AP Inter 2nd Year Physics Important Questions Chapter 10 Alternating Current

Students get through AP Inter 2nd Year Physics Important Questions 10th Lesson Alternating Current which are most likely to be asked in the exam.

AP Inter 2nd Year Physics Important Questions 10th Lesson Alternating Current

Very Short Answer Questions

Question 1.
A transformer converts 200 V ac into 2000 V ac. Calculate the number of turns in the ‘ secondary if the primary has 10 turns. [T.S. Mar. 16]
Answer:
\(\frac{\mathrm{V}_{\mathrm{s}}}{\mathrm{V}_{\mathrm{p}}}=\frac{\mathrm{N}_{\mathrm{s}}}{\mathrm{N}_{\mathrm{p}}}\)
Vp = 200V, Vs = 2000V, Np = 10
Ns = \(\frac{\mathrm{V}_{\mathrm{s}}}{\mathrm{V}_{\mathrm{p}}}\) × Np = \(\frac{2000}{200}\) × 10
Ns = 100.

Question 2.
What type of transformer is used in a 6V bed lamp ? [A.P. Mar. 17]
Answer:
Step down transformer is used in 6V bed lamp.

AP Inter 2nd Year Physics Important Questions Chapter 10 Alternating Current

Question 3.
What is the phenomenon involved in the working of a transformer ? [Mar. 16(A.P.) Mar. 14]
Answer:
Transformer works on the principle of mutual induction.

Question 4.
What is transformer ratio ?
Answer:
The ratio of secondary e.m.f to the primary e.m.f. (or) number of turns in secondary to the number of turns in the primary is called the transformer ratio.
Transformer ratio = \(\frac{\mathrm{V}_{\mathrm{s}}}{\mathrm{V}_{\mathrm{p}}}=\frac{\mathrm{N}_{\mathrm{s}}}{\mathrm{N}_{\mathrm{p}}}\)

Question 5.
Write the expression for the reactance of i) an inductor and (ii) a capacitor.
Answer:

  1. Inductive reactance (XL) = ωL
  2. Capacitive reactance (XC) = \(\frac{1}{\omega C}\)

Question 6.
What is the phase difference between A.C emf and current in the following: Pure resistor, pure inductor and pure capacitor. [T.S. Mar. 15]
Answer:

  1. In pure resistor A.C. e.m.f and current are in phase with each other.
  2. In pure inductor, current lags behind the e.m.f. by an angle of \(\frac{\pi}{2}\) (or) 90°.
  3. In pure capacitor, current leads the e.m.f by an angle \(\frac{\pi}{2}\) (or) 90°.

AP Inter 2nd Year Physics Important Questions Chapter 10 Alternating Current

Question 7.
Define power factor. On which factors does power factor depend ?
Answer:
The ratio of true power and apparent power (virtual power) in an a.c circuit is called as power factor of the circuit.
Power factor (cosΦ) = \(\frac{\mathrm{P}}{\mathrm{P}_{\mathrm{rms}}}\) [∵ Prms = Vrms Irms]
Power factor depends on r.m.s voltage, r.m.s current and average power (P).

Question 8.
What is meant by wattless component of current ?
Answer:
Average power (P) = Vrms(Irms sinΦ) cos\(\frac{\pi}{2}\)
The average power consumed in the circuit due to (Irms sinΦ) component of current is zero. This component of current is known as wattless current. (Irms sinΦ) is the wattless component of current.

Question 9.
When does a LCR series circuit have minimum impedance ?
Answer:
In LCR series circuit, Impendence (Z) = \(\sqrt{R^2+\left(\frac{1}{\omega C}-\omega L\right)^2}\)
At a particular frequency, ωL = \(\frac{1}{\omega C}\)
The impedance is minimum (Z = R)
This frequency is called resonant frequency.

AP Inter 2nd Year Physics Important Questions Chapter 10 Alternating Current

Question 10.
What is the phase difference between voltage and current when the power factor in LCR series circuit is unity ?
Answer:
In LCR series circuit power factor (cosΦ) = 1
Phase difference between voltage and current is zero. (Φ = 0)

Short Answer Questions

Question 1.
State the principle on which a transformer works. Describe the working of a transformer with necessary theory.
Answer:
Transformer is a device to convert a low alternating current of high voltage into high alternating current of low voltage and vice versa.
Principle : It works on the principle of mutual induction between two coils.
Working : When an alternating emf is applied across the primary coil, the input voltage changes with time. Hence the magnetic flux through the primary also changes with time.
AP Inter 2nd Year Physics Important Questions Chapter 10 Alternating Current 1
This changing magnetic flux will be linked with secondary through the core. An emf is induced in the secondary.

Theory: Let N1 and N2 be the number of turns in the primary and secondary. Let VP and VS be the emf s across the primary and secondary.
\(\frac{V_S}{V_p}=\frac{\text { Output emf }}{\text{Input emf}}=\frac{-N_2 \frac{d \phi}{d t}}{-N_1 \frac{d \phi}{d t}}=\frac{N_2}{N_1}\)
∴ \(\frac{\mathrm{V}_{\mathrm{S}}}{\mathrm{V}_{\mathrm{P}}}=\frac{\mathrm{N}_2}{\mathrm{~N}_1}\) = Transformer ratio
Efficiency of transformer :
It is the ratio of output power to the input power.
η = \(\frac{\text { Outputpower }}{\text { Input power }}\) × 100

Problems

Question 1.
A light bulb is rated at 100W for a 220 V supply. Find
(a) the resistance of the bulb;
(b) the peak voltage of the source; and
(c) the rms current through the bulb. [A.P. Mar. 15]
Solution:
(a) We are given P = 100 W and V = 220V The resistance of the bulb is
R = \(\frac{\mathrm{V}^2}{\mathrm{P}}=\frac{(220 \mathrm{~V})^2}{100 \mathrm{~W}}\) = 484 Ω

(b) The peak voltage of the source is υm = \(\sqrt{2}\)V = 311 V

(c) Since, P = 1 V
I = \(\frac{\mathrm{P}}{\mathrm{V}}=\frac{100 \mathrm{~W}}{220 \mathrm{~V}}\) = 0.450 A.

AP Inter 2nd Year Physics Important Questions Chapter 10 Alternating Current

Question 2.
A pure inductor of 25.0 mH is connected to a source of 220 V. Find the inductive reactance and rms current in the circuit if the frequency of the source is 50 Hz.
Solution:
The inductive reactance,
XL = 2πvL = 2 × 3.14 × 50 × 25 × 10-3 = 7.85 Ω
The rms current in the circuit is I = \(\frac{\mathrm{V}}{\mathrm{X}_{\mathrm{L}}}=\frac{220 \mathrm{~V}}{7.85 \Omega}\) = 28 A

Question 3.
The instantaneous current and instantaneous voltage across a series circuit containing resistance and inductance are given by i = \(\sqrt{2}\) sin (100t – π/4)A and υ = 40 sin (100t) V. Calculate the resistance ?
Solution:
i = \(\sqrt{2}\) sin (100t – π/4)A (∵i = i0sin(ωt – Φ))
υ = 40 sin(100t)V (∵ V = V0sin(ωt ))
i0 = \(\sqrt{2}\) , V0 = 40, ω = 100, Φ = π/4
R = \(\frac{\mathrm{V}_0}{\mathrm{i}_0}\) cosΦ = \(\frac{40}{\sqrt{2}}\) cos\(\frac{\pi}{4}\), R = \(\frac{40}{\sqrt{2}} \times \frac{1}{\sqrt{2}}\), R = 20 Ω

Question 4.
In an AC circuit, a condenser, a resistor and a pure inductor are connected in series across an alternator (AC generator). If the voltages across them are 20 V, 35 V and 20 V respectively, find the voltage supplied by the alternator.
Solution:
VC = 20V, VR = 35V, VL = 20V
V = \(\sqrt{V_R^2+\left(V_L^2-V_C^2\right)}\) ; V = \(\sqrt{(35)^2+\left(20^2-20^2\right)}\) ; V = \(\sqrt{35^2}\); V = 35 Volt.

AP Inter 2nd Year Physics Important Questions Chapter 10 Alternating Current

Question 5.
What is step up transformer ? How it differs from step down transformer ?
Solution:
The ratio of number of turns in the secondary coil to the number of turns in the primary coil is called transformer ratio.
T = \(\frac{N_S}{N_p}=\frac{\text {No. of turns in the secondary }}{\text {No. of turns in the primary }}\)
If NS > NP, then the transformer is called step up transformer.
If NS < NP, then the transformer is called step down transformer.

Textual Examples

Question 1.
A light bulb is rated at 100W for a 220 V supply. Find
(a) the resistance of the bulb;
(b) the peak voltage of the source; and
(c) the rms current through the bulb. [A.P. Mar. 15]
Solution:
(a) We are given P = 100 W and V = 220V.
The resistance of the bulb is
R = \(\frac{\mathrm{V}^2}{\mathrm{P}}=\frac{(220 \mathrm{~V})^2}{100 \mathrm{~W}}\) = 484 Ω

(b) The peak voltage of the source is υm = \(\sqrt{2}\)V = 311 V

(c) Since, P = 1 V
I = \(\frac{\mathrm{P}}{\mathrm{V}}=\frac{100 \mathrm{~W}}{220 \mathrm{~V}}\) = 0.450 A.

Question 2.
A pure inductor of 25.0 mH is connected to a source of 220 V. Find the inductive reactance and rms current in the circuit if the frequency of the source is 50 Hz.
Solution:
The inductive reactance,
XL = 2πvL = 2 × 3.14 × 50 × 25 × 10-3 = 7.85 Ω
The rms current in the circuit is I = \(\frac{\mathrm{V}}{\mathrm{X}_{\mathrm{L}}}=\frac{220 \mathrm{~V}}{7.85 \Omega}\) = 28 A

AP Inter 2nd Year Physics Important Questions Chapter 10 Alternating Current

Question 3.
A lamp is connected in series with a capacitor. Predict your observations for dc and ac connections. What happens in each case if the capacitance of the capacitor is reduced ?
Solution:
When a dc source is connected to a capacitor, the capacitor gets charged and after charging no current flows in the circuit and the lamp will not glow. There will be no change even if C is reduced. With ac source, the capacitor offers capacitative reactance (1/ωC) and the current flows in the circuit. Consequently, the lamp will shine. Reducing C will increase reactance and the lamp will shine less brightly than before.

Question 4.
A 15.0 μF capacitor is connected to a 220 V, 50 Hz source. Find the capacitive reactance and the current (rms and peak) in the circuit. If the frequency is doubled, what happens to the capacitive reactance and the current ?
Solution:
The capacitive reactance is
Xc = \(\frac{1}{2 \pi v C}=\frac{1}{2 \pi(50 \mathrm{~Hz})\left(15.0 \times 10^{-6} \mathrm{~F}\right)}\)
= 212 Ω
The rms current is I = \(\frac{\mathrm{V}}{\mathrm{X}_{\mathrm{C}}}=\frac{220 \mathrm{~V}}{212 \Omega}\)
= 1.04 A
The peak current is
im = \(\sqrt{2}\)I = (1.41)(1.04A) = 1.47A
This current oscillates between + 1.47A and – 1.47 A, and is ahead of the voltage by π/2.
If the frequency is doubled, the capacitive reactance is halved and consequently, the current is doubled.

Question 5.
A light bulb and an open coil inductor are connected to an ac source through a key as shown in the figure.
AP Inter 2nd Year Physics Important Questions Chapter 10 Alternating Current 2
The switch is closed and after sometime, an iron rod is inserted into the interior of the inductor. The glow of the light bulb (a) increases; (b) decreases; (c) is unchanged, as the iron rod is inserted. Give your answer with reasons.
Solution:
As the iron rod is inserted, the magnetic field inside the coil magnetizes the iron increasing the magnetic field inside it. Hence, the inductance of the coil increases. Consequently, the inductive reactance of the coil increases. As a result, a larger fraction of the applied ac voltage appears across the inductor, leaving less voltage across the bulb. Therefore, the glow of the light bulb decreases.

AP Inter 2nd Year Physics Important Questions Chapter 10 Alternating Current

Question 6.
A resistor of 200Ω and a capacitor of 15.0 μF are connected in series to a 220V, 50 Hz ac source,
(a) Calculate the current in the circuit;
(b) Calculate the voltage (rms) across the resistor and the capacitor. Is the algebraic sum of these voltages more than the source voltage ? If yes, resolve the paradox.
Solution:
Given
R = 200Ω. C = 15.0 μF = 15.0 × 10-6F
V = 220V, v = 50Hz
(a) In order to calculate the current, we need the impedance of the circuit. It is
Z = \(\sqrt{\mathrm{R}^2+\mathrm{X}_{\mathrm{C}}^2}=\sqrt{\mathrm{R}^2+(2 \pi v C)^{-2}}\)
= \(\sqrt{(200 \Omega)^2+\left(2 \times 3.14 \times 50 \times 10^{-6} \mathrm{~F}\right)^{-2}}\)
= \(\sqrt{(200 \Omega)^2+(212 \Omega)^2}\) = 291.5Ω
Therefore, the current in the circuit is
I = \(\frac{\mathrm{V}}{\mathrm{Z}}=\frac{220 \mathrm{~V}}{291.5 \Omega}\) = 0.755A

(b) Since the current is the same throughout the circuit, we have
VR = IR = (0.755 A) (200Ω) = 151V
VC = IXC = (0.755A) (212.3Ω) = 160.3V
The algebraic sum of the two voltages, VR and VC is 311.3 V which is more than the source voltage of 220 V. How to resolve this paradox ? As you have learnt in the text, the two voltages are not in the same phase. Therefore, they cannot be added like ordinary numbers. The two voltages are out of phase by ninety degrees. Therefore, the total of these voltages must be obtained using the Pythagorean theorem:
VR+C = \(\sqrt{\mathrm{V}_{\mathrm{R}}^2+\mathrm{V}_{\mathrm{C}}^2}\) = 220 V
Thus, if the phase difference between two voltages is properly taken into account, the total voltage across the resistor and the capacitor is equal to the voltage of the source.

Question 7.
a) For circuits used for transporting electric power, a low power factor implies large power loss in transmission. Explain.
b) Power factor can often be improved by the use of a capacitor of appropriate capacitance in the circuit. Explain.
Solution:
a) We know that P = IV cosΦ where cosΦ is the power factor. To supply a given power at a given voltage, if cosΦ is small, we have to increase current accordingly. But this will lead to large power loss (IR) in transmission.

b) Suppose in a circuit, current I lags the voltage by an angle Φ.
Then power factor cosΦ = R/Z
We can improve the power factor (tending to 1) by making Z tend to R. Let us understand, with the help of a phasor diagram in the figure how this can be achieved. Let us resolve I into two components, IP
AP Inter 2nd Year Physics Important Questions Chapter 10 Alternating Current 3
along the applied voltage V and Iq perpendicular to the applied voltage. Iq is called the wattless component since corresponding to this component of current, there is no power loss. IP is known as the power component because it is in phase with the voltage and corresponds to power loss in the circuit.

It’s clear from this analysis that if we want to improve power factor, we must completely neutralize the lagging wattless current Iq by an equal leading wattless current I’q. This can be done by connecting a capacitor of appropriate value in parallel so that Iq and I’q cancel each other and P is effectively IP V.

AP Inter 2nd Year Physics Important Questions Chapter 10 Alternating Current

Question 8.
A sinusoidal voltage of peak value 283 V , and frequency 50 Hz is applied to a series LCR circuit in which R = 3Ω. L = 25.48 mH. and C = 796μF. Find
(a) the impedance of the circuit;
(b) the phase difference between the voltage across the source and the current;
(c) the power dissipated in the circuit; and
(d) the power factor.
Solution:
a) To find the impedance of the circuit, we first calculate XL and XC.
XL = 2πvL
= 2 × 3.14 × 50 × 25.48 × 10-3Ω = 8Ω
XC = \(\frac{1}{2 \pi v \mathrm{C}}\)
= \(\frac{1}{2 \times 3.14 \times 50 \times 796 \times 10^{-6}}\) = 4Ω
Therefore,
z = \(\sqrt{\mathrm{R}^2+\left(\mathrm{X}_{\mathrm{L}}-\mathrm{X}_{\mathrm{C}}\right)^2}=\sqrt{3^2+(8-4)^2}\)
= 5Ω

b) Phase difference, Φ = tan-1\(\frac{\mathrm{X}_{\mathrm{C}}-\mathrm{X}_{\mathrm{L}}}{\mathrm{R}}\)
= tan-1\(\left(\frac{4-8}{3}\right)\) = -53.1°

c) The power dissipated in the circuit is
P = I2R
AP Inter 2nd Year Physics Important Questions Chapter 10 Alternating Current 4
Therefore, P = (40A)2 × 3Ω = 4800W
= 4.8 kW

d) Power factor = cos Φ = cos 53.1° = 0.6.

Question 9.
Suppose the frequency of the source in the previous example can be varied,
(a) What is the frequency of the source at which resonance occurs ?
(b) Calculate the impedance, the current, and the power dissipated at the resonant condition.
Solution:
(a) The frequency at which the resonance occurs is
ω0 = \(\frac{1}{\sqrt{\mathrm{LC}}}=\frac{1}{\sqrt{25.48 \times 10^{-3} \times 796 \times 10^{-6}}}\)
= 222.1 rad/s
vr = \(\frac{\omega_0}{2 \pi}=\frac{221.1}{2 \times 3.14}\) Hz = 35.4Hz

b) The impedance Z at resonant condition is equal to the resistance
Z = R = 3Ω
The rms current at resonance is ,
as V = \(\frac{v_{\mathrm{m}}}{\sqrt{2}}\)
I = \(\frac{\mathrm{V}}{\mathrm{Z}}=\frac{\mathrm{V}}{\mathrm{R}}=\left(\frac{283}{\sqrt{2}}\right) \frac{1}{3}\) = 66.7 A
The power dissipated at resonance is
P = I2 × R = (66.7)2 × 3 = 13.35 kW
You can see that in the present case, power dissipated at resonance is more than the power dissipated in Example 8.

AP Inter 2nd Year Physics Important Questions Chapter 10 Alternating Current

Question 10.
At an airport, a person is made to walk through the doorway of a metal detector, for security reasons. If she/he is carrying anything made of metal, the metal detector emits a sound. On what principle does this detector work ?
Solution:
The metal detector works on the principle of resonance in ac circuits. When you walk through a metal detector, you are, in fact, walking through a coil of many turns. The coil is connected to a capacitor tuned so that the circuit is in resonance. When you walk through with metal in your pocket, the impedance of the circuit changes – resulting in significant change in current in the circuit. This change in current is detected and the electronic circuitry causes a sound to be emitted as an alarm.

Question 11.
Show that in the free oscillations of an LC circuit, the sum of energies stored in the capacitor and the inductor is constant in time.
Solution:
Let q0 be the initial charge tin a capacitor. Let the charged capacitor be connected to an inductor of inductance L. this LC circuit will sustain an oscillation with frequency
\(\omega\left(2 \pi v=\frac{1}{\sqrt{\mathrm{LC}}}\right)\)
At an instant t, charge q on the capacitor and the current i are given by :
q(t) = q0 cos ωt
i(t) = -q0 co sin ωt
Energy stored in the capacitor at time ‘t’ is
UE = \(\frac{1}{2}\) C V2 = \(\frac{1}{2} \frac{\mathrm{q}^2}{\mathrm{C}}=\frac{\mathrm{q}_0^2}{2 \mathrm{C}}\) cos2 (ωt)
Energy stored in the inductor at time ‘t’ is
UM = \(\frac{1}{2}\) L i2
= \(\frac{1}{2}\) L q02 ω2sin2 (ωt)
= \(\frac{Q_0^2}{2 C} \sin ^2(\omega t)\) [∵  ω = 1/ \(\sqrt{L C}\)]
Sum of energies
UE + UM = \(\frac{\mathrm{q}_0^2}{2 \mathrm{C}}\) [cos2 ωt + sin2ωt) = \(\frac{\mathrm{q}_0^2}{2 \mathrm{C}}\)
This sum is constant in time as q0 and C, both are time-independent.

AP Inter 2nd Year Physics Important Questions Chapter 11 Electromagnetic Waves

Students get through AP Inter 2nd Year Physics Important Questions 11th Lesson Electromagnetic Waves which are most likely to be asked in the exam.

AP Inter 2nd Year Physics Important Questions 11th Lesson Electromagnetic Waves

Very Short Answer Questions

Question 1.
Give any one use of infrared rays. [T.S. Mar. 17; A.P. Mar. 16]
Answer:

  1. Infrared radiation plays an important role in maintaining the Earth warm.
  2. Infrared lamps are used in physical therapy.
  3. Infrared detectors are used in Earth Satellites.
  4. These are used in taking photographs during the conditions of fog, smoke etc.

Question 2.
How are infrared rays produced ? How they can be detected ?
Answer:
Infrared rays can be produced by vibrations of atoms and molecules. These waves can be detected by Thermopile, Bolometer, IR photographic film.

AP Inter 2nd Year Physics Important Questions Chapter 11 Electromagnetic Waves

Question 3.
How are radio waves produced ? How can they detected ?
Answer:
Radio waves can be produced by rapid acceleration and deceleration of electrons in aerials (conductors). These can be detected by receivers of aerials.

Question 4.
If the wave length of E.M radiation is doubled, what happens to the energy of photon ? [IPE 2016 (TS)]
Answer:
If the wave length of electromagnetic radiation is doubled, then energy will be halved because energy is inversely proportional to.wavelength of electromagnetic waves.
E = hυ = hc/λ ⇒ E ∝ 1/λ (∵ hc is a constant)

Question 5.
What is the principle of production of electromagnetic waves ?
Answer:
If the charge is accelerated both the magnetic field and electric field will change with Space and time, then electromagnetic waves are produced.

Question 6.
What is the ratio of speed of infrared rays and ultraviolet rays in vaccum ?
Answer:
The ratio of speed of infrared rays and ultraviolet rays in vacuum is 1 : 1.
All electromagnetic waves travel with same speed 3 × 108 m /s in vaccum.

AP Inter 2nd Year Physics Important Questions Chapter 11 Electromagnetic Waves

Question 7.
What is the relation between the amplitudes of the electric and magnetic fields in free space for an electromagnetic wave ?
Answer:
E0 = CB0
Where E0 = Amplitude of electric field.
B0 = Amplitude of magnetic field.
C = velocity of light.

Question 8.
What are the applications of microwaves ? [A.P. Mar. 17; T.S. Mar. 15]
Answer:

  1. Microwaves are used in Radars.
  2. Microwaves are used for cooking purposes.
  3. A radar using microwave can help in detecting the speed of automobile while in motion.

Question 9.
Microwaves are used in Radars, why ? [Mar. 14]
Answer:
As microwaves are of smaller wavelengths, hence they can be transmitted as a beam signal in a particular direction. Microwaves do not bend around the comers of any obstacle coming in their path.

Question 10.
Give two uses of infrared rays.
Answer:

  1. Infrared rays are used for producing dehydrated fruits.
  2. They are used in the secret writings on the ancient walls.
  3. They are used in green houses to keep the plants warm.

AP Inter 2nd Year Physics Important Questions Chapter 11 Electromagnetic Waves

Question 11.
How are microwaves produced ? How can they detected ? [A.P. Mar. 16; IPE 15]
Answer:
Microwaves can be produced using Klystron valve or Magnetrons.
Microwaves can be detected using point contact diodes.

Question 12.
The charging current for a capacitor is 0.6 A. What is the displacement current across its plates ?
Answer:
i = charging current for a capacitor = 0.6 A
i = id = ε0 = \(\frac{\mathrm{d} \phi_{\mathrm{B}}}{\mathrm{dt}}\)
∴ i = id = 0.6 A
∴ Displacement current (id) = 0.6 A.

Question 13.
What physical quantity is the same for X-rays of wavelength 10-10m, red light of wavelength 6800 Å and radiowaves of wavelength 500in ?
Answer:
The speed in vaccum is same for all the given wavelengths, which is 3 × 108 m/s.

Question 14.
A radio can tune into any station in the 7.5 MHz to 12MHz band. What is the corresponding wavelength band ?
Answer:
λ1 = \(\frac{3 \times 10^8}{7.5 \times 10^6}\) = 40 m
λ2 = \(\frac{3 \times 10^8}{12 \times 10^6}\) = 25 m
Thus wavelength band is 40m to 25m.

AP Inter 2nd Year Physics Important Questions Chapter 11 Electromagnetic Waves

Question 15.
The amplitude of the magnetic field part of a harmonic electromagnetic wave in vaccum is B0 = 510 nT. What is the amplitude of the electric field part of the wave ?
Answer:
Here, B0 = 510nT = 510 × 10-9T
E0 = CB0 = 3 × 108 × 510 × 10-9 = 153 NC-1.

Question 16.
Define displacement current
Answer:
Displacement current (Id) is equal to ε0 times to the rate of change of electric flux. Displacement current is not the current produced due to charge carried. But it is due to varying electric flux. It is the current in the sense that it produces a magnetic field.
Id = ε0 \(\frac{\mathrm{d} \phi_{\varepsilon}}{\mathrm{dt}}\)

Short Answer Questions

Question 1.
State six characteristics of electromagnetic waves.
Answer:
Characteristics of electromagnetic waves :

  1. Electromagnetic waves are produced by accelerated charges.
  2. Electromagnetic waves are transverse in nature.
  3. Electromagnetic waves donot require material medium for their propagation.
  4. Electromagnetic waves obey principle of superposition of waves.
  5. Velocity of E.M waves in vaccum depends on permittivity and permeability of free space.
  6. Electromagnetic waves carry energy and momentum.
  7. Electromagnetic waves exert pressure when they strike a surface.

AP Inter 2nd Year Physics Important Questions Chapter 11 Electromagnetic Waves

Question 2.
What is Greenhouse effect and its contribution towards the surface temperature of earth ?
Answer:
Green house effect: Temperature of the earth increases due to the radiation emitted by
the earth is trapped by atmospheric gases like CO2, CH4, N2, chlorofluoro carbons etc., is called green house effect.

  1. Radiation from the sun enters the atmosphere and heat the objects on the earth. These heated objects emit infrared rays.
  2. These rays are reflected back to Earth’s surface and trapped in the Earth’s atmosphere. Due to this temperature of the earth increases.
  3. The layers of carbon dioxide (CO2) and low lying clouds prevent infrared rays to escape Earth’s atmosphere.
  4. Since day-by-day the amount of carbondioxide in the atmosphere increases, more • infrared rays are entrapped in the atmosphere.
  5. Hence the temperature of the Earth’s surface increases day by day.

Problems

Question 1.
A plane electromagnetic wave travels in vaccum along z-direction. What can you say about the directions of its electric and magnetic field vectors ? If the frequency of the wave is 30 MHz. What is its wavelength ?
Answer:
Electric and magnetic fields \(\overline{\mathrm{E}}\) and \(\overline{\mathrm{B}}\) of an electromagnetic wave must be perpendicular to the propagation of electromagnetic wave. Hence they lie in X – Y plane mutually perpendicular to each other.
Frequency of wave, v = 30MHz = 30 × 106Hz.; Velocity of light, C = 3 × 108m/s
Wavelength of the wave, λ = \(\frac{C}{V}=\frac{3 \times 10^8}{30 \times 10^6}\) = 10m

AP Inter 2nd Year Physics Important Questions Chapter 11 Electromagnetic Waves

Question 2.
A charged particle oscillates about its mean equilibrium position with a frequency of 109 Hz. What is the frequency of the electromagnetic waves produced by the oscillator ?
Answer:
According to Maxwell, a charged particle oscillating with a frequency produces electro-magnetic waves of same frequency. Hence frequency of EM waves produced is, 109Hz.

Textual Examples

Question 1.
A parallel plate capacitor with circular plates of radius 1 m has a capacitance of 1 nF. At t = 0, It is connected for charging in series with a resistor R = 1 M Q across a 2V battery (fig)- Calculate the magnetic field at a point P halfway between the centre and the periphery of the plates, after t = 10-3 s. (The charge on the capacitor at time τ is q (t) = CV [1 – exp (-t/τ)], where the time constant τ is equal to CR.).
Solution:
The time constant of the CR circuit is τ = CR = 10-3s. Then, we have
q(t) = CV [1 – exp (-t/τ) ]
= 2 × 10-9 [1 – exp (-t /10-3)
The electric field in between the plates at time t is
E = \(\frac{q(t)}{\varepsilon_0 A}=\frac{q}{\pi \varepsilon_0}\); A = π (1)2 m2 = area of the plates.
AP Inter 2nd Year Physics Important Questions Chapter 11 Electromagnetic Waves 1
Consider now a circular loop of radius (1/2)m parallel to the plates passing through R The magnetic field B at all points on the loop is along the loop arid of the same value.
The flux ΦE through this loop is
The flux ΦE = E × area of the loop
= E × π × (\(\frac{1}{2}\))2 = \(\frac{\pi \mathrm{E}}{4}=\frac{\mathrm{q}}{4 \varepsilon_0}\)
The displacement current
id = e0 \(\frac{\mathrm{d} \phi_{\mathrm{E}}}{\mathrm{dt}}\) = \(\frac{1}{4} \frac{\mathrm{dq}}{\mathrm{dt}}\) = 0.5 × 10-6 exp (-1)
at t = 10-3s. Now, applying Ampere-Maxwell law to the loop, we get
B × 2π × (\(\frac{1}{2}\)) = m0.(ic + id) = m0(0 + id) = 0.5 × 10-6 m0 exp(-1)
or, B = 0.74 × 10-13 T.

Question 2.
A plane electromagnetic wave of frequency 25 MHz travels in free space along the x – direction. At a particular point in space and time, E = \(6.3 \hat{\mathbf{j}}\) V/m. What is B at the point ?
Solution:
Using Eq. B0 = [E0/c] the magnitude of B is
B = \(\frac{\mathrm{E}}{\mathrm{c}}\)
= \(\frac{6.3 \mathrm{~V} / \mathrm{m}}{3 \times 10^8 \mathrm{~m} / \mathrm{s}}\) = 2.1 × 10-8 T
To find the direction, we note that E is along y-direction and the wave propagates along x- axis. Therefore, B should be in a direction perpendicular to both x- and y-axes. Using vector algebra, E × B should be along x-direction. Since, \((+\hat{\mathrm{j}})\) × \((+\hat{\mathrm{k}})\) = i, B is along the z-direction. Thus. B = 2.1 × 10-8 \(\hat{\mathrm{k}}\) T

AP Inter 2nd Year Physics Important Questions Chapter 11 Electromagnetic Waves

Question 3.
The magnetic field in a plane electromagnetic wave is given by
By = 2 × 10-7 sin (0.5 × 103 × + 1.5 × 1011 t)T.
a) What is the wavelength and frequency of the wave ?
b) Write an expression for the electric field.
Solution:
a) Comparing the given equation with
By = B0 sin \(\left[2 \pi\left(\frac{\mathrm{x}}{\lambda}+\frac{\mathrm{t}}{\mathrm{T}}\right)\right]\)
We get, λ = \(\frac{2 \pi}{0.5 \times 10^3}\) m = 1.26 cm,
and \(\frac{1}{\mathrm{~T}}\) = v = (1.5 × 1011)/2π = 23.9 GHz

b) E0 = B0C = 2 × 10-7 T × 3 × 108 m/s = 6 × 101 V/m
The electric field component is perpendicular to the direction of propagation and the di-rection of magnetic field. Therefore, the electric field component along the z-axis is obtained as Ez = 60 sin (0.5 × 103x + 1.5 × 1011t) V/m.

Question 4.
Light with an energy flux of 18 W/cm2 falls on a non reflecting surface at normal incidence. If the surface has an area of 20 cm2 find the average force exerted on the surface during a 30 minute time span.
Solution:
The total energy falling on the surface is
U = (18 W/cm2) × (20 cm2) × (30 × 60)
= 6.48 × 105 J
Therefore, the total momentum delivered (for complete absorption) is
P = \(\frac{\mathrm{U}}{\mathrm{C}}=\frac{6.48 \times 10^5 \mathrm{I}}{3 \times 10^8 \mathrm{m} / \mathrm{s}}\) = 2.16 × 10-3 kg m/s
The average force exerted on the surface is
F = \(\frac{\mathrm{p}}{\mathrm{t}}=\frac{2.16 \times 10^{-3}}{0.18 \times 10^4}\) = 1.2 × 10-6 N

AP Inter 2nd Year Physics Important Questions Chapter 11 Electromagnetic Waves

Question 5.
Calculate the electric and magnetic fields produced by the radiation coming from a 100 W bulb at a distance of 3m. Assume that the efficiency of the bulb is 2.5% and it is a point source.
Solution:
The bulb, as a point source, radiates light in all directions uniformly. At a distance of 3m, the surface area of the surrounding sphere is A = 4 πr2 = 4π (3)2 = 113m2
The intensity at this distance is
I = \(\frac{\text { Power }}{\text { Area }}=\frac{100 \mathrm{~W} \times 2.5 \%}{113 \mathrm{~m}^2}\) = 0.022 W.m2
Half of this intensity is provided by the electric field and half by the magnetic field.
\(\frac{1}{2} I=\frac{1}{2}\left(\varepsilon_0 \mathrm{E}_{\mathrm{rms}}^2 \mathrm{C}\right)\)
= \(\frac{1}{2}\) (0.022 W/m2)
Erms = \(\sqrt{\frac{0.022}{\left(8.85 \times 10^{-12}\right)\left(3 \times 10^8\right)}}\) V/m = 2.9 V/m
The value of E found above is the root mean square value of the electric field. Since the electric field in a light beam is sinusoidal, the peak electric field, E0 is
E0 = \(\sqrt{2}\)Erms = \(\sqrt{2}\) × 2.9 V/m
= 4.07 V/m
Thus, you see that the electric field strength of the light that you use for reading is fairly large. Compare it with electric field strength of TV or FM waves, which is of the order of a few microvolts per metre.
Now, let us calculate the strength of the magnetic field. It is
Brms = \(\frac{E_{\mathrm{rms}}}{\mathrm{C}}=\frac{2.9 \mathrm{Vm}^{-1}}{3 \times 10^8 \mathrm{~ms}^{-1}}\) = 9.6 × 10-9 T.
Again, since tbs field in the light beam is sinusoidal, the peak magnetic field is B0 = \(\sqrt{2}\) Brms = 1.4 × 10-8 T. Note that although the energy in the magnetic field is equal to the energy in the electric field, the magnetic field strength is evidently very weak.

AP Inter 2nd Year Physics Important Questions Chapter 12 Dual Nature of Radiation and Matter

Students get through AP Inter 2nd Year Physics Important Questions 12th Lesson Dual Nature of Radiation and Matter which are most likely to be asked in the exam.

AP Inter 2nd Year Physics Important Questions 12th Lesson Dual Nature of Radiation and Matter

Very Short Answer Questions

Question 1.
What are “cathode rays” ? [A.P. Mar. 17]
Answer:
Cathode rays are streams of fast moving electrons or negatively charged particles.

Question 2.
What important fact did Millikan’s experiment establish?
Answer:
Millikan’s experiment established that electric charge is quantised. That means the charge on any body (oil drop) is always an integral multiple of charge of an electron, i.e., Q = ne.

AP Inter 2nd Year Physics Important Questions Chapter 12 Dual Nature of Radiation and Matter

Question 3.
What is “work function” ? [T.S. Mar. 17, 15]
Answer:
The minimum energy required to liberate an electron from photometal surface is called work function, Φ0.

Question 4.
What is “photoelectric effect” ?
Answer:
When light of sufficierit energy is incident on the photometal surface, electrons are emitted. This phenomenon is called photoelectric effect.

Question 5.
Give examples of “photosensitive substances”. Why are they called so ?
Answer:
Examples of photosensitive substances are Li, Na, K, Rb and Cs etc.
The work function of alkali metals is very low. Even the ordinary visible light, alkali metals can produce photoelectric emission. Hence they are called photosensitive substances.

Question 6.
Write down Einstein’s photoelectric equation. [A.P. Mar. 15]
Answer:
Einstein’s photoelectric equation is given by Kmax = \(\frac{1}{2}\) mv2max = hυ – Φ0

AP Inter 2nd Year Physics Important Questions Chapter 12 Dual Nature of Radiation and Matter

Question 7.
Write down de-Broglie’s relation and explain the terms therein. [T.S. & A.P. Mar. 16]
Answer:
The de-Broglie wave length (λ) associated with a moving particle is related to its momentum (p) is λ = \(\frac{h}{p}=\frac{h}{m v}\), where h is planck’s constant.

Question 8.
State Heisenberg’s Uncertainity Principle. [Mar. 14]
Answer:
Uncertainity principle states that “it is impossible to measure both position (∆x) and momentum of an electron (∆p) [or any other particle] at the same time exactly”, i.e., ∆x . ∆p ≈ h where ∆x is uncertainty in the specification of position and ∆p is uncertainty in the specification of momentum.

Question 9.
The photoelectric cut off voltage in a certain experiment is 1.5 V. What is the maximum kinetic energy of photoelectrons emitted ? [Mar. 11]
Answer:
Cut off voltage, V0 = 1.5 V; Maximum kinetic energy, (KE)max = eV0 = e × 1.5 = 1.5eV

Question 10.
An electron, an α-particle and a proton have the same kinetic energy. Which of these particles has the shortest de Broglie wavelength ? [T.S. Mar. 15]
Answer:
For a particle,
de Broglie wavelength, λ = h/p
Kinetic energy, K = p2/2m
Then, λ = h/\(\sqrt{2 \mathrm{mK}}\)
For the same kinetic energy K, the de Broglie wavelength associated with the particle is inversely proportional to the square root of their masses.A proton AP Inter 2nd Year Physics Important Questions Chapter 12 Dual Nature of Radiation and Matter 1 is 1836 times massive than an electron and an a-particle \(\left({ }_2^4 \mathrm{He}\right)\) four times that of a proton.
Hence, α-particle has the shortest de Broglie wavelength.

AP Inter 2nd Year Physics Important Questions Chapter 12 Dual Nature of Radiation and Matter

Qeustion 11.
What is-the de Broglie wavelength associated with an electron, accelerated through a potential difference of 100 volts ? [A.P. Mar. 15]
Solution:
Accelerating potential V = 100 V The de Broglie wavelngth λ is
λ = h/p = \(\frac{1.227}{\sqrt{\mathrm{V}}}\) nm
λ = \(\frac{1.227}{\sqrt{100}}\) nm = 0.123 nm
The de Broglie wavelength associated with an electron in this case is of the order of X-ray wavelengths.

Question 12.
If the wave length of electro magnetic radiation is doubled, what happens to energy of photon ? [IPE 2015 (TS)]
Answer:
E = \(\frac{\mathrm{hc}}{\lambda}\) ⇒ E ∝ \(\frac{1}{\lambda}\), since wave length of photon is doubled, its energy becomes halved.

Question 13.
In an experiment on photoelectric effect, the slope of the cut-off voltage versus frequency of incident light is found to be 4.12 × 10-15 V s. Calculate the value of Planck’s constant.
Solution:
Given, slope of graph tan θ = 4.12 × 10-15 V – s;
AP Inter 2nd Year Physics Important Questions Chapter 12 Dual Nature of Radiation and Matter 2
For slope of graph, tan θ = \(\frac{\mathrm{V}}{\mathrm{v}}\)
We know that hv = eV
\(\frac{\mathrm{V}}{\mathrm{v}}=\frac{\mathrm{h}}{\mathrm{e}} \Rightarrow \frac{\mathrm{h}}{\mathrm{e}}\) = 4.12 × 10-15; h = 4.12 × 10-15 × 1.6 × 10-19 = 6.592 × 10-34 J – s.

Question 14.
The energy flux of sunlight reaching the surface of the earth is 1.388 × 103 W/m2. How many photons (nearly) per square metre are incident on the Earth per second ? Assume that the photons in the sunlight have an average wavelength of 550 nm.
Solution:
Given, P = 1.388 × 103 W/m2; λ = 550 nm = 550 × 10-9 m ,
h = 6.63 × 10-34 J-s; c = 3 × 108 m/s
Energy of each photon E = \(\frac{\text { hc }}{\lambda}=\frac{6.63 \times 10^{-34} \times 3 \times 10^8}{550 \times 10^{-9}}\) = 3.616 × 10-19 J
No. of photons incident on the earth’s surface, N = \(\frac{\mathrm{P}}{\mathrm{E}}=\frac{1.388 \times 10^3}{3.66 \times 10^{-19}}\)
∴ N = 3.838 × 1021 photons/m2 – s.

AP Inter 2nd Year Physics Important Questions Chapter 12 Dual Nature of Radiation and Matter

Question 15.
Show that the wavelength of electromagnetic radiation is equal to the de-Brogile wavelength of its quantum (photon). [Mar. 14]
Answer:
Wave length of electromagnetic wave of frequency v and velocity C is given by,
λ = \(\frac{\mathrm{C}}{\mathrm{v}} \Rightarrow \lambda=\frac{\mathrm{C}}{\mathrm{v}} \times \frac{\mathrm{h}}{\mathrm{h}}=\frac{\mathrm{h}}{\left(\frac{\mathrm{hv}}{\mathrm{C}}\right)}=\frac{\mathrm{h}}{\mathrm{p}}\)    (∵ \(\frac{\mathrm{hv}}{\mathrm{C}}\) = p)
Hence, we can say wavelength of electromagnetic radiation is equal to the de-Brogile wavelength.

Sample Problem :

Question 1.
Calculate the (a) momentum and (b) dE-Brogile wavelength of the electrons accelerated through a potential difference of 56 V. [Mar. 14]
Answer:
a) Mass of the electron, m = 9 × 10-31 kg;
Potential difference, V = 56V
Momentum of electron, mv = \(\sqrt{2 \mathrm{eVm}}=\sqrt{2 \times\left(1.6 \times 10^{-19}\right) \times 56 \times 9 \times 10^{-31}}\) = 4.02 × 10-24kg ms-1

b) de-Broglie wavelength, λ = \(\frac{\mathrm{h}}{\mathrm{p}}=\frac{\mathrm{h}}{\mathrm{mv}}=\frac{6.62 \times 10^{-34}}{4.04 \times 10^{-24}}\) = 1.64 × 10-10m

Short Answer Questions

Question 1.
Describe an experiment to study the effect of frequency of incident radiation on ‘stopping potential’.
Answer:
Experimental study of the effect of frequency of incident radiation on stopping potential:

  1. The experimental set up is shown in fig.
    AP Inter 2nd Year Physics Important Questions Chapter 12 Dual Nature of Radiation and Matter 3
  2. Monochromatic light of sufficient energy (E = hv) from source ‘s’ is incident on photosensitive plate ‘C’ (emitter), electrons are emitted by it.
  3. The electrons are collected by the plate A (collector), by the electric field created by the battery.
  4. The polarity of the plates C and A can be reversed by a commutator.
  5. For a particular frequency of incident radiation, the minimum negative (retarding) potential V0 given to the plate A for which the photo current stops or becomes zero is called stopping potential.
  6. The experiment is repeated with different frequencies, and their different stopping potential are measured with voltmeter.
  7. From graph, we note that
    AP Inter 2nd Year Physics Important Questions Chapter 12 Dual Nature of Radiation and Matter 4

    • The values of stopping potentials are different for different frequencies.
    • The value of stopping potential is more negative for radiation of highef incident frequency.
    • The value of saturation current depends on the intensity of incident radiation but it is independent of the frequency of the incident radiation.

Question 2.
What is the deBroglie wavelength of a ball of mass 0.12 Kg moving with a speed of 20 ms-1? What can we infer from this result ?
Answer:
Given, m = 0.12 kg; υ = 20 m/s; h = 6.63 × 10-34 J-s;
λ = \(\frac{\mathrm{h}}{\mathrm{mv}}=\frac{6.63 \times 10^{-34}}{0.12 \times 20}=\frac{6.63 \times 10^{-34}}{2.4}\)
∴ λ = 2.762 × 10-34 m = 2762 × 10-21 Å.
The wave length of ball is very very small. Hence, its motion can be observed.

AP Inter 2nd Year Physics Important Questions Chapter 12 Dual Nature of Radiation and Matter

Question 3.
What is the effect of (i) intensity of light (ii) potential on photoelectric current ?
Answer:
(i) Effect of intensity of light on photoelectric current:
1) When the intensity (I) of incident light, with frequency greater than the threshold frequency (υ > υ0) is increased then the number of photoelectrons emitted increases i.e., the value of photoelectric current (i) increases, ie., i ∝ I.
AP Inter 2nd Year Physics Important Questions Chapter 12 Dual Nature of Radiation and Matter 5
ii) The effect of potential on photoelectric current:

  1. On increasing the positive potential on collecting electrode, the photoelectric current increases. At a particular positive potential, the photocurrent becomes maximum which is known as saturated current.
    AP Inter 2nd Year Physics Important Questions Chapter 12 Dual Nature of Radiation and Matter 6
  2. On increasing the value of negative potential on collecting electrode, the photoelectric current gradually goes on decreasing. At a particular negative potential the value of photoelectric current becomes zero. This is known as stopping potential.
  3. Stopping potential does not depend on the intensity of incident light. On increasing intensity, the value of saturated current increases, whereas the stopping potential remains unchanged.

Sample Problem :

Question 1.
The work function of caesium metal is 2.14 eV When light of frequency 6 × 1014 Hz is incident on the metal surface, photoemission of electrons occurs. What is the (a) maximum kinetic energy of the emitted electrons, (b) stopping potential and (c) maximum speed of the emitted photoelectrons ? [A.P. Mar. 16]
Solution:
Given, Φ0 = 2.14 eV; v = 6 × 1014 Hz
a) KEmax = hv – Φ0 = \(\frac{6.63 \times 10^{-34} \times 6 \times 10^{14}}{1.6 \times 10^{-19}}\) – 2.14
∴ KEmax = 0.35 eV

b) KEmax = eV0 ⇒ 0.35 eV = eV0
∴ V0 = 0.35 V

AP Inter 2nd Year Physics Important Questions Chapter 12 Dual Nature of Radiation and Matter 7

Long Answer Questions

Question 1.
How did Einstein’s photoelectric equation explain the effect of intensity and potential on photoelectric current ? How did this equation account for the effect of frequency of ‘ incident light on stopping potential ?
Answer:

  1. Einstein postulated that a beam of light consists of small energy packets called photons or quanta.
  2. The energy of photon is E = hv. Where ‘h1 is Planck’s constant; v is frequency of incident light (or radiation).
  3. If the absorbed energy of photon is greater than the work function (Φ0 = hυ0), the electron is emitted with maximum kinetic energy i.e., kmax = \(\frac{1}{2} \mathrm{mv}_{\max }^2\) = eV0 = hv – Φ0. This equation is known as Einstein’s photoelectric equation.
  4. Effect of intensity of light on photoelectric current:
    When the intensity (I) of incident light, with frequency greater than the threshold frequency (υ > υ0) is increased then the number of photoelectrons emitted decreases i.e., the value of photoelectric current (i) increases, ie., i ∝ I.
    AP Inter 2nd Year Physics Important Questions Chapter 12 Dual Nature of Radiation and Matter 5
  5. The effect of potential on photoelectric current:
    • On increasing the positive potential on collecting electrode, the photoelectric current increases. At a particular positive potential, the photocurrent becomes maximum which is known as saturated current.
      AP Inter 2nd Year Physics Important Questions Chapter 12 Dual Nature of Radiation and Matter 6
    • On increasing the value of negative potential on collecting electrode, the photoelectric current gradually goes on decreasing. At a particular negative potential the value of photoelectric current becomes zero. This is known as stopping potential (v0).
    • Stopping potential does not depend on the intensity of incident light. On increasing intensity, the value of saturated current increases, whereas the stopping potential remains unchanged.
  6. The effect of frequency of incident radiation on stopping potential:
    On increasing the frequency of incident light, the value of stopping potential goes on increasing gradually as shown in fig. That means kmax increases eV0 also increases.
    AP Inter 2nd Year Physics Important Questions Chapter 12 Dual Nature of Radiation and Matter 8
  7. From the graph, we note that
    • For a given photosensitive metal, the cut off potential (v0) varies linearly with the frequency of the incident radiation.
    • For a given photosensitive metal, there is a certain minimum cut off frequency υ0 (called threshold frequency) for which the stopping potential is zero.
      AP Inter 2nd Year Physics Important Questions Chapter 12 Dual Nature of Radiation and Matter 9
  8. From the graph we note that
    • The value of cut-off potential is different for radiation of different frequency.
    • The value of stopping potential is more negative for radiation of higher incident frequency.
  9.  From above experiments, it is found that, if the incident radiation is of higher frequency than that of threshold frequency, the photoelectric emission is. possible.

Textual Examples

Question 1.
Monochromatic light of frequency 6.0 × 1014 Hz is produced by a laser. The power emitted is 2.0 × 10-3 W. (a) What is the energy of a photon in the light beam ? (b) How many photons per second, on an average, are emitted by the source ?
Solution:
a) Each photon has an energy
E = hv = (6.63 × 10-34 J s) (6.0 × 1014 Hz)
= 3.98 × 10-19 J

b) If N is the number of photons emitted by the source per second, the power P transmitted in the beam equals N times the energy per photon E, so that P = NE.
Then N = \(\frac{\mathrm{P}}{\mathrm{E}}=\frac{2.0 \times 10^{-3} \mathrm{~W}}{3.98 \times 10^{-19} \mathrm{~J}}\)
= 5.0 × 1015 photons per second.

AP Inter 2nd Year Physics Important Questions Chapter 12 Dual Nature of Radiation and Matter

Question 2.
The work function of caesium is 2.14 eV. Find (a) the threshold frequency for caesium, and (b) the wavelength of the incident light if the photocurrent is brought to zero by a stopping potential of 0.60 V. [A.P. Mar. 16]
Solution:
a) For the cut-off or threshold frequency, the energy hv0 of the incident radiation must be equal to work function Φ0, so that
v0 = \(\frac{\phi_0}{\mathrm{~h}}=\frac{2.14 \mathrm{eV}}{6.63 \times 10^{-34} \mathrm{Js}}\)
= \(\frac{2.14 \times 1.6 \times 10^{-19} \mathrm{~J}}{6.63 \times 10^{-34} \mathrm{~J} \mathrm{~s}}\) = 5.16 × 1014 Hz
Thus, for frequencies less than this threshold frequency, no photoelectrons are ejected.

b) Photocurrent reduces to zero, when maximum kinetic energy of the emitted photoelectrons equals the potential energy e V0 by the retarding potential V0. Einstein’s Photoelectric equation is
eV0 = hv – Φ0 = \(\frac{\mathrm{hc}}{\lambda}\) – Φ0
or λ = hc/(eV0 + Φ0)
= \(\frac{\left(6.63 \times 10^{-34} \mathrm{~J} \mathrm{~s}\right) \times\left(3 \times 10^8 \mathrm{~m} / \mathrm{s}\right)}{(0.60 \mathrm{eV}+2.14 \mathrm{eV})}\)
= \(\frac{19.89 \times 10^{-26} \mathrm{~J} \mathrm{~m}}{(2.74 \mathrm{eV})}\)
λ = \(\frac{19.89 \times 10^{-26} \mathrm{~J} \mathrm{~m}}{2.74 \times 1.6 \times 10^{-19} \mathrm{~J}}\) = 454 nm

Question 3.
The wavelength of light in the visible region is about 390 nm for violet colour, about 550 nm (average wavelength) for yellow-green colour and about 760 nm for red colour.
(a) What are the energies of photons in (eV) at the (i) violet end, (ii) average wavelength, yellow-green colour, and (iii) red end of the visible spectrum ? (Take h = 6.63 × 10-34 J s and 1 eV = 1.6 × 10-19 J.)
(b) From which of the photosensitive materials with work functions listed in table and using the results of (i), (ii) and (iii) of (a) can you build a photoelectric device that operates with visible light ?
AP Inter 2nd Year Physics Important Questions Chapter 12 Dual Nature of Radiation and Matter 10
Solution:
a) Energy of the incident photon,
E = hv = hc/λ
E = (6.63 × 10-34 J s) (3 × 108 m/s)/λ.
= \(\frac{1.989 \times 10^{-25} \mathrm{~J} \mathrm{~m}}{\lambda}\)
i) For violet light,
λ1 = 390 nm (lower wavelength end)
Incident photon energy,
E1 = \(\frac{1.989 \times 10^{-25} \mathrm{~J} \mathrm{~m}}{390 \times 10^{-9} \mathrm{~m}}\)
5.10 × 10-19 J = \(\frac{5.10 \times 10^{-19} \mathrm{~J}}{1.6 \times 10^{-19} \mathrm{~J} / \mathrm{eV}}\)
= 3.19 eV

ii) For yellow-green light,
λ2 = 550 nm (average wavelength)
Incident photon energy,
E2 = \(\frac{1.989 \times 10^{-25} \mathrm{~J} \mathrm{~m}}{550 \times 10^{-9} \mathrm{~m}}\)
= 3.62 × 10-19 J = 2.26 eV.

iii) For red light,
λ3 = 760 nm (higher wavelength end)
Incident photon energy,
E3 = \(\frac{1.989 \times 10^{-25} \mathrm{~J} \mathrm{~m}}{760 \times 10^{-9} \mathrm{~m}}\)
= 2.62 × 10-19 J = 1.64 eV

AP Inter 2nd Year Physics Important Questions Chapter 12 Dual Nature of Radiation and Matter

b) For a photoelectric device to operate, we require incident light energy E to be equal to or greater than the work function Φ0 of the material. Thus, the photoelectric device will operate with violet light (with E = 3.19 eV) photosensitive material Na (with Φ0 = 2.75 eV), K (with Φ0 = 2.30 eV) and Cs (with Φ0 = 2.14 eV). It will also operate with yellow-green light (with E = 2.26 eV) for Cs (with Φ0 = 2.14 eV) only. However, it will not operate with red light (with E = 1.64 eV) for any of these photosensitive materials.

Question 4.
What is the de Broglie wavelength associated with (a) an electron moving with a speed of 5.4 × 106 m/s, and (b) a ball of mass 150 g travelling at 30.0 m/s ?
Solution:
a) For the electron :
Mass m= 9.11 × 10-31 kg, speed υ = 5.4 × 106 m/s. Then, momentum
P = mυ = 9.11 × 10-31 (kg) × 5.4 × 106 (m/s)
P = 4.92 × 10-24 kg m/s
de Broglie wavelength, λ = h/p
= \(\frac{6.63 \times 10^{-34} \mathrm{~J} \mathrm{~s}}{4.92 \times 10^{-24} \mathrm{~kg} \mathrm{~m} / \mathrm{s}}\)
λ = 0.135 nm

b) For the ball:
Mass m’ = 0.150 kg, speed υ’ = 30.0 m/s.
Then momentum
p’ = m’υ’ = 0.150 (kg) × 30.0 (m/s)
p’ = 4.50 kg m/s
de Broglie wavelength λ’ = h/p’.
= \(\frac{6.63 \times 10^{\pm 34} \mathrm{Js}}{4.50 \times \mathrm{kg} \mathrm{m} / \mathrm{s}}\)
λ’ = 1.47 × 10-34 m
The de Broglie wavelength of electron is comparable with X-ray wavelengths. However, for the ball it is about 10-19 times the size of the proton, quite beyond experimental measurement.

Question 5.
An electron, an a-particle and a proton have the same kinetic energy. Which of these particles has the shortest de Broglie wavelength ? [T.S. Mar. 15]
Solution:
For a particle,
de Broglie wavelength, λ = h/p
Kinetic energy, K = p2/2m
Then, λ = h /\(\sqrt{2 \mathrm{mK}}\)
For the same kinetic energy K, the de Broglie wavelength associated with the particle is inversely proportional to the square root of their masses. A proton \(\left({ }_1^1 \mathrm{He}\right)\) is 1836 times massive than an electron and an a-particle \(\left({ }_2^4 \mathrm{He}\right)\) four times that of a proton
Hence, α-particle has the shortest de Broglie wavelength.

AP Inter 2nd Year Physics Important Questions Chapter 12 Dual Nature of Radiation and Matter

Question 6.
A particle is moving three times as fast as an electron. The ratio of the de Broglie wavelength of the particle to that of the electron is 1.813 × 10-4. Calculate the particle’s mass and identify the particle.
Solution:
de Broglie wavelength of a moving particle, having mass m and velocity υ :
λ = \(\frac{\mathrm{h}}{\mathrm{p}}=\frac{\mathrm{h}}{\mathrm{mv}}\)
Mass, m = h/A.
For an electron, mass me = h/λe υe
Now, we have υ/υe = 3 and
λ/λe = 1.813 × 10-4
Then, mass of the particle,
m = me \(\left(\frac{\lambda_{\mathrm{e}}}{\lambda}\right)\left(\frac{v_{\mathrm{e}}}{v}\right)\)
m = (9.11 × 10-31 kg) × (1/3) × (1/1.813 × 10-4)
m = 1.675 × 10-27 kg.
Thus, the particle, with this mass could be a proton or a neutron.

Question 7.
What is the de Broglie wavelength associated with an electron, accelerated through a potential difference of 100 volts? [A.P. Mar. 15]
Solution:
Accelerating potential V = 100 V The de Broglie wavelength λ is
λ = h/p = \(\frac{1.227}{\sqrt{\mathrm{V}}}\) nm
λ = \(\frac{1.227}{\sqrt{100}}\) nm = 0.123 nm
The de Broglie wavelength associated with an electron in this case is of the order of X-ray wavelengths.

AP Inter 2nd Year Physics Study Material Chapter 1 Waves

Andhra Pradesh BIEAP AP Inter 2nd Year Physics Study Material 1st Lesson Waves Textbook Questions and Answers.

AP Inter 2nd Year Physics Study Material 1st Lesson Waves

Very Short Answer Questions

Question 1.
What does a wave represent ?
Answer:
A wave represent the transport of energy through a medium from one point to another without translation of the medium.

Question 2.
Distinguish between transverse and longitudinal waves.
Answer:
Transverse waves

  1. The particles of the medium vibrate perpendicular to the direction of wave propagation.
  2. Crests and troughs are formed alternatively.

Longitudinal waves

  1. The particles of the medium vibrate parallel to the direction of wave propagation.
  2. Compressions and rare fractions are formed alternatively.

AP Inter 2nd Year Physics Study Material Chapter 1 Waves

Question 3.
What are the parameters used to describe a progressive harmonic wave ?
Answer:
Progressive wave equation is given y = a sin (ωt – kx)
Where ω = 2πv = \(\frac{2 \pi}{T}\)

Parameters :

  1. a = Amplitude
  2. λ = Wavelength
  3. T = Time period
  4. v = Frequency
  5. k = Propagation constant
  6. ω = Angular frequency.

Question 4.
Obtain an expression for the wave velocity in terms of these parameters.
Answer:
Let ‘v’ be the velocity of a wave, ‘v’ be frequency and ‘λ’ be the wavelength. If T is the time period, then v = \(\frac{1}{\mathrm{~T}}\)
The distance travelled by the wave in the time T = λ.
Distance travelled in one second = \(\frac{\lambda}{T}\)
which is equal to wave velocity v = \(\frac{\lambda}{T}\).
∴ v = vλ

Question 5.
Using dimensional analysis obtain an expression for the speed of transverse waves in a stretched string.
Answer:
Wave velocity v ∝ Ta µb ⇒ V = K Ta µb ——-> (1)
Dimensions of v = M0L1 T-1, Tension T = M1L1T-2,
Linear mass µ = M1L-1, Constant K = M0L0T0
Now (1) becomes M0L1T-1 = [M1L1T-2]a [M1L-1]b
M0L1T1 = Ma + bLa-bT-2a Comparing the powers of same physical quantity.
-1 = -2a ⇒ a = \(\frac{1}{2}\)
a + b = 0 ⇒ b = –\(\frac{1}{2}\)
⇒ v = (1)\(T^{\frac{1}{2}} \mu^{\frac{1}{2}}\) [∵ K = 1 Practically]
∴ v = \(\sqrt{\frac{\mathrm{T}}{\mu}}\)

Question 6.
Using dimensional analysis obtain an expression for the speed of sound waves in a medium. .
Answer:
Speed of sound v ∝ Ba ρb ⇒ v = K Ba ρb ——–> (1)
Dimensions of v = M0L1T-1,
Elasticity of medium,
B = M1L-1T-2, density ρ = M1L-3, constant K = M0L0T0.
Now (1) becomes M0L1T-1 = M0L0T0 [M1L-1T-2]a [M1L-3]b
0 = a + b
1 = -a – 3b
-1 = -2a ⇒ a = \(\frac{1}{2}\)
b = –\(\frac{1}{2}\)
v = K \(B^{\frac{1}{2}} \rho^{\frac{1}{2}}\)
∴ v = \(\sqrt{\frac{B}{\rho}}\) [∵ K = 1, practically]

AP Inter 2nd Year Physics Study Material Chapter 1 Waves

Question 7.
What is the principle of superposition of waves ?
Answer:
When two or more waves, are acting simultaneously on the particle of the medium, the resultant displacement is equal to the algebraic sum of individual displacements of all the waves. This is the principle of superposition of waves.

If y1, y2, …… yn be the individual displacements of the particles,then resultant displacement y = y1 + y2 + ……. + yn

Question 8.
Under what conditions will a wave be reflected ?
Answer:

  1. When the medium ends abruptly at any point.
  2. If the density and rigidity modulus of the medium changes at any point.

Question 9.
What is the phase difference between the incident and reflected waves when the wave is reflected by a rigid boundary.
Answer:
π Radian or 180°.

Question 10.
What is a stationary or standing wave ?
Answer:
When two identical progressive (Transverse or longitudinal) waves travelling opposite directions in a medium along the same straight line, which are superimposed then the resultant wave is called stationary waves or standing wave.

Question 11.
What do you understand by the terms ‘node ‘and’ antinode’?
Answer:
Node : The points at which the amplitude is zero, are called nodes.
Antinodes : The points at which the amplitude is maximum, are called antinodes.

Question 12.
What is the distance between a node and an antinode in a stationary wave ?
Answer:
The distance between node and antinode is \(\frac{\lambda}{4}\)

Question 13.
What do you understand by ‘natural frequency’ or ‘normal mode of vibration’ ?
Answer:
When a body is set into vibration and then left to itself, the vibrations made by it are called natural or free vibrations. Its frequency is called natural frequency or normal mode of vibration.

Question 14.
What are harmonics ?
Answer:
The frequencies in which the standing waves can be formed are called harmonics.
(Or)
The integral multiple of fundamental frequencies are called harmonics.

Question 15.
A string is stretched between two rigid supports. What frequencies of vibration are possible in such a string ?
Answer:
The possible frequencies of vibrations in a stretched string between two rigid supports is given by
vn = (n + \(\frac{1}{2}\))\(\frac{v}{21}\) where n = 0, 1, 2, 3, ……

Question 16.
The air column in a long tube, closed at one end, is set in vibration. What harmonics are possible in the vibrating air column ?
Answer:
The possible harmonics in the vibrating air column of a long closed tube is given by
vn = [2n + 1]\(\frac{v}{4 l}\) where n = 0, 1, 2, 3, ……..

Question 17.
If the air column in a tube, open at both ends, is set in vibration; what harmonics are possible?
Answer:
The possible harmonics in vibrating air column of a long open tube is given by
Vn = \(\frac{n v}{21}\)
where n = 1, 2, 3, ……….

AP Inter 2nd Year Physics Study Material Chapter 1 Waves

Question 18.
What are ‘beats’ ?
Answer:
Beats : When two sound notes of nearly frequency travelling in the same direction and interfere to produce waxing and waning of sound at regular intervals of time is called “Beats”.

Question 19.
Write down an expression beat frequency and explain the terms there in.
Answer:
Expression of beat frequency, Δv = v1 ~ v2
where v1 and v2 are the frequencies of two waves.

Question 20.
What is ‘Doppler effect’? Give an example.
Answer:
Doppler effect: The apparent change in the frequency heard by the observer due to relative motion between source of sound and observer is called “Doppler effect”.

E.g.: When the whistling railway engine approaches the stationary observer on the platform, the frequency of sound appears to increase above the actual frequency. When it moves away from the observer, the apparent frequency decreases.

Question 21.
Write down an expression for the observed frequency when both source and observer are moving relative to each other in the same direction.
Answer:
Apparent frequency of sound heard by an observer,
v’ = \(\left[\frac{v-v_0}{v-v_s}\right] v\)
AP Inter 2nd Year Physics Study Material Chapter 1 Waves 1
where v = frequency of sound
v = velocity of sound
v0 = velocity of observer
vS = velocity of source

Short Answer Questions

Question 1.
What are transverse waves ? Give illustrative examples of such waves.
Answer:
Transverse waves: In a wave motion, the vibration of the particles and the direction of the propagation of the waves are perpendicular to each other, the waves are said to be transverse waves.

Illustration:

  1. Waves produced in the stretched strings are transverse.
  2. When a stretched string is plucked, the waves travel along the string.
  3. But the particles in the string vibrate in the direction perpendicular to the propagation of the wave.
  4. They can propagate only in solids and on the surface of the liquids.
  5. Ex : Light waves, surface water waves.

Question 2.
What are longitudinal waves ? Give illustrative example of such waves.
Answer:
Longitudinal waves : In a wave motion, the direction of the propagation of the wave and vibrations of particles are in the same direction, the waves are said to be longitudinal waves.
Illustration:

  1. Longitudinal waves may be easily illustrated by releasing a compressed spring.
  2. A series of compressions and rarefactions (expansions) propagate along the spring.
    AP Inter 2nd Year Physics Study Material Chapter 1 Waves 2
    C = Compression; R = Rarefaction.
  3. They can travel in solids, liquids and gases.
  4. Ex : Sound waves.

Question 3.
Write an expression for a progressive harmonic wave and explain the various parameters used in the expression.
Answer:
The expression of a progressive harmonic wave is written as y = a sin(ωt – \(\frac{2 \pi}{\lambda} x\))
or y = a sin(ωt – kx) where ω = 2πv, k = \(\frac{2 \pi}{\lambda}\)

Parameters:

  1. Amplitude (a) : It is the maximum displacement of a vibrating particle from its mean position.
  2. Frequence (v): It is the number of complete vibrations made by a vibrating body in one second.
  3. Wave length (λ) : It is defined as the distance covered by a wave while it completes one vibration, (or) It is the distance between two consecutive points in the same phase.
  4. Phase of vibration (ϕ) : The phase of vibration of a vibrating particle gives its state of displacement at a given instant. It is generally given by the phase angle.

AP Inter 2nd Year Physics Study Material Chapter 1 Waves

Question 4.
Explain the modes of vibration of a stretched string with examples.
Answer:
Modes of vibrations of a stretched string :

  1. In sitar or Guitar, a stretched string can vibrate, in different frequencies and form stationary waves. This mode of vibrations are known as harmonics.
  2. If it vibrates in one segment, which is known as fundamental harmonic. The higher harmonics are called the overtones.
    AP Inter 2nd Year Physics Study Material Chapter 1 Waves 3
  3. It vibrates in two segments then the second harmonic is called first overtone. Similarly the pattern of vibrations are shown fig.
  4. If a stretched string vibrates with P ’Seg’ ments (loop) then frequency of vibration v = \(\frac{\mathrm{P}}{2 l} \sqrt{\frac{\mathrm{T}}{\mu}}\) where T = tension in the string, µ = linear density = \(\frac{\text { mass }}{\text { length }}\)
  5. In first mode of vibration, P = 1, then v = \(\frac{1}{2l} \sqrt{\frac{T}{\mu}}\) (1st hamonic (or) fundamental frequency)
  6. second mode of vibration, P = 2, then v1 = \(\frac{2}{2l} \sqrt{\frac{T}{\mu}}\) = 2v (2nd harmonic (or) 1st overtone)
  7. In third mode of vibration, P = 3, then v2 = \(\frac{3}{2l} \sqrt{\frac{T}{\mu}}\) = 3v (3rd harmonic (or) 2nd overtone)
    The ratio of the frequency of Harmonics are, v : v1 : v2 = v : 2v : 3v = 1 : 2 : 3

Question 5.
Explain the modes of vibration of an air column in an open pipe.
Answer:
Modes of vibration of an air column in an open pipe :
1) For a open pipe both the ends are open. So antinodes will be formed at both the ends. But two antinodes cannot exist without a node between them.
2) The possible harmonics in vibrating air column of a open pipe is given by AP Inter 2nd Year Physics Study Material Chapter 1 Waves 4.
Where n = 1, 2, 3
(1st harmonic or fundamental frequence)
AP Inter 2nd Year Physics Study Material Chapter 1 Waves 5
3) In first normal Mode of vibrating air column in a open pipe v1 = \(\frac{v}{2l}\) = v
(2nd harmonic 1st overtone)
AP Inter 2nd Year Physics Study Material Chapter 1 Waves 6

4) In second normal Mode of vibrating air column in a open pipe, v2 = \(\frac{2 v}{2l}\) = 2v

5) In third, normal Mode of vibrating air column in a open pipe, v3 = \(\frac{3 v}{21}\) = 3u
(3rd harmonic 2nd overtone)
AP Inter 2nd Year Physics Study Material Chapter 1 Waves 7

6) In open pipe the ratio of frequencies of harmonics is
v1 : v2 : v3 = v : 2v : 3v = 1 : 2 : 3

Question 6.
What do you understand by ‘resonance’ ? How would you use resonance to determine the velocity of sound in air ?
Answer:
Resonance: If the natural frequency of a vibrating body is equal to the frequency of external periodic force then the two bodies are said to be in resonance. At resonance the bodies will vibrate with increasing amplitude.

Determination of velocity of sound in air using resonance :

1) In resonance tube, an air column is made to vibrate by means of vibrating fork. At certain length of air column, the air column would have the same frequency as that of the fork. Then the air column vibrates with the maximum amplitude and the intense sound is produced.
AP Inter 2nd Year Physics Study Material Chapter 1 Waves 8

2) The vibrating fork of known frequency (v) is placed above the open end of the tube.

3) The length of air column is gradually increased until the booming sound can be heard at two different lengths of air column.

4) In first resonance, l1 be the length of air column, then \(\frac{\lambda}{4}\) = l1 + C …….. (1)
Where λ is the wavelength of sound emitted by the fork and C is the end correction of the tube.
AP Inter 2nd Year Physics Study Material Chapter 1 Waves 9
5) In second resonance, l2 be the length of air column, then \(\frac{3 \lambda}{4}\) = l2 + C …… (2)
(2) – (1) ⇒ \(\frac{\lambda}{2}\) = l2 – l1
AP Inter 2nd Year Physics Study Material Chapter 1 Waves 10
λ = 2 (l2 – l1)
Speed of sound is given by
υ = vλ = v[2(l2 – l1)]
∴ υ = 2v (l2 – l1)

6) Hence by knowing v, l1, l2 the speed of sound is calculated.

AP Inter 2nd Year Physics Study Material Chapter 1 Waves

Question 7.
What are standing waves ? Explain how standing waves may be formed in a stretched string.
Answer:
Standing wave or stationary: When two identical progressive (Transverse or longitudinal) waves travelling opposite directions in a medium along the same straight line, which are super-imposed then the resultant wave is called standing wave.

Formation of standing wave in a stretched string : –

  1. If a string of length ‘l’ is stretched between two fixed points and set into vibration, a transverse progressive wave begins to travel along the string.
  2. The wave is get reflected at the other fixed end.
  3. The incident and reflected waves interfere and produce a stationary wave.
  4. The stationary wave with nodes and antinodes is shown below.
    AP Inter 2nd Year Physics Study Material Chapter 1 Waves 11

Question 8.
Describe a procedure for measuring the velocity of sound in a stretched string.
Answer:
The velocity of a transverse wave travelling along a stretched string in fundamental mode is given by v = 2vl, where v = frequency, l = resonating length.

Measurement of velocity of sound in a stretched string using sonometer :

  1. The wire is subjected to a fixed tension with suitable load.
  2. A tuning fork of known frequency (v), is excited and the stem is held against the sono – meter box.
    AP Inter 2nd Year Physics Study Material Chapter 1 Waves 12
  3. The distance between the two bridges is adjusted such that a small paper rider at the middle of B1 B2 vibrates vigorously and flies off due to resonance.
  4. The resonating length ‘l’ can be measured between two bridges with scale.
  5. By knowing v and l; we can find the velocity of a wave using υ = 2vl.

Question 9.
Explain, using suitable diagrams, the formation of standing waves in a closed pipe. How may this be used to determine the frequency of a source of sound ?
Answer:
Formation of standing waves in a closed pipe :

  1. In closed pipe one end is closed and the other end is open. So antinode is formed at open end and antinode is formed at closed end.
  2. The possible harmonics in vibrating air column in a closed pipe vn = \(\frac{(2 n+1) v}{4 l}\) where v = 0, 1, 2, 3, …….
  3. In first normal mode of vibrating air column in a closed pipe, v1 = \(\frac{v}{41}\)
    [first harmonic (or) fundamental frequency]
    AP Inter 2nd Year Physics Study Material Chapter 1 Waves 13
  4. In second normal mode of vibrating air column in a closed pipe,
    v3 = \(\frac{3 \mathrm{v}}{4l}\) [Third harmonic (or) first overtone]
  5. In third normal mode of vibrating air column in a closed pipe,
    v5 = \(\frac{5 \mathrm{v}}{4 \mathrm{l}}\) [Fifth harmonic (or) second overtone]

Determination of frequency of a source of sound :

  1. The vibrating fork of unknown frequency (v) is placed above the open end of the tube.
    AP Inter 2nd Year Physics Study Material Chapter 1 Waves 14
  2. Reservoir is slowly lowered, until a large booming sound is heard. Measure 1st resonating, air column length l1.
  3. Further lower the reservoir, until second time a large booming sound is heard. Measure 2nd resonating air column length l2.
  4. Velocity of a wave at 0°C is v = 331 m/s.
  5. By knowing v, l1 and l2 we can find unknown frequency of a tuning fork using
    v = \(\frac{v}{2\left(l_2-l_1\right)}\)

AP Inter 2nd Year Physics Study Material Chapter 1 Waves

Question 10.
What are ‘beats’ ? When do they occur ? Explain their use, if any.
Answer:
Two sound waves of nearly same frequency are travelling in the same direction and interfere to produce a regular waxing (maximum) and waning (minimum) in the intensity of the resultant sound waves at regular intervals of time is called beats.

It two vibrating bodies have slightly difference in frequencies, beats can occur.
No. of beats can be heard Δυ = υ1 – υ2

Importance :

  1. It can be used to tune musical Instruments.
  2. Beats are used to detect dangerous gases

Explanation-for tuning musical instruments with beats :
Musicians use the beat phenomenon in tuning their musical instruments. If an instrument is sounded against a standard frequency and tuned until the beats disappear. Then the instrument is in tune with the standard frequency.

Question 11.
What is ‘Doppler effect’? Give illustrative examples.
Answer:
Doppler effect: The apparent change in the frequency heard by the observer due to relative motion between the observer and the source of sound is called doppler effect.

Examples:

  1. The frequency of whistling engine heard by a person standing on the platform appears to increase, when the engine is approaching the platform and it appears to decrease when the engine is moving away from the platform.
  2. Due to Doppler effect the frequency of sound emitted by the siren of an approaching ambulance appears to increase. Similarly the frequency of sound appears to drop when it is moving away.

Long Answer Questions

Question 1.
Explain the formation of stationary waves in stretched strings and hence deduce the laws of transverse wave in stretched strings. (A.P. Mar. ’19)
Answer:
A string is a metal wire whose length is large when compared to its thickness. A stretched string is fixed at both ends, when it is plucked at mid point, two reflected waves of same amplitude and frequency at the ends are travelling in opposite direction and overlap along the length. Then the resultant waves are known as the standing waves (or) stationary waves.

Let two transverse progressive waves of same amplitude a, wave length λ and frequency ‘v’, travelling in opposite direction be given by
y1 = a sin (kx – ωt) and y2 = a sin (kx + ωt)
where ω’ = 2πv and k = \(\frac{2 \pi}{\lambda}\)
AP Inter 2nd Year Physics Study Material Chapter 1 Waves 15
The resultant wave is given by y = y1 + y2
y = a sin (kx – ωt) + a sin (kx + ωt)
y = (2a sin kx) cos ωt
2a sin kx = Amplitude of resultant wave.
It depends on ‘kx’. If x = 0, \(\frac{\lambda}{2}\), \(\frac{2 \lambda}{2}\), \(\frac{3 \lambda}{2}\),……. etc, the amplitude = zero
These positions are known as “Nodes”.
If x = \(\frac{\lambda}{4}\), \(\frac{2 \lambda}{2}\), \(\frac{3 \lambda}{2}\) ……… etc, the amplitude = zero

The positions are known as “Nodes”
If x = \(\frac{\lambda}{4}\), \(\frac{3 \lambda}{4}\), \(\frac{5 \lambda}{4}\) ……. etc, the amplitude = maximum (2a).
These positions are called “Antinodes”.

If the string vibrates in ‘P’ segments and ‘l’ is its length then length of each segment = \(\frac{l}{p}\)
Which is equal to \(\frac{\lambda}{2}\)
∴ \(\frac{l}{\mathrm{p}}\) = \(\frac{\lambda}{2}\) ⇒ λ = \(\frac{2 l}{\mathrm{P}}\)
Harmonic frequency v = \(\frac{v}{\lambda}=\frac{v \mathrm{P}}{2 l}\)
v = \(\frac{v P}{2 l}\) ——- (1)
If ‘ T’ is tension (stretching force) in the string and ‘μ’ is linear density then velocity of transverse wave (v) in the string is v = \(\sqrt{\frac{\mathrm{T}}{\mu}}\) —– (2)
From the Eqs (1) and (2)
Harmonic frequency v = \(\frac{p}{2 l} \sqrt{\frac{T}{\mu}}\)
If P = 1 then it is called fundamental frequency (or) first harmonic frequency
∴ Fundamental Frequency v = \(\frac{1}{2 l} \sqrt{\frac{\mathrm{T}}{\mu}}\) —— (3)

Laws of Transverse Waves Along Stretched String:

Fundamental frequency of the vibrating string v = \(\frac{1}{2 l} \sqrt{\frac{\mathrm{T}}{\mu}}\)

First Law : When the tension (T) and linear density (μ) are constant, the fundamental frequency (v) of a vibrating string is inversely proportional to its length.
∴ v ∝ \(\frac{1}{l}\) ⇒ vl = constant, when T and ‘μ’ are constant.

Second Law: When the length (I) and its, linear density (m) are constant the fundamental frequency of a vibrating string is directly proportional to the square root of the stretching force (T).
∴ v ∝ \(\sqrt{\mathrm{T}}\) ⇒ \(\frac{v}{\sqrt{T}}\) = constant, when ‘l’ and ‘m’ are constant.

Third Law: WHien the length (J) and the tension (T) are constant, the fundamental frequency of a vibrating string is inversely proportional to the square root of the linear density (m).
∴ v ∝ \(\frac{1}{\sqrt{\mu}}\) ⇒ \(v \sqrt{\mu}\) = constant, when ‘l’ and ‘T’ are constant.

AP Inter 2nd Year Physics Study Material Chapter 1 Waves

Question 2.
Explain the formation of stationary waves in an air column enclosed in open pipe. Derive the equations for the frequencies of the harmonics produced. (T.S. Mar. ’16, A.P. Mar. ’15)
Answer:
A pipe, which is opened at both ends is called open pipe. When a sound wave is sent through a open pipe, which gets reflected by the earth. Then incident and reflected waves are in same frequency, travelling in the opposite directions are super – imposed stationary waves are formed.
AP Inter 2nd Year Physics Study Material Chapter 1 Waves 16
Harmonics in open pipe : To form the stationary wave in open pipe, which has two antinodes at two ends of the pipe with a node between them.
∴ The vibrating length (l) = half of the wavelength \(\left(\frac{\lambda_1}{2}\right)\)
l = \(\frac{\lambda_1}{2}\) ⇒ λ1 = 2l
fundamental frequency v1 = \(\frac{\mathrm{v}}{\lambda_1}\) where v is velocity of sound in air v1 = \(\frac{v}{2 l}\) = v —— (1)
For second harmonic (first overtone) will have one more node and antinode than the fundamental.
If λ2 is wavelength of second harmonic l = \(\frac{2 \lambda_2}{2}\) ⇒ λ2 = \(\frac{2l}{2}\)
If ‘v2‘ is frequency of second harmonic then v2 = \(\frac{v}{\lambda_2}\) = \(\frac{v \times 2}{2 l}\) = 2v
v2 = 2v —– (2)

Similarly for third harmonic (second overtone) will have three nodes and four antinodes as shown in above figure.
If λ3 is wave length of third harmonic l = \(\frac{3 \lambda_3}{2}\)
λ3 = \(\frac{2l}{3}\)
If ‘v2’ is frequency of third harmonic then
v3 = 3v —– (3)

Similarly we can find the remaining or higher harmonic frequencies i.e v3, v4 etc, can be determined in the same way.
Therefore the ratio of the harmonic frequencies in open pipe can be written as given below.
v : v1 : v2 = 1 : 2 : 3 ………

Question 3.
How are stationary waves formed in closed pipes ? Explain the various modes of vibrations and obtain relations for their frequencies. (A.P. & T.S. Mar. ’15)
Answer:
A pipe, which is closed at one end and the other is opened is called closed pipe. When a sound wave is sent through a closed pipe, which gets reflected at the closed end of the pipe. Then incident and reflected waves are in same frequency, travelling in the opposite directions are superimposed stationary waves are formed.
AP Inter 2nd Year Physics Study Material Chapter 1 Waves 17

To form the stationary wave in closed pipe, which has atleast a node at closed end and antinode at open end of the pipe, it is known as first harmonic in closed pipe. Then length of the pipe (l) is equal to one fourth of the wave length.
∴ l = \(\frac{\lambda_1}{4}\) ⇒ λ1 = 4l
If ‘v1‘ is fundamental frequency then
v1 = \(\frac{v}{\lambda_1}\) where ‘υ’ is velocity of sound in air
v1 = \(\frac{v}{4l}\) = v ——- (1)

To form the next harmonic in closed pipe, two nodes and two antinodes should be formed. So that there is possible to form third harmonic in closed pipe. Since one more node and antinode should be included.
Then length of the pipe is equal to \(\frac{3}{4}\) of the wavelength.
∴ l = \(\frac{3 \lambda_3}{4}\) where ‘λ3‘ is wave length of third harmonic.
λ3 = \(\frac{4l}{3}\)
If ‘v3‘ is third harmonic frequency (first overtone)
∴ v3 = \(\frac{v}{\lambda_3}\) = \(\frac{3 v}{41}\)
v3 = 3v ——- (2)

Similarly the next overtone in the close pipe is only fifth harmonic it will have three nodes and 3 antinodes between the closed end and open end.
Then length of the pipe is equal to \(\frac{5}{4}\) of wave length (λ5)
∴ l = \(\frac{5 \lambda_5}{4}\) where ’λ5‘ is wave length of fifth harmonic. .
λ5 = \(\frac{4l}{5}\)
If ‘V5‘ is frequency of fifth harmonic (second overtone)
V5 = \(\frac{v}{\lambda_5}=\frac{5 v}{4 I}\)
v5 = 5v —– (3)

∴ The frequencies of higher harmonics can be determined by using the same procedure. Therefore from the Eq (1), (2) and (3) only odd harmonics are formed.
Therefore the ratio of the frequencies of harmonics in closed pipe can be written as
v1 : v3 : v5 = v : 3v : 5v
v1 : v3 : v5 = 1 : 3 : 5

AP Inter 2nd Year Physics Study Material Chapter 1 Waves

Question 4.
What are beats ? Obtain an expression for the beat frequency ? Where and how are beats made use of ?
Answer:
Beats : Two sound waves of nearly same frequency are travelling in the same direction and interfere to produce a regular waxing and waning in the intensity of the resultant sound waves at regular intervals of time are called Beats.
If v1 and v2 are the frequencies of two sound notes superimposed along the same direction, no of beats heard per second = Δv = v1 – v2.
Maximum no. of beats heard per sec is 10 due to persistence of hearing.

Expression for the beat frequency :

  1. Consider the two wave trains of equal amplitude but of nearly equal frequencies.
  2. Let the frequencies of the waves be v1 and v2. Say v1 is slightly greater than v2.
  3. Let the beat period be T seconds.
  4. No.of vibrations, made by the first wave train in T seconds – v1T
    [∵ no.of oscillations in 1 sec = v]
    [∵ no.of oscillations in T sec = vt]
  5. No.of vibrations, made by the second wave train in T seconds = v2T
  6. During the time interval T, the first wave train would have completed one vibration more than the second wave train.
  7. Hence, v1T – v2T = 1 or v1 – v2 = \(\frac{1}{\mathrm{~T}}\)
  8. Since, T is the beat period, no.of beats per seconds = \(\frac{1}{\mathrm{~T}}\)
  9. Hence the beat frequency = \(\frac{1}{\mathrm{~T}}\) = v1 – v2 = Δv
  10. That is the beat frequency is the difference between the frequencies of the two wave trains.

Practical applications of beats:

  1. Determination of an unknown frequency: Out of two tuning forks, one is loaded with wax and the other is filed. The excited tuning forks are close together and no.of beats can be heard. Then after unknown frequencies of them will be found practically.
  2. For tuning musical instruments : Musicians use the beat phenomenon in tuning their musical instruments.
  3. For producing colourful effects in music: Sometimes, a rapid succession of beats is knowingly introduced in music. This produces an effect similar to that of human voice and is appreciated by the audience.
  4. For detection of Marsh gas (dangerous gases) in mines.

Question 5.
What is Doppler effect? Obtain an expression for the apparent frequencý of sound heard when the source is in motion with respect to an observer at rest. (A.P. Mar. ’16, Mar. ’14)
Answer:
Doppler effect : The apparent change in the frequency heard by the observer due to the relative motion between the observer and the source of sound is called doppler effect.

When a whistling railway engine approaches an observer standing on the platform, the frequency of sound appears to increase. When it moves away the frequency appear to decrease.
AP Inter 2nd Year Physics Study Material Chapter 1 Waves 18
Expression for apparent frequency when source is in motion and listener at rest:
Let S = Source of sound
O = listener

Let ‘S be the source, moving with a velocity ‘υs‘ towards the stationary listener.
The distance travelled by the source in time period T’ = υs. T
Therefore the successive compressions and rarefactions are drawn closer to listener.
∴ Apparent wavelength = λ’ = λ – υsT.
AP Inter 2nd Year Physics Study Material Chapter 1 Waves 19

If ‘v’ “is apparent frequency heard by the listener
then v’ = \(\frac{v}{\lambda^{\prime}}\) where ‘υ’ is velocity of sound in air
v’ = \(\frac{v . V}{v-v_s}\)
Therefore the apparent frequency is greater than the actual frequency.
Similarly, if the source is away from the stationary listener then apparent frequency
v’ = \(\frac{v . V}{v+v_s}\), which is less than the actual frequency.

Limitation : Doppler effect is applicable when the velocities of the source and listener are much less than that of sound velocity

Question 6.
What is Doppler shift? Obtain an expression for the apparent frequency of sound heard when the observer Is In motion with respect to a source át rest.
Answer:
Doppler Shift: Due to the relative motion, when the source comes closer to listener, the apparent frequency is greater than actual frequency and source away from listener; the apparent frequency is less than actual frequency So the difference in apparent and actual frequencies is known as Doppler shift.
Expression for the apparent frequency heard by a moving observer:

Case (1) : When observer Is moving towards source:
Let ‘υ0’ be velocity of listener ‘O’, moving towards the stationary source ‘s’ as shown in figure. So observer will receive more number of waves in each second.
The distance travelled by observer in one second = υ0

The number of extra waves received by the observer = \(\frac{v_0}{\lambda}\)
We know v = vλ ⇒ λ = \(\frac{v}{v}\)
AP Inter 2nd Year Physics Study Material Chapter 1 Waves 20
Where υ = Velocity of sound
v = Frequency of sound
If ‘v’ is apparent frequency heard by him then
AP Inter 2nd Year Physics Study Material Chapter 1 Waves 21
Therefore the apparent frequency is greater than actual frequency.

Case (2) : When observer Is moving away from rest source:
If the observer is moving away from the stationary source then he loses the number of waves \(\frac{v_0}{\lambda}\)
∴ Apparent frequency v’ = v – \(\frac{v_0}{\lambda}\) = v – \(\frac{v_0 \cdot v}{v}\)
v’ = \(\left[\frac{v-v_0}{v}\right] \cdot v\)
Hence the apparent frequency is less than actual frequency.

Problems

Question 1.
A stretched wire of length 0.6m is observed to vibrate with a frequency of 30Hz in the fundamental mode. If the string has a linear mass of 0.05 kg/m find
(a) the velocity of propagation of transverse waves in the string
(b) the tension in the string.
Solution:
v = 30 Hz; l = 0.6 m ; μ = 0.05 kg m-1 υ = ? ; T = ?
a) υ = 2vl = 2 × 30 × 0.6 = 36 m/s
b) T = υ2μ = 36 × 36 × 0.05 = 64.8 N

Question 2.
A steel cable of diameter 3 cm is kept under a tension of 10kN. The density of steel is 7.8 g/cm3. With what speed would transverse waves propagate along the cable ?
Solution:
T = 10 kN = 104 N
D = 3 cm; r = \(\frac{\mathrm{D}}{2}\) = \(\frac{3}{2}\)cm
= \(\frac{3}{2}\) × 10-2m;
A = πr² = \(\frac{22}{7} \times\left[\frac{3}{2} \times 10^{-2}\right]^2\)
AP Inter 2nd Year Physics Study Material Chapter 1 Waves 22

Question 3.
Two progressive transverse waves given by y1 = 0.07 sinπ(12x-500t) and y2 = 0.07 sinπ(12x + 500t) travelling along a stretched string from nodes and antinodes. What is the displacement at the (a) nodes (b) antinodes ? What is the wavelength of the standing wave ?
Solution:
A1 = 0.07; A2 = 0.07; K = 12π
a) At nodes, displacement
y = A1 – A2 = 0.07 – 0.07 = 0
b) At antinodes, displacement
y = A1 + A2 = 0.07 + 0.07 = 0.14 m
c) Wavelength λ = \(\frac{2 \pi}{\mathrm{K}}=\frac{2 \pi}{12 \pi}\) = 0.16 m

AP Inter 2nd Year Physics Study Material Chapter 1 Waves

Question 4.
A string has a length of 0.4m and a mass of 0.16g. If the tension in the string is 70N, what are the three lowest frequencies it produces when plucked ?
Solution:
I = 0.4 m; M = 0.16g = 0.16 × 10-3 kg;
AP Inter 2nd Year Physics Study Material Chapter 1 Waves 23

Question 5.
A metal bar when clamped at its centre resonantes in its fundamental frequency with longitudinal waves of frequency 4 kHz. If the clamp is moved to one end. What will be its fundamental resonance frequency ?
Solution:
When a metal bar of length l is clamped in the middle, it has one node in the middle and two antinodes at its free ends. In the fundamental mode. l = \(\frac{\lambda}{2}\) ⇒ λ = 21
AP Inter 2nd Year Physics Study Material Chapter 1 Waves 24
In fundamental mode of frequency of bar
= frequency of wave = 4 kHz.
∴ v = \(\frac{\mathrm{v}}{\lambda}\) = \(\frac{\mathrm{v}}{2l}\) = 4kHz —– (1)
When clamp is moved to one end,
l = \(\frac{\lambda^{\prime}}{4}\) ⇒ λ’ = 4l
∴ V1 = \(\frac{\mathrm{v}}{\lambda}\) = \(\frac{\mathrm{v}}{4 \mathrm{l}}\) = \(\frac{4 \mathrm{kHz}}{2}\) = 2kHz
[∵ from (1)]

Question 6.
A closed organ pipe 70 cm long is sounded. If the velocity of sound is 331 m/s, what is the fundamental frequency of vibration of the air column ?
Solution:
I = 70 cm = 70 × 10-2m; v = 331 m/s ; v = ?
v = ?
v = \(\frac{v}{4 l}\) = \(\frac{331}{4 \times 70 \times 10^{-2}}\) = 118.2 Hz

Question 7.
A vertical tube is made to stand in water so that the water level can be adjusted. Sound waves of frequency 320 Hz are sent into the top of the tube. If standing waves are produced at two successive water levels of 20 cm and 73 cm, what is the speed of sound waves in the air in the tube ?
Solution:
v = 320 Hz; l1 = 20cm = 20 × 10-2
l2 = 73 cm = 73 × 10-2m; v = ?
v = 2v(l2 – l1)
= 2 × 320 (73 × 10-2 – 20 × 10-2)
∴ v = 339 m/s .

Question 8.
Two organ pipes of lengths 65cm and 70cm respectively, are sounded simultaneously. How many beats per second will be produced between the fundamental frequencies of the two pipes ? (Velocity of sound = 330 m/s).
Solution:
l1 = 65 cm = 0.65 m
2 = 70 cm = 0.7 m
v = 330 m/s
No. of beats per second ∆υ = υ1 – υ2
= \(\frac{v}{2 h}\) – \(\frac{\mathrm{v}}{2 l_2}\) = \(\frac{330}{2 \times 0.65}\) – \(\frac{330}{2 \times 0.7}\)
∴ ∆v = 253.8 – 235.8 = 18Hz

Question 9.
A train sounds its whistle as it approaches and crosses a level crossing. An observer at the crossing measures a frequency of 219 Hz as the train, approaches and a frequency of 184 Hz as it leaves. If the speed of sound is taken to be 340 m/s, find the speed of the train and the frequency of its whistle.
Solution:
When a whistling train approaches to rest observer,
AP Inter 2nd Year Physics Study Material Chapter 1 Waves 25
v’ = \(\left[\frac{v}{v-v_s}\right] v\) ——– (1)
When a whistling train away from rest observer
AP Inter 2nd Year Physics Study Material Chapter 1 Waves 26
v” = \(\left[\frac{v}{v+v_{\mathrm{S}}}\right] v\) —— (2)
Here v’ = 219 Hz; V” = 184 Hz;
v = 340 m/s
\(\frac{(1)}{(2)}\) ⇒ \(\frac{v^{\prime}}{v^{\prime \prime}}\) = \(\frac{\left(v+v_s\right)}{\left(v-v_s\right)}\)
\(\frac{219}{184}\) = \(\frac{340+v_{\mathrm{s}}}{340-v_{\mathrm{s}}}\)
219(340 – υs) = 184(340 + υs)
219 × 340 – 219 υs = 184 × 340 + 184 υs
403 υs = 35 × 340
∴ Velocity of train υs = 29.5 m/s
Frequency of whistle, v = v’ × \(\left[\frac{v-v_{\mathrm{S}}}{v}\right]\)
= 219 × \(\left[\frac{340-29.5}{340}\right]\)
= 199.98
∴ v = 200 Hz.

Question 10.
Two trucks heading in opposite directions with speeds of 60 kmph and 70 kmph respectively, approach each other. The driver of the first truck sounds his horn of frequency 400Hz. What frequency does the driver of the second truck hear ? (Velocity of sound = 330 m/ s). After the two trucks have passed each other, what frequency does the driver of the second truck hear ?
Solution:
vs = 60 kmph = 60 × \(\frac{5}{18}\) m/s = \(\frac{300}{18}\) m/s
v0 = 70 kmph = 70 × \(\frac{5}{18}\) m/s = \(\frac{350}{18}\) m/s
v = 400 Hz
When two trucks approach each other
AP Inter 2nd Year Physics Study Material Chapter 1 Waves 27
When two trucks crossed each other,
AP Inter 2nd Year Physics Study Material Chapter 1 Waves 28

Textual Exercises

Question 1.
A string of mass 2.50 kg is under a tension of 200 N. The length of the stretched string is 20.0 m. If the transverse jerk is caused at one end of the string, how long does the disturbance take to reach the other end ?
Answer:
Here M = 2.50 kg, T = 200 N, l = 20.0M
Mass per unit length; μ = \(\frac{\mathrm{M}}{l}\) = \(\frac{2.5}{20.0}\)
= 0.125 kg/m
Velocity V = \(\sqrt{\frac{\mathrm{T}}{\mu}}=\sqrt{\frac{200}{0.125}}\) = 40 m/s
Time taken by disturbance to reach the other end
t = \(\frac{l}{\mathrm{~V}}\) = \(\frac{20}{40}\) = 0.5s.

AP Inter 2nd Year Physics Study Material Chapter 1 Waves

Question 2.
A stone dropped from the top of a tower of height 300 m high splashes into the pond of water near the base of the tower. When is the splash heard at the top given that the speed of sound in air is 340 m s-1, (g = 9.8m s-2)
Answer:
Here, h = 300m, g = 9.8 m/s2, V = 340 m/s. If t1 = time taken by stone to strike the surface of water in the pond, then from
S = ut + \(\frac{1}{2}\) at2
300 = 0 + \(\frac{1}{2}\) × 9.8 \(\mathrm{t}_1^2\)
t1 = \(\sqrt{\frac{300}{4.9}}\) = 7.82s.
Time taken by sound to reach the top of tower t2 = \(\frac{\mathrm{h}}{\mathrm{v}}\)
= \(\frac{300}{400}\) = 0.88s
Total time after which splash of sound is heard = t1 + t2 = 7.82 + 0.88 = 8.70s

Question 3.
A steel wire has a length of 12.0 m and a mass of 2.10 kg. What should be the tension in the wire so that speed of a transverse wave on the wire equals the speed of sound in dry air at 20° C = 343 m s-1.
Answer:
Here, l = 12.0M, M = 2.10kg, T = ?
V = 343 m/s
Mass per unit length μ = \(\frac{\mathrm{M}}{l}\) = \(\frac{2.10}{12.0}\)
= 0.175 kg/m
As V = \(\sqrt{\frac{\mathrm{T}}{\mu}}\)
T = V2 . μ = (343)2 × 0.175 = 2.06 × 104N.

Question 4.
Use the formula υ = \(\sqrt{\frac{\gamma P}{\rho}}\) to explain why the speed of sound in air
a) is independent of pressure,
b) increases with temperature,
c) increases with humidity.
Answer:
a) Effect of pressure:
The speed of sound in gases, υ = \(\sqrt{\frac{\gamma \mathrm{P}}{\rho}}\)
At constant temperature, PV = constant
P\(\frac{m}{\rho}\) = constant ⇒ \(\frac{\mathrm{P}}{\rho}\) = constant
If P increases, ρ also increases. Hence speed of sound in air is independent of pressure.

b) Effect of temperature:
As PV = nRT, P\(\frac{\mathrm{m}}{\rho}\) = \(\frac{m}{M} R T\)
⇒ \(\frac{P}{\rho}\) = \(\frac{\mathrm{RT}}{\mathrm{M}}\)
∴ υ = \(\sqrt{\frac{R T}{M}}\)
Since R, M are constants υ ∝ \(\sqrt{\mathrm{T}}\)
∴ Velocity of sound in air depends on temperature.

c) Effect of humidity:
As υ = \(\sqrt{\frac{\gamma \mathrm{P}}{\rho}}\) ∴ υ ∝ \(\frac{1}{\sqrt{\rho}}\)
As the density of water vapour is less than density of dry air at STP. So the presence of moisture in air decreases the
density of air. Since the speed of sound is inversely proportional to the square root of density. So sound travels faster in moist air than dry air. Hence velocity of sound
V ∝ humidity

Question 5.
You have learnt that a travelling wave in one dimension is represented by a function y = f(x, t) where x and t must appear in the combination x – υt or x + υt, i.e. y = f(x ± υ t). Is the converse true? Examine if the following functions for y can possibly represent a travelling wave:
Answer:
No, the converse is not true. The basic requirement for a wave function to represent a travelling wave is that for all values of x & t, wave function must have a finite value.

Out of the given functions y, no one satisfies this condition therefore, none can represent a travelling wave.

Question 6.
A bat emits ultrasonic sound a frequency 1000 kHz in air. If the sound meets a water surface, what is the wavelength of
(a) the reflected sound,
(b) the transmitted sound?
Speed of sound in air is 340 m s-1 and in water 1486 m s-1.
Answer:
Here V = 100 KHz = 105Hz, Va = 340m/s, Vw = 1486 m/s-1
Wavelength of reflected sound, λa = \(\frac{\mathrm{V}^{\mathrm{a}}}{\mathrm{V}}\)
= \(\frac{340}{10^5}\) = 3.4 × 10-3 m
Wavelength of transmitted sound,
λw = \(\frac{V_w}{V}\) = \(\frac{1486}{10^5}\) = 1.486 × 10-2 m

AP Inter 2nd Year Physics Study Material Chapter 1 Waves

Question 7.
A hospital uses an ultrasonic scanner to locate tumours in a tissue. What is the wavelength of sound in the tissue in which the speed of sound is 1.7 km s-1? The operating frequency of the scanner is 4.2 MHz.
Answer:
λ = ? υ = 1.7 Km/s = 1700 ms-1
y = 4.2 MHz = 4.2 × 106Hz
λ = \(\frac{v}{v}\) = \(\frac{1700}{4.2 \times 10^6} \mathrm{~m}\) = 0.405 × 10-3 m
= 0.405 mm

Question 8.
A transverse harmonic wave on a string is described by y(x, t) = 3.0 sin (36 t + 0.018 x + π/4) where x and y are in cm and t in s. The positive direction of x is from left to right.
a) Is this a travelling wave or a stationary wave?
If it is travelling, what are the speed and direction of its propagation?
b) What are Its amplitude and frequency?
C) What is the initial phase at the origin?
d) What is the least distance between two successive crests in the wave?
Answer:
Compare the given equation with that of plane progressive wave of amplitude r, travelling with a velocity V from right to left.
y(x, t) = rsin\(\left[\frac{2 \pi}{\lambda}(v t+x)+\phi_0\right]\) ……… (1)
We find that
a) The given equation represents a transvërse harmonic wave travelling from right to left. It is ñot a stationary wave.

b) The given equation can be written as
Y(x, t) = 3.0sin[0.018(\(\frac{36}{0.018}\) + x) + \(\frac{\pi}{4}\)] ……… (2)
equating coefficient of t in the two
(1) & (2) we get. :
V = \(\frac{.36}{0.018}\) = 2000 cm/sec.
Obviously, r = 3.0 cm
Also, \(\frac{2 \pi}{\lambda}\) = 0.018
λ = \(\frac{2 \pi}{0.018} \mathrm{~cm}\)
Frequency, v = \(\frac{v}{\lambda}\) = \(\frac{2000}{2 \pi}\) × 0.018
= 5.7351.

c) Intial phase, φ0 = \(\frac{\pi}{4}\) radian.

d) Least distance between two successive crests of the wave =
Wave length, λ = \(\frac{2 \pi}{0.018 \mathrm{~cm}}\) = 349 cm,

Question 9.
For the wave described in the last problem plot the displacement (y) versus (t) graphs for x = 0.2 and 4 cm. What are the shapes of these graphs? In which aspects does the oscillatory motion in travelling wave differ from one point to another: amplitude, frequency or phase?
Answer:
The transverse harmonic wave is
y(x,t) = 3.0 sin[36t + 0.018x + \(\frac{\pi}{4}\)]
For x = 0
y(0, t) = 3.0 sin(36t + \(\frac{\pi}{4}\)) —— (i)
Here w = \(\frac{2 \pi}{T}\) = 36, T = \(\frac{2 \pi}{36}\) = \(\frac{\pi}{18}\)-sec.

For different values of t, we calculate y using eq(i). These values are tabulated below.
AP Inter 2nd Year Physics Study Material Chapter 1 Waves 29
On plotting y versus t graph, we obtain a sinusoidal curve as shown in fig.
AP Inter 2nd Year Physics Study Material Chapter 1 Waves 30
Similar graphs are obtained for x = 2cm & x = 4cm. The oscillary motion in travelling wave differs from one point to another only in terms of phase. Amplitude and frequency of oscillatory motion remain the same in all the three areas.

Question 10.
For the travelling harmonic wave
y(x, t) = 2.0 cos 2 π (10t – 0.0080 x + 0.35)
where x and y are in cm and t in s. Calculate the phase difference between oscillatory motion of two points separated by a distance of
a) 4 m,
b) 0.5 m,
c) λ/2,
d) 3λ/4,
Answer:
The given equation be written as
y = 2.0 cos[2π(10t – 0.0080x) + 2π × 0.35]
y = 2.0 cos[2π × (0.0080(\(\frac{10 \mathrm{t}}{0.0080}\) – x) + 0.7π]
Compare it with the standard equation of a travelling harmonic, we have
y = r.cos[\(\frac{2 \pi}{\lambda}(v t-x)+\phi_0\)
We get, \(\frac{2 \pi}{\lambda}\) = 2π × 0.0080
Further we know that phase diff. φ = \(\frac{2 \pi}{\lambda} \mathrm{x}\)
a) When x = 4m = 400 cm
φ = \(\frac{2 \pi}{\lambda}\) . x = 2π × 0.0080 × 400
= 6.4 π rad.

b) When x = 0.5 = 50 cm
φ = \(\frac{2 \pi}{\lambda}\) . x = 2π × 0.0080 × 50
= 0.8π rad.

c) When x = \(\frac{\lambda}{2}\)
φ = \(\frac{2 \pi}{\lambda} \times \frac{\lambda}{2}\) = π rad.

d) When x = \(\frac{3 \lambda}{4}\)
φ = \(\frac{2 \pi}{\lambda} \times \frac{3 \lambda}{4}\) = \(\frac{3 \lambda}{2} \mathrm{rad}\)

Question 11.
The transverse displacement of a string (clamped at its both ends) is given by y(x, t) = 0.06 sin\(\left(\frac{2 \pi}{3} x\right)\)cos (120 πt)
where x and y are in m and t in s. The length of the string is 1.5 m and its mass is 3.0 × 10-2 kg
Answer the following:
a) Does the function represent a travelling wave or a stationary wave?
b) Interpret the wave as a superposition of two waves travelling in opposite directions. What is the wavelength, frequency, and speed of each wave?
c) Determine the tension in the string.
Answer:
The given equation is
y(x, t) = 0.06 sin \(\frac{2 \pi}{3} \mathrm{x} \cos 120 \pi \mathrm{t}\)

a) As the equation involves harmonic functions of x and t seperately, it represents a stationary wave.

b) We know that when a wave pulse
y1 = r sin \(\frac{2 \pi}{\lambda}(v t+x)\) —– (i)
travelling along + direction of x-axis is super imposed by the reflected wave
y = y1 + y2 = -2rsin\(\frac{2 \pi}{\lambda}\) xcos \(\frac{2 \pi}{\lambda}\) vt is formed. ——- (ii)
Comparing (i) & (ii) we find that
\(\frac{2 \pi}{\lambda}\) = \(\frac{2 \pi}{3}\) ⇒ λ = 3m.
Also \(\frac{2 \pi}{\lambda} v\) = 120π (Or)
V = 60λ = 60 × 3 = 180m/s.
frequency, v = \(\frac{v}{\lambda}\) = \(\frac{180}{3}\) = 60 hertz.
Note that both the waves have same wave length, same frequency and same speed.

c) Velocity of transverse waves is
υ = \(\sqrt{\frac{T}{\mu}}\) (or) υ2 = T/μ
T = V2 × μ where μ = \(\frac{3 \times 10^{-2}}{1.5}\)
= 2 × 10-2 kg/m
T = (180)2 × 2 × 10-2 = 648 N.

Question 12.
i) For the wave on a string described in previous problem do all the points on the string oscillate with the same (a) frequency, (b) phase, (c) amplitude? Explain your answers.
(ii) What is the amplitude of a point 0.375 m away from one end?
Answer:
i) All the points on the string
a) have the same frequency except at the nodes (where frequency is cos θ),
b) have the same phase every where in one loop except at the nodes,
c) however, the amplitude of vibration at different points is different.

ii) From y(x, t) = 0.06 sin\(\left(\frac{2 \pi}{3} x\right)\) cos (120 πt)
The amplitude at x = 0.375 m is 0.06
sin \(\frac{2 \pi}{3} x \times 1\) = 0.06 × sin \(\frac{2 \pi}{3} \times 0.375\)
= 0.06sin\(\frac{\pi}{4}\) = \(\frac{0.06}{\sqrt{2}}\) = 0.042 M

AP Inter 2nd Year Physics Study Material Chapter 1 Waves

Question 13.
Given below are some functions of x and t to represent the displacement (transverse or longitudinal) of an elastic wave State which of these represent
(i) a travelling wave,
(ii) a stationary wave or
(iii) none at all:
a) y = 2 cos (3x) sin (10t)
b) y = \(2 \sqrt{x-v t}\)
c) y = 3 sin(5x – 0.5t) + 4 cos(5x – 0.5t)
d) y = cos x sin t + cos 2x sin 2t
Answer:
a) It represents a stationary wave as harmonic functions of x & t are contained separetely in the equation.
b) It cannot represent any type of wave.
c) It represents a progressive / travelling harmonic wave.
d) This equation is sum of two functions each representing a stationary wave. Therefore it represents superposition of two stationary waves.

Question 14.
A wire stretched between two right supports vibrates in its fundamental mode with a frequency of 45 Hz. The mass of the wire is 3.5 × 10-2 kg and its linear mass density is 4.0 × 10-2 kg m-1. What is (a) the speed of transverse wave on the string, and (b) the tension in the string?
Answer:
Here, v = 45Hz, μ = 3.5 × 10-2 kg
Mass/length = μ = 4.0 × 10-2 kg/m
l = \(\frac{\mu}{\mu}\) = \(\frac{3.5 \times 10^{-2}}{4.0 \times 10^{-2}}\) = \(\frac{7}{8}\)
As \(\frac{\lambda}{2}\) = l = \(\frac{7}{8}\) ∴ λ = \(\frac{7}{4}\)m = 1.75m
a) The speed of transverse wave
υ = vλ = 45 × 1.75 = 78.75 m/s.

b) As υ = \(\sqrt{\frac{\mathrm{T}}{\mu}}\)
∴ T = υ2 × μ = (78.75)2 × 4.0 × 10-2
= 248.06 N.

Question 15.
A metre-long tube open at one end, with a movable piston at the other end, shows resonance with a fixed frequency source (a tuning fork of frequency 340 Hz) when the tube length Is 25.5 cm or 79.3 cm. Estimate the speed of sound in air at the temperature of the experiment. The edge effect may be neglected.
Answer:
As there is a piston at one end of the tube, it behaves as a closed organ pipe, which produces odd harmonics only. Therefore the pipe is in resonance with the fundamental note at the third harmonic (79.3 cm is about 3 times 25.5 cm)
In the fundamental note = \(\frac{\lambda}{4}\) = l1 = 25.5
λ = 4 × 25.5 = 102 cm = 1.02m
Speed of sound in air.
υ = vλ = 340 × 1.02
= 346.0 m/s

Question 16.
A steel rod 100 cm long is clamped at its middle. The fundamental frequency of longitudinal vibrations of the rod are given to be 2.53 kHz. What is the speed of sound in steel?
Answer:
Here, l = 100 cm = 1 m, y = 2.53 KHz
= 2.53 × 103 Hz
When the rod is clamped at the middle, then in the fundamental mode of vibration of the rod, anode is formed at the middle and antinode is formed at each end.
Therefore, it is clear from fig
AP Inter 2nd Year Physics Study Material Chapter 1 Waves 31
l = \(\frac{\lambda}{4}\) + \(\frac{\lambda}{4}\) + \(\frac{\lambda}{2}\)
λ = 2l = 2m
As v = λl
v = 2.53 × 103 × 2
= 5.06 × 103 ms-1

Question 17.
A pipe 20 cm long is closed at one end. Which harmonic mode of the pipe is resonantly excited by a 430 Hz source? Will the same source be in resonance with the pipe if both ends are open? (speed of sound in air is 340 m s-1).
Answer:
Here l = 20 cm = 0.2m, vn = 430 Hz,
υ = 340 m/s
The frequency of nth normal mode of vibration of closed pipe is
vn = (2n – 1)\(\frac{v}{4l}\)
∴ 430 = (2n – 1)\(\frac{340}{4 \times 0.2}\)
2n – 1 = \(\frac{430 \times 4 \times 0.2}{340}\) = 1.02
2n = 2.02, n = 1.01
Hence it will be the 1st normal mode of vibration. In a pipe, open at both ends we have
vn = n × \(\frac{\mathrm{v}}{2l}\) = \(\frac{\mathrm{n} \times 340}{2 \times 0.2}\) = 430.
∴ n = \(\frac{430 \times 2 \times 0.2}{340}\) = 0.5
As n has to be an integer, therefore open organ pipe cannot be in resonance with the source.

Question 18.
Two sitar strings A and B playing the note ‘Ga’ are slightly out of tune and produce beats of frequency 6 Hz. The tension in the string A is slightly reduced and the beat frequency is found to reduce to 3 Hz. If the original frequency of A is 324Hz. What is the frequency of B?
Answer:
Let original frequency of sitar string A be na & original frequeny of sitar string B be nb.
As number of beats / sec = 6
∴ nb = na ± 6 = 330 (or) 318Hz.
When tension in A is reduced, its frequency reduces (∴ n ∝ \(\sqrt{T}\))
As number of beats/sec decreases to 3 therefore, frequency of B = 324 – 6
= 318Hz.

Question 19.
Explain why (or how):
a) in a sound wave, a displacement node is a pressure antinode and vice versa,
b) bats can ascertain distances, directions, nature and sizes of the obstacles without any ‘eyes”.
c) a violin note and sitar note may have the same frequency, yet we can distinguish between the two notes,
d) Soils can support both longitudinal and transverse waves, but only longitudinal waves can propagate in gases and
e) The shape of a pulse gets distorted during propagation In a dispersive medium.
Answer:
a) Node (N) is a point where the amplitude of oscillation is 0. (and pressure is maximum)
Antinode (A) is a point where the amplitude of oscillation is maximum (and pressure is min).
These nodes & antinodes do not coincide with pressure nodes & antinodes.
Infact, N coincides with pressure antinode and A coincides with pressure node, as is clear from the definitions.

b) Bats emit ultrasonic wave of large frequencies, when these waveš are reflected from the obstacles in their path,
they give them the idea about the distance, direction, size & nature of the obstacle.

c) Though the violin note and sitar note have the same frequency, yet the over tones produced and their reactive strengths are different in the two flotes that is why we can distinguish between the two notes.

d) This is because solids have both, the elasticity of volume and elasticity of shape where as gases have only the volume elasticity.

e) A sound pulse is a combination of waves of different wavelengths. As waves of different wavelengths travel in a disperse medium with different velocities, therefore the shape of the pulse gets distorted.

Question 20.
A train, standing at the outer signal of a railway station blows a whistle of frequency 400 Hz in still air.
i) What is the frequency of the whistle for a platform observer when the train
(a) approaches the platform with a speed of 10ms-1,
(b) recedes from the platform with a speed of 10 m s-1 ?
ii) What is the speed of sound in each case ? The speed of sound in still air can be taken as 340 m s-1.
Answer:
i) Here, y = 400 Hz, υ = 340 m/s
a) Train approaches the platform
υs = 10m/s
v’ = \(\frac{v}{v-v_s}\) = \(\frac{340 \times 400}{340-10}\) = 412.12 Hz.

b) Train recedes from the platform
υs = 10m/s
v’ = \(\frac{v \times v}{v \times v_s}\) = \(\frac{340 \times 400}{340+10}\)
= 388.6Hz

ii) The speed of sound in each case is the same = 340 m/s

AP Inter 2nd Year Physics Study Material Chapter 1 Waves

Question 21.
A train, stañdingin a station-yard, blows a whistle of frequency 400 Hz in still air. The wind starts blowing in the direction from the yard to the station with a speed of 10m s-1. What are the frequency, wavelength and speed of sound for an observer standing on the station’s platform ? Is the situation exactly identical to the case when the air is still and the observer runs towards the yard at a speed of 10 m s-1 ? The speed of sound in still air can be taken as 340 m-1
Answer:
Here y = 400 Hz, υm = 10ms-1, υ = 340m/s
As the wind is blowing in the direction of sound, therefore effective speed of sound
= υ + υm = 340 + 10 = 350m/s
As the source & Iistner both are at rest, therefore, frequency remains unchanged
i.e. v = 400 Hz.
Wavelength, λ = \(\frac{v+v_m}{v}\) = \(\frac{350}{400}\)
= 0.875 M.
When air is still, υm = 0
Speed of observer υ1 = 10m/s υs = 0
As observer moves toward the source
υ’ = \(\frac{\left(v+v_l\right)}{v} \times v\) = \(\frac{(340+10)}{340} \times 400\)
= 411.76 Hz.
As source is at rest, wavelength does not change
i.e, λ’ = λ = 0.875M.
Also, speed of sound = υ + υm = 340 + 0
= 340 m/s
The situations in the two cases are entirely different.

Additional Exercises

Question 1.
A travelling harmonic wave on a string is described by
y(x, t) = 7.5 sin (0.0050x + 12t + π/4)
a) What are the displacement and velocity of oscillation of a point at x = 1 cm, and t = 1 s? Is this velocity equal to the velocity of wave propagation?
b) Locate the points of the string which have the same transverse displacements and velocity as the x = 1cm
point at t = 2 s, 5 s and 11 s.
Answer:
a) The travelling harmonic wave is y(x, t)
AP Inter 2nd Year Physics Study Material Chapter 1 Waves 32
From (1), y(1, 1) = 7.5 sin(732.55°)
= 7.5 sin (720 + 12.55°)
= 7.5 sin 12.55° = 7.5 × 0.2173 = 1.63 cm
velocity of oscillation, v = \(\frac{d y}{d t} y(1,1)\)
= \(\frac{\mathrm{d}}{\mathrm{dt}}\)[7.5 sin (0.005x + 12t + \(\frac{\pi}{4}\))
= 7.5 × 12 cos (0.005x + 12t + \(\frac{\pi}{4}\))
At x = 1 cm, t = 1 sec
= 7.5 × 12 cos (o.oo5 + 12 + \(\frac{\pi}{4}\))
= 90 cos (732.55°).
= 90 cos(720 + 12.55°)
= 90 cos (12.55°)
= 90 × 0.9765
= 87.89 cm/s.
Comparing the given equation with the standard form
y(x, t) = r sin\(\left[\frac{2 \pi}{\lambda}(v t+x)+\phi_0\right]\)
We get r = 7.5 cm, \(\frac{2 \pi v}{\lambda}\) = 12 (or) 2πV = 12
V = \(\frac{6}{\pi}\)
2\(\frac{\pi}{\lambda}\) = 0.005
∴ λ = \(\frac{2 \pi}{0.005}\) = \(\frac{2 \times 3.14}{0.005}\) = 1256 cm
= 12.56 m.
Velocity of wave propagation, υ = Vλ
= \(\frac{6}{\pi}\) × 12.56 .
= 24 m/s.
We find that velocity at x = 1 cm t = 1 sec is not equal to velocity of wave propagation.

b) Now, all points which are at a distance of ±λ, ± 2λ, ± 3λ from x = 1 cm will have same transverse displacement and velocity. As λ = 12.56 m, therefore, all points at distances ± 12.6m, ± 25.2 m displacement and velocity As λ = 12.56m, therefore all points at distances ± 12.6m, ± 25.2m, ± 37.8m from x = 11m will have same displacement & velocity at x = 1 cm point at t = 25.55 & 115s.

Question 2.
A narrow sound pulse (for example, a short pip by a whistle) is sent across a medium. (a) Does the pulse have a definite (I) frequency, (ii) wavelength, (iii) speed of propagation? (b) If the pulse rate is 1 after every 20 s (that is the whistle blown for a split of second after every 20 s), is the frequency of the note produced by the whistle equal to 1/20 or 0.05 Hz?
Answer:
a) A short pip by a whistle has neither a definite wavelength nor a definite frequency. However its speed of propagation is fixed, being equal to speed of sound in air.

b) No, frequency of the note produced by whistle is not 1/20 = 0.05 Hz. Rather 0.05 Hz is the frequency of repetition of the short pipe of the whistle.

Question 3.
One end of a long string of linear mass density 8.0 × 10-3 kg m-1 is connected to an electrically driven tuniúg fork of frequency 256 Hz. The other end passes over a pulley and is tied to a pan containing a mass of 90 kg. The pulley end absorbs all the Incoming energy so that reflected waves at this end have negligible amplitude. At t = 0, the left end (fork end) of the string x = 0 has zero transverse displacement (y = 0) and is moving along positive y-direction. The amplitude of the wave is 5.0 cm. Write down the transverse displacement y as function of x and t that describes the wave on the string.
Answer:
Here, μ = 8.0 × 10-3 kg/m, y = 256 Hz, T = 90kg = 90 × 9.8 = 882N.
Amplitude of wave, r = 5.0 1m = 0.05m.
As the wave propagation along the string is a transverse travelling wave, the velocity of the wave is given by
AP Inter 2nd Year Physics Study Material Chapter 1 Waves 33

As the wave is propagating along x direction, the equation of the wave is
y(x, t) = r sin (ωt – kx)
= 0.05 sin (1.61 × 103t – 4.84x)
Here x, y are in mt & t in sec

AP Inter 2nd Year Physics Study Material Chapter 1 Waves

Question 4.
A SONAR system fixed in a submarine operates at a frequency 40.0 kHz. An enemy submarine moves towards the SONAR with a speed of 360 km h-1. What is the frequency of sound reflected by the submarine ? Take the speed of sound in water to be 1450 m s-1.
Answer:
Here, frequency of SONAR,
v = 40 KHz = 40 × 103 Hz.
Speed of observer / enemy’s submarine
υ1 = 360km/h .
= 360 × \(\frac{5}{18}\)m/s = 100m/s.
Speed of sound wave in water; υ = 1450 m/s.
As the source is at rest & observer is moving towards the source, therefore, apparent frequency received by enemy submarine
v’ = \(\frac{\left(v+v_1\right) v}{v}\)
= \(\frac{(1450+100) 40 \times 10^3 \mathrm{~Hz}}{1450}\)
= 4.27 × 104 Hz.
This frequency is reflected by enemy submarine (source) and is observed by SONAR. Therefore in this case,
υs = 360 km/s = 100 m/s, υ1 = 0
∴ Apparent frequency, v11 = \(\frac{v \times v}{v_i-v_s}\)
= \(\frac{1450 \times 4.276 \times 10^4}{1450-10}\)
= 4.59 × 104 Hz = 45.9 Hz.

Question 5.
Earthquakes generate sound waves inside the earth. Unlike a gas, the earth can experience both transverse (S) and longitudinal (P) sound waves. Typically the speed of S wave is about 4.0 km s-1 and that of P wave is 8.0 km s-1. A seismograph records P and S waves from an earthquake. The first P wave arrives 4 min before the first S wave. Assuming the waves travel in straight line, at what distance does the earthquake occur?
Answer:
Let υ1, υ2 be the velocity of S waves & P waves & t1, t2 be the time taken by these waves to travel to the position of seismograph. If l is the distance of occurrence of earth quake from the seismograph, then
l = υ1t1 = υ2t2 ——- (i)
now υ1 = 4 km/s & υ2 = 8 km/s .
∴ 4t1 = 8t2 (or) t1 = 2t2 ——- (ii)
Also t1 – t2 = 4min = 240s.
using (iii), 2t2 – t2 = 240s, t2 = 240s
t1 = 2t2 = 2 × 240 = 480s.
Now from (i) l = υ1t1 = 4 × 480 = 1920 km.
Hence earthquake occurs 1920 km away from the seismograph.

Question 6.
A bat is flitting about in a cave, navigating via ultrasonic beeps. Assume that the sound emission frequency of the bat is 40 kHz. During one fast swoop directly toward a flat wall surface, the bat is moving at 0.03 times the speed of sound in air. What frequency does the bat hear reflected off the wall?
Answer:
Here, the frequency of sound emitted by the bat, v = 40 kHz.
velocity of bat, υs = 0.03υ, where υ is velocity of sound.
Apparent frequency of sound striking the wall.
v’ = \(\frac{v \times v}{v-v_s}\) = \(\frac{v}{v-0.03 v}\) × 40 kHxz
= \(\frac{40}{0.97}\) kHZ.
This frequency is reflected by the wall & is received by the bat moving towards the wall, So υs = 0.
υ1 = 0.03 υ
v’ = \(\frac{\left(v+v_1\right) v^{\prime}}{v}\) = \(\frac{(v+0.03 v)}{v}\left(\frac{40}{0.97}\right)\)
= \(\frac{1.03}{0.97} \times 40 \mathrm{kHz}\)
= 42.47 kHz

AP Inter 2nd Year Physics Important Questions Chapter 13 Atoms

Students get through AP Inter 2nd Year Physics Important Questions 13th Lesson Atoms which are most likely to be asked in the exam.

AP Inter 2nd Year Physics Important Questions 13th Lesson Atoms

Very Short Answer Questions

Question 1.
What is the physical meaning of negative energy of an electron’ ?
Answer:
The ‘negative energy of an electron’ indicates that the electron is bound to the nucleus due to force of attraction.

Question 2.
Sharp lines are present in the spectrum of a gas. What does this indicate ?
Answer:
Sharp lines in the spectrum of gas, indicates bright lines against dark background.

AP Inter 2nd Year Physics Important Questions Chapter 13 Atoms

Question 3.
Name a physical quantity whose dimensions are the same as those of angular momentum.
Answer:
Planck’s constant.

Question 4.
What is the difference between α – particle and helium atom ?
Answer:
Alpha particle

  1. It is a + 2e charged Helium nucleus.
  2. It contains 2 protons and 2 neutrons.

Helium atom

  1. It has no charge.
  2. It contains 2 protons, 2 electrons and 2 neutrons.

Question 5.
Among alpha, beta and gamma radiations, which get affected by the electric field ?
Answer:
Alpha and Beta radiations are get affected by the electric field.

Question 6.
What do you understand by the phrase ground state atom ?
Answer:
If the electron is present in the ground state, it is called ground state atom.

AP Inter 2nd Year Physics Important Questions Chapter 13 Atoms

Question 7.
Why does the mass of the nucleus not have any significance in scattering in Rutherford’s experiment ?
Answer:
The size of the atom is 10-10 m and size of the nucleus is 10-15 m. Hence atom has large empty space. So the mass of nucleus has no significance in Rutherford’s scattering experiment.

Question 8.
The Lyman series of hydrogen spectrum lies in the ultraviolet region. Why ? [A.P. Mar. 15]
Answer:
The calculated values of wavelengths lie in the ultraviolet region of the spectrum will agree with the values of wavelengths observed experimentally by Lyman.

Question 9.
Write down a table giving longest and shortest wavelengths of different spectral series.
Answer:
Wavelength limits of some spectral series of hydrogen.
AP Inter 2nd Year Physics Important Questions Chapter 13 Atoms 1

Question 10.
Give two drawbacks of Rutherford’s atomic model.
Answer:
Drawbacks of Rutherford’s atomic model:

  1. As the revolving electron loses energy continuously, it must spiral inwards and eventually fall into the nucleus. But matter .is stable, we cannot expect the atom collapse.
  2. The atoms should emit continuous spectrum, but what we observe is only a line spectrum.

AP Inter 2nd Year Physics Important Questions Chapter 13 Atoms

Question 11.
If the kinetic energy of revolving electron in an orbit is K, what is its potential energy and total energy ?
Answer:
For an electron revolving round the nucleus, total energy is always negative and it is numerically equal to kinetic energy.
∴ Total energy = -Kinetic energy = -K
Potential energy is always negative and PE = 2 × TE = -2K

Short Answer Questions

Question 1.
What is impact parameter and angle of scattering ? How are they related to each other ?
Answer:

  1. Impact parameter (b) : Impact parameter is defined as the perpendicular distance of the initial velocity vector of the alpha particle from the central line of the nucleus, when the particle is far away from the nucleus of the atom.
  2. Scattering angle (θ) : The scattering angle (θ) is the angle between the asymtotic direction of approach of the α – particle and the asymtotic direction in which it receeds.
  3. The relation between b and θ is b = \(\frac{1}{4 \pi \varepsilon_0} \frac{\mathrm{Ze}^2}{\mathrm{E}}\) cot \(\frac{\theta}{2}\) where E = K.E of α – particle = \(\frac{1}{2}\) mυ2.

Question 2.
Explain the distance of closest approach and impact parameter.
Answer:
Distance of closest approach :

  1. Suppose an α-particle with initial kinetic energy (K.E) is directed towards the centre of the nucleus of an atom.
  2. On account of Coulomb’s repulsive force between nucleus and alpha particle, kinetic energy of alpha particle goes on decreasing and in turn, electric potential energy of the particle goes on increasing.
  3. At certain distance d’ from the nucleus, K. E of α-particle reduces to zero. The particle stops and it cannot go closer to the nucleus. It is repelled by the nucleus and therefore it retraces its path, turning through 180°.
  4. Therefore, the distance d is known as the distance of closest of approach.
    AP Inter 2nd Year Physics Important Questions Chapter 13 Atoms 2
    The closest distance of approach,
    d = \(\frac{1}{4 \pi \varepsilon_0} \times \frac{Z e^2}{\left(\frac{1}{2} m v^2\right)}\)
  5. Impact parameter (b) : Impact parameter is defined as the ⊥r distance of the initial velocity vector of the α – particle from the central line of the nucleus, when the particle is far away from the nucleus of the atom.
    AP Inter 2nd Year Physics Important Questions Chapter 13 Atoms 3

AP Inter 2nd Year Physics Important Questions Chapter 13 Atoms

Question 3.
Describe Rutherford atom model. What are the draw backs of this model ?
Answer:
Rutherford atom model: The essential features of Rutherford’s nuclear model of the atom or planetary model of the atom are as follows :

  1. Every atom consists of tiny central core, called the atomic nucleus, in which the entire positive charge and almost entire mass of the atom are concentrated.
  2. The size of nucleus is of the order of 10-15m, which is very small as compared to the size of the atom which is of the order of 10-10 m.
  3. The atomic nucleus is surrounded by certain number of electrons. As atom on the whole is electrically neutral, the total negative charge of electrons surrounding the nucleus is equal to total positive charge on the nucleus.
  4. These electrons revolve around the nucleus in various circular orbits as do the planets around the sun. The centripetal force required by electron for revolution is provided by the electrostatic force of attraction between the electrons and the nucleus.

Drawbacks : According to classical E.M. theory,

  1. the revolving electron loses energy continuously, it must spiral inwards and eventually fall into the nucleus. As matter is stable, we cannot expect the atoms to collapse.
  2. since the frequency of radiation emitted is the same as the frequency of revolution, the atom should radiate a continuous spectrum, but what we observe is only a line spectrum.

Question 4.
What are the limitations of Bohr’s theory of hydrogen atom ? [A.P. Mar. 17; Mar. 14]
Answer:
Limitations of Bohr’s theory of Hydrogen atom :

  1. This theory is applicable only to simplest atom like hydrogen, with z = 1. The theory fails in case of atoms of other elements for which z > 1.
  2. The theory does not explain why orbits of electrons are taken as circular, while elliptical orbits are also possible.
  3. Bohr’s theory does not say anything about the relative intensities of spectral lines.
  4. Bohr’s theory does not take into account the wave properties of electrons.

AP Inter 2nd Year Physics Important Questions Chapter 13 Atoms

Question 5.
Write a short note on Debroglie’s explanation of Bohr’s second postulate of quantization. [T.S. Mar. 17]
Answer:
Debroglie’s explanation of Bohr’s second postulate of quantization :

  1. The second postulate of Bohr atom model says that angular momentum of electron orbiting around the nucleus is quantized i.e., mυr = \(\frac{\mathrm{nh}}{2 \pi}\) where n = 1, 2, 3,….
  2. According to Debroglie, the electron in its circular orbit, as proposed by Bohr, must be seen as a particle wave.
  3. When a string fixed at two ends is plucked, a large number of wavelengths are excited and standing wave is formed.
  4. It means that in a string, standing waves form when total distance travelled by a wave down the string and back is an integral number of wavelengths.
  5. According to Debroglie, a stationary orbit is that which contains an integral number of Debrogile waves associated with the revolving electron.
  6. For an electron revolving in nth circular orbit of radius rn, total distance covered = circumference of the orbit = 2πrn
    ∴ For permissible orbit, 2πrn = nλ
  7. According to Debrogile, λ = \(\frac{h}{m v_n}\)
    Where υn is speed of electron revolving in nth orbit
    ∴ 2πrn = \(\frac{\mathrm{nh}}{\mathrm{m} v_{\mathrm{n}}}\)
    nrn = \(\frac{\mathrm{nh}}{2 \pi}=\mathrm{n}\left(\frac{\mathrm{h}}{2 \pi}\right)\)
    i.e., angular momentum of electron revolving in nth orbit must be an integral multiple of \(\frac{\mathrm{h}}{2 \pi}\), which is the quantum condition proposed by Bohr in second postulate.

AP Inter 2nd Year Physics Important Questions Chapter 13 Atoms

Question 6.
Explain the different types of spectral series in hydrogen atom. [T.S. Mar. 16, A.P. Mar. 15]
Answer:
The atomic hydrogen emits a line spectrum consisting of five series.

  1. Lyman series : v = Rc \(\left(\frac{1}{1^2}-\frac{1}{n^2}\right)\) where n = 2, 3, 4
  2. Balmer series : v = Rc \(\left(\frac{1}{2^2}-\frac{1}{n^2}\right)\) where n = 3, 4, 5.
  3. Paschen series : v = Rc \(\left(\frac{1}{3^2}-\frac{1}{n^2}\right)\) where n = 4, 5, 6
  4. Brackett series: v = Rc \(\left(\frac{1}{4^2}-\frac{1}{n^2}\right)\) where n = 5, 6, 7
  5. Pfund series: v = Rc \(\left(\frac{1}{5^2}-\frac{1}{n^2}\right)\) where n = 6, 7, 8

Long Answer Questions

Question 1.
State the basic postulates of Bohr’s theory of atomic spectra. Hence obtain an expression for the radius of orbit and the energy of orbital electron in a hydrogen atom. [A.P. Mar. 16]
Answer:
a) Basic postulates of Bohr’s theory are

  1. The electron revolves round a nucleus in an atom in various orbits known as stationary orbits. The electrons cannot emit radiation when moving in their own stationary levels.
  2. The electron can revolve round the nucleus only in allowed orbits whose angular momentum is the integral multiple of \(\frac{\mathrm{h}}{2 \pi}\)
    i.e., mυnrn = \(\frac{\mathrm{nh}}{2 \pi}\) …………… (1)
    where n = 1, 2, 3 …………..
    If an electron jumps from higher energy (E2) orbit to the lower energy (E1) orbit, the difference of energy is radiated in the form of radiation.
    i.e., E = hv = E2 – E1 ⇒ v = \(\frac{E_2-E_1}{h}\) …………… (2)

b) Energy of emitted radiation: In hydrogen atom, a single electron of charge – e, revolves around the nucleus of charge e in a circular orbit of radius rn.
1) K.E. of electron : For the electron to be in circular orbit, centripetal force = The electrostatic force of attraction between the electron and nucleus.
AP Inter 2nd Year Physics Important Questions Chapter 13 Atoms 4
3) Radius of the oribit: Substituting the value of (6) in (2),
\(\frac{\mathrm{m}}{\mathrm{r}_{\mathrm{n}}}\left(\frac{\mathrm{n}^2 \mathrm{~h}^2}{4 \pi^2 \mathrm{r}_{\mathrm{n}}^2 \mathrm{~m}^2}\right)=\frac{\mathrm{Ke}^2}{\mathrm{r}_{\mathrm{n}}^2}\)
rn = \(\frac{n^2 h^2}{4 \pi^2 m K e^2}\) ……………. (1)
∴ rn = 0.53 n2

4) Total energy (En) : Revolving electron possess K.E. as well as P.E.
AP Inter 2nd Year Physics Important Questions Chapter 13 Atoms 5

Problems

Question 1.
The radius of the first electron orbit of a hydrogen atom is 5.3 x HTum. What is the radius of the second orbit ?
Solution:
∴ rn ∝ n2
\(\) ⇒ r2 = 4r1
∴ r2 = 4 × 5.3 × 10-11 = 2.12 × 10-10m.

Question 2.
The total energy of an electron in the first excited state of the hydrogen atom is -3.4eV. What is the potential energy of the electron in this state?
Solution:
In 1st orbit, E = -3.4 eV
Total energy E = \(\frac{\mathrm{KZe}^2}{2 \mathrm{r}}-\frac{\mathrm{KZe}^2}{\mathrm{r}}\)
\(\frac{\mathrm{KZe}^2}{\mathrm{r}}\) = U(say)
E = \(\frac{\mathrm{U}}{2}\) – u = \(\frac{\mathrm{-U}}{2}\)
U = -2E
∴ U = -2 × -3.4 = 6.8 eV.

AP Inter 2nd Year Physics Important Questions Chapter 13 Atoms

Question 3.
The total energy of an electron in the first excited state of hydrogen atom is -3.4eV. What is the kinetic energy of the electron in this state ?
Solution:
In Hydrogen like atom, we know that
K = – Total energy E
Here E = – 3.4eV
∴ K = -(-3.4) = 3.4 eV . . ,

Question 4.
Prove that the ionisation energy of hydrogen atom is 13.6 eV.
Solution:
n = 1 corresponds to ground state.
E = \(\frac{-13.6}{\mathrm{n}^2}\) eV
E = \(\frac{-13.6}{\mathrm{1}^2}\) eV
E = -13.6 eV
The minimum energy required to free the electron from the ground state of hydrogen atom = 13.6 eV.
∴ Ionisation energy of hydrogen atom = 13.6 eV

Question 5.
Calculate the ionization energy for a lithium atom.
Solution:
For 3Li7 atom, Z = 3, n = 2 [∵Li = 1s2 2s1]
En = \(\frac{13.6 \mathrm{Z}^2}{\mathrm{n}^2}\) eV
= \(\frac{13.6 \times(3)^2}{4}\) = 30.6 eV
∴ Ionization energy of Lithium = 30.6eV

Question 6.
The wavelength of the first member of Lyman series is 1216 A. Calculate the wavelength of second member of Balmer series.
Solution:
AP Inter 2nd Year Physics Important Questions Chapter 13 Atoms 6

Textual Examples

Question 1.
In the Rutherford’s nuclear model of the atom, the nucleus (radius about 10-15m) is analogous to the sun about which the electron move in orbit (radius ≈ 10-10m) like the earth orbits around the sun. If the dimensions of the solar system had the same proportions as those of the atom, would the earth be closer to or farther away from the sun than actually it is ? The radius of earth’s orbit is about 1.5 × 1011m. The radius of sun is taken as 7 × 108 m.
Solution:
The ratio of the radius of electron’s orbit to the radius of nucleus is (10-10 m)/(10-15 m) = 105, that is, the radius of the electron’s orbit is 105 times larger than the radius of nucleus. If the radius of the earth’s orbit around the sun were 105 times larger than the radius of the sun, the radius of the earth’s orbit would be 105 × 7 × 108 m = 7 × 1013 m. This is more than 100 times greater than the actual orbital radius of earth. Thus, the earth would be much farther away from the sun.

It implies that an atom contains a much greater fraction of empty space than our solar system does.

AP Inter 2nd Year Physics Important Questions Chapter 13 Atoms

Question 2.
In a Geiger-Marsden experiment, what is the distance of closest approach to the nucleus of a 7.7 MeV α-particle before it comes momentarily to rest and reverses its direction?
Solution:
The key idea here is that throughout the scattering process, the total mechanical energy of the system consisting of an α-particle and a gold nucleus is conserved. The system’s initial mechanical energy is Ei; before the particle and nucleus interact, and it is equal to its mechanical energy Ef when the α-particle momentarily stops. The initial energy Et is just the kinetic energy K of the incoming α-particle. The final energy Ef is just the electric potential energy U of the system. The potential energy U can be calculated from Equation.
F = \(\frac{1}{4 \pi \varepsilon_0} \frac{(2 \mathrm{e})(\mathrm{Ze})}{\mathrm{r}^2}\)
Let d be the centre-to-centre distance between the α-particle and the gold nucleus when the α-particle is at its stopping point. Then we can write the conservation of energy
Ei = Ef as
K = \(\frac{1}{4 \pi \varepsilon_0} \frac{(2 \mathrm{e})(\mathrm{Ze})}{\mathrm{d}}=\frac{2 \mathrm{Ze}^2}{4 \pi \varepsilon_0 \mathrm{~d}}\)
Thus the distance of closest approach d is given by
d = \(\frac{2 \mathrm{Ze}^2}{4 \pi \varepsilon_0 \mathrm{~K}}\)
The maximum kinetic energy found in α-particles of natural origin is 7.7 MeV or 1.2 × 10-12 J. Since 1/4πε0 = 9.0 × 109N m2/C2. Therefore with e = 1.6 × 10-19C, we have,
d = \(\frac{(2)\left(9.0 \times 10^9 \mathrm{Nm}^2 / \mathrm{C}^2\right)\left(1.6 \times 10^{-19} \mathrm{C}\right)^2 \mathrm{Z}}{1.2 \times 10^{-12} \mathrm{~J}}\)
= 3.84 × 10-16 Zm
The atomic number of foil material gold is Z = 79, so that
d (Au) = 3.0 × 10-14m = 30 fm. (1 fm (i.e. fermi) = 10-15m.)
The radius of gold nucleus is, therefore, less than 3.0 × 10-14 m. This is not in very good agreement with the observed result as the actual radius of gold nucleus is 6 fm. The cause of discrepancy is that the distance of closest approach is considerably larger than the sum of the radii of the gold nucleus and the a-particle. Thus, the α-particle reverses its motion without ever actually touching the gold nucleus.

Question 3.
It is found experimentally that 13.6 eV energy is required to separate a hydrogen atom into a proton and an electron. Compute the orbital radius and the velocity of the electron in a hydrogen atom.
Solution:
Total energy of the electron in hydrogen atom is – 13.6 eV = -13.6 × 1.6 × 10-19J
= -2.2 × 10-18 J.
Thus from equation, E = – \(\frac{\mathrm{e}^2}{8 \pi \varepsilon_0 \mathrm{r}}\) we have
– \(\frac{\mathrm{e}^2}{8 \pi \varepsilon_0 \mathrm{r}}\) = 2.2 10-18 J
This gives the orbital radius
r = – \(\frac{\mathrm{e}^2}{8 \pi \varepsilon_0 \mathrm{r}}\) = \(\frac{\left(9 \times 10^9 \mathrm{Nm}^2 / \mathrm{C}^2\right)\left(1.6 \times 10^{-19} \mathrm{C}\right)^2}{(2)\left(-2.2 \times 10^{-18} \mathrm{~J}\right)}\)
= 5.3 × 10-11 m.
The velocity of the revolving electron can be computed from Equation r = –\(\frac{e^2}{4 \pi \varepsilon_0 r m v^2}\)
with
m = 9.1 × 10-31kg,
υ = \(\frac{\mathrm{e}}{\sqrt{4 \pi \varepsilon_0 \mathrm{mr}}}\) = 2.2 × 106 m/s

AP Inter 2nd Year Physics Important Questions Chapter 13 Atoms

Question 4.
According to the classical electromagnetic theory, calculate the initial frequency of the light emitted by the electron revolving around a proton in hydrogen atom.
Solution:
From Examjple 3 we know that velocity of electron moving around a proton in hydrogen atom in an orbit of radius 5.3 × 10-11 m is 2.2 × 10-6 m/s. Thus, the frequency of the electron moving around the proton is
v = \(\frac{v}{2 \pi \mathrm{r}}=\frac{2.2 \times 10^6 \mathrm{~m} \mathrm{~s}^{-1}}{2 \pi\left(5.3 \times 10^{-11} \mathrm{~m}\right)}\)
≈ 6.6 × 1015 Hz.
According to the classical electromagnetic theory we know that the frequency of the electromagnetic waves emitted by the revolving electrons is equal to the frequency of its revolution around the nucleus. Thus the initial frequency of the light emitted is 6.6 × 1015Hz.

Question 5.
A 10 kg satellite circles earth once every 2 h in an orbit having a radius of 8000 km. Assuming that Bohr’s angular momentum postulate applies to satellites just as it does to an electron in the hydrogen atom, find the quantum number of the orbit of the satellite.
Solution:
From equation, we have
nrn = nh/2π
Here m = 10 kg and rn = 8 × 106 m. We have the time period T of the circling satellite as 2h. That is T = 7200 s.
Thus the velocity υn = 2π rn/T
The quantum number of the orbit of satellite
n = (2π rn)2 × m(T × h)
Substituting the values,
n = (2π × 8 × 106m)2 × 10/(7200 s × 6.64 × 10-34 J s)
= 5.3 × 1045
Note that the quantum number for the satellite motion is extremely large. In fact for such large quantum numbers the results of quantisation conditions tend to those of classical physics.

AP Inter 2nd Year Physics Important Questions Chapter 13 Atoms

Question 6.
Using the Rydberg formula, calculate the wavelengths of the first four spectral lines in the Lyman series of the hydrogen spectrum.
Solution:
The Rydberg formula is
hc/λif = \(\left(\frac{1}{n_{\mathrm{f}}^2}-\frac{1}{\mathrm{n}_{\mathrm{i}}^2}\right)\)
The wavelengths of the first four lines in the Lyman series correspond to transitions from ni = 2, 3, 4, 5 to nf = 1. We know that
\(\frac{\mathrm{me}^4}{8 \varepsilon_0^2 \mathrm{~h}^2}\) = 13.6 eV = 21.76 × 10-19 J
Therefore,
AP Inter 2nd Year Physics Important Questions Chapter 13 Atoms 7
Substituting, ni = 2, 3, 4, 5 we get λ21 = 1218 Å, λ31 = 1028 Å, λ41 = 974.3 Å, and λ51 = 951.4 Å.

AP Inter 2nd Year Physics Important Questions Chapter 14 Nuclei

Students get through AP Inter 2nd Year Physics Important Questions 14th Lesson Nuclei which are most likely to be asked in the exam.

AP Inter 2nd Year Physics Important Questions 14th Lesson Nuclei

Very Short Answer Questions

Question 1.
The half life of 58Co is 72 days. Calculate Its average life. [Board model Paper]
Answer:
Thalf = 0.693 × TMean ⇒ TMean = \(\frac{T_{\text {half }}}{0.693}=\frac{72}{0.693}\) = 103.8 days.

Question 2.
Why do all electrons emitted during p-decay not have the same energy?
Answer:
When a neutron is converted into a proton, an electron and neutron are emitted along with it.
\({ }_1^1 \mathrm{n} \longrightarrow{ }_1^1 \mathrm{H}+{ }_{-1}^0 \mathrm{e}+\mathrm{v}\)
In β – decay proton remains in the nucleus, but electron and neutron are emitted with constant energy. The energy of neutron is not constant. So, ail electrons do not have same energy.

AP Inter 2nd Year Physics Important Questions Chapter 14 Nuclei

Question 3.
Neutrons are the best projectiles to produce nuclear reactions. Why ?
Answer:
Neutrons are uncharged particles. So they do not get deflected by the electric and magnetic fields. Hence Neutrons are considered as best projectiles in nuclear reaction.

Question 4.
Neutrons cannot produce ionization. Why ?
Answer:
Because neutrons are uncharged particles and cannot produce ionization.

Question 5.
What are delayed neutrons ?
Answer:
Neutrons are emitted in the fission products after, sometime are called delayed neutrons.

Question 6.
What are thermal neutrons ? What is their importance ?
Answer:
Neutrons having kinetic energies approximately 0.025 eV are called as slow neutrons or thermal neutrons. 235U undergoes fission only when bombarded with thermal neutrons.

AP Inter 2nd Year Physics Important Questions Chapter 14 Nuclei

Question 7.
What is the value of neutron multiplication factor in a controlled reaction and in an uncontrolled chain reaction ?
Answer:
In controlled chain reaction K = 1
In uncontrolled chain reaction K > 1

Question 8.
What is the role of controlling rods in a nuclear reactor ?
Answer:
In nuclear reactor, controlling rods are used to absorb the neutrons. Cadmium, boron materials are used in the form of rods in reactor. These control the fission rate.

Question 9.
Why are nuclear fusion reactions called thermo nuclear reactions ?
Answer:
Nuclear fusion occurs at very high temperatures. So it is called as thermo nuclear reaction.

Question 10.
Define Becquerel and Curie.
Answer:
Becquerel: 1 disintegration or decay per second is called Becquerel. It is SI unit of activity.
1 disintegration or decay
i.e., Becquerel = \(\frac{1 \text { disintegration or decay }}{\text { second }}\)
Curie : 3.7 × 1010 decays per second is called Curie.
1 Curie : 1 Ci = \(\frac{3.7 \times 10^{10} \text { decays }}{\text { second }}\) = 3.7 × 1010Bq.

Question 11.
What is a chain reaction ?
Answer:
Chain reaction : The neutrons produced in the fission of a nucleus can cause fission in other neighbouring nuclei producing more and more neutrons to continue the fission until the whole fissionable material is disintegrated. This is called chain reaction.

AP Inter 2nd Year Physics Important Questions Chapter 14 Nuclei

Question 12.
What is the function of moderator in a nuclear reactor ?
Answer:
They are used to slow down the fast moving neutrons produced during the fission process.
e.g.: Heavy water, Beryllium.

Question 13.
What is the energy released in the fusion of four protons to form a helium nucleus ?
Answer:
26.7 MeV energy is released.

Short Answer Questions

Question 1.
Write a short note on the discovery of neutron.
Answer:

  1. Bothe and Becker found that when beryllium is bombarded with a – particles of energy 5 MeY which emitted a highly penetrating radiation.
  2. The equation for above process can be written as
    \({ }_4^9 \mathrm{Be}+{ }_2^4 \mathrm{He} \rightarrow{ }_6^{13} \mathrm{C}\) + γ – (radiation energy)
  3. The radiations are not effected by electric and magnetic fields.
  4. In 1932, James Chadwick, had subjected nitrogen and argon to the beryllium radiation. He interpreted the experimental results by assuming that the radiation is of a new kind of particles which has no charge and its mass is equal to proton. These neutral particles were named as neutrons. Thus the neutron was discovered.
  5. The experimental results can be represented by the following equation.
    \({ }_4^9 \mathrm{Be}+{ }_2^4 \mathrm{He} \rightarrow{ }_6^{12} \mathrm{C}+{ }_0^1 \mathrm{n}+\mathrm{Q}\)
    AP Inter 2nd Year Physics Important Questions Chapter 14 Nuclei 1

Question 2.
What are nuclear forces ? Write their properties.
Answer:
The forces which hold the nucleons together in nucleus are called nuclear forces.
Properties of Nuclear forces:

  1. Nuclear forces are attractive forces between proton and proton (P – N, proton and neutron (P – N) and neutron and neutron (N – N).
  2. Nuclear forces are independent of charge. It was found that force between proton and proton is same as force between neutron and neutron.
  3. These forces are short range forces i.e., these forces will act upto a small distance only. Generally the range of nuclear forces is upto few Fermi (10-15 m).
  4. These forces are non central forces, i.e., they do not act along the line joining the two nucleons.
  5. These forces are exchange forces. The force between two nucleons is due to exchange of n-mesons.
  6. These forces are spin dependent. These forces are strong when the spin of two nucleons are in same direction and they are weak when they are in opposite direction.
  7. Nuclear forces are saturated forces i.e., the force between nucleons will extend upto the immediate neighbouring nucleons only.
  8. These are the strongest forces in nature. They are nearly 1038 times stronger than gravitational forces and nearly 100 times stronger than Coulombic forces.

AP Inter 2nd Year Physics Important Questions Chapter 14 Nuclei

Question 3.
Define half life period and decay constant for a radioactive substance. Deduce the relation between them.
Answer:
Half life period (T) : Time taken for the number of radio active nuclei to disintegrate to half of its original number of nuclei is called Half life period.

Decay constant (λ) : The ratio of the rate of radioactive decay to the number of nuclei present at that instant.
It is a proportional constant and is denoted by ‘λ’.
∴ λ = \(\frac{-\left(\frac{d N}{d t}\right)}{N}\)

Relation between half the period and decay constant:

  1. The radioactive decay law N = N0 e-λt states that the number of radioactive nuclei in a radioactive sample decreases exponentially with time. Here A is called decay constant.
  2. If N0 is the number of nuclei at t = 0 and N is the radioactive nuclei at any instant of time’t’.
  3. Substituting N = \(\frac{\mathrm{N}_0}{2}\) at t = T in N = N0 e-λt
  4. \(\frac{\mathrm{N}_0}{2}\) = N0 e-λT
    eλT = 2
    λT= ln 2
    T = \(\frac{\ln 2}{\lambda}=\frac{2.303 \log _{10}^2}{\lambda}\)
    ∴ T = \(\frac{0.693}{\lambda}\)

Question 4.
What is nuclear fission ? Give an example to illustrate it.
Answer:
Nuclear fission : The process of dividing a heavy nucleus into two or more smaller and stable nuclei due to nuclear reaction is called nuclear fission.
Ex: The fission reaction is \({ }_{92}^{235} \mathrm{U}+{ }_0^1 \mathrm{n} \rightarrow{ }_{56}^{141} \mathrm{Ba}+{ }_{36}^{92} \mathrm{Kr}+3{ }_0^1 \mathrm{n}+\mathrm{Q}\)
Where Q is the energy released.
Q = (Total mass of reactants – Total mass of product) C2
= [(Mass of \({ }_{92}^{235} \mathrm{U}\) + Mass of \({ }_0^1 \mathrm{n}\)) – (Mass of \({ }_{56}^{141} \mathrm{Ba}\) + Mass of \({ }_{36}^{92} \mathrm{Kr}\) + Mass of three neutrons)] C2
= (235.043933 – 140.9177 – 91.895400 – 2 × 1.008665) amu × C2.
= 0.2135 × 931.5 MeV = 198.9 MeV = 200 MeV

AP Inter 2nd Year Physics Important Questions Chapter 14 Nuclei

Question 5.
What is nuclear fusion ? Write the conditions for nuclear fusion to occur.
Answer:
Nuclear fusion : The process of combining lighter nuclei to produce a larger nucleus is known as nuclear fusion.
E.g : Hydrogen nuclei (1H1) are fused together to form heavy Helium (2He4) along with 25.71 MeV energy released.
Conditions for nuclear fusion :

  1. Nuclear fusion occurs at very high temperatures such as 107 kelvin and very high pressures. These are obtained under the explosion of an atom bomb.
  2. Higher density is also desirable so that collisions between light nuclei occur quite frequently.

Question 6.
Distinguish between nuclear fission and nuclear fusion.
Answer:
Nuclear fission

  1. In this process heavy nucleus is divided into two fragments along with few neutrons.
  2. These reactions will takes place even at room temperature.
  3. To start fission atleast one thermal neutron from outside is compulsory.
  4. Energy released per unit mass of participants is less.
  5. In this process neutrons are liberated.
  6. This reaction can be controlled.
    Ex: Nuclear reactor.
  7. Atom bomb works on principle of fission reaction.
  8. The energy released in fission can be used for peaceful purpose.
    Ex: Nuclear reactor and Atomic power stations.

Nuclear fusion

  1. In this process lighter nuclei will join together to produce heavy nucleus.
  2. These reactions will takes place at very high temperature such as 107 Kelwin.
  3. No necessary of external neutrons.
  4. Energy released per unit mass of participants is high. Nearly seven times more than fission reaction.
  5. In this process positrons are liberated.
  6. There is no control on fusion reaction.
  7. Hydrogen bomb works on the principle of fusion reaction.
  8. The energy released in fusion cannot be used for peaceful purpose.

Long Answer Questions

Question 1.
Define mass defect and binding energy. How does binding energy per nucleon vary with mass number ? What is its significance ?
Answer:
1. Mass defect (∆M) : The difference in mass of a nucleus and its constituents is called the mass defect. The nuclear mass M is always less than the total mass, Em, of its constituents.
Mass defect, (∆M) = [Zmp + (A – Z)mn – M]

2. Binding energy: The energy required to break the nucleus into its constituent nucleons is called the binding energy.
Binding Energy, (Eb) = ∆MC2 = [Zmp + (A – Z)mn – M] 931.5 MeV
Nuclear binding energy is an indication of the stability of the nucleus.
Nuclear binding energy per nucleon Ebn = \(\frac{E_{\mathrm{b}}}{\mathrm{A}}\)

3. The following graph represents how the binding energy per nucleon varies with the mass number A.
AP Inter 2nd Year Physics Important Questions Chapter 14 Nuclei 2
4. From the graph that the binding energy is highest in the range 28 < A < 138. The binding energy of these nuclei is very close to 8.7 MeV

5. With the increase in the mass number the binding energy per nucleon decreases and consequently for the heavy nuclei like Uranium it is 7.6 MeV

6. In the region of smaller mass numbers, the binding energy per nucleon curve shows the characteristic minima and maxima.

7. Minima are associated with nuclei containing an odd number of protons and neutrons such as \({ }_3^6 \mathrm{Li},{ }_5^{10} \mathrm{~B},{ }_7^{14} \mathrm{~N}\) and the maxima are associated with nuclei having an even number of protons and neutrons such as \({ }_2^4 \mathrm{He},{ }_6^{12} \mathrm{C},{ }_8^{16} \mathrm{O}\).

Significance:
8. The curve explains the relationship between binding energy per nucleon and stability of the nuclei.

9. Uranium has a relatively low binding energy per nucleon as 7.6 MeV Hence to attain greater stability Uranium breaks up into intermediate mass nuclei resulting in a phenomenon called fission.

10. On the other hand light nuclei such as hydrogen combine to form heavy nucleus to form helium for greater stability, resulting in a phenomenon called fusion.

11. Iron is the most stable having binding energy per nucleon 8.7 MeV, and it neither undergoes fission per fusion.

AP Inter 2nd Year Physics Important Questions Chapter 14 Nuclei

Question 2.
What is radioactivity ? State the law of radioactive decay. Show that radioactive decay is exponential in nature. [T.S. Mar. 16]
Answer:
1. Radioactivity : The nuclei of certain elements disintegrate spontaneously by emitting alpha (α), beta (β) and gamma (δ) rays. This phenomenon is called Radioactivity or Natural radioactivity.

2. Law of radioactivity decay : “The rate of radioactive decay \(\left(\frac{d N}{d t}\right)\) (or) the number of
nuclei decaying per unit time at any instant, is directly proportional to the number of nuclei (N) present at that instant is called law of radioactivity decay”.

3. Radioactive decay is exponential in nature : Consider a radioactive substance. Let the number of nuclei present in the sample at t = 0, be N0 and let N be the radioactive nuclei remain at an instant t.
\(\frac{\mathrm{dN}}{\mathrm{dt}}\) ∝ N ⇒ \(\frac{\mathrm{dN}}{\mathrm{dt}}\) = – λN
dN = – λ Ndt …………………….. (1)
The proportionality constant λ is called decay constant or disintegration constant. The negative sign indicates the decrease in the number of nuclei.

4. From eq. (1) \(\frac{\mathrm{dN}}{\mathrm{N}}\) = – λ dt ……………… (2)

5. Integrating on both sides
\(\int \frac{\mathrm{dN}}{\mathrm{N}}=-\lambda \int \mathrm{dt}\)
ln N = – λt + C …………….. (3)
Where C = Integration constant.

6. At t = O; N = N0. Substituting in eq. (3), we get ln N0 = C
∴ ln N = -λt + ln N0
ln N – ln N0 = – λt
ln (\(\frac{\mathrm{N}}{\mathrm{N}_0}\)) = – λt
∴ N = N0 e-λt
The above equation represents radioactive decay law.

7. It states that the number of radioactive nuclei in a radioactive sample decreases exponentially with time.

Sample Problems

Question 1.
The half life of radium is 1600 years. How much time does lg of radium take to reduce to 0.125g. [IPE 2016 (TS)]
Answer:
Half – life of radium = 1600 years;
\(\frac{N}{N_0}=\frac{1}{2^n} \Rightarrow \frac{0.125}{1}=\frac{1}{2^n} \Rightarrow \frac{125}{1000}=\frac{1}{2^n} \Rightarrow \frac{1}{8}=\frac{1}{2^n} \Rightarrow \frac{1}{2^3}=\frac{1}{2^n} \Rightarrow n=3\)
∴ Time taken = Half life × no. of Half lives = 1600 × 3 = 4800 years

AP Inter 2nd Year Physics Important Questions Chapter 14 Nuclei

Question 2.
Plutonium decays with a half-life of 24,000 years. If plutonium is stored for 72,000 years, what fraction of it remains ?
Answer:
Half life of plutonium, T = 24000 years; Stored time of plutonium, t = 72000 years
no. of half lives, n = \(\frac{t}{T}=\frac{72000}{24000}\) = 3; Fraction of plutonium remains = \(\frac{N}{N_0}=\frac{1}{2^n}=\frac{1}{2^3}=\frac{1}{8}\)

Question 3.
Explain the principle and working of a nuclear reactor with the help of a labelled diagram. [A.P. & T.S. Mar. 17, A.P. Mar. 16 15, Mar. 14]
Answer:
Principle : A nuclear reactor works on the principle of achieving controlled chain reaction in natural Uranium 238U enriched with 235U, consequently generating large amounts of heat.
A nuclear reactor consists of (1) Fuel (2) Moderator (3) Control rods (4) Radiation shielding (5) Coolant.
AP Inter 2nd Year Physics Important Questions Chapter 14 Nuclei 3
1. Fuel and clad : In reactor the nuclear fuel is fabricated in the form of thin and long cylindrical rods. These group of rods treated as a fuel assembly. These rods are surrounded by coolant, which is used to transfer of heat produced in them. A part of the nuclear reactor which is used to store the nuclear fuel is called the core of the reactor. Natural uranium, enriched uranium, plutonium and uranium – 233 are used as nuclear fuels.

2. Moderator : The average energy of neutrons released in fission process is 2 MeV. They are used to slow down the velocity of neutrons. Heavy water or graphite are used as moderating materials in reactor.

3. Control Rods : These are used to control the fission rate in reactor by absorbing the neutrons. Cadmium and boron are used as controlling the neutrons, in the form of rods.

4. Shielding : During fission reaction beta and gamma rays are emitted in addition to neutrons. Suitable shielding such as steel, lead, concrete etc., are provided around the reactor to absorb and reduce the intensity of radiations to such low levels that do not harm the operating personnel.

5. Coolant : The heat generated in fuel elements is removed by using a suitable coolant to flow around them. The coolants used are water at high pressures, molten sodium etc.

Working : Uranium fuel rods are placed in the aluminium cylinders. The graphite moderator is placed in between the fuel cylinders. To control the number of neutrons, a number of control rods of cadnium or beryllium or boron are placed in the holes of graphite block. When a few 235U nuclei undergo fission fast neutrons are liberated. These neutrons pass through the surrounding graphite moderator and loose their energy to become thermal neutrons. These thermal neutrons are captured by 235U. The heat generated here is used for heating suitable coolants which is turn heat water and produce steam. This steam is made to rotate steam turbine and there by drive a generator of production for electric power.

AP Inter 2nd Year Physics Important Questions Chapter 14 Nuclei

Question 4.
Explain the source of stellar energy. Explain the carbon – nitrogen cycle, proton – proton cycle occuring in stars.
Answer:
Scientists proposed two types of cyclic processes for the sources of energy in the sun and stars. The first is known as carbon-nitrogen cycle and the second is proton-proton cylce.

1. Carbon-Nitrogen Cycle : According to Bethe carbon-nitrogen cycle is mainly responsible for the production of solar energy. This cycle consists of a chain of nuclear reactions in which hydrogen is converted into Helium, with the help of Carbon and Nitrogen as catalysts. The nuclear reactions are as given below.
AP Inter 2nd Year Physics Important Questions Chapter 14 Nuclei 7

2. Proton – Proton Cycle: A star is formed by the condensation of a large amount of matter at a point in space. Its temperature rises to 2,00,000°C as the matter contracts under the influence of gravitational attraction. At this temperature the thermal energy of the protons is sufficient to form a deuteron and a positron. The deuteron then combines with another proton to form lighter nuclei of helium \({ }_2^3 \mathrm{He}\). Two such helium nuclei combine to form a helium nucleus latex]{ }_2^4 \mathrm{He}[/latex] and two protons releasing a total amount of energy 25.71 MeV The nuclear fusion reactions are given below.
AP Inter 2nd Year Physics Important Questions Chapter 14 Nuclei 4

Problems

Question 1.
Compare the radii of the nuclei of mass numbers 27 and 64.
Solution:
A1 = 27; Asub>2 = 64
\(\frac{\mathrm{R}_1}{\mathrm{R}_2}=\left[\frac{\mathrm{A}_1}{\mathrm{~A}_2}\right]^{1 / 3}\) [∵ R = R0A1/3]
\(\frac{\mathrm{R}_1}{\mathrm{R}_2}=\left[\frac{27}{64}\right]^{\frac{1}{3}}=\frac{3}{4}\)
∴ R1 : R2 = 3 : 4

Question 2.
Find the energy required to split \({ }_8^{16} \mathrm{O}\) nucleus into four α-particles. The mass of an a-particle is 4.002603u and that of oxygen is 15.994915u.
Solution:
The energy required to split O = [Total mass of the products – Total mass of the reactants] c2
Mass of four \({ }_2^4 \mathrm{He}\) – Mass of \({ }_8^{16} \mathrm{O}\)] × c2
= [(4 X 4.002603) – 15.994915] u × c2
= [16.010412 – 15.994915] u × c2
= (0.015497) 931.5 MeV = 14.43 MeV

AP Inter 2nd Year Physics Important Questions Chapter 14 Nuclei

Question 3.
A certain substance decays to 1/232 of its initial activity in 25 days. Calculate its half-life.
Solution:
Fraction of substance decays
= \(\frac{\text { Quantity remains }}{\text { Initial quantity }}\)
= \(\frac{1}{2^n}=\frac{1}{32}=\frac{1}{2^5}\)
∴ n = 5
Duration of time = 25 days
We know (n) = \(\frac{\text { Duration of time }}{\text { Half life time }}\)
∴ Half life time = \(\frac{\text { Duration of time }}{\mathrm{n}}\)
\(\frac{25}{5}\) = 5 days

Question 4.
One gram of radium is reduced by 2 milli- gram in 5 years by a-decay. Calculate the half¬life of radium.
Solution:
Initial (original) mass (N0) = 1 gram
Reduced mass – 2 mg = 0.002 grams
Final mass (N)= 1 – 0.002 = 0.998 grams
t = 5 years
e-λt = \(\frac{\mathrm{N}}{\mathrm{N}_0}\) ⇒ eλt = \(\frac{\mathrm{N}_0}{\mathrm{~N}}\) ⇒ λt = loge[latex]\frac{\mathrm{N}_0}{\mathrm{~N}}[/latex]
λt = 2.303 log [latex]\frac{\mathrm{N}_0}{\mathrm{~N}}[/latex]
λt = 2.303 log [latex]\frac{1}{0.998}[/latex]
= 2.303 log (1.002)
= 2.303 × 0.000868
= 0.001999
λ = \(\frac{0.001999}{5}\) = 0.0003998
T = \(\frac{0.693}{\lambda}=\frac{0.693}{0.0003998}\) = 1733.3 years

Question 5.
If one microgram of \({ }_92^{235} \mathrm{U}\) is completely destroyed in an atomhomb, how much energy will be released ?
Solution:
m = 1 μg = 1 × 10-6 g = 1 × 10-6 × 10-3 kg
= 10-9 kg
c = 3 × 108 m/s
E = mc2 = 1 × 10-9 × 9 × 106 = 9 × 107 J

AP Inter 2nd Year Physics Important Questions Chapter 14 Nuclei

Question 6.
200 MeV energy is released when one nucleus of 235U undergoes fission. Find the number of fissions per second required for producing a power of 1 megawatt.
Solution:
E = 200 MeV
P = 1 × 106 W
P = \(\frac{\mathrm{nE}}{\mathrm{t}} \Rightarrow \frac{\mathrm{n}}{\mathrm{t}}=\frac{\mathrm{P}}{\mathrm{E}}=\frac{10^6}{200 \times 10^6 \times 1.6 \times 10^{-19}}\)
= \(\frac{1}{32}\) × 1018
∴ P = 0.03125 × 1018
= 3.125 × 106

Textual Examples

Question 1.
Given the mass of iron nucleus as 55.85u and A = 56, find the nuclear density ?
Solution:
mFe = 55.85, u = 9.27 × 10-26 kg
Nuclear density = \(\frac{\text { mass }}{\text { volume }}\)
= \(\frac{9.27 \times 10^{-26}}{(4 \pi / 3)\left(1.2 \times 10^{-15}\right)^3} \times \frac{1}{56}\)
= 2.29 × 1017 kg m-3
The density of matter in neutron stars (an astrophysical object) is comparable to this density. This shows that matter in these objects has been compressed to such an extent that they resemble a big nucleus.

Question 2.
Calculate the energy equivalent of 1 g of substance.
Solution:
Energy, E = 10-3 × (3 × 108)2 J
E = 10-3 × 9 × 1016 = 9 × 1013 J
Thus, if one gram of matter is converted to energy, there is a release of an enormous amount of energy.

AP Inter 2nd Year Physics Important Questions Chapter 14 Nuclei

Question 3.
Find the energy equivalent of one atomic mass unit, first in Joules and then in MeV. Using this, express the mass defect of \({ }_8^{16} \mathrm{O}\) in MeV/c2.
Solution:
1 u = 1.6605 × 10-27 kg
To convert it into energy units, we multiply it by c2 and find, that energy. equivalent
= 1.6605 × 10-27 × 2.9979 × 108 kg m2/s2
= 1.4924 × 10-10 J
= \(\frac{1.4924 \times 10^{-10}}{1.602 \times 10^{-19}}\) eV
= 0.9315 × 109 eV = 931.5 MeV
or, 1 u = 931.5 MeV/c2.
For \({ }_8^{16} \mathrm{O}\), ∆M = 0.13691 u = 0.13691 × 931.5 MeV/c2 = 127.5 MeV/c2
The energy needed to separate \({ }_8^{16} \mathrm{O}\) into hs constituents is thus 127.5 MeV/c2.

Question 4.
The half-life of \({ }_{92}^{238} \mathrm{U}\) undergoing a – decay is 4.5 × 109 years. What is the activity of 1 g sample of \({ }_{92}^{238} \mathrm{U}\) ?
Solution:
T1/2 = 4.5 × 109 y
= 4.5 × 109 y × 3.16 × 107 s/y.
= 1.42 × 1017 s
One k mol of any isotope contains Avogadro’s number of atoms, and so 1 g of \({ }_{92}^{238} \mathrm{U}\) contains
\(\frac{1}{238 \times 10^{-3}}\) k mol × 6.025 × 1026 atoms/kmol
= 25.3 × 1020 atoms.
The decay rate R is
R = λN
= \(\frac{0.693}{\mathrm{~T}_{1 / 2}}\) N = \(\frac{0.693 \times 25.3 \times 10^{20}}{1.42 \times 10^{17}}\) S-1
= 1.23 × 104 S-1
= 1.23 × 104 Bq

AP Inter 2nd Year Physics Important Questions Chapter 14 Nuclei

Question 5.
Tritium has a half-life of 12.5 y undergoing beta decay. What fraction of sample of pure tritium will remain undecayed after 25 y.
Solution:
By definition of half-life, half of the initial i sample will remain undecayed after 12.5 y. In the next 12.5 y, one-half of these nuclei would have decayed. Hence, one fourth of the sample of the initial pure s tritium will remain undecayed.

Question 6.
We are given the following atomic masses:
\({ }_{92}^{238} \mathrm{U}\) = 238.05079 u
\({ }_{2}^{4} \mathrm{He}\) = 4.00260 u
\({ }_{90}^{234} \mathrm{Th}\) = 234.04363 u
\({ }_{1}^{1} \mathrm{H}\) = 1.00783 u
\({ }_{91}^{237} \mathrm{Pa}\) = 237.05121 u
Here the symbol Pa is for the element protactinium (Z = 91).
a) Calculate the energy released during the alpha decay of \({ }_{92}^{238} \mathrm{U}\).
b) Show that \({ }_{92}^{238} \mathrm{U}\) cannot spontaneously emit a proton.
Solution:
a) The alpha decay of \({ }_{92}^{238} \mathrm{U}\) is given by
AP Inter 2nd Year Physics Important Questions Chapter 14 Nuclei 6
The energy released in this process is given by
Q = (MU – MTh – MHe) c2
Substituting the atomic masses as given in the data, we find
Q = (238.05079 – 234.04363 – 4.00260)u × c2
= (0.00456 u) c2
= (0.00456 u) (931.5 MeV/u)
= 4.25 MeV

AP Inter 2nd Year Physics Important Questions Chapter 14 Nuclei

b) If \({ }_{92}^{238} \mathrm{U}\) spontaneously emits a proton, the decay process would be
\({ }_{92}^{238} \mathrm{U} \rightarrow{ }_{91}^{237} \mathrm{~Pa}+{ }_1^1 \mathrm{H}\)
The Q for this process to happen is
= (MU – MPa – MH)c2
(238.05079 – 237.05121 – 1.00783) u × c2
=(- 0.00825 u) c2
= – (0.00825 u) (931.5 MeV/u)
= -7.68 MeV
Thus, the Q of the process is negative and therefore it cannot proceed spontaneously. We will have to supply an energy of 7.68 MeV to a \({ }_{92}^{238} \mathrm{U}\) nucleus to make it emit a proton.

AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics: Material, Devices and Simple Circuits

Students get through AP Inter 2nd Year Physics Important Questions 15th Lesson Semiconductor Electronics: Material, Devices and Simple Circuits which are most likely to be asked in the exam.

AP Inter 2nd Year Physics Important Questions 15th Lesson Semiconductor Electronics: Material, Devices and Simple Circuits

Very Short Answer Questions

Question 1.
What is an n-type semiconductor ? What are the majority and minority charge carriers in it ?
Answer:

  • If a pentavalent impurity is added to pure tetravalent semiconductor, it is called n-type semiconductor.
  • In n-type semiconductor majority charge carriers are electrons and minority charge carriers are holes.

Question 2.
What are intrinsic and extrinsic semiconductors ? [A.P. Mar. 15]
Answer:

  • Pure form of semiconductors are called intrinsic semiconductors.
  • When impure atoms are added to increase their conductivity, they are called extrinsic semiconductors.

AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics: Material, Devices and Simple Circuits

Question 3.
What is a p-type semiconductor ? What are the majority and minority charge carriers in it ? [A.P. & T.S. Mar. 17]
Answer:
If a trivalent impurity is added to a tetravalent semiconductor, it is called p-type semi-conductor.
AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics Material, Devices and Simple Circuits 1
In p-type semiconductor majority charge carriers are holes and minority charge carriers are electrons.

Question 4.
What is a p-n junction diode ? Define depletion’ layer.
Answer:
When an intrinsic semiconductor crystal is grown with one side doped with trivalent element and on the other side doped with pentavalent element, a junction is formed in the crystal. It is called p-n junction diode.

A thin narrow region is formed on either side of the p-n junction, which is free from charge carriers is called depletion layer.

Question 5.
How is a battery connected to a junction diode in i) forward and ii) reverse bias ?
Answer:
i) In p-n junction diode, if p-side is connected to positive terminal of a cell and n-side to negative terminal, it is called forward bias.
AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics Material, Devices and Simple Circuits 2
ii) In a p-n junction diode, p-side is connected to negative terminal of a cell and n-side to positive terminal, it is called reverse bias.
AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics Material, Devices and Simple Circuits 3

Question 6.
What is the maximum percentage of rectification in half wave and full wave rectifiers ?
Answer:

  1. The percentage of rectification in half-wave rectifier is 40.6%.
  2. The percentage of rectification in full-wave rectifiers is 81.2%.

AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics: Material, Devices and Simple Circuits

Question 7.
What is Zener voltage (Vz) and how will a Zener diode be connected in circuits generally ?
Answer:

  1. When Zener diode is in reverse biased, at a particular voltage current increases suddenly is called Zener (or) break down voltage.
  2. Zener diode always connected in reverse bias.

Question 8.
Write the expressions for the efficiency of a full wave rectifier and a half wave rectifier.
Answer:

  1. Efficiency of half wave rectifier (η) = \(\frac{0.406 \mathrm{R}_{\mathrm{L}}}{\mathrm{r}_{\mathrm{f}}+\mathrm{R}_{\mathrm{L}}}\)
  2. Efficiency of full wave rectifier (η) = \(\frac{0.812 \mathrm{R}_{\mathrm{L}}}{\mathrm{r}_{\mathrm{f}}+\mathrm{R}_{\mathrm{L}}}\)

Question 9.
What happens to the width of the depletion layer in a p-n junction diode when it is
i) forward-biased and ii) reverse biased ?
Answer:
When a p-n junction diode is forward bias, thickness of depletion layer decreases and in reverse bias, thickness of depletion layer increases.

Question 10.
Draw the circuit symbols for p-n-p and n-p-n transistors. [Mar. 14]
Answer:
AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics Material, Devices and Simple Circuits 4

Question 11.
Define amplifier and amplification factor.
Answer:

  1. Rising the strength of a weak signal is known as amplification and the device is called amplifier.
  2. Amplification factor is the ratio between output voltage to the input voltage.
    A = \(\frac{\mathrm{V}_0}{\mathrm{~V}_{\mathrm{i}}}\)

AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics: Material, Devices and Simple Circuits

Question 12.
In which bias can a Zener diode be used as voltage regulator ?
Answer:
In reverse bias Zener diode can be used as voltage regulator.

Question 13.
Which gates are called universal gates ? [T.S. Mar. 15]
Answer:
NAND gate and NOR gates are called universal gates.

Question 14.
Write the truth table of NAND gate. How does it differ from AND gate ?
Answer:
AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics Material, Devices and Simple Circuits 5
AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics Material, Devices and Simple Circuits 6

Short Answer Questions

Question 1.
What is a rectifier ? Explain the working of half wave and full Wave rectifiers with diagrams. [A.P. Mar. 17]
Answer:
Rectifier: It is a circuit which converts ac into d.c. A p-n junction diode is used as a rectifier.
Half-wave rectifier:
AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics Material, Devices and Simple Circuits 7

  1. A half wave rectifier can be constructed with a single diode. The ac input signal is fed to the primary coil of a transformer. The output signal is taken across the load resistance RL.
  2. During positive half cycle, the diode is forward biased and current flows through the diode.
  3. During negative half cycle, diode is reverse biased and no current flows through the load resistance.
  4. This means current flows through the diode only during positive half cycles and negative half cycles are blocked. Hence in the output we get only positive half cycles.
  5. Rectifier efficiency is defined as the ratio of output dc power to the input ac power.
    η = \(\frac{\mathrm{P}_{\mathrm{dc}}}{\mathrm{P}_{\mathrm{ac}}}=\frac{0.406 \mathrm{R}_{\mathrm{L}}}{\mathrm{r}_{\mathrm{f}}+\mathrm{R}_{\mathrm{L}}}\)
    Where rf = Forward resistance of a diode; RL = Load resistance
    The maximum efficiency of half wave rectifier is 40.6%.

Full wave rectifier : The process of converting an alternating current into a direct current is called rectification.
The device used for this purpose is called rectifier.

  1. A full wave rectifier can be constructed with the help of two diodes D1 and D2.
  2. The secondary transformer is centre tapped at C and its ends are connected to the p regions of two diodes D1 & D2.
  3. The output voltage is measured across the load resistance RL.
  4. During positive half cycles of ac, the diode D1 is forward biased and current flows through the load resistance RL. At this time D2 will be reverse biased and will be in switch off position.
  5. During negative half cycles of ac, the diode D2 is forward biased and the current flows through RL. At this time D1 will be reverse biased and will be in switch off position.
  6. Hence positive output is obtained for all the input ac signals.
  7. The efficiency of a rectifier is defined as the ratio between the output dc power to the input ac power.
    η = \(\frac{\mathrm{P}_{\mathrm{dc}}}{\mathrm{P}_{\mathrm{ac}}}=\frac{0.812 \mathrm{R}_{\mathrm{L}}}{\mathrm{r}_{\mathrm{f}}+\mathrm{R}_{\mathrm{L}}}\)
    The maximum efficiency of a full wave rectifier is 81.2%.

AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics: Material, Devices and Simple Circuits

Question 2.
What is a junction diode ? Explain the formation of depletion region at the junction. Explain the variation of depletion region in forward and reverse-biased condition.
Answer:
p-n junction diode: When a p-type semiconductor is suitablyjoined to n-type semiconductor, a p-n junction diode is formed. ’
The circuit symbol of p-n junction diode is shown in figure.
AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics Material, Devices and Simple Circuits 8
Formation of depletion layer at the junction: When p-n junction is formed, the free electrons on n-side diffuse over to p-side, they combine with holes and become neutral. Similarly holes on p-side diffuse over to n-side and combihe with electrons become neutral.

This results in a narrow region formed on either side of the junction. This region is called depletion layer. Depletion layer is free from charge carriers.

The n-type material near the junction becomes positively charged due to immobile donor ions and p-type material becomes negatively charged due to immobile acceptor ions. This creates an electric field near the junction directed from n-region to p-region and cause a potential barrier.
AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics Material, Devices and Simple Circuits 9
The potential barrier stops further diffusion of holes and electrons across the* junction. The value of the potential barrier depends upon the nature of the crystal, its temperature and the amount of doping.

Forward bias :
“When a positive terminal of a battery is connected to p-side and negative terminal is connected to n-side; then p-n junction diode is said to be forward bias”.
AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics Material, Devices and Simple Circuits 10
The holes in the p-region are repelled by the positive polarity and move towards the junction. Similarly electrons in the n-region are repelled by the negative polarity and move towards the junction.

As a result, the width of the depletion layer decreases. The charge carriers cross the junction apd electric current flows in the circuit.
Hence in forward bias resistance of diode is low. This position is called switch on position.
Reverse bias:
AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics Material, Devices and Simple Circuits 11
“When the negative terminal of the battery is connected to p-side and positive terminal of the battery is connected to n-side of the p-n junction, then the diode is said to be reverse bias”.

The holes in the p-region are attracted towards negative polarity and move away from the junction. Similarly the electrons in the n-region are attracted towards positive polarity and move away from the junction.

So, width of the depletion layer and potential barrier increases. Hence resistance of p-n junction diode increases. Thus the reverse biased diode is called switch off position.

AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics: Material, Devices and Simple Circuits

Question 3.
What is a Zener diode ? Explain how it is used as a voltage regulator.
Answer:
Zener diode : Zener diode is a heavily doped germanium (or) silicon p-n junction diode. It works on reverse bias break down region.
The circuit symbol of zener diode is shown in figure.
AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics Material, Devices and Simple Circuits 12
Zener diode can be used as a voltage regulator. In- general zener diode is connected in reverse bias in the circuits.
AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics Material, Devices and Simple Circuits 13

  1. The zener diode is connected to a battery, through a resistance R. The battery reverse biases the zener diode.
  2. The load resistance RL is connected across the terminals of the zener diode.
  3. The value of R is selected in such away that In the absence of load RL maximum safe current flows in the diode.
  4. Now consider that load is connected across the diode. The load draws a current.
  5. The current through the diode falls by the same amount but the voltage drops across the load remains constant.
  6. The series resistance ‘R’ absorbs the output voltage fluctuations, so as to maintain constant voltage across the load.
  7. The voltage across the zener diode remains constant even if the load RL varies. Thus, zener diode works as voltage regulator.
  8. If I is the input current, IZ and IL are zener and load currents.
    I = IZ + IL; Vin = IR + VZ
    But V0ut = VZ
    ∴ Vout = Vin – IR

Question 4.
Explain the working of LED and what are its advantages over conventional incandescent low power lamps.
Answer:
Light emitting diode (LED) : It is a photoelectronic device which converts electrical energy into light energy.

It is a heavily doped p-n junction diode which under forward bias emits spontaneous radiation. The diode is covered with a transparent cover so that the emitted light may not come out. Working : When p-n junction diode is forward biased, the movement of majority charge carriers takes place across the junction. The electrons move from n-side to p-side through the junction and holes move from p-side to n-side.

As a result of it, concentration of majority carriers increases rapidly at the junction.

When there is no bias across the junction, therefore there are excess minority carriers on either side of the junction, which recombine with majority carriers near the junction.

On recombination of electrons and hole, the energy is given out in the form of heat and light.
Advantages of LED’s over incandescent lamp :

  1. LED is cheap and easy to handle.
  2. LED has less power and low operational voltage.
  3. LED has fast action and requires no warm up time.
  4. LED can be used in burglar alarm system.

AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics: Material, Devices and Simple Circuits

Question 5.
Define NAND and NOR gates. Give their truth tables. [T.S. Mar. 17]
Answer:
NAND gate : NAND gate is a combination of AND gate and NOT gate.
AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics Material, Devices and Simple Circuits 14
NAND gate can be obtained by connecting a NOT gate in the output of an AND gate. NAND gates are called universal gates.

  1. If both inputs are low, output is high.
    A = 0, B = 0, X = 1
  2. If any input is low, output is high.
    A = 0, B = 1, X = 1
    A = 1, B = 0, X =1
  3. If both inputs are high, output is low.
    A = 1, B = 1, X = 0
    AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics Material, Devices and Simple Circuits 15

NOR gate : NOR gate is a combination of OR gate and NOT gate when the output of OR gate is connected to NOT gate. It has two (or) more inputs and one output.
AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics Material, Devices and Simple Circuits 16

  1. If both inputs are low, output is high.
    A = 0, B = 0, X = 1
  2. If any input is high, the output is low.
    A = 0, B = 1, X = 0
    A = 1, B = 0, X = 0
  3. If both inputs are high, the output is low.
    A = 1, B = 1, X = 0
    NOR gate is also a universal gate.
    AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics Material, Devices and Simple Circuits 17

AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics: Material, Devices and Simple Circuits

Question 6.
Explain the working of a solar cell and draw its I-V characteristics.
Answer:
Solar cell is a p-n junction device which converts solar energy into electric energy.
It consists of a silicon (or) gallium – arsenic p-n junction diode packed in a can with glass window on top. The upper layer is of p- type semiconductor. It is very thin so that the incident light photons may easily reach the p-n junction.
AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics Material, Devices and Simple Circuits 18
Working : When light (E = hv) falls at the junction, electron – hole pairs are generated near the junction. The electrons and holes produced move in opposite directions due to junction field. They will be collected at the two sides of the junction giving rise to a photo voltage between top and bottom metal electrodes. Top metal acts as positive electrode and bottom metal acts as a negative electrode. When an external load is connected across metal electrodes a photo current flows.
AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics Material, Devices and Simple Circuits 19
I-V characteristics : I – V characteristics of solar cell is drawn in the fourth quadrant of the coordinate axes. Because it does not draw current.
AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics Material, Devices and Simple Circuits 20
Uses : They are used in calculators, wrist watches, artificial satellites etc.

Question 7.
Explain the operation of a NOT gate and give its truth table. [IPE 15, T.S.]
Answer:
NOT gate: NOT gate is the basic gate. It has one input and one output. The NOT gate is also called an inverter. The circuit symbol of NOT gate is shown in figure.
AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics Material, Devices and Simple Circuits 21

  1. If input is low, output is high.
    A = 0, X = \(\overline{0}\) = 1
  2. If input is high, output is low.
    A = 1, X = \(\overline{1}\) = 0
    AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics Material, Devices and Simple Circuits 22

AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics: Material, Devices and Simple Circuits

Question 8.
Draw an OR gate using two diodes and explain its operation. Write the truth table and logic symbol of OR gate.
Answer:
OR Gate : It has two input terminals and one output terminal. If both inputs are low the output is low. If one of the inputs is high, or both the inputs are high then the output of the gate is high. The truth tables of OR gate.
AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics Material, Devices and Simple Circuits 23
In truth table logic function is written as A or B ‘OR’ logic function is represented by the symbol ‘plus’.
Q = A + B
Logic gate ‘OR’ is shown given below.

Implementation of OR gate using diodes :
Let D1 and D2 be two diodes.
AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics Material, Devices and Simple Circuits 24
A potential of 5V represents the logical value 1.
A potential of OV represents the logical value 0.
AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics Material, Devices and Simple Circuits 25
When A = 0, B = 0 both the diodes are reverse biased and there is no current through the resistance. So, the potential at Q is zero i.e., Q = 0. When A = 0 or B = 0 and the other equal to a potential of 5 V the diode whose anode is at a potential of 5 V is forward – biased and that diode behaves like a closed switch. The output potential then becomes 5 V i.e., Q = 1. When both A and B are 1, both the diodes are forward-biased and the potential at Q is same as that at A and B which is 5 V i.e., Q = 1. The output is same as that of the OR gate.

Question 9.
Sketch a basic AND circuit with two diodes and explain its operation. Explain how doping increases the conductivity in semiconductors ?
Answer:
AND gate : It has two input terminals and one output terminal.

  • If both the inputs are low (or) one of the inputs is low.
    • The output is low in an AND gate.
  • If both the inputs .are high
    • The output of the gate is high.
  • Note : If A and B are the inputs of the gate and the output is ‘Q’ then ‘Q’ is a logical function of A and B.
    AND gate Truth Tables
    AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics Material, Devices and Simple Circuits 26
    The logical function AND is represented by the symbol dot so that the output, Q = A.B and the circuit symbol used for the logic gate AND is shown in Fig.
    AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics Material, Devices and Simple Circuits 27
    The logical function AND is similar to the multiplication.

Implementation of AND gate using diodes : Let D1 and D2 represents two diodes. A potential of 5 V represents the logical value 1 and a potential of 0 V represents the logical value zero (0). When A = 0, B = ,0 both the diodes D1 and D2 are forward-biased and they behave like closed switches. Hence, the output Q is same as that A or B (equal to zero.) When A = 0 or B = 0, D1 or D2 is forward – biased and Q is zero. When A = 1 and B = T both the diodes are reverse – biased and they behave like open switches. There is no current through the resistance R making the potential at Q equal to 5V i.e., Q = 1. The output is same as that of an AND gate.

Doping increases the conductivity in Semiconductors: If a pentavalent impurity (Arsenic) is added to a pure tetravalent semiconductor it is called n-type semiconductor. Arsenic has 5 valence electrons, but only 4 electrons are needed to form covalent bonds with the adjacent Germanium atoms. The fifth electron is very loosely bound and become a free electron. Therefore excess electrons are available for conduction and conductivity of semiconductor increases.

Similarly when a trivalent impurity Indium is added to pure Germanium it is called p-type semi-conductor. In this excess holes in addition to those formed due to thermal energy are available for conduction in the valence band and the conductivity of semiconductor increases.

AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics: Material, Devices and Simple Circuits

Question 10.
Explain how transistor can be used as a switch ?
Answer:
To understand the operation of transistor as a switch.

  1. As long as Vi is low and unable to forward bias the transistor, V0 is high (at Vcc).
  2. If Vi is high enough to drive the transistor into saturation then V0 is low, very near to zero.
    AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics Material, Devices and Simple Circuits 28
  3. When the transistor is not conducting it is said to be switched off and when it is driven into saturation it is said to be switched on.
  4. We can say that a low input to the transistor gives a high output and high input gives a low output.
  5. When the transistor is used in the cutoff (or) saturation state it acts as a switch.
    AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics Material, Devices and Simple Circuits 29

Problems

Question 1.
In a half wave rectifier, a p-n junction diode with internal resistance 20 ohm is used. If the load resistance of 2 ohm is used in the circuit, then find the efficiency of this half wave rectifier.
Solution:
Internal resistance (rf) = 20Ω
RL = 2kΩ = 2000 Ω
η = \(\frac{0.406 \mathrm{R}_{\mathrm{L}}}{\mathrm{r}_{\mathrm{f}}+\mathrm{R}_{\mathrm{L}}}=\frac{0.406 \times 2000}{20+2000} \times 100=\frac{812 \times 100}{2020}\)
η = 40.2%.

Question 2.
A full wave p-n junction diode rectifier uses a load resistance of 1300 ohm. The internal resistance of each diode is 9 ohm. Find the efficiency of this full wave rectifier.
Solution:
Given that RL = 1300 Ω
rf = 9Ω
η = \(\frac{0.812 \mathrm{R}_{\mathrm{L}}}{\mathrm{r}_{\mathrm{f}}+\mathrm{R}_{\mathrm{L}}}=\frac{0.812 \times 1300}{9+1300} \times 100 ; \eta=\frac{8120 \times 13}{1309}\)
η = 80.64%.

AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics: Material, Devices and Simple Circuits

Question 3.
Calculate the current amplification factor β(beta) when change in collector current is 1mA and change in base current is 20μA.
Solution:
Change in collector current (∆IC) = 1mA = 10-3 A
Change in base current (∆IB) = 20 μA = 20 × 10-6 A
β = \(\frac{\Delta \mathrm{I}_{\mathrm{C}}}{\Delta \mathrm{I}_{\mathrm{B}}}=\frac{10^{-3}}{20 \times 10^{-6}}\); β = \(\frac{1000}{20}\)
β = 50

Question 4.
For a transistor amplifier, the collector load resistance RL = 2k ohm and the input resistance Ri = 1 k ohm. If the current gain is 50, calculate voltage gain of the amplifier.
Solution:
RL = 2kΩ = 2 × 103
Ri = 1kΩ = 1 × 103
β = 50.
Voltage gain (AV) = β × \(\frac{\mathrm{R}_{\mathrm{L}}}{\mathrm{R}_{\mathrm{i}}}=\frac{50 \times 2 \times 10^3}{1 \times 10^3}\)

AV = 100.

Textual Examples

Question 1.
C, Si and Ge have same lattice structure. Why is C insulator while Si and Ge intrinsic semiconductors ?
Solution:
The 4 bonding electrons of C, Si or Ge lie, respectively, in the second, third and fourth orbit. Hence, energy required to take out and electron from these atoms(i.e., ionisation energy Eg) will be least for Ge, following by Si and highest for C. Hence, number of free electrons for conduction in Ge and Si are significant but negligibly small for C.

Question 2.
Suppose a pure Si crystal has 5 × 1028 atoms m-3. It is doped by 1 ppm concentration of pentavalent As. Calculate the number of electrons and holes. Given that n. = 1.5 × 1016 m-3.
Solution:
Note that thermally generated electrons (ni ~ 1016 m-3) are negligibly small as compared to those produced by doping.
Therefore, ni ≈ ND.
Since nenh = ni2, The number of holes, nh = (1.5 × 1016)2 / 5 × 1028 × 16-6
nh = (2.25 × 1032)/(5 × 1022) ~ 4.5 × 109m-3

AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics: Material, Devices and Simple Circuits

Question 3.
Can we take one slab of p-type semiconductor and physically join it to another n-type semiconductor to get p-n junction ?
Solution:
No ! Any slab, howsoever flat, will have roughness much large than the inter-atomic crystal spacing(~2 to 3 A°) and hence continuous contact at the atomic level will not be possible. The junction will behave as a discontinuity for the flowing charge carriers.

Question 4.
The V-I characteristics of a silicon diode are shown in the Fig. Calculate the resistance of the diode at (a) ID = 15 mA and (b) VD = -10 V.
AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics Material, Devices and Simple Circuits 30
Solution:
Considering the diode characteristics as a straight line between I = 10 mA to I = 20 mA passing through the origin, we can calculate the resistance using Ohm’s law.
a) From the curve, at I = 20 mA, V = 0.8 V, I = 10 mA, V = 0.7 V
rfb = ∆V/∆I = 0.1V/10mA = 10Ω
b) From the curve at V = -10 V, I = -1 µA.
Therefore rrb = 10V/1µA = 1.0 × 107

Question 5.
In a Zener regulated power supply a Zener diode’with Vz = 6.0 V is used for regulation. The load current is to be 4.0 mA and the unregulated input is 10,0 V. What should be the value of series resistor Rs?
Solution:
The value of Rs should be such that the current through the Zener diode is much larger than the load current. This is to have good load regulation. Choose Zener current as five times the load current, i.e., Iz = 20 mA. The total current through RS is , therefore, 24 mA. The voltage drop across RS is 10.0 – 6.0 = 4.0 V This gives RS = 4.0V/(24 × 10-3) A = 167Ω. The nearest value of carbon resistor is 150 Ω. So, a series resistor of 150 Ω is appropriate. Note that slight variation in the value of the resistor does not matter, what is important is that the current IZ should be sufficiently larger than IL.

AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics: Material, Devices and Simple Circuits

Question 6.
The current in the forward bias is known to be more (~mA) than the current in the reverse bias (~µA). What is the reason then to operate the photodiodes in reverse bias ?
Solution:
Consider the case of an n-type semiconductor. Obviously, the majority carrier density (n) is considerably larger than the minority hole density p (i.e., n> > p). On illumination, let the excess eletrons and holes generated be ∆n and ∆p, respectively,
n’ = n + ∆n
p’ = p + ∆p
Here n’ and p’ are the electron and hole concentrations at any particular illumination and n and p are carriers concentration when there is no illumination. Remember ∆n = ∆p and n > > p. Hence, the fractional change in the majority carriers (i.e., ∆n/n) would be much less than that in the minority carriers (i.e,, ∆p/p). In general, we can state that the fractional change due to the photo-effects on the minority carrier dominated reverse bias current is more easily measurable than the fractional change in the forward bias current. Hence, photodiodes-are preferably used in the reverse bias condition for measuring light intensity.

Question 7.
Why are Si and GaAs are preferred materials for solar cells ?
Solution:
The solar radiation spectrum received by us is shown in figure.
AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics Material, Devices and Simple Circuits 31
The maxima is near 1.5 eV. For photo-excitation, hv > Eg. Hence, semiconductor with band gap ~1.5 eV or lower is likely to give better solar conversion efficiency. Silicon has Eg ~ 1.1 eV while for GaAS it is ~ 1.53 eV. In fact, GaAs is better On spite of its higher band gap) than Si because of its relatively higher absorption coefficient. If we choose materials like CdS or CdSe(Eg ~ 2.4 eV), we can use only the high energy component of the solar energy for photo-conversion and a significant part of energy will be of no use.

The question arises: why we do not use material like pbS(Eg ~ 0.4 eV) which satisfy the condition hv > Eg for v maxima corresponding to the solar radiation spectra ? if we do so, most of the solar radiation will be absorbed on the top-layer of solar cell mid will not reach in or near the depletion region. For effective electron-hole separation, due to the junction field, we want the photo-generation to occur in the junction region only.

Question 8.
From the output charactristics shown in fig, calculate the values of βac and βdc of the transistor when VCE is 10 V and IC = 4.0 mA.
AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics Material, Devices and Simple Circuits 32
Solution:
βac = \(\left(\frac{\Delta \mathrm{I}_{\mathrm{C}}}{\Delta \mathrm{I}_{\mathrm{B}}}\right)_{\mathrm{V}_{\mathrm{CE}}}\) ; βdc = \(\frac{\mathrm{I}_{\mathrm{C}}}{\mathrm{I}_{\mathrm{B}}}\)
For determining βac and βdc at the stated values of VCE and IC one can proceed as follows. Consider any two characteristics for two values of IB which lie above and below the given value of IC. Here IC = 4.0 mA. (Choose characteristics for IB = 30 and 20μA.) At V CE = 10V
we read the two values of Ic from the graph. Then
∆IB = (30 – 20)μA = 10μA. ∆IC
= (4.5 – 3.0) mA = 1.5 mA
Therefore, βac =1.5 mA/ 10μA = 150
For determining βdc either estimate the value of IB corresponding to IC = 4.0 mA at VCE = 10V or calculate the two values of βdc for the two characteristics chosen and find their mean.
Therefore, for IC = 4.5 mA and IB = 30 μA
βdc = 4.5 mA/30 μA = 150 and for IC = 3.0 mA/ and IB = 20 μA
βdc = 3.0 mA / 20 μA= 150
Hence, βdc = (150 + 150)/ 2 = 150

AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics: Material, Devices and Simple Circuits

Question 9.
In Fig. the VBB supply can be varied from OV to 5.0 V. The Si transistor has βdc = 250 and RB = 100 kΩ, RC = 1 KΩ, VCC = 5.0V. Assume that when the transistor is saturated, VCE = 0V and VBE = 0.8V. Calculate
(a) the minimum base current, for which the transistor will reach saturation. Hence,
(b) determine V1 for when the transistor is ‘switched on’,
(c) find the ranges of V1 for which the transistor is ‘switched of and ‘switched on’.
AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics Material, Devices and Simple Circuits 33
Solution:
Given at stauration
VCE = OV, VBE = 0.8V
VCE = VCC – ICRC
IC = VCC / RC = 5.0V/1/0 kΩ = 5.0mA
Therefore, IB = IC
= 5.0 mA/250 = 20μA
The input voltage at which the transistor will go into saturation is given by
VIH = VBB = IBRB + VBE
= 20μA × 100 kΩ + 0.8V = 2.8V
The value of input voltage below which the transistor remains cutoff is given by
VIL = 0.6V, VIH = 2.8V
Between 0.0V and 0.6V the transistor will be in the ‘switched off-state. Between 2.8V and 5.0V, it will be in ‘switched on’ state.
Note that the transistor is in active state when IB varies from 0.0mA to 20mA. In this range, IC = βIB is valid. In the saturation range IC ≤ βIB.

Question 10.
For a CE transistor amplifier, the audio signal voltage across the collector resistance of 2.0 kΩ is 2.0 V. Suppose the current amplification factor of the transistor is 100, what should be the value of RB in series with VBB supply of 2.0 V if the dc base current has to be 10 times the signal current. Also calculate the dc drop across the collector resistance. (Refer to Fig)
AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics Material, Devices and Simple Circuits 34
Solution:
The output ac voltage is 2.0 V. So, the ac collector current iC= 2.0/2000 = 1.0 mA. The signal current through the base is, therefore given by iB = iC/β = 1.0 mA/100 = 0.010 mA. The dc base current has to be 10 × 0.010 = 0.10 mA
From VBB = VBE+ IB RB RB = (VBB – VBE)/IB. Assuming VBE = 0.6V
RB = (2.0 – 0.6)/0.10 = 14kΩ
The dc collector current IC = 100 × 0.10 = 10 mA.

Question 11.
Justify the output waveform (Y) of the OR gate for the following inputs A and B given in fig.
AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics Material, Devices and Simple Circuits 35
Solution:
Note the following :

  • At t ≤ t1; A = 0, B = 0; Hence Y = 0
  • For t1 to t2; A = 1, B = 0; Hence Y = 1
  • For t2 to t3; A = 1, B = 1; Hence Y = 1
  • For t3 to t4; A = 0, B = 1; Hence Y = 1
  • For t4 to t5; A = 0, B = 0; Hence Y = 0
  • For t5 to t6; A = 1, B = 0; Hence Y = 1
  • For t > t6; A = 0,B = 1; Hence Y = 1

Therefore the waveform Y will be as shown in the Fig.
AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics Material, Devices and Simple Circuits 36

AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics: Material, Devices and Simple Circuits

Question 12.
Take A and B input waveforms similar to that in Example 11. Sketch the output waveform obtained from AND gate.
Solution:

  • For t ≤ t1; A = 0, B = 0; Hence Y = 0
  • For t1 to t2; A = 1, B = 0; Hence Y = 0
  • For t2 to t3; A = 1, B = 1; Hence Y = 1
  • For t3 to t4; A = 0, B = 1; Hence Y = 0
  • For t4 to t5; A = 0, B = 0; Hence Y = 0
  • For t5 to t6; A = 1, B = 0; Hence Y = 0
  • For t > t6; A = 0,B = 1; Hence Y = 0

Based on the above, the output waveform for AND gate can be drawn as given below.
AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics Material, Devices and Simple Circuits 37

Question 13.
Sketch the output Y from a NAND gate having inputs A and B given below :
Solution:

  • For t ≤ t1; A = 1, B = 1; Hence Y = 0
  • For t1 to t2; A = 0, B = 0; Hence Y = 1
  • For t2 to t3; A = 0, B = 1; Hence Y = 1
  • For t3 to t4; A = 1, B = 0; Hence Y = 1
  • For t4 to t5; A = 1, B = 1; Hence Y = 0
  • For t5 to t6; A = 0, B = 0; Hence Y = 1
  • For t > t6; A = 0,B = 1; Hence Y = 1
    AP Inter 2nd Year Physics Important Questions Chapter 15 Semiconductor Electronics Material, Devices and Simple Circuits 38

AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry

AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry

Andhra Pradesh BIEAP AP Inter 1st Year Chemistry Study Material 5th Lesson Stoichiometry Textbook Questions and Answers.

AP Inter 1st Year Chemistry Study Material 5th Lesson Stoichiometry

Very Short Answer Questions

Question 1.
How many number of moles of glucose are present in 540 gms of glucose ?
Answer:
Formulae:
No. of moles = \(\frac{\text { Weight }}{\text { G.M. Wt }}\)
= \(\frac{540}{180}\)
= 3
Given
Weight of glucose = 540 gms
G.M. wt = 180 (C6H12O6)
No. of moles of glucose in 540 gms = 3 moles.

Question 2.
Calculate the weight of 0.1 mole of sodium carbonate.
Answer:
Formulae:
No. of moles = \(\frac{\text { Weight }}{\mathrm{GMWt}}\), 0.1 = \(\frac{\text { Weight }}{106}\)
Weight = 10.6 × 0.1 = 10.6 gms
Given
No. of moles = 0.1
G.M. Wt of Na2CO3 = 106

AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry

Question 3.
How many molecules of glucose are present in 5.23g of glucose (Molecular weight of glucose 180 u).
Answer:
Formulae:
No. of molecules = No. of moles × Avagadro no
No. of moles = \(\frac{\text { Weight }}{\text { GM Wt }}\) = \(\frac{5.23}{180}\) = 0.02906 moles
No. of molecules = 0.02906 × 6.023 × 1023
= 1.75 × 1022 molecules
Given
Wt of glucose = 5.23 gms
G.M. Wt = 180

Question 4.
Calculate the number of molecules present in 1.12 × 10-7 c.c. of a gas at STP (c.c. cubic centimeters = cm3).
Answer:
1 mole of any gas occupies 22400 cc of volume
1 mole of any gas contains 6.023 × 1023 molecules
∴ 22400 cc of gas at STP contains 6.023 × 1023 molecules
1.12 × 10-7 cc of a gas at STP contains ——–?
\(\frac{1.12 \times 10^{-7} \times 6.023 \times 10^{23}}{22400}\) = 3.015 × 1012 molecules

Question 5.
The empirical formula of a compound is CH2O. Its molecular weight is 90. Calculate the molecular formula of the compound. (A.P. Mar. ‘16) (Mar. 13)
Answer:
Formulae:
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 1
∴ Molecular formula = 3(CH2O) = Cr (NO3)3(aq) + Pb(s).
Given
Molecular wt = 90
Emperical formula = CH3O
∴ Emperical wt = 30

AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry

Question 6.
Balance the following equation by the oxidation number method
Cr(S) + Pb(NO3)2(aq) → Cr (NO3)3(aq) + Pb(s).
Answer:
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 2
So the balanced equation is 2Cr + 3Pb(NO3)2 → 2Cr(NO3)3 + 3Pb

Question 7.
What volume of H2 at STP is required to reduce 0.795 g of CuO to give Cu and H2O.
Answer:
Balance equation is
CuO + H2 → Cu + H2O
79.5 gms → 1 mole H2 gas for reduction → 22.4 lit of volume at STP
∴ 79.5 gms of CuO require 22.4 lit of H2 at STP
Then 0.795 gms of CuO require ——-?
\(\frac{0.795 \times 22.4}{79.5}\) = 0.224 lit.

AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry

Question 8.
Calculate the volume of O2 at STP required to completely burn 100 ml of acetylene.
Answer:
Balanced chemical equation for combustion of acetylene is
2C2H2 + 5O2 → 4CO2 + 2H2O
2 moles of C2H2 require 5 moles of O2 for complete combustion at STP
2 × 22400 ml of C2H2 require 5 × 22400 ml of O2 at STP
Then 100 ml of C2H2 require ————?
\(\frac{100 \times 5 \times 22400}{2 \times 22400}\) = \(\frac{500}{2}\) = 250 ml

Question 9.
Now-a-days it is thought that oxidation is simply decrease in electron density and reduction is increase in electron density. How would you justify this?
Answer:
Oxidation : The process of removing electrons from an element is called oxidation. It is also called
as de-electronation.

  • Oxidation means decrease in electron density
    e.g. : 2Cl → Cl2 + 2\(e^{\ominus}\)
  • Oxidation process occurs at anode in the electrolysis process.
  • Reduction : The process of addition of electrons to an element is called reduction. It is also called as electronation.
  • Reduction means increase in electron density.
    e.g.: Mg2+ + 2e → Mg
  • Reduction process occurs of cathode in the electrolysis process.

Question 10.
What is a redox concept? Give an example.
Answer:
Redox reactions are the reactions in which reduction and oxidation both takes place.
e.g.:
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 3

Question 11.
Calculate the mass per cent of the different elements present in sodium sulphate (Na2SO4).
Answer:
Given compound is sodium sulphate (Na2SO4)
Molecular weight of the compound = 2(23) + 1(32) + 4(16)
= 46 + 32 + 64
= 142

Step – I:
Mass percent of Na
142 gms of Na2SO4 → 46 gms of Na
100 gms of Na2SO4 → ?
\(\frac{100 \times 46}{142}\) = 32.39%

Step – II:
Mass percent of ‘S’
142 gms of Na2SO4 → 32 gms of ‘S’
100 gms of Na2SO4 → ?
\(\frac{100 \times 32}{142}\) = 22.53%

Step-III:
Mass percent of ‘0’
142 gms of Na2SO4 → 64 gms of oxygen
100 gms of Na2SO4 → ?
\(\frac{100 \times 64}{142}\) = 45.07%
∴ Mass percents of Na, S, O are 32.39, 22.53, 45.07.

Question 12.
What do you mean by significant figures?
Answer:
The meaningful digits which are known with certainty are called significant figures.

  • The uncertainty in the experimental (or) calculated values is indicated by mentioning the number of significant figures.

AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry

Question 13.
If the speed of light is 3.0 × 108 ms-1. Calculate the distance covered by light in 2.00 ns.
Answer:
Given the speed of light 3 × 108 m/sec.
i.e., 1 sec → 3 × 108 m
In 2 nano seconds → ?
2 × 10-9 sec
\(\frac{2 \times 10^{-9} \times 3 \times 10^8}{1}\) = 6 × 10-1 = 0.6 m

Short Answer Questions

Question 1.
The approximate production of sodium carbonate per month is 424 × 108 g, while that of methyl alcohol is 320 × 106 g. Which is produced more in terms of moles?
Answer:
Given that weight of Na2CO3 produced per month = 424 × 106 gms
No. of moles of Na2CO3 = \(\frac{w t}{\text { GMwt }}\) = \(\frac{424 \times 10^6}{106}\) = 4 × 106 moles per month
Given weight of CH3OH produced per month = 320 × 106 gms
No. of moles of CH3OH = = \(\frac{320 \times 10^6}{32}\) = 107 moles per month
Methyl Alcohol is produced more (per month) in terms of moles.

Question 2.
How much minimum volume of CO at STP is needed to react completely with 0.112 L of O2 at 1.5 atm pressure and 127°C to give CO2.
Answer:
Formula \(\frac{P_1 V_1}{T_1}\) = \(\frac{P_2 V_2}{T_2}\)
At STP
P1 = 1 atm
V1 = ?
T1 = 0°C = 273 K
Given P2 = 1.5 at m
V2 = 0.112 lit
T2 = 127°C = 400 K
\(\frac{1 \times V_1}{273}\) = \(\frac{1.5 \times 0.112}{400}\)
V1 = \(\frac{1.5 \times 0.112 \times 273}{400}\)
= 0.1147 lit
Chemical equation
2CO + O2 → 2CO2
2 moles of CO → 1 mole of O2 at STP
2 × 22.4 lit → 1 × 22.4 lit
? ← 0.1147 lit
\(\frac{0.1147 \times 2 \times 22.4}{22.4}\) = 0.2294 lit
The volume of CO required at STP = 0.2294 lit = 229.4 ml

Question 3.
Chemical analysis of a carbon compound gave the following percentage composition by weight of the elements present, carbon = 10.06%, hydrogen = 0.84%. chlorine = 89.10%. Calculate the empirical formula of the compound.
Answer:
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 4
∴ Emperical formula of given compound = C1H1Cl3
= CHCl3

Question 4.
A carbon compound on analysis gave the following percentage composition, carbon 14.5%, hydrogen 1.8%, chlorine 64.46%, oxygen 19.24%. Calculate the empirical formula of the compound.
Answer:
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 5
∴ The formula of the compound = C1 H1.5 Cl1.5 O1
Emperical formula of the compound = C2 H3 Cl3 O2

Question 5.
Calculate the empirical formula of a compound having percentage composition : Potassium (K) = 26.57, chromium (Cr) = 35.36; oxygen (0) 38.07. (Given the atomic weights of K. Cr and O as 39; 52 and 16 respectively).
Answer:
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 6
Formula of the compound = K1 Cr1 O3.5
Emperical formula of the compound = K2 Cr2 O7

Question 6.
A carbon compound contains 12.8% carbon, 2.1% hydrogen, 85.1% bromine. The molecular weight of the compound is 187.9. Calculate the molecular formula.
Answer:
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 7
∴ Emperical formula of the compound = C1 H2 Br
Molecular formula = n (Emperical formula)
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 8
Given molecular wt = 187.9
Emperical wt = 94 (CH2 Br)
∴ Molecular formula = 2 (CH2Br)
= C2H4Br2

Question 7.
0.188 g of an organic compound having an empirical formula CH2Br displaced 24.2 cc. of air at 14 °C and 752 mm pressure. Calculate the molecular formula of the compound. (Aqueous tension at 14°C is 12mm).
Answer:
Given Emperical formula CH2Br
Wt. of compound = 0.188 gms
Volume of displaced air = 24.2 CC
Temperature = 14°C = 287 K
Pressure = 752 mm.
STP conditions
P1 = 760 mm
v1 = ?
T1 = 273 K
Given conditions
P2 = Pressure – aqueous tension
= 752 – 12mm = 740 mm
V2 = 24.2K
T2 = 287 K

Formula

\(\frac{P_1 V_1}{T_1}\) = \(\frac{P_2 V_2}{T_2}\) ⇒ \(\frac{760 \times V_1}{273}\) = \(\frac{740 \times 24.2}{287}\)
V1 = \(\frac{740 \times 24.2}{287} \times \frac{273}{760}\) = 22.414ml (or) 22.414 CC

∴ 0.188 gms of organic compound displaced 22.414 CC of air gms of organic compound displaced 22400 CC of air
\(\frac{22400}{22.414} \times 0.188\) = 188 ∴ Gram molecular wt of compound = 188
Molecular formula = n (Emperical formula)
= n (CH2Br)
n = \(\frac{\text { Molecular wt }}{\text { Emperical wt }}\) = \(\frac{188}{94}\) = 2 ∴ Molecular formula = C2 H4 Br2

Question 8.
Calculate the amount of 90% H2S04 required for the preparation of 420 kg HCl.
2 NaCl + H2SO4 → Na2SO4 + 2 HCl
Answer:
Given equation is
2NaCl + H2SO4 → 4 Na2SO4 + 2HCl
1 Mole H2SO4 → 2 Moles HCl
98 gms of H2SO4 → 2 × 36.5 = 73 gms of HCl
? → 420 × 103 gms of HCl
\(\frac{420 \times 10^3 \times 98}{73}\) = 563.84 × 103 gms
For 100% H2SO4 → 563.84 × 103
For 90% H2SO4 → 563.84 × 103 × \(\frac{100}{90}\)
= 626.5 × 103 gms of H2SO4
∴ 626.5 × 103 gms of 90% H2SO4 is required to produce 420 × 103 gms of HCl

Question 9.
An astronaut receives the energy required in his body by the combustion of 34g of sucrose per hour. How much oxygen he has to carry along with him for his energy requirement in a day ?
Answer:
Given sucrose quantity required per hour = 34 gms
Sucrose quantity required per a day = 34 × 24 gms
Chemical equation of combustion of sucrose
C12 H22 O11 + 12O2 → 12CO2 + 11 H2O + Energy
1 mole sucrose → 12 moles of Oxygen
342 gms of sucrose → 12 × 32 gms of oxygen
34 × 24 gs of sucrose → ?
\(\frac{34 \times 24}{342}\) × 12 × 32 = 916.21 gms
∴ The astronaut require 916.21 gms of Oxygen per a day

Question 10.
What volume of CO2 is obtained at STP by heating 4g of CaCO3 ?
Answer:
Chemical Equation is
CaCO3 AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 9 CaO + CO2
1 mole CaCO3 → 1 mole CO2 at STP
100 gms of CaCO3 → 22.4 lit of CO2 at STP
4gms of CaCO3 → ?
\(\frac{4 \times 22.4}{100}\) = 0.896 lit

Question 11.
When 50g of a sample of sulphur was burnt in air 4% of the sample was left over. Calculate the volume of air required at STP containing 21% oxygen by volume.
Answer:
Given that when a 50 gms of sample of sulphur was burnt 4% of the sample was left over.
∴ 50 gms of S → 48 gms of ‘S’ burnt.
Chemical equation is S + O2 → SO2
32 gm of ‘S’ 22.4 lit of O2 at STP
48 gms of S →
100 ml air contains 21 ml of oxygen
100 lit of air contains 21 lit of oxygen
33.6 lit of oxygen is present in \(\frac{33.6 \times 100}{21}\) = 160 lit
∴ The volume of air needed for combustion = 160 lit

Question 12.
Calculate the volume of oxygen gas required at STP conditions for the complete combustion of 10 cc of methane gas at 20°C and 770 mm pressure.
Answer:
Given 10 cc of methane (CH4) undergoes combustion at 20°C and 770 mm pressure
STP Conditions
P1 = 760 mm
V1 = ?
T1 = 273 k
Given conditions
P2 = 770 mm
V2 = 10 cc
T2 = 20°C = 293 k

AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry

Formulae
\(\frac{P_1 V_1}{T_1}\) = \(\frac{\mathrm{P}_2 \mathrm{~V}_2}{\mathrm{~T}_2}\)
\(\frac{760 \times V_1}{273}\) = \(\frac{770 \times 10}{293}\)
V1 = \(\frac{770 \times 10 \times 273}{293 \times 760}\) = 9.44 cc ∴ Volume of methane (CH4) at STP = 9.44 cc.

Chemical equation is
CH4 + 2O2 → CO2 + 2H2O
1 mole CH4 → 2 moles of oxygen
22,400 cc of CH4 → 2 × 22400 cc of oxygen
9.44 cc of CH4 → ?
\(\frac{9.44}{22400}\) × 2 × 22400 = 18.88 cc ∴ Volume of oxygen gas required STP = 18.88 cc.

Question 13.
Calculate the volume of H2 liberated at 27°C and 760 mm of Hg pressure by action by 0.6g of magnesium with excess of dil HCl.
Answer:
Chemical equation is
Mg + 2HCl → MgCl2 +H2
24 gms. of Mg → 1 mole of H2 at STP
= 22.4 lit at STP
0.6 gms. of Mg → ?
\(\frac{0.6 \times 22.4}{24}\) = 0.56 lit = 560 ml

Formula
\(\frac{P_1 V_1}{T_1}\) = \(\frac{P_2 V_2}{T_2}\)
\(\frac{760 \times V_1}{300}\) = \(\frac{760 \times 560}{273}\)
V1 = \(\frac{560 \times 300}{273}\) = 615.4 ml
∴ volume of H2 liberated at 27°C and 760 mm of presence = 615.4 ml = 0.6154 Lit.

Given conditions
P1 = 760 mm
V1 = ?
T1 = 27° mm

STP conditions
P2 = 760 mm
V2 = 560 ml
T2 = 0° C = 273 k

Question 14.
Explain the role of redox reactions in titrimetre processes and galvanic cells.
Answer:
a) Redox reactions in titrimetric quantitative analysis : In titrimetric analysis the substance of known concentration is called the titrant and the substance being titrated is called the titrand. The standard solution is generally added from a long graduated tube called burette. The process of adding the standard solution until the reaction is just complete is called titration. The substance to be estimated is titrated. The point at which the titrand just completely reacts is called the equivalence point (or) the theoretical point (or) stoichiometric end point. In the redox reactions the completion of the titration is detected by a suitable method like

  1. Observing a physical change, for example, the light pink colour of KMnO4 titrations.
  2. By using a reagent known as indicator which gives a clear visual change like colour change, formation of turbidity etc. The point at wh’ch this is observed is called the end point of the titration which should coincide with theoretical end point.

Examples:

  1. In Cr2\(\mathrm{O}_7^{2-}\) titrations dippenyl amine is used as indicator and at the end poirtt it produces intense blue colour due to oxidation by Cr2\(\mathrm{O}_7^{2-}\).
  2. In the titration of Cu+2 with I
    2Cu2+ (aq) + 4I (aq) → Cu2I2 (s) + I2 (aq)
    A redox reaction I2 gives a deep blue colour with starch solution.
  3. In the titration of I2 (aq) S2\(\mathrm{O}_3^{-2}\) as per the stoichiometric reduce equation.
    In this way redox reactions are taken as the basis for titrimetric analysis with Mn\(\mathrm{O}_4^{-}\), Cr2\(\mathrm{O}_7^{-2}\) etc., as oxidising agents and S2\(\mathrm{O}_3^{-}\) etc., as a reducing agents,

b) Redox Reactions – Galvanic cells : Redox reaction (i.e) cell reaction that takes place in gavamic cell is
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 10
The process of transfer of electrons from Zn(s) to Cu+2 take place directly. To make this transfer indirectly, Zn rod is kept in ZnSO4 solution in one beaker and in the other beaker CuSO4 solution is taken and a copper rod is dipped in it. Now the redox reaction takes place in either of the beakers. Each beakers contains both oxidised and reduced forms of the respective species in the beakers containing CuSO4 solution & Cu rod at the interface Cu and Cu+2 and in the other beaker at the interface Zn and Zn+2. The two forms i.e., oxidised and reduced forms of a species participating in oxidation and reduction half reactions is called redox couple. Both the beakers contains each a redox couple. The oxidised form and reduced form are separated by a vertical line (or) a slash that represent an interface, e.g.: Zn (s) / Zn+2 (aq)

In the above arrangement, the two redox couples are represented by Zn+2 / Zn and Cu+2 / Cu.
halvamic cell is represented as : Zn/Zn2+//Cu+2/Cu.

Question 15.
Define and explain molar mass.
Answer:
Molar mass : The mass of one mole of any substance in gms is called its molar mass,
e.g. :

  1. Molar mass of sulphuric acid = 98 g
  2. Molar mass of hydrogen is one gram for gram atomic mass and two grams for gram molecular mass.

Question 16.
What are disproportionation reactions ? Give example. (T.S. Mar. ’16)
Answer:
Disproportionation Reaction : Reactions in which same element in the given form to undergo both oxidation and reduction simultaneously.
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 11
Here Cl2 undergoes both oxidation and reduction reactions. So, above reaction is disproportionation reaction.

Question 17.
What are comproportionation reactions ? Give example.
Answer:
Comproportionation reactions : In these reactions, two species with the same element in two different oxidation states form a single product in which the element is in an intermediate oxidation state. Reverse of disproportionation is comproportionation.
e.g.:
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 12

Question 18.
Determine the empirical formula of an oxide of iron which has 69.9% iron and 30.1% dioxygen by mass.
Answer:
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 13
∴ Formula of compound = Fe1O1.5
Emperical formula = Fe2O3

AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry

Question 19.
Calculate the mass of sodium acetate (CH3COONa) required to make 500 ml of 0.375 molar aqueous solution. Molar mass of sodium acetate is 82.0245 g mol-1.
Answer:
Formula
Molarity = \(\frac{\mathrm{Wt}}{\mathrm{GMwt}}\) × \(\frac{1000}{V(m l)}\)
Molarity = 0.375M
V = 500 ml
GMwt of CH3COONa = 82.0245
0.375 = \(\frac{W t}{82.0245}\) × \(\frac{1000}{500}\)
Wt = \(\frac{82.0245 \times 0.375}{2}\)
= 15.3795 gms
∴ Mass of CH3COONa required = 15.3795 gms

Question 20.
What is the concentration of sugar (C12H22O11) in mol L-1 if 20g are dissolved in enough water to make a final volume up to 2L ?
Answer:
Formula
Molarity = \(\frac{W t}{G M W t} \times \frac{1}{V(\text { lit })}\)
= \(\frac{20}{342} \times \frac{1}{2}\)
= 0.02924 M
Wt = 20 gms
GMWt = 342
[C12H22OH]
V = 2 lit

Question 21.
How many significant figures are present in the following ?

  1. 0.0025
  2. 208
  3. 5005
  4. 126,000
  5. 500.0
  6. 2.0034

Answer:

  1. 0.0025 has 2 significant figures
  2. 208 has 3 significant figures
  3. 5005 has 4 significant figures
  4. 126000 has 3 significant figures
  5. 500.0 has 4 significant figures
  6. 2.0034 has 5 significant figures

Question 22.
Round up the following upto three significant figures :

  1. 34.216
  2. 10.4107
  3. 0.04597
  4. 2808

Answer:

  1. 34.216 becomes 34.2
  2. 10.4107 becomes 10.4
  3. 0.04597 becomes 0.046
  4. 2808 becomes 281

Question 23.
Calculate the molarity of a solution of ethanol in water in which the mole fraction of ethanol is 0.040 (assume the density of water to be one)
Answer:
The relation between Molarity and mole fraction is given by
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 14
Given data, x2 = 0.040, M = ?
M1 = 18(H2O), M2 = 46 (C2H5OH), Cl = 1
0040 = \(\frac{M \times 18}{M(18-46)+1000 \times 1}\) = \(\frac{M \times 18}{-28 M+1000}\)
18 M = -0.040 × M × 28 + 0.040 × 1000
18 M = -1.12M + 40
(18 + 1.12)M = 40
19.12 M = 40
M = \(\frac{40}{19.12}\) = 2 09 M ∴ Molarity of Ethanol = 2.09 M

Given data
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 15

Question 24.
A welding fuel gas contains carbon and hydrogen only. Burning a small sample of it in oxygen gives 3.38 g carbon dioxide, 0.690 g of water and no other products. A volume of 10.0 L (measured at STP) of this welding gas is found to weigh 11.6 g. calculate
(i) empirical formula,
(ii) molar mass of the gas and
(iii) molecular formula.
Answer:
Assuming 1 gm. of gas is burnt
The weight % of ‘C’ = \(\frac{3.38 \times 12}{44}\) × 100 = 92.18%
The weight % of ‘H’ = \(\frac{0.69 \times 2}{18}\) = 7.67%
i)
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 16
Experical formula of compound = C1H1

ii) Given 10 lit of gas at STP weighs – 11.6 gas
22.44 lit of gas at STP weights = \(\frac{22.4 \times 11.6}{10}\) = 25.984
∴ Molecular weight of given gas = 25.984

iii) Molecular formula = n(emperical formula)
n = \(\frac{\text { Mol. wt }}{\text { Emp. wt }}\) = \(\frac{25.984}{13}\) = 2 ∴ Molecular formula = 2(CH) = C2H2.

Question 25.
Calcium carbonate reacts with aqueous HCl to give CaCl2 and CO2 according to the reaction CaCO3(5) + 2 HCl(aq) → CaCl2(aq) + CO2(g) + H2O (L). What mass of CaCO3 is required to react completely with 25 mL of 0.75 M HCl?
Answer:
Given chemical reaction
CaCO3(s) + 2HCl(aq) → CaCl2(aq) + CO2(g) + H2O(l)
Molarity = \(\frac{\mathrm{Wt}}{\mathrm{GMWt}} \times \frac{1000}{\mathrm{~V}(\mathrm{~m})}\)
0.75 = \(\frac{\mathrm{Wt}}{36.5} \times \frac{1000}{25}\)
M = 0.75 M
W = ?
Wt = \(\frac{0.75 \times 36.5 \times 25}{1000}\)
= 0.6844 gms of HCl
GMwt = 36.5
V = 25 ml
From the equation
1 mole of CaCO3 – 2 moles of HCl
100 gms of CaCO3 – 2 × 36.5 gms of HCl
? – 0.6844 gms of HCl
∴ Mass of CaCO3 = \(\frac{0.6844}{2 \times 36.5}\) × 100 = 0.937 5 gms

AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry

Question 26.
Chlorine is prepared in the laboratory by treating manganese dioxide (MnO2) with aqueous hydrochloric acid according to the reaction
4 HCl (aq) + MnO2(s) → 2 H2O (l) + MnCl2(aq) + Cl2(g)
How many grams of HCl react with 5.0 g of manganese dioxide?
Answer:
Given chemical reaction is
4HCl(aq) + MnO2(s) → MnCl2(aq) + 2H2O(l) + Cl2(g)
4 moles of HCl – 1 mole of MnO2
4 × 36.5 gms. of HCl – 87 gms. of MnO2
∴ Weight of HCl required = \(\frac{5 \times 4 \times 36.5}{87}\) = 8.39 gms.

Question 27.
To 50 ml of 0.1 N Na2CO3 solution 150 ml of H2O is added. Then calculate a normality of resultant solution.
Answer:
Formulae N1V1 = N2V2
0.1 × 50 = N2 × 200
N2 = \(\frac{0.1}{4}\)
= 0.025 N
N1 = 0.1 N
V1 = 50 ml
N2 = ?
V2 = 150 + 150 = 200 ml

Question 28.
Calculate the volume of 0.1 N H2SO4 required to neutralize 200 ml of 0.2 N NaOH solution. It is an acid base neutralisation reaction. Hence, at the neutralisation point. Number of equivalents of acid = Number of equivalents of base.
Answer:
Formula N1V1 = N2V2
0.1 × V1 = 0.2 × V2
V2 = 2 × 200
= 400 ml
N1 = 0.1 N
V1 = ?
V2 = 200 ml
N2 = 0.2 N

Question 29.
Calculate normality of H2SO4 solutions if 50 ml of it completely neutralises 250 ml of 0.1 N Ba(OH)2 solutions.
Solution:
Formulae N1V1 = N2V2
N × 50 = 0.1 × 250
N1 = 0.5 N
N1 = ?
V1 = 50 ml
N2 = 0.1 N
V2 = 250 ml

AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry

Question 30.
Calculate the volume of 0.1 M KMnO4 required to react with 100 ml of 0.1 M H2C2O4.2H2O solution in presence of H2SO4.
Answer:
Balanced chemical equation for the given data is
Formulae
\(\frac{M_1 V_1}{n_1}\) = \(\frac{M_2 V_2}{n_2}\)
\(\frac{0.1 \times V_1}{2}\) = \(\frac{0.1 \times 100}{8}\)
V1 = 40 ml

KMnO4
M1 = 0.1 M
V1 = ?
n1 = 2

Oxalic acid

M2 = 0.1 M
V2 = 100 ml
n2 = 5

Question 31.
Assign oxidation number to the underlined elements in each of the following species:
a) NaH2PO4
b) NaHSO4
c) H4P2O7
d) K2MnO4
e) CaO2
f) NaH4
g) H2S2O7
h) KAl(SO4)2.12 H2O
Answer:
a) NaH2PO4
1(+1) + 2(+1) + x + 4(-2) = 0
1 + 2 + x – 8 = 0
x – 5 = 0
x= + 5
Oxidation no. of ‘P in NaH2PO4 = + 5

b) NaHSO4
1(+1) + 1(+1) + x + 4(-2) = 0
1 + 1 + x – 8 = 0
x = +6
Oxidation no. of ‘S in NaHSO4 = + 6

c) H4P2O7
4(+1) + 2x + 7(-2) = 0
4 + 2x – 14 = 0
2x – 10 = 0.
x = +5
Oxidation no. of ‘P’ in H4P2O7 is +5

d) K2MnO4
2(+1) + x + 4(-2) = 0
2 + x – 8 = 0
x = + 6
Oxidation no of Mn is K2MnO4 = +6

e) CaO2
+ 2 + 2x = 0
2x = -2
x = -1
Oxidation no. of oxygen in CaO2 = -1

f) NaBH4
1(+1) + x + 4 (-1) = 0
1 + x – 4 = 0
x = +3
Oxidation no. of ‘B’ in NaBH4 = + 3
But B’ most probably exhibits – 3 oxidation state.

g) H2S2O7
2(1) + 2x + 7(-2) = 0
2 + 2x – 14 = 0
2x – 12 = 0
x = +6
Oxidation state of ‘S’ in H2S2O7 = + 6

h) K Al(SO4)2 12H2O :
General formula of above compound is K2SO4 Al2 (SO4)3 2 + H2O (Potash alum)
Consider Al2(SO4)3 from the above double salt
2x + 3 (-2) = 0
2x – 6 = 0
x = + 3

Question 32.
What are the oxidation number of the underlined elements in each of the following and how do you rationalise your results ?
a) KI3
b) H2SO4O6
c) Fe3O4
Answer:
a) KI3
It is formed by the combining KI, I2
∴ Oxidation no. of ‘I’ in KI = – 1
[1 + x = 0
x = – 1]
In I2 oxidation no. of. I = 0

b) H2S4O6
According to H2S4O6 structure
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 17
2(1) + 4x + 6(-2) = 0 ⇒ 4x – 10 = 0
x = 2.5
‘S’ average Ox. No = 2.5

c) Fe3O4
3x + 4(-2) = 0
3x – 8 = 0
x = \(\frac{8}{3}\)
In general Fe3O4 obtained by FeO + Fe2O3
∴ In FeO → x = + 2 [x – 2 = 0 ⇒ x = + 2]
In Fe2O3 → x = + 3 [2x – 6 = 0 ⇒ x = + 3]

Question 33.
Justify that the following reactions are redox reactions :
Answer:
a) CuO(s) + H2(g) → Cu(s) + H2O(g)
b) Fe2O3(s) + 3CO(g) → 2 Fe(s) + 3 CO2(g)
c) 4 BCl3(g) + 3 LiAlH4(s) → 2 B2H6(g) + 3 LiCl(s) + 3 AlCl3(s)
d) 2 K(s) + F2(g) → 2 K+F(s)
e) 4 NH3(g) + 5 O2(g) → 4 NO(g) + 6 H2O(g)
a)
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 18
Hence it is a Redox reaction.

b)
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 19
Hence it is a Redox reaction.

c) \(4 \mathrm{BCl}_{3_{(g)}}\) + \(3 \mathrm{LiAlH}_{4_{(s)}}\) → \(2 \mathrm{~B}_2 \mathrm{H}_{6_{(g)}}\) + \(3 \mathrm{LiCl}_{(s)}\) + \(4 \mathrm{AlCl}_{3_{(\mathrm{s})}}\)
Here in the above reaction the oxidation states of all elements doesnot changed. So it is not a redox reaction.

d)
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 20
Hence it is a redox reaction.

e)
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 21
Hence it is a redox reaction.

AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry

Question 34.
Fluorine reacts with ice and results in the change
H2O(S) + F2(g) → HF(g) + HOF(g)
Justify that this reaction is a redox reaction.
Answer:
Given chemical equation
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 22
Hence it is a redox reaction.

Question 35.
Calculate the oxidation number of sulphur, chromium and nitrogen in H2SO5, Cr2\(O_7^{-2}\) and \(\mathrm{NO}_3^{-}\). Suggest structure of these compounds.
Answer:
a) H2SO5 Structure:
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 23
2(1) + x + 2(-1) + 3(-2) = 0 (One peroxy linkage)
2 + x – 2 – 6 = 0
x = +6

b) Cr2\(\mathrm{O}_7^{-2}\) Structure:
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 24
2(x) + 7(-2) = -2
2x – 14 = -2
x = +6

c) \(\mathrm{NO}_3^{-}\) Structure:
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 25
x + 3(-2) = -1
x – 6 = -1
x = +5

Question 36.
Write formulas for the following compounds: .
a) Mercury (II) chloride
b) Nickel (II) sulphate
c) Tin (IV) oxide
d) Thallium (I) sulphate
e) Iron (III) sulphate
f) Chromium (III) oxide
Answer:
Formula
a) Mercury (II) Chloride — HgCl2
b) Nickel (II) Sulphate — NiSO4
c) Tin (IV) Chloride — SnCl4
d) Thallium (I) Sulphate — Tl2SO4
e) Iron (III) Sulphate — Fe2(SO4)3
f) Chromium (III) Oxide — Cr2O3

AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry

Question 37.
Suggest a list of the substances where carbon can exhibit oxidation states from -4 to +4 and nitrogen from -3 to +5.
Answer:
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 26
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 27

Question 38.
While sulphur dioxide and hydrogen peroxide can act as oxidising as well as reducing agents in their reactions, ozone and nitric acid act only as oxidants. Why ?
Answer:
SO3 and H2O2 both acts as oxidising as well as reducing agents.

  • Ozone acts as strong oxidising agent in acidic medium (Ozone acts as reducing agent also in some cases).
  • HNO3 acts as strong oxidising agent because in HNO3. He H+ ion readily released and the oxidation state of ‘N’ is +5 in HNO3. So it is a good oxidising agent.

Question 39.
Consider the reactions :
a) 6 CO2(g) + 6 H2O (l) → C6H12O6 (aq) + 6 O2(g)
b) O3 (g) + H2O2 (l) → H2O () + 2 O2 (g)
Why it is more appropriate to write these reactions as :
a) 6 CO2 (g) + 12 H2O(l) → C6H12O6(aq) + 6 H2O(l) + 6 O2(g)
b) O3(g) + H2O2(l) → H2O2(l) + O2(g) + O2(g)
Also suggest a technique to investigate the path of the above (a) and (b) redox reactions.
Answer:
a) 6CO2 + 12 H2O → C6H12O6 + 6 H2O + 6 O2
This equation is more appropriate writing equation of photosynthesis.
Because of evolution of oxygen from H2O but not from CO2.

b) O3(g) + H2O2(l) → H2O(l) + O2(g) + O2(g)
This is the more appropriate equation to write because in this reaction clearly mentioned that which is oxidised and which is reduced.

Question 40.
The compounds AgF2 is unstable compound. However, if formed, the compound acts as a very strong oxidising agent. Why ?
Answer:

  • AgF2 is an unstable compound.
  • If suppose it is formed it acts as good oxidising agent.

Reason :
AgF2 releases fluorine gas which is a powerful oxidising agent.
∴ AgF2 is a good oxidising agent.

Question 41.
Whenever a reaction between an oxidising agent and a reducing agent is carried out, a compound of lower oxidation state is formed if the reducing agent is in excess and a compound of higher oxidation state is formed if the oxidising agent is in excess. Justify this statement giving three illustrations.
Answer:
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 28
The above equations illustrates the given statement.

Question 42.
How do you count for the following observations ?
a) Though alkaline potassium permagnate and acidic potassium permanganate both are used as oxidants, yet in the manufacture of benzoic acid from toluene we use alcoholic potassium permanganate as an oxidant. Why ? Write a balanced redox equation for the reaction.
b) When concentrated sulphuric acid is added to an inorganic mixture containing chloride, we get colourless pungent smelling gas HCl, but if the mixture contains bromide then we get red vapour of bromine. Why ?
Answer:
a) Balanced equation of KMn04/H+ (Acidic medium)
Mn\(\mathrm{O}_4^{-}\) + 8 H+ + 5 e → Mn+2 + 4 H2O
Balanced equation of KMnO4/ H+ (Basic medium)
Mn\(\mathrm{O}_4^{-}\) + 2 H2O + 3 e → MnO2 + 4 OH

  • Toluene is oxidised to benzoic acid in alcoholic medium.
  • Toluene is easily oxidised in presence of alcoholic KMnO4
    AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 29

b) When Conc. H2SO4 reacts with NaCl then HCl vapours are evolved.
2 NaCl + H2SO4 → Na2SO4 + 2 HCl
When Conc. H4SO4 reacts with KBr than HBr vapours are evolved which on further oxidation gives reddish brown Br2 vapours.
2 KBr + H2SO4 → Na2SO4 + 2 HBr
2MBr + H2SO4 → 2 H2O + SO2 + Br2 (Reddish brown)

Question 43.
Identify the substance oxidised, reduced, oxidising agent and reducing agent for each of the following reactions :
a) 2 AgBr (s) + C6H6O2 (aq) → 2 Ag (s) + 2 HBr (aq) + C6H6O2 (aq)
b) HCHO(l) + 2[Af(NH3)2]+(aq) + 3 OH(aq) → 2 Ag(s) + HCOO(aq) + 4 NH3(aq) + 2 H2O(l)
c) HCHO(l) + 2 Cu2+(aq) + 5 OH(aq) → Cu2O(s) + HCOO(aq) + 3 H2O(l)
d) N2H4 (l) + 2 H2O2(l) → N2(g) + 4 H2O(l)
e) Pb(s) + PbO2(s) + 2 H2SO4(aq) → 2 PbSO4(s) + 2 H2O(l)
Answer:
a) Given equation is
2 AgBr(s) + C6H6O2(aq)

  • 2 Ag(s) + 2 HBr(aq) + C6H4O2 (aq)
  • C6H6O2 oxidised to C6H4O2
  • Ag+Br reduced to Ag
  • Oxidising agent is Ag+
  • Reducing agent is C6H6O2

b) Given equation is HCHO(l) + \(2\left[\mathrm{Ag}\left(\mathrm{NH}_3\right)_2\right]_{(\mathrm{aq})}^{+}\) + \(3 \mathrm{OH}_{(\mathrm{aq})}^{-}\)

  • 2 Ag(s) + \(3 \mathrm{OH}_{\text {(aq) }}^{-}\) + 4NH3(aq) + 2 H2O(l)
  • HCHO oxidised to HCOO
  • [Ag(NH3)2]+ reduced to Ag
  • [Ag(NH3)2]+ is oxidising agent
  • HCHO is reducing agent

c) Given equation is
HCHO(l) + \(2 \mathrm{Cu}_{(\mathrm{aq})}^{+2}\) + 5OH → Cu2O(s) + \(\mathrm{HCOO}_{(a q)}^{-}\) + 3H2O(l)

  • HCHO oxidised to HCOO
  • Cu+2 reduced to Cu+(in Cu2O)
  • Oxidising agent is Cu+2 ions (In Fehling’s reagent)
  • Reducing agent is HCHO

d) Given equation is
N2H4(l) + 2K2O2(l) → N2(g) + 4H2O(l)

  • \(N^{-2}\) oxidised to N2
  • \(\mathrm{O}_2^{-2}\) reduced to O-2
  • Oxidising agent is H2O2
  • Reducing agent is N2H4
    Pb(s) + pbO4(s) + 2H2SO4(aq)

e)

  • 2PbSO4(s) + 2H2O
  • Pb oxidised to Pb+2
  • PbO2 reduced to Pb+2
  • Oxidising agent is PbO2
  • Reducing agent is Pb

Question 44.
Consider the reactions :
2S2\(\mathrm{O}_3^{-2}\) (aQ) + I2 (s) → S4\(\mathrm{O}_6^{-2}\) (aQ) + 2I (aQ)
S2\(\mathrm{O}_3^{-2}\) (aQ) + I2(s) + 5H2O(l) → 2S\(\mathbf{O}_4^{-2}\) (aq) + 4\(\mathrm{Br}_{(\mathrm{aq})}^{-}\)(aq) + 10\(H_{(aq)}^{+}\)
Why does the same reductant, thiosulphate react differently with iodine and bromine ?
Answer:

  • Thio sulphate ion is not a strong reducing agent.
  • I2 is not a strong oxidising agent.
  • The reaction between I2 and thio sulphate form tetrathionate (S4\(\mathrm{O}_6^{-2}\)) ion.
    AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 30
  • The above reaction has high rate of reaction.
  • The reaction between Thiosulphate and Bromine involves the formation of sulphate ion.
    AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 31
  • Br2 is some what better oxidising agent than I2.
    Hence the difference observed in the above reactions.

Question 45.
Justify giving reactions that among halogens, fluorine is the best oxidant and among hydrohalic compounds, hydroiodic acid is the best reductant.
Answer:
a)

  • The oxidising capacity of any substance depends on the net result of several contributing energy factors like enthalpy change in a reaction, standard Electrode potential.
  • With reference to Born-Haber type energy cycle the change in enthalpy value is greater for fluorine.
  • Greater the magnitude of negative change in enthalpy greater is the oxidising power.

Supporting reaction :

  • Fluorine reacts with carbon while the other elements of group do not combine even under drastic conditions.
    C + 2 F2 → CF4
  • Fluorine is also called as superhalogen.

b)

  • Hydrohalic compounds acts as reducing agents. Their stability order is
    HF >> HCl > HBr > HI
  • HI is least stable and acts as strons reducing agent among these halides.
    H2 + I2 \(\rightleftharpoons\) 2 HI

Question 46.
Why does the following reaction occur ?
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 32
What conclusion about the compound Na4XeO6(of which AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 75 is a part) can be drawn from the reaction.
Answer:
Given equation is
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 34

  • In the above reaction ‘Xe’ under go reduction (+8 to +6)
  • F oxidised to F2 here perxenate ion acts as powerful oxidising agent.
  • Perxenates are stable in alkali solutions.
  • Na4XeO6 is a powerful oxidant.

AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry

Question 47.
Consider the reactions :
(a) H3PO2(aq) + 4 AgNO3(aq) + 2 H2O(1) → H3PO4(aq) + 4Ag(s) + 4HNO3(aq)
(b) H3PO2(aq) + 2 CuSO4(aq) + 2 H2O(1) → H3PO4(aq) + 2Cu(s) + H2SO4(aq)
(c) C6H5CHO(l) + 2[Ag (NH3)2]+(aq) + 3OH(aq) → C6H5COO(aq) + 2Ag(s) + 4NH3(aq) + 2H2O(l)
(d) C6H5CHO(l) + 2Cu2+(aq) + 5OH(aq) → No Change observed.
what inference do you draw about the behaviour of Ag+ and Cu2+ from these reactions ?
Answer:

  • H3PO2 is a strong reducing agent hence it reduce Ag+ to Ag and Cu+2 to Cu.
  • C6H5CHO is also a reducing agent it just reduced.
    Ag+ to Ag in tollen’s reagent but it does not reduced Cu+2 in alkaline solution.

Question 48.
Balance the following redox reactions by ion – electron method :
(a) Mn\(\mathrm{O}_4^{-}\) (aq) + I (aq) → MnO2(s) + I2 (s) (in basic medium)
(b) Mn\(\mathrm{O}_4^{-}\) (aq) + SO2 (g) → Mn2+ (aq) + HSO4 (aq) (in acidic solution)
(c) H2O2 (aq) + Fe2+ (aq) → Fe3+ (aq) + H2O(l) (in acidic solution)
(d) Cr2\(\mathrm{O}_7^{2-}\)SO2 (g) → Cr3+ (aq) + S\(\mathrm{O}_4^{2-}\) (aq) (in acidic solution)
Answer:
a) Mn\(\mathrm{O}_{4 \text { (aq) }}^{-}\) + \(\mathbf{I}_{(\mathrm{aq})}^{-}\) → MnO2(s) + I2(s) (Basic medium)

Reduction half cell
\(\mathrm{MnO}_4^{-}\) → MnO2
\(\mathrm{MnO}_4^{-}\) → MnO2 + 2H2O (Oxygens balanced)
\(\mathrm{MnO}_4^{-}\) + 4H2O → MnO2 + 2H2O + 4OH
(Hydrogens balanced)
\(\mathrm{MnO}_4^{-}\) + 2H2O → MnO2 + 4OH
\(\mathrm{MnO}_4^{-}\) + 2H2O + 3e → MnO2 +4OH]
(Charge balanced)
2 × [Mn\(\mathrm{MnO}_4^{-}\) + 2H2O + 3e → MnO2 + 4OH-1]
3 × [2 I → I2 + 2e]
Oxidation half cell
I → I2
2 I → I2 (Iodines balanced)
2 I → I2 + 2 e (Charge balanced)
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 35
The above equation is balanced equation.

b)
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 36
Above equation is balanced chemical equation.

c)
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 37

d)
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 38
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 39

Question 49.
Balance the following equations in basic medium by ion-electron method and oxidation number methods and identify the oxidising agent and the reducting agent
(a) P4(s) + OH(aq) → PH3(g) + HP\(\mathrm{O}_2^{-}\)
(b) N2H4(l) + Cl\(\mathrm{O}_3^{-}(\mathrm{aq})\)(aq) → NO(g) + Cl(g)
(c) Cl2O7 (g) + H2O2 (aq) → ClO2 (aq) + O2 (g) + H+
Answer:
a) P4 + OH → PH3 + HP\(\mathrm{O}_2^{-}\)
Reduction half cell
P4 → PH3
P4 → 4PH3
P4 + 12H2O → 4PH3 + 12 OH (H balanced)
P4 + 12H2O + 12e → 4PH3 + 12 OH
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 40

Oxidation half cell
P4 → H2P\(\mathrm{O}_2^{-}\)
P4 → 4H2P\(\mathrm{O}_2^{-}\) (‘P’ balanced)
P4 + 8H2O → 4H2\(\mathrm{PO}_2^{-}\) (Oxygen balanced)
P4 + 8H2O + 8 OH → 4H2\(\mathrm{PO}_2^{-}\) + 8H2O
(Hydrogen balanced)
P4 + 8OH → 4H2\(\mathrm{PO}_2^{-}\) + 4e
(Charge balanced)
Here P4 is oxidising agent as well reducing agent.

b)
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 41
\(\mathrm{ClO}_3^{-}\) is oxidant
N2H4 is reductant

c)
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 42

AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry

Question 50.
What sorts of informations can you draw from the following reaction ?
(CN)2(g) + 20H(aq → CN(aq) + CNO(aq) + H2O(1)
Answer:
(CN)2(g) + 2(OH) → \(\mathrm{CN}_{(\mathrm{aq})}^{-}\) + \(\mathrm{CNO}_{(\mathrm{aq})}^{-}\) + \(\mathrm{H}_2 \mathrm{O}_{(l)}\)
(CN)2 + 2e → 2 CN (Reduction)
(CN)2 + 2H2O → 2 CNO + 4H+ + 2e (Oxidation)
Here (CN2) undergo oxidation as well as reduction.
It is a dis proportionation reaction.

Question 51.
The Mn3+ ion is unstable in solution and undergoes disproportionation to give Mn2+. MnO2+ and H+ ion. Write a balanced ionic equation for the reaction.
Answer:
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 43

  • Above equation is a dis proportionation reaction

AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 44

Question 52.
Consider the elements :
Cs, Ne, I and F
(a) Identify the element that exhibits only negative oxidation state.
(b) Identify the element that exhibits only positive oxidation state.
(c) Identify the element that exhibits both postive and negative oxidation states.
(d) Identify the element which exhibits neither the negative nor does the positive oxidation state.
Answer:
a) ‘F’ Exhibits negative oxidation state i.e. ‘-1’
b) ‘Cs’ Exhibits only positive oxidation state
c) ‘I’ Exhibits positive as well as negative oxidation states
d) ‘Ne’ doesnot exhibit any oxidation state in ground state
(Neither positive nor negative)

Question 53.
Chlorine is used to purify drinking water. Excess of chlorine is harmful. The excess of chlorine is removed by treating with sulphur dioxide. Present a balanced equation for this redox change taking place in water.
Answer:
Balanced equations for the given data is
Cl2 + SO2 + H2O → SO3 + 2H+ + Cl

Question 54.
Refer to the periodic table given in your book and now answer the following questions :
a) Select the possible non metals that can show disproportionation reaction.
b) Select three metals that can show disproportionation reaction.
Answer:
a) Non metals
Chlorine, Bromine, Oxygen, Sulphur, Phosphorous, Iodine undergo disproportionations
b) Metals Cr, Mn and pb undergo disproportionation.

Question 55.
In Ostwald’s process for the manufacture of nitric acid, the first step involves the oxidation of ammonia gas by oxygen gas to give nitric oxide gas and steam. What is the maximum weight of nitric oxide that can be obtained starting only with 10.00 g, of ammonia and 20.00 g of oxygen ?
Answer:
Chemical reaction is
4NH3 + 5O2 → 4NO + 6H2O
4 moles NH3 – 5 moles of O2
Given 10 gms of NH3
∴ No. of moles = \(\frac{10}{17}\) = 0.588 moles
Given 20 gms of NH3
∴ No. of moles = \(\frac{20}{32}\) = \(\frac{5}{8}\) = 0.4
4 moles of ammonia reacts with 5 moles of O2
0.588 moles of NH3 …. ?
\(\frac{0.588}{4}\) × 5 = 0.735 moles
5 moles of O2 reacts with 4 moles of NH3
0.4 moles of O2…. ?
\(\frac{0.4}{5}\) × 4 = \(\frac{1.6}{5}\) = 0.32

  • Here O2 is not present in sufficient amount NH3 has sufficient in amount

5 moles of O2 → 4 moles of NO.
5 × 32 gms of O2 → 4 × 30 gms of NO.
20 gms of O2 → ?
\(\frac{20 \times 4 \times 30}{5 \times 32}\) = \(\frac{20 \times 24}{32}\) = \(\frac{120}{8}\) = 15 gms

AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry

Question 56.
Arrange the following metals in the order in which they displace each other from the solution of their salts.
Al, Cu, Fe, Mg and Zn.
Answer:
According to electro chemical series electrode potentials of given elements as follows
Al – -1.66 V
Cul – + 0.34 V
Fe – -0.40 V
Mg – -2.37 V
Zn – -0.76 V
∴ The order of the above metals in which they displace each other from the solution of their salts. Mg > AZ > Zn > Fe > Cu

Long Answer Questions

Question 1.
Wtire the balanced ionic equation which represents the oxidation of iodide (I-) ion by permanganate ion in basic medium at give iodine (I) and managanese dioxide (MnO2)
Answer:
Basic equation is
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 45
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 46
The above equation is balanced equation.

Question 2.
Write the balanced equation for the oxidation of suiphite ions to sulphate ions in acid medium by permanganate ion.
Answer:
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 47

AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry

Question 3.
Oxalic acid is oxidised by permanganate ion is acid medium of Mn2+ balance the reaction by ion — electron method.
Answer:
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 48

AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry

Question 4.
Phosphorus when heated with NaOH solution gives phosphine (PH3) and H2P\(\mathrm{O}_2^{-}\). Give the balanced equation.
Answer:
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 49
Here P4 is oxidising agent as well reducing agent.

Question 5.
Balance the following equation
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 50
Answer:
Oxidation half reaction Cr(OH)3 → \(\mathrm{CrO}_4^{2-}\)
(The ox. no of Cr changes from +3 to +6)
Reduction half reaction \(\mathrm{IO}_3^{-}\) → I (The ox. no. of ‘I’ changes from +5 to —1)
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 51
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 52

Question 6.
Balance the following equation by the oxidation number method.
Mn\(\mathrm{O}_4^{-2}\) + Cl2 → Mn\(\mathrm{O}_4^{-2}\) + Cl
Answer:
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 53

Question 7.
Exaplain the different types of redox reactions.
Answer:
Redox reaction : “The reaction that involves loss of electrons is called an oxidation reaction and that involving gain of electrons is called a reduction reaction. The overall reaction is called as oxidation – reduction reaction” or simply ‘redox reaction’.

Types of redox reactions:

a) Chemical combination reactions: In these reactions one species combine with another species to form product. In this conversion one species undergo oxidation and other species undergo reduction.
e.g.:
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 54
As the oxidation state of carbon increases from ‘0’ to ‘+4’ in the reaction. So, carbon undergoes oxidation.
Oxygen changes from ‘0’ to ‘-2’ oxidation state. So, it undergoes reduction. Therefore, the above overall combination reaction is redox reaction.

b) Decomposition reactions : Chemical compounds chemically split into two or more simpler substances during decomposition reactions. These are again redox reactions.
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 55
Here, water decomposes. H2 undergoes reduction from +1 to 0 in the reaction. Oxygen on the other hand changes its oxidation state from -2 to 0 (or) it undergoes oxidation.
Therefore, decomposition of H20 is a redox reaction.

c) Displacement reactions : In these reactions the place of one species in its compound is taken up by other species.
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 56
‘Zn’ displaces ‘Cu’ from CuSO4 solution. In this reaction Zn undergoes oxidation and Cu undergoes reduction. Therefore overall reaction is a redox reaction.

d) Disproportionation reactions : These reactions involve the same element in the given form to undergo both oxidation and reduction simultaneously.
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 57
In the above reaction ‘Cl2‘ undergoes both oxidation and reduction. It is a special type of redox reaction.

e) Comproportionation reactions : Reverse of disproportionation is comproportionation. In comproportionation reactions, two species with the same element in two different oxidation states form a single product in which the element is in an intermediate oxidation state.
e.g.:
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 58

AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry

Question 8.
Calculate the amount of carbon dioxide that could be produced when
(i) 1 mole of carbon is burnt in air.
(ii) 1 mole of carbon is burnt in 16 g of dioxygen,
(iii) 2 moles of carbon are burnt in 16 g of dioxygen.
Answer:
i) C + O2 → CO2
1 mole of ‘C’ → 1 mole of CO2
∴ 44 gms of CO2

ii) C + O2 → CO2
Given 1 mole of O2 burnt in 16 gms of di oxygen

iii) C + O2 → CO2
2 moles of ‘C’ burnt in 16 gms of O2
∴ 1 mole of O2 → 44 gms of CO2
32 gms of O2 → 44 gms of CO2
16 gms of O2 → 22 gms of CO2

AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry

Question 9.
Dinitrogen and dihydrogen react with each other to produce ammonia according to the following chemical equation :
N2(g) + H2(g) → 2NH3 (g)
(i) Calculate the mass of ammonia produced if 2.00 × 103 g dinitrogen reacts with 1.00 × 103 g of dihydrogen.
(ii) Will any of the two reactants remain unreacted ?
(iii) If yes, which one and what would be its mass ?
Answer:
i) N2 + 3H2 → 2NH3
Given Nitrogen = 2 × 103 gms
No. of moles = \(\frac{2000}{28}\) = 71.4285
Given hydrogen = 1 × 103 gms
No. of moles = \(\frac{1000}{2}\) = 500 moles
N2 is limiting agent
∴28 gms of N2 → 2 × 17gms NH3
2000 gms of N2
\(\frac{2000 \times 17}{28}\) = 2428.57 gms

ii) Used amount of hydrogen
28 → 6 gms
100 gms → ?
\(\frac{1000 \times 6}{28}\) = 214.285 gms

iii) Remaining amount of hydrogens
= 1000 – 2140.285
= 785.715 gms

AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry

Question 10.
Assign oxidation number to the underlined elements in each of the following epecies :
(a) NaH2PO4
b) NaHSO4
c) H4P2O7
d) K2MnO4
(e) CaO2
(f) NaBH4
(g) H2S2O7
(h) KA1(SO4)2.12 H2O
Answer:
(a) NaH2PO4
1(+1) + 2(+1) + x + 4 (-2) = 0
1 + 2 + x – 8 = 0
x – 5 = 0
x = +5
Oxidation no. of ‘P’ in NaH2PO4 = + 5

b) NaHSO4
1(+1) +1(+1) + x + 4(-2) = 0
1 + 1 + x – 8 = 0
x = + 6
Oxidation no. of ‘S’ in NaHSO4 = + 6

c) H4P2O7
4(+1) + 2x + 7(-2) = 0
4 + 2x – 14 = 0
2x – 10 = 0
x = + 5
Oxidation no. of ‘P’ in H4P2O7 is +5

d) K2MnO4
2(+1) + x + 4(-2) = 0
2 + x + 8 = 0
x = + 6
Oxidation no of Mn is K2MnO4 = + 6

e) CaO2
+ 2 + 2x = 0.
2x = -2
x = -1
Oxidation no. of oxygen in CaO2 = -1

f) NaBH4
1(+1) + x+ 4(-1) = 0
1 + x – 4 = 0
x = +3
Oxidation no. of ‘B’ in NaBH4 = +3
But ‘B’ most probably exhibits -3 oxidation state.

g) H2S2O7
2(1) + 2x + 7(-2) = 0
2 + 2x – 14 = 0
2x – 12 = 0
x = +6
Oxidation state of ‘S’ in H2S2O7 = + 6

h) k Al(SO4)2 12H2O :
General formula of above compound is
k2SO4 Al2 (SO4) 24 H2O (Potash alum)
Consider Al2(SO4)3 from the above double salt
2x + 3(-2) = 0
2x – 6 = 0
x = +3

Question 11.
What are the oxidation number of the underlined elements in each of the following and how do you rationalise your results?
(a) H2S4O6
(b) Fe3O4
(c) CH3CH2OH
(d) CH3COOH
Answer:
a) 45 (b)

b) 45 (c)

c) CH3 CH2 – OH
C2H6O
2x + 6(1) + (-2) = 0
2x + 6 – 2 = 0
2x + 4 = 0
x = -2

d) CH3COOH
C2H4O2
2x + 4(+1) + 2(-2) = 0
2x + 4 – 4 = 0
x = 0

Solved Problems

Question 1.
Calculate molecular mass of glucose (C6H12O6) molecule.
Solution:
Molecular mass of glucose (C6H12O6)
= 6(12.011 u) + 12(1.008 u) + 6(16.00.u)
= (72.066 u) + (12.096 u) + (96.00 u)
= 180.162 u

Question 2.
A compound contains 4.07% hydrogen, 24.27% carbon and 71.65% chlorine. Its molar mass is 98.96 g. What are its empirical and molecular formulae ?
Solution:
Step 1.
Conversion of mass per cent to grams :
Since we are having mass percent, it is convenient to use 100 g of the compound as the starting material. Thus, in the 100 g sample of the above compound, 4.07 g hydrogen; 24.27 g carbon; and 71.65 g chlorine are present.

Step 2.
Convertion into number of moles of each element:
Divide the masses obtained above by respective atomic masses of various elements.
Moles of hydrogen = \(\frac{4.07 \mathrm{~g}}{1.008 \mathrm{~g} \mathrm{~mol}^{-1}}\) = 4.04
Moles of carbon = \(\frac{24.27 \mathrm{~g}}{12.01 \mathrm{~g} \mathrm{~mol}^{-1}}\) = 2.021 Moles of chlorine = \(\frac{71.65 \mathrm{~g}}{35.453 \mathrm{~g} \mathrm{~mol}^{-1}}\) = 2.021

Step 3.
Divide the mole value obtained above by the smallest number:
Since 2.021 is smallest value, division by it gives a ratio of 2:1:1 for H:C:Cl.

In case the ratios are not whole numbers, then they may be converted into whole number by multiplying by the suitable coefficient.

Step 4.
These numbers indicate the relative number of atoms of the elements. Write empirical formula by mentioning the numbers after writing the symbols of respective elements :
CH2Cl is thus, the empirical formula of the above compound.

Step 5.
Writing molecular formula :

a) Determine empirical formula mass.
Add the atomic masses of various atoms present in the empirical formula.
For CH2Cl, empirical formula mass is 12.01 + 2 × 1.008 + 35.453 = 49.48 g

b) Divide Molar mass by empirical formula mass
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 59
= 2 = (n)

c) Multiply empirical formula by n obtained above to get the molecular formula.
Empirical formula = CH2Cl, n = 2. Hence molecular formula is C2H4Cl2.

AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry

Question 3.
Calculate the amount of water (g) produced by the combustion of 16 g of methane.
Solution:
The balanced equation for combustion of methane is :
CH4(g) + 2O2(g) → CO2(g) + 2H2O(g)

  1. 16 g of CH4 corresponds to one mole.
  2. From the above equation, 1 mol of CH4(g) gives 2 mol of H2O(g).
    2 mol of water (H2O) = 2 × (2 + 16)
    = 2 × 18 = 36 g

1 mol H2O = 18g H2O ⇒ \(\frac{18 \mathrm{~g} \mathrm{H}_2 \mathrm{O}}{1 \mathrm{~mol} \mathrm{H}_2 \mathrm{O}}\) = 1
Hence 2 mol H2O × \(\frac{18 \mathrm{~g} \mathrm{H}_2 \mathrm{O}}{1 \mathrm{~mol} \mathrm{H}_2 \mathrm{O}}\)
= 2 × 18 g H2O = 36 g H2O

Question 4.
How many moles of methane are required to produce 22 g CO2(g) after combustion ?
Solution:
According to the chemical equation
CH4(g) + 2O2(g) → CO2(g) + 2H2O(g)
44 g CO2 (g) is obtained from 16 g CH4(g).
[∵ 1 mol CO2(g) is obtained from 1 mol of CH4(g)]
Mole of CO2(g)
= 22 g CO2(g) × \(\frac{1 \mathrm{~mol} \mathrm{CO}_2(\mathrm{~g})}{44 \mathrm{gCO}_2(\mathrm{~g})}\)
= 0.5 mol CO2(g)
Hence 0.5 mol CO2(g) would be obtained from 0.5 mol CH4(g) or 0.5 mol of CH4(g) would be required to produce 22g CO2(g).

Question 5.
50.0 kg of N2(g) anfd 10.0 kg of H2(g) are mixed to produce NH3(g). Calculate the NH3 (g) formed. Identify the limiting reagent in the production of NH3 in this situation.
Solution:
A balanced equation for the above reaction is written as follows :
Calculation of moles:
N2 + 3H2(g) ⇌ 2HN3(g)
Moles of N2
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 60

According to the above equation, 1 mol N2(g) requires 3 mol H2(g), for the reaction. Hence, for 17.86 × 102 mol N2, the moles of H2(g) required would be
17.86 × 102 mol N2 × \(\frac{3 \mathrm{~mol} \mathrm{H}_2(\mathrm{~g})}{1 \mathrm{molN}_2(\mathrm{~g})}\)
= 5.36 × 103 mol H2

But we have only 4.96 × 103 mol H2. Hence, dihydrogen is the limiting reagent in
this case. So NH3(g) would be formed only from that amount of available dihydrogen i.e.,
4.96 × 103 mol.
Since 3 mol H2(g) gives 2 mol NH3(g)
4.96 × 103 mol H2 (g) × \(\frac{2 \mathrm{~mol} \mathrm{NH}_3(\mathrm{~g})}{3 \mathrm{molH}_2(\mathrm{~g})}\)
= 3.30 × 103 mol NH3 (g)
3.30 × 103 mol NH3 (g) is obtained.

If they are to be converted to grams, it is done as follows :
1 mol NH3(g) = 17.0 g NH3(g)
3.30 × 103 mol NH3 (g) × \(\frac{17.0 \mathrm{~g} \mathrm{NH}_3(\mathrm{~g})}{1 \mathrm{~mol} \mathrm{NH}_3(\mathrm{~g})}\)
= 3.30 × 103 × 17 g NH3 (g)
= 56.1 × 103 g NH3 (g)
= 56.1 kg NH3

Question 6.
A solution ¡s prepared by adding 2 g of a substance A to 18 g water. Calculate the mass percent of the solute.
Solution:
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 61

Question 7.
Calculate the molarity of NaOH in the solution prepared by dissolving 4 g in enough water to form 250 mL of the solution.
Solution:
Since molarity
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 62
Note that molarity of a solution depends upon temperature because volume of a solution is temperature dependent.

AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry

Question 8.
The density of 3 M solution of NaCl is 1.25 g mL-1. Calculate molality of the solution.
Solution:
M = 3 mol L-1
Mass of NaCl in 1 L solution
= 3 × 58.5 = 175.5 g
Mass of
1 L solution = 1000 × 1.25 = 1250 g
(Since density = 1.25 gmL-1)
Mass of water of solution = 1250 – 175.5
= 1074.5 g = 1.0745 kg
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 63

Often in a chemistry laboratory, a solution of a desired concentration is prepared by diluting a solution of known higher concentration. The solution of higher concentration is also known as stock solution. Note that molality of a solution does not change with temperature since mass remains unaffeoted with temperature.

Question 9.
Calculte the normality of oxalic acid solutions containing 6.3g of H2C2O4.2H2O in 500 ml of solutions.
Solution:
Weight of solute = 6.3 g
GEW of solute = \(\frac{126}{2}\) = 63
∴ Normality (N) = \(\frac{\omega}{\mathrm{GEW}} \times \frac{1000}{\mathrm{~V}(\mathrm{~mL})}\)
Normality (N) = \(\frac{6.3}{63} \times \frac{1000}{500}\) = 0.2 N

Question 10.
Calculate the mass of Na2CO3 required to prepare 250 ml of 0.5 N solution.
Solution:
Normality of required solution = 0.5 N
Volume of required solution = 250 ml
Equivalent wt.of. Na2CO3 = \(\frac{106}{2}\) = 53
Normality (N) = \(\frac{\mathrm{W}}{\mathrm{GEW}} \times \frac{1000}{\mathrm{~V}(\mathrm{~mL})}\)
wt. of solute = N × GEW × \(\frac{V(\mathrm{ml})}{1000}\)
= 0.5 × 53 × \(\frac{250}{1000}\) = \(\frac{53}{8}\) = 6.62 g

Question 11.
In the reactions given below, identify the species undergoing oxidation and reduction :
(i) H2S (g) + Cl2 (g) → 2HC (g) + S (s)
(ii) 3Fe3O4 (s) + 8Al (s) → 9Fe (s) + 4Al2O3 (s)
(iii) 2Na (s) + H2 (g) → 2NaH (s)
Solution:
(i) H2S is oxidised because a more electronegative element, chlorine is added to hydrogen (or a more electropositive element, hydrogen has been removed from S). Chlorine is reduced due to addition of hydrogen to it.

(ii) Aluminium is oxidised because oxygen is added to it. Ferrous ferric oxide (Fe3O4) is reduced because oxygen has been removed from it.

(iii) This reaction is very interesting. It may be said from the above definitions that the reaction is only a reduction reaction as it involves the addition of sodium (electropositive metal) or hydrogen whereas sodium undergoes oxidation and hydrogen undergoes reduction. It is to say that the above definitions for oxidation and reduction can not explain this and they have limitations like this. Therefore a new concept has to be considered for the oxidation and the reduction.

Question 12.
Justify that the reaction :
2Na (s) + H2 (g) → 2NaH (s) is a redox change.
Solution:
Since in the above reaction the compound formed is an ionic compound, which may also be represented as Na+H (s), this suggests that one half reaction in this process is :
2Na (s) → 2Na+ (g) + 2e and the other half reaction is H2 (g) + 2e→ 2H (g)
This splitting of the reaction under examination into two half reactions automatically reveals that here sodium is oxidised and hydrogen is reduced, therefore, the complete reaction is a redox change.

Question 13.
Using stock notation, represent the following compounds : HAuCl4, Tl2O, FeO, Fe2O3, Cul, CuO, MnO and MnO2.
Solution:
By applying various rules of calculating the oxidation number of the desired element in a compound, the oxidation number of each metallic element in its compound is as follows:
HAuCl4 → Au has 3
Tl2O → Tl has 1
FeO → Fe has 2
Fe2O3 → Fe has 3
Cul → Cu has 1
CuO → Cu has 2
MnO → Mn has 2
MnO2 → Mn has 4
Therefore, these compounds may be represented as :
HAu (III) Cl4, Tl2(I)O, Fe(II)O, Fe2(III)O3, Cu(I)I, Cu(II)O, Mn(II)O, Mn(IV)O2.

Question 14.
Justify that the reaction :
2Cu2O(S) + Cu2S(s) → 6Cu(s) + SO2(g) is a redox reaction. Identify the species oxidised/reduced, which acts as an oxidant and which acts as a reductant.
Solution:
Let us assign oxidation number to each of the species in the reaction under examination. This results into :
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 64
We therefore, conclude that in this reaction copper is reduced from +1 state to zero oxidation state and sulphur is oxidised from -2 state to +4 state. The above reaction is thus a redox reaction.

Further, Cu2O helps sulphur in Cu2S to increase its oxidation number, therefore Cu (I) is an oxidant; and sulphur of Cu2S helps copper both in Cu2S itself and Cu2O to decrease its oxidation number ; therefore, sulphur of Cu2S is reductant.

Question 15.
Which of the following species, do not show disproportionation reaction and why ? ClO, Cl\(\mathrm{O}_2^{-}\), ClO3 and ClO4. Also write reaction for each of the species that disproportionates.
Solution:
Among the oxoanions of chlorine listed above, ClO4 does not disproportionate because in this oxoanion chlorine is present in its highest oxidation state that is, +7. The disproportionation reactions for the other three oxoanions of chlorine areas follows :
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 65

Question 16.
Suggest a scheme of classification of
the following redox reactions
(a) N2 (g) + O2 (g) → 2NO (g)
(b) 2Pb(NO3)2(s) → 2PbO(s) + 2NO2(g) + 1/2 O2(g)
(c) NaH(s) + H2O (l) → NaOH (aq) + H2(g)
(d) 2NO2(g) + 2OH(aq) → N\(\mathrm{O}_2^{-}\) (aq) + N\(\mathrm{O}_3^{-}\) (aq) + H2O(l)
Solution:
In reaction
(a) the compound nitric oxide is formed by the combination of the elemental substances, nitrogen and oxygen ; therefore, this is an example of combination redox reactions. The reaction.

(b) involves the breaking down of lead nitrate into three components, therefore, this is categorised under decomposition redox reaction. In reaction.

(c) hydrogen of water has been displaced by hydride ion into dihydrogen gas. There-fore, this may be called as displacement redox reaction.

(d) The reaction involves disproportionation of NO2 (+4 state) into NO2 (+3 state) and NO3 (+5 state). Therefore reaction (d) is an example of disproportionation redox reaction.

Question 17.
Why do the following reactiolns proceed differently ?
Pb3O4 + 8HCl → 3PbCl2 + Cl2 + 4H2O and
Pb3O4 + 4HNO3 → 2Pb(NO3)3 + PbO2 + 2H2O
Solution:
Pb3O4 is actually a stoichiometric mixture of 2 mol. of PbO and 1 mol. of PbO2. In PbO2, lead is present in +4 oxidation state, whereas the stable oxidation state of lead in PbO is +2. PbO2 thus can act as an oxidant (oxidising agent) and, therefore, can oxidise Cl ion of HCl into chlorine. We may also keep in mind that PbO is a basic oxide. Therefore, the reaction
Pb3O4 + 8HCl → 3PbCl2 + Cl2 + 4H2O can be splitted into two reactions namely :
2 PbO + 4HCl → 2PbCl2 + 2H2O (acid base reaction)
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 66

Question 18.
Write the net ionic equation for the reaction of potassium dichromate (VI) K2Cr2O7 with sodium sulphite, Na2SO3 in an acid solution to give chromium (III) ion and the sulphate ion.
Solution:
Step – 1 :
The skeletal ionic equation is :
Cr2\(\mathrm{O}_7^{2-}\) (aq) + S\(\mathrm{O}_3^{-2}\) (aq) → Cr3+ (aq) + S\(\mathrm{O}_4^{2-}\) (aq)

Step – 2 :
Assign oxidation numbers for Cr and S
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 67
This indicates that the dichromate ion is the oxidant (it oxidises sulphite ion to sulphate ion) and the sulphite ion is the reductant. (it reduces dichromate ion to chromium (III).

Step – 3 :
Calculate the increase and decrease of oxidation numbers of respective species and make them equal.
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 68
As reduction is total 6 units due to two Cr3+ formed, oxidation also must be 6 units. This is obtained by multiplying S\(\mathrm{O}_3^{-2}\) with 3.

Step – 4 :
Adjust the coefficients of the products accordingly
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 69

Step – 5 :
a) Add H+ ions in acid medium or H2O molecules in basic medium in the required number to hydrogen atoms deficient side.
b) Add H2O molecules in acid medium or OH ions in basic medium in the required number to oxygen atoms diffident side, (a) and (b) may be repeated any number of times by hit and trial method until hydrogen and oxygen atoms are same in number on both the sides of the redox reaction.
The given reaction is in acid medium
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 70

AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry

Question 19.
Permanganate ion reacts with bromide ion in basic medium to give manganese dioxide and bromate ion. Write the balanced ionic equation for the reaction.
Solution:
Step – 1 : The skeletal ionic equation is :
Mn\(\mathrm{O}_4^{-}\) (aq) + Br (aq) → MnO2 (s) + Br\(\mathrm{O}_3^{-}\) (aq)

Step – 2 : Assign oxidation numbers for Mn and Br.
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 71
This indicates that permanganate ion is the oxidant and bromide ion is the reductant.

Step – 3 :
Calculate the increase or decrease in the oxidation number per atom and for the entire molecule/ion in which it occurs. The number units of oxidations change must be equal to the number units of reduction change. If this is not observed then multiply the oxidising agent with the number of units of oxidation and the reducing agent with the number of units of reduction i.e., multiply oxidant MnO4 by 2 and reductant Br by I.
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 72
Know number of units oxidation number of units reduction.

Step – 4:
Adjust the coefficients of products accordingly.
2Mn\(\mathrm{O}_4^{-}\) (aq) + Br (aq) → 2MnO2 (s) + BrO3 (aq)

Step – 5 :
a) Add H+ ions in acid medium or H2O molecules in basic medium in the required number to hydrogen atoms deficient side.
b) Add H2O molecules in acid medium or OH ions in basic medium in the required number to oxygen atoms dificient side, (a) and (b) may be repeated any number of times by hit and trial method until hydrogen and oxygen atoms are same in number on both the sides of the redox reaction. This reaction is basic medium.
2Mn\(\mathrm{O}_4^{-}\) (aq) + Br (aq) + H2O (I)

  • 2MnO2 (s) + Br\(\mathrm{O}_3^{-}\) (aq) + 2OH (aq)

Question 20.
Permanganate (VII) ion, MnO4 in basic solution oxidises iodide ion. I to produce molecular iodine (I2) and manganese (IV) oxide (MnO2). Write a balanced ionic equation to represent this redox reaction.
Solution:
Step – 1 :
First we write the skeletal ionic equation, which is
Mn\(\mathrm{O}_4^{-}\) (aq) + I (aq) → MnO2 (s) + I2 (s)

Step – 2 :
The two half-reactions are :
-1        0
Oxidation half :
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 73
Reduction half : Mn\(\mathrm{O}_4^{-}\) (aq) → MnO2 (s)

Step – 3 :
To balance the I atoms in the oxidation half reaction, we rewrite it as:
2I (aq) → I2 (s)

Step – 4:
As the reaction takes place to basic medium, to balance the 0 atoms in the reduction half reaction, we add OH ions in required number.
Mn\(\mathrm{O}_4^{-}\) (aq) → MnO2 (s) + 2HO (I)
To balance the H atoms, we add two H2O molecules on the left.
Mn\(\mathrm{O}_4^{-}\) (aq) + 2H2O (aq) → MnO2 (s) + 2 HO (I)
Balance H and 0 atoms by bit and trial method even if it requires morethan once we haye to do it.
The resultant equation is :
Mn\(\mathrm{O}_4^{-}\) (aq) + 2H2O (I) → MnO2 (s) + 4OH (aq)
Note : While balanceing H atom 0 atoms do not disturb the coefficients of other species (oxidant and reduCtant and the products).

Step-5 :
In this step we balance the charges of the two half-reactions in the manner depicted as:
2I (aq) → I2 (s) + 2e
Mn\(\mathrm{O}_4^{-}\) (aq) + 2H2O (I) + 3e → MnO2 (s) + 4OH (aq)

Now to equalise the number of electrons, in the two half reactions we multiply the oxidation hatf-reaction by 3 and the reduction half-reaction by 2.
6I (aq) → 3I2 (s) + 6e
2MnO4 (aq) + 4H4O (I) + 6e → 2MnO2 (s) + 8OH(aq)

Step-6:
Add two half-reactions to obtain the net reaction and canceel the electrons on both sides.
6I (aq) + 2Mn\(\mathrm{O}_4^{-}\) (aq) + 4H2O (I) → 3I2 (s) + 2MnO2 (s) + 8 OH (aq)

Step – 7:
A final verification shows that the equation is balanced in respect of the number of atoms and charges on both sides.

Question 21.
Calculate the normality of H2So4 solution. If 50 ml of it completely neutralised 250 ml of 0.2 N sodium hydroxide (NaoH) solution.
Answer:
AP Inter 1st Year Chemistry Study Material Chapter 5 Stoichiometry 74
∴ Normality of H2So4 solution = 1 N.

AP Inter 2nd Year Physics Important Questions Chapter 16 Communication Systems

Students get through AP Inter 2nd Year Physics Important Questions 16th Lesson Communication Systems which are most likely to be asked in the exam.

AP Inter 2nd Year Physics Important Questions 16th Lesson Communication Systems

Very Short Answer Questions

Question 1.
What are the basic blocks of a communication system ?
Answer:
Basic blocks in communication system are

  1. Transmitter
  2. Receiver
  3. Channel

Question 2.
What is ‘World Wide Web” (WWW) ?
Answer:
Tern Berners -Lee invented the World Wide Web.
It is an encyclopedia of knowledge accessible to every one round the clock through out the year.

AP Inter 2nd Year Physics Important Questions Chapter 16 Communication Systems

Question 3.
Mention the frequency range of speech signals.
Answer:
Speech signals frequency range is 300 Hz to 3100 Hz.

Question 4.
What is sky wave propagation ?
Answer:
In the frequency range from a MHz upto about 30 MHz, long distance communication can be achieved by ionospheric reflection of radio waves back towards the earth. This mode of propagation is called sky wave propagation.

Question 5.
Mention the various parts of the ionosphere ?
Answer:
Parts of ionosphere are

  1. D Part of stratosphere (65-70 km day only),
  2. E Part of stratosphere (100 km day only),
  3. F1 Part of mesosphere (170 km – 190 km),
  4. F2 Part of thermosphere [300 km at night 250 – 400 km during day time].

AP Inter 2nd Year Physics Important Questions Chapter 16 Communication Systems

Question 6.
Define modulation. Why is it necessary ? [A.P. 17; A.P., T.S. Mar. 16, T.S. Mar. 15, Mar. 14]
Answer:
Modulation : The process of combining low frequency audio signal with high frequency carrier wave is called modulation.
Necessary: Low frequency signals cannot transmit directly. To reduce size of the antenna and to avoid mixing up of signal from different transmitters modulation is necessary.

Question 7.
Mention the basic methods of modulation. [T.S. Mar. 15, 17, A.P. Mar. 16]
Answer:
The basic methods of modulation are :

  1. Amplitude modulation (AM)
  2. Frequency modulation (FM)
  3. Phase modulation (PM)

Question 8.
Which type of communication is employed in Mobile Phones ? [A.P. Mar. 15]
Answer:
Space wave mode of propagation is employed in mobile phones.

Textual Examples

Question 1.
A transmitting antenna at the top of a tower has a height 32 m and the height of the receiving antenna Is 50 m. What is the maximum distance between them for satisfactory communication in LOS mode? Given radius of earth 6.4 × 106 m.
Solution:
dm = \(\sqrt{2 \mathrm{Rh}_{\mathrm{T}}}+\sqrt{2 \mathrm{Rh}_{\mathrm{R}}}\)
dm = \(\sqrt{2 \times 64 \times 10^5 \times 32}+\sqrt{2 \times 64 \times 10^5 \times 50 \mathrm{~m}}\)
=64 × 102 × \(\sqrt{10}\) +8 × 103 × \(\sqrt{10}\)m = 144 × 102 × \(\sqrt{10}\)m
= 45.5 km.

AP Inter 2nd Year Physics Important Questions Chapter 16 Communication Systems

Question 2.
A message signal of frequency 10 kHz and peak voltage of 10 volts Is used to modulate a carrier of frequency 1 MHz and peak voltage of 20 volts. Determine
(a) modulation index
(b) the side bands produced.
Solution:
a) Modulation index = \(\frac{10}{2}\) = 0.5
b) The side bands are at (1000 + 10 kHz) = 1010 kHz and (1000 – 10 kHz) = 990 kHz.

AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics

AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics

Andhra Pradesh BIEAP AP Inter 1st Year Chemistry Study Material 6th Lesson Thermodynamics Textbook Questions and Answers.

AP Inter 1st Year Chemistry Study Material 6th Lesson Thermodynamics

Very Short Answer Questions

Question 1.
What is the information given by the term thermodynamics ?
Answer:

  • Thermodynamics gives information of energy changes involved in a chemical reaction (or) process.
  • The laws of thermodynamics deal with energy changes of macroscopic systems involving a large no.of molecules.

Question 2.
What is the relationship between the laws of thermodynamics and equilibrium state ?
Answer:

  • Laws of thermodynamics apply only when a system is in equilibrium (or) moves from one equilibrium state another equilibrium state.
  • Macroscopic properties like pressure and temperature do not change with time for a system in equilibrium state.

AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics

Question 3.
Define a system. Give an example.
Answer:
System : A small part of the universe chosen for thermodynamic study is called system.

Question 4.
The wall is adiabatic and AU = Wad. What do you understand about the heat and work with respect to the system ?
Answer:

  • In adiabatic process there is no transfer of heat between system and surroundings.
  • Work done adiabatic process is the change in internal energy of the system.
    Work done in adabatic process Wad = U2 – U1 = ∆U
  • Here work done on the system is positive and work done by the system is negative.

Question 5.
The system loses ‘q’ amount of heat though no work is done on the system. What type of wall does the system have ?
Answer:

  • Here heat is transferred through a thermally conducting walls.
  • In this case charge in internal energy ∆U = q can be measured interms of temperature difference.
    q = TB – TA
  • Here no work is done at constant volume,
    q = +ve → heat transferred from surrounding to system.
    q = -ve → heat transferred from system to surrounding.

AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics

Question 6.
Work is done by the system and ‘q’ amount at heat is supplied to the system. What type of system would it be ?
OR
Work is done by the system and heat (q) is supplied to the system. What type of system would it be ?
Answer:
The system is closed system .
Mathematical form
∆U = q – W
∆U = Change in internal energy
q = heat supplied
W = Workdone by the system.

Question 7.
What is the workdone in the free expansion of an ideal gas in reversible and irreversible processes ?
Answer:
In case of free expansion of an ideal gas pressure becomes zero, (into vaccum) .
∴ No workdone during free expansion of an ideal gas in case to reversible as well as irreversible process.

Question 8.
From the equation ∆U = q-pex∆V, if the volume is constant what is the value of ∆U.
Answer:
Suppose volume is constant change in volume ∆V = 0
∆U = q = pex∆V
∆U = q – 0
∆U = q
∴ Internal energy change = heat supplied.

Question 9.
In isothermal free expansion of an ideal gas find the value of q and ∆U.
Answer:
∆U = q + W
In case of free expansion of a gas in isothermal process.
W = 0, q = 0, ∆U = 0
Work done W = 0 (∵ Pexp = 0)
∴ q = 0, U = 0

AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics

Question 10.
In isothermal irreversible change of ideal gas what is value of q ?
Answer:
In case of an isothermal irreversible change,
q = W = 0
q = -W
= pexp(Vf – Vi)

Question 11.
In isothermal reversible change of an ideal gas, what is the value of q ?
Answer:
In case of isothermal irreversible change
q = W
q = nRTln \(\frac{V_f}{V_i}\)
q = 2.303 nRT log \(\frac{V_f}{V_i}\)

Question 12.
For an adiabatic change is an ideal gas what is the relationship between its ∆U and W (adiabatic) ?
Answer:
For an adiabatic change q = 0
∆U = q + W
∆U = 0 + W
∆U = W
Change in internal energy = Work done in adiabatic process.

Question 13.
State the first law of the thermodynamics.
Answer:
First law of thermodynamics :
The law of conservation of energy is taken as the first law of thermodynamics.
Statements :

  1. “Energy can neither be created nor destoyed, although it can be transformed from one form to another”.
  2. “It is impossible to construct perpetual motion machine of first kind”.
  3. “The total energy of the system, and surroundings in constant”.

AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics

Question 14.
What are the sign conventions of the work done on the system and work done by the system ?
Answer:

  • Work done on the system uses positive sign.
  • Work done by the system uses negative sign.

Question 15.
Volume (V), pressure (P) and Temperature (T) are state functions. Is the statement true ?
Answer:
Volume (V), Pressure (P) and Temperature (T) are state functions. These do not depends upon path or reaction but only depends upon state of reaction.

Question 16.
What are the heat (q) sign conventions when heat is transferred from the surroundings to the system and that transferred from system to the surrounding ?
Answer:
q = positive → Heat transferred from surrounding to the system.
q = negative → Heat transferred from system to the surroundings.

Question 17.
No heat is absorbed by the system from the surroundings, but work (w) is done on the system. What type of wall does the system have ?
Answer:
No heat is absorbed by the system from the surroundings but work (w) is done on the system. It is adiabatic change. Hence the wall is called adiabatic wall.

AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics

Question 18.
No work is done on the system, but heat (q) is taken out from the system by the surroundings. What type of wall does the system have ?
Answer:
No work is done on the system, but heat (q) is taken out from the system by the surroundings. Then the containing wall is thermally conducting wall.

Question 19.
q = w = -Pext (vf – vi) is for irreversible ………… change.
Answer:
q = w = -Pext (vf – Vi) is for irreversible isothermal change.

Question 20.
q = -w = nRT in (vf/vi) is for isothermal change.
Answer:
q = -w = nRT/n \(\frac{v_f}{v_i}\) is for isothermal reversible change.

Question 21.
What are the ‘∆H’ sign conventions for exothermic and endothermic reactions ? [T.S. Mar. ’16]
Answer:
Compounds which liberate heat in their formation reaction from the constituent elements are called exothermic compounds.
For this reactions, ∆Hf = negative
Compounds which absorb heat in the formation reaction from the constituent elements are called endothermic compounds.
For this reactions, ∆Hf = positive.

Question 22.
What are intensive and extensive properties ?
Answer:
Measurable (or) macroscopic properties such as mass, pressure, volume, temperature, surface tension, viscosity etc., can be subdivided into two categories as below :
i) Extensive properties : The properties whose magnitude depends upon the quantity of matter present in the system are called extensive properties. Examples of such properties are mass, volume, heat capacity, internal energy, entropy, heat content, Gibbs free energy etc. These properties change with quantity of matter present in the system. These properties are additive in nature.

ii) Intensive properties : The properties which do not depend upon the quantity of matter present in the system are called intensive properties. Examples of such properties are density, molar volume, molar entropy, molar heat capacity, surface tension, viscosity, specific heat, refractive index, pressure, temperature, boiling point, freezing point, vapour pressure etc. These properties depends only on the nature of the substance.

AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics

Question 23.
In the equation q = c. m. ∆T. If ∆T is change in temperature’m1 mass of the substance and ‘q1 is heat required, what is ‘c’ ?
Answer:
Given q = c × m × ∆T.
In the above equation c = Specific heat.
Specific heat: The amount of heat required to raise the temperature of 1gm of substance by 1°C is called specific heat.

Question 24.
Give the equation that gives the relationship between ∆U and ∆H.
Answer:
The equation that gives the relationship between ∆U and ∆H.
∆H = ∆U + ∆nRT
∆H = Change in Enthalpy
∆U = Change in Internal energy
∆n = nP – nR, R = Universal gas constant :
T = Temperature

Question 25.
What is the relationship between Cp and Cv ?
Answer:
Cp – Cv = R
Cp = Heat capacity at constant pressure
Cv = Heat capacity at constant volume
R = Universal gas constant.

Question 26.
1g of graphite is burnt in a bomb calorimeter in excess of O2 at 298K and 1 atm. pressure according to the equation.
C(graphite) + O2(g) → CO2(g)
During the reaction the temperature rises from 298K to 299K. Hat capacity of the bomb calorimeter is 20.7kJK-1. What is the enthalpy change for the above reaction at 298 K and 1 atm ?
Answer:
Given equation C(graphite) + O2(g) → CO2(g)
∆T = 299 – 298 = 1K
q = Cv × ∆T = -20.7 × 1 = -20.7 KJ
Given 1 gm of graphite ∴ ∆H = ∆U (∵ ∆h = 0)
∴ ∆U = – 20.7 KJ/k
For 1 mole
∆U = 12 × -20.7 = -248 KJ/mole

AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics

Question 27.
For the above reaction what is the internal energy change, ∆U ?
Answer:
For the above reactions
Internal Energy change ∆U = – 20.7 KJ/k

Question 28.
What is ∆rH for CH4(g) + 2O2(g) → CO2(g) + 2H2O(l) interms of molar enthalpies of the respective reactants and products ?
Answer:
rH = Sum of enthalpies of products – sum of enthalpies of reactants
= ∆Hp – ∆Hr
rH = (H(CO2) + 2 × H(H2O)) – (H(CH4) + 2 × H(O2))

Question 29.
Enthalpy decrease is not the criterion for spontaneity. Why ?
Answer:
Change in enthalpy “(∆H) = negative” may be a condition but not a necessary and sufficient condition for the spontaneous nature of a reaction.

Question 30.
Is increase of entropy the criterion for spontaneity ? Why ?
Answer:
Change in entropy “(∆S) = positive” is a condition but is not necessary and sufficient condition for the spontaneous nature of a reaction.

Question 31.
Explain the relationship between Gibbs energy change and equilibrium constant.
Answer:
∆G° =2.303 RT logk
∆G = Change in Gibbs free energy
K = Equilibrium constant

AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics

Question 32.
If we measure ∆Hθ and ∆Sθ it is possible to estimate ∆Gθ. Is it true ? why ?
Answer:
∆G° = ∆H° – T ∆S°
If we can measure ∆H° and ∆°S° then we can estimate ∆G° by the above equation, the given statement is true.

Question 33.
Equilibrium constant ‘K’ is measured accurately in the laboratory at given temperature. Is it possible to calculate ∆G° at any other temperature ? How ?
Answer:
It is possible to calculate ∆G° from the measurement of ∆H° and ∆S° and we can calculate K at any temperature.
∆G° = ∆H° – T ∆S°

Question 34.
Comment on the thermodynamic stability of NO(g) given that
\(\frac{1}{2}\) N2(g) + \(\frac{1}{2}\) O2(g) → NO(g); ∆rH° = 90kJ mol-1
NO(g) + \(\frac{1}{2}\) O2(g) → NO2(g); ∆rH° = -74kJ mol-1
Answer:
\(\frac{1}{2}\) N2(g) + \(\frac{1}{2}\) O2(g) → NO(g); ∆rH° = 90kJ mol-1
NO(g) + \(\frac{1}{2}\) O2(g) → NO2(g); ∆rH° = -74kJ mol-1
NO(g) formation is endothermic (∆H = + ve)
NO2(g) formation is exothermic (∆H = -ve)
∴ NO(g) is thermodynamically unstable.

Question 35.
Calculate the entropy change in surroundings when 1.00 mole of H2O(1) is formed under standard conditions ∆fH° = – 286k J mol-1.
Answer:
Formula ∆S = \(\frac{\Delta_{\mathrm{f}} \mathrm{H}^{\circ}}{\mathrm{T}}=\frac{-286}{273}\) = -1.0476 KJ/mole.K
∆Ssurrounding = 1.0476 KJ/mole.k

AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics

Question 36.
The equilibrium constant for a reaction is 10. What will be value of ∆G°? R = 8.314JK-1 mol-1,T = 300K.
Answer:
Formulae ∆G° = —RTlnk
K = 10
R = 8.314 J/K.mole
T = 300 K
= – 2.303 RT logK
= – 2.303 × 8.314 × 300 × log 10 = – 5774.14 J/mole

Question 37.
State the third law of thermodynamics. [T.S. Mar. 16]
Answer:
Third law of thermodynamics: “The entropy of a pure and perfectly crystalline substance is zero at the absolute zero of a temperature (- 273° C)”.
Slim T → 0 = 0
This law imposes a limitation on the value of entropy.
ST = \(\int_0^T \frac{C_p}{T} \cdot d T\)

Short Answer Questions

Question 1.
What are open, closed and isolated systems ? Give one example for each.
Answer:
Types of systems : The systems are classified into three types. These are :
a) Open system : A system which can exchange matter as well as energy with the surroundings is called an open system.
Ex :Evaporation of water from a beaker presents as open system.
AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics 2
Here vapours of water (matter) go into the atmosphere and heat (energy) required is absorbed by water from the surroundings.

b) Closed system : A system which may exchange energy but not matter with surroundings is called a closed system.
Ex : Consider boiling of water in a closed metallic vessel. Here, heat is transferred from the burner (surroundings) to the system. Steam remains inside the vessel. Thus, matter is not exchanged.
AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics 3

c) Isolated system : A system which can neither exchange matter nor energy with the surroundings is called an isolated system.
Ex : Take some water in an insulated vessel and put a small piece of sodium metal in it. An exothermic reaction takes place.
AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics 4
Neither hydrogen gas (matter) nor heat (energy) is transferred to the surroundings.

AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics

Question 2.
Define the state function and state variables, Give examples.
Answer:
State function : Properties of the systems which depends on the initial and the final states of the system but are independent of the path are called state functions.
Eg. : Energy, Volume, Enthalpy, Gibbs Energy.
State variables : Variables like P, V, T are called state variables. These are used to describe a system completely.

Question 3.
“Internal energy is a state function”. Explain.
Answer:
Internal energy (E) : ‘The total energy stored in the substance at constant temperature and pres-sure is called its internal energy (E)”.
It is a state function and is an extensive property.
Internal energy change (∆E) is considered as the difference between the internal energies of products and reactants.
∆E = (Ep – Er)
In any path taken (Q – W) is equal to ∆E.
Where Q = heat and W = Work.
Internal energy is a state function it depends on initial state and find state of the system.

Question 4.
“Work is not a state function”. Explain.
Answer:
Work is not a state function.

  • One state of a system can be changed to another in a several ways. Work involved in different paths differ in their quantity.
  • Work depends on the path of the reaction but not on the state of the system.
  • Work is dependent on the capacity factor and intensity factor.

AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics

Question 5.
What is heat ? Explain.
Answer:
The internal energy of a system can be change by transfer of heat from the surroundings to system (or) vice versa with out consumption of work.
This exchange of energy which results of temperature difference is called heat (q).

Question 6.
Derive the equation for ‘Wrev‘ in isothermal reversible process.
Answer:
Derivation of equation for Wrev in isothermal reversible process :
Let the work done against the pressure Sw = -P.dv
Assume that expansion isothermally from volume V1 to volume V2
[P = pressure (External)
dv = change in volume]
AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics 1

Question 7.
Two litres of an ideal gas at a pressure of 10 atm expands isothermally into a vacuum until its total volume is 20 litres. How much heat is absorbed and how much work is done in the expansion ?
Answer:
Heat absorbed = 0
‘P‘ of gas = 10 at m
PExp = 0
V1 = 2 lit
V2 = 20 lit
Because the given process is isothermal
∴ ∆U = 0
q + w = 0
q = -w
= pex (V2 – V2)
= Pex(20 – 2) = 0
= 0 (18) = 0
q = 0 and work done also zero
∴ w = 0

AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics

Question 8.
If the ideal gas given in the problem 45 expands against constant external pressure of 1 atm what is the q value ?
Answer:
Formula
q = – W
pex = 1 atm
pgas = 10 atm
V1 = 2 lit
V2 = 20 lit
= pex(VV2 – V2)
= 1 (20 – 2)
= 18 lit. atm

Question 9.
If the ideal gas given in the problem 45 expands to a final volume of 10L conducted reversibly what is q value ?
Answer:
V1 = 2 lit, V2 = 10 lit.
The expression for reversible isothermal expansion; q = 2.303 log \(\frac{\mathrm{v_2}}{\mathrm{v_1}}\)
q = -w = 2.303 × log \(\frac{20}{2}\) = 2.303 × 20 log 10
= 46.06 lit. atm.

Question 10.
Explain the state function enthalpy, H’. What is the relationship between ∆U and ∆H ?
Answer:
Enthalpy (H) : “Enthalpy is the amount of heat exchanged by a system with its surroundings at constant pressure and temperature”.
The energy change taking place at constant pressure and at a constant temperature is called enthalpy change (∆H).
Mathematically, ∆H = ∆U + P. ∆N. When ∆U = Internal energy change.
Enthalpy is state function. Thus, the magnitude of enthalpy change depends only upon the enthalpies in the initial and the final states.
∆H = [HProducts – HReactants]
For gaseous reactions, ∆H = ∆U + ∆nRT

AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics

Question 11.
Show that ∆H = ∆U + ∆n(g), RT
Answer:
The energy change taking place at constant pressure and at a constant temperature is called enthalpy change (∆H).
Mathematically, ∆H = ∆U + P. ∆V. When ∆U = Internal energy change.
Enthalpy is state function. Thus, the magnitude of enthalpy change depends only upon the enthalpies in the initial and the final states.
∆H = [HProducts – Hreactants]
For gaseous reaction, ∆H = ∆U + ∆nRT
Consider PV1 = n1 RT
PV2 = n2 RT
PV2 – PV1 = n2 RT – n1 RT
P(V2 – V1) = (n2 – n11)RT
P × ∆V = ∆ ng RT
[We know that ∆H = ∆U + P∆V]
∆H = ∆U + ∆ng RT

Question 12.
If Water vapour is assumed to be a perfect gas, molar enthalpy change for vapourisation of 1 mole of water at 1 bar and 100° C is 41 kJ mol-1. Calculate the internal energy change when
a) 1 mol of water is vapourised at 1 bar and 100° C
b) 1 mol of water liquid is converted into ice.
Answer:
∆H = 41kJ
∆ng = 1
a) H2O ⇌ H2O(g)
∆H = ∆U + ∆n(g) RT
41 = ∆U + 8.314 × 10-3 × 373
∆U = 41 – 3.101 = 37.899 KJ / mole

b) H2O ⇌ H2O(ice)
∆H = ∆U + ∆n(g) RT
∆H = ∆U [∵ ∆n = 0]
∆H = ∆U = 41 kJ /mole

AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics

Question 13.
Explain extensive and intensive properties.
Answer:
Measurable (or) macroscopic properties such as mass, pressure, volume, temperature, surface tension, viscosity etc., can be subdivided into two categories as below :
i) Extensive properties: The properties whose magnitude depends upon the quantity of matter present in the system are called extensive properties. Examples of such properties are mass, volume, heat capacity, internal energy, entropy, heat content, Gibbs free energy etc. These properties change with quantity of matter present in the system. These properties are additive in nature.

ii) Intensive properties: The properties which do not depend upon the quantity of matter present in the system are called intensive properties. Examples of such properties are density, molar volume, molar entropy, molar heat capacity, surface tension, viscosity, specific heat, refractive index, pressure, temperature, boiling point, freezing point, vapour pressure etc. These properties depends only on the nature of the substance.

Question 14.
Define heat capacity. What are Cp and Cv ? Show that Cp – Cv = R.
Answer:
Heat capacity (C) : “Heat capacity (G) of a substance is defined as the amount of heat required to raise its temperature through one degree”.
(or)
It is the ratio of heat absorbed (q) to the resulting increase in temperature (dT)
∴ C = \(\frac{\mathrm{q}}{\mathrm{dT}}\)
Heat capacity (C) is a state function. Hence to evaluate ‘C’, the conditions such as volume or constant pressure have to be specified in order to define the path. Thus there are two different types of heat capacities. These are :
i) Heat capacity at constant volume (Cv)
ii) Heat capacity at constant pressure (Cp).
From first law of thermodynamicsl,
q = dE + W = dE + P.dV (∵ W = PV)
∴ C = \(\frac{\mathrm{q}}{\mathrm{dT}}\) = \(\frac{\mathrm{dE}+\mathrm{P} \cdot \mathrm{dV}}{\mathrm{dT}}\)
If the absorption of heat takes place at constant volume, Cv (i.e.,) dV = 0
∴ Cv = \(\frac{\mathrm{q}_{\mathrm{V}}}{\mathrm{dT}}=\frac{\partial \mathrm{E}+\mathrm{O}}{\partial \mathrm{T}}=\left(\frac{\partial \mathrm{E}}{\partial \mathrm{T}}\right)_V\)
∴ Cv = \(\left(\frac{\partial E}{\partial T}\right)_V\)
Definition for “Cv“: Heat capacity of a system at constant volume (Cv) may be defined as rate of change of internal energy with temperature at constant volume.
If heat is absorbed by the system at constant pressure, heat capacity is denoted by ‘Cp‘ and is called heat capacity at constant pressure.
Cp = \(\frac{\mathrm{q}_{\mathrm{p}}}{\mathrm{dT}}=\left(\frac{\partial \mathrm{E}}{\partial \mathrm{T}}\right)_{\mathrm{p}}+\mathrm{P}\left(\frac{\partial \mathrm{V}}{\partial \mathrm{T}}\right)_{\mathrm{p}}\) ………… (1)
But H = E + PV
Differentiating this equation w.r.to T at constant pressure, we get
\(\left(\frac{\partial \mathrm{H}}{\partial \mathrm{T}}\right)_p=\left(\frac{\partial \mathrm{E}}{\partial \mathrm{T}}\right)_p+\left(\frac{\partial \mathrm{V}}{\partial \mathrm{T}}\right)_p\) ………………. (2)
From equation (1) and (2), we get
Cp = \(\left(\frac{\partial \mathrm{H}}{\partial \mathrm{T}}\right)_p\)
Definition for Cp: Heat capacity of a system at constant pressure (Cp) may be defined as rate of change of enthalpy with temperature at constant pressure.
Derivation of Cp – Cv = R
H = E + PV (For an ideal gas)
differentiating w.r.t. T.
\(\frac{\mathrm{dH}}{\mathrm{dT}}=\frac{\mathrm{dE}}{\mathrm{dT}}+\frac{\mathrm{d}(\mathrm{PV})}{\mathrm{dT}}\)
⇒ \(\frac{\mathrm{dH}}{\mathrm{dT}}=\frac{\mathrm{dE}}{\mathrm{dT}}+\frac{\mathrm{d}(\mathrm{RT})}{\mathrm{dT}}\) [∵ PV = RT]
\(\frac{\mathrm{dH}}{\mathrm{dT}}=\frac{\mathrm{dE}}{\mathrm{dT}}+\mathrm{R}\)
Cp – Cv = R

AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics

Question 15.
Explain the determination of ‘∆U’ of a reaction calorimetric
Answer:

  • The technique of measuring heats of reactions is called calorimetres. The apparatus used is calorimeter.
  • Calorimeter is a vessel properly insulated and contains water as the calorimetric liquids.
  • A combustioble substance is burnt in oxygen.
  • Heat evolved during the reaction is measured from the rise in temperature.
  • Here there is no worke done as ∆V = 0
  • By the raise of temperature of water bath and calorimeter the amount of heat evolved can be determined.
    AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics 5
    Internal Enegy ∆U = Q. ∆t × \(\frac{M}{m}\)
    Q = heat capacity
    ∆t = t2 – t1
    M = molar mass of substance
    m = mass of the substance

Question 16.
Explain the determination of ∆H of a reaction calorimetrically.
Answer:
Heat change at constant pressure can be measured in a calorimeter at constant pressure ∆H = qp
∴ For Exothermic reaction ∆H = -ve
For Endothermic reaction ∆H = +ve
AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics 6

Question 17.
What is enthalpy of a reaction ? Explain the standard enthalpy of a reaction.
Answer:
Enthalpy (H) : “Enthalpy is the amount of heat exchanged by a system with its surroundings at constant pressure and temperature”.
The energy change taking place at constant pressure and at a constant temperature is called enthalpy change (∆H).
Mathematically, ∆H = ∆U + P. ∆V. When ∆U = Internal energy change.
Enthalpy is state function. Thus, the magnitude of enthalpy change depends only upon the enthalpies in the initial and the final states.
∆H = [HProducts – HReactants]
For gaseous reactions, ∆H = ∆U + ∆nRT

AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics

Question 18.
What is the standard enthalpy of formation ? Explain it with example.
Answer:
The heat change involved in the reaction of formation of one mole of a compound is called enthalpy of formation.
If all the substances are in the standard states the enthalpy of formation is called standard enthalpy of formation.
AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics 7
In the above reactions the substances are in standand states

Question 19.
Define and explain enthalpy of phase transformation.
Answer:
Enthalpy of phase transformation : The heat change involved in the change of phase (or) physical state of one mole of compound at atmospheric pressure.
a) Standard molar enthalpy of fusion (∆fusH0) : The enthalpy change that accompanies melting of one mole of a soild substance in standard state is called Standard molar enthalpy of fusion.
Eg : H2O(s) → H2O(l)fusH0 = 6.00 KJ mol-1

b) Standard molar enthalpy of vapourisation (∆vapH0): Amount of heat required to vapourise 1 mole of a liquid at constant temperature and under standard pressure (1 bar) is called its standard molar enthalpy of vapourisation (∆vapH0).
Eg : H2O(l) → H2O(g)vapH0 = + 40.79 KJ mol-1

c) Standard enthalpy of sublimation (∆subH0): It is the change in the enthalpy when are mole of a solid substance sublimes at a constant temperature under standard pressure.
AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics 8

Question 20.
Define and explain the standard enthalpy of vapourisation (Molar enthalpy of vapourisation)
Answer:
Standard molar enthalpy of fusion (∆fusH0) : The enthalpy change that accompanies melting of one mole of a soild substance in standard state is called Standard molar enthalpy of fusion.
Eg : H2O(s) → H2O(l)fusH0 = 6.00 KJ mol-1

Question 21.
Define and explain the standard enthalpy of vapourisation (Molar entghalpy of vapourisation)
Answer:
Standard molar enthalpy of vapourisation (∆vapH0) : Amount of heat required to vapourise 1 mole of a liquid at constant temperature and under standard pressure (1 bar) is called its standard molar enthalpy of vapourisation (∆vapH0).
Eg : H2O(l) → H2O(g)vapH0 = + 40.79 KJ mol-1

AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics

Question 22.
Define and explain the standard enthalpy sumlimation.
Answer:
Standard enthalpy of sublimation ∆subH0 : It is the change in the enthalpy when are mole of a solid substance sublimes at a constant temperature under standard pressure.
AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics 8

Question 23.
Define and explain the standard enthalpy of formation (∆f Hθ).
Answer:
The heat change involved in the reaction of formation of one mole of a compound is called Enthalpy of formation
If all the substances are in the standard states the enthalpy of formation is called standard enthalpy of formation.
AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics 7

Question 24.
State and explain the Hess’s law of constant Heat summation. [Mar. ’13]
Answer:
Hess’s law states that the total amount of heat evolved or absorbed in a chemical reaction is always same whether the reaction is carried out in one step (or) in several steps.
Illustration : This means that the heat of reaction depends only on the initial and final stages and not on the intermediate stages through which the reaction is carried out. Let us consider a reaction in which A gives D. The reaction is brought out in one step and let the heat of reaction be ∆H.
AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics 9
A → D; ∆H
Suppose the same reaction is brought out in three stages as follows :
A → B : ∆H1
B → C : ∆H2
C → D : ∆H3
The net heat of reaction is ∆H1 + ∆H2+ ∆H3.
According to Hess law ∆H = ∆H1 + ∆H2 + ∆H3.

Ex : Consider the formation of CO2. It can be prepared in two ways.
1) Direct method : By heating carbon in excess of 02.
C(s) + O2(g) → CO2(g); ∆H = – 393.5 kJ

2) Indirect method : Carbon can be converted into CO2 in the following two steps.
C(s) + \(\frac{1}{2}\)O2(g) → CO2(g); ∆H1 = – 110.5 kJ
CO (g) + \(\frac{1}{2}\) O2(g) → CO2(g); ∆H2 = – 283.02 kJ
Total ∆H = -393.52 kJ (∆H1 + ∆H2)
The two ∆H values are same.

AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics

Question 25.
Define and exlain the enthalpy of combustion (∆cHθ).
Answer:
Standard enthalpy of combustion is defined as the enthalpy change per mole (or per unit amount) of a substance, when it undergoes combustion and all the reactants and products being in their standard states at the specified temperature.

Cooking gas in cylinders contains mostly butane (C4 H10). During complete combustion of one mole of butane, 2658 kJ of heat is released. We can write the thermochemical reactions ofr this as :
C4 H10 (g) + \(\frac{13}{2}\) O2 (g) → 4CO2 (g) + 5H2O(1);
cHθ = -2658.0 kJ mol-1

Question 26.
Define and explain enthalpy of atomisation (∆aHθ).
Answer:

  • The heat required to dissociate 1 mole of a simple molecule in the gaseous state into its constituent atoms.
  • This is an endothermic process.
    Ex: O2(g) → 2O(g) ∆H = 489.5 KJ
    N2(g) → 2N(g) ∆H = 937.4 KJ

Question 27.
Define and explain the Bond enthalpy (∆bondHθ)
Answer:
Bond Enthalpy : The change in enthalpy when one mole of covalent bonds of a gaseous covalent compond is broken to from gaseous products is called Bond dissociation Enthalpy.
Ex : H2(g) → 2H(g) ∆H = 435.9 KJ / mole

Question 28.
What is the bond enthalpy of C-H bond of CH4 ?
Answer:
The bond dissociation enthalpy for step wise dissociation of CH4 molecule are CH.
CH4(g) → CH3(g) + H(g) ; ∆ H0 = 427.0 kJ
CH3(g) → CH2(g) + H(g) ; ∆ H0 = 418.4kJ
CH2(g) → CH(g) + H(g) ; ∆ H0 = 460.2 kJ
CH(g) → C(g) + H(g) ; ∆ H0 343.1 kJ
The enthalpy for each step is different due to the reason that in each step of dissociation different fragments of CH4 are involved.

But the average bond dissociation enthalpy is 1/4 of the total energy required to break one mole of CH4 into C(g) + 4H (g).
C-H bond dissociation Enthalpy = \(\frac{427.0+418.4+460.2+343.1}{4}\)
= \(\frac{1648.7}{4}\) = 412.2 kJ

AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics

Question 29.
Define heat of solution (∆solHθ) and heat of dilution.
Answer:
i) Enthalpy of solution : Definition : The amount of heat evolved or absorbed when 1 mole of a solute is dissolved completely in large excess of solvent is called Enthalpy of solution.
(Or)
The change of enthalpy when a solution containing 1 mole of a solute is diluted from one concentration to another.
Ex : KCl (s) + aq → KCl (aq), DH = + 19.75 kJ
This is measured per mole of solute. If solvent is water then it is called enthalpy of dilution.

ii) Enthalpy of dilution : The change of enthalpy when a solution containing 1 mole of a solute is diluted from one concentration to another.
Ex : HCl (g) + aq → HCl (aq), ∆H = – 75.4 kJ

Question 30.
Define ionisation enthalpy and electron affinity.
Answer:
Ionisation Enthalpy : The enthalpy change observed by the removal of an electron from the neutral isolated gaseous atom is called ionisation enthalpy.
AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics 10
Electron affinity : The enthalpy evolved by the addition of electron to neutral gaseous atom is called electron affinity.
X(g) + e → X(g) + EA
Observe the following spontaneous reactions
1) H2(g) + 1/2 O2(g) H2O = -285.8 kJ / mole
2) c(graphite) + 2S(g) → CS2(g) ∆ H0 = 91.91 kJ / mole
In reaction (1) ∆H = -ve, in reaction (2) ∆H = +ve
∴ The decrease in Enthalpy may be a condition but not a necessary and sufficient condition for spontaneity of a reaction.

Ionization Energy and Electron Affinity : Ionization energy and electron affinity are defined at absolute zero. At any other temperature, heat capacities for the reactants and the products have to be taken into account. Enthalpies of reactions for
M(g) → M+(g) + e (for ionization)
M(g) + e → M(g) (for electron gain) at temperature, T is
rHΘ(T) = ∆rHΘ(0) + \(\int_0^T \Delta_r C_p^{\Theta} d t\)
The value of Cp for each species in the above reaction is 5/2 R (Cv = 3/2R)
So, ∆rCpΘ = 5/2 R (for ionization)
rCpΘ = -5/2 R (for electron gain)
Therefore,
rHΘ = (Ionization enthalpy)
= E0 (ionization energy) + 5/2 RT
∆rHΘ = (electron gain enthalpy)
= -A (electron affinity) = 5/2 RT

AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics

Question 31.
Explain the spontaneity of a process.
Answer:
Spontaneous reaction :
A process is said to be spontaneous if it occurs on its own without intervention of any external agency of any kind.
All natural processes are spontaneous.
Entropy increases in all spontaneous processes.
Change in entropy “(∆S) = positive” is a condition but is not necessary and sufficient condition for the spontaneous nature of a reaction.
Change in enthalpy “(∆H) = negative” may be a condition but not a necessary and sufficient condition for the spontaneous nature of a reaction.

At this juncture, “Gibbs” introduced another thermodynamic function which involves both enthalpy (H) and entropy (S) functions. This is known as free energy function (G). This is given by the equation.
G = H – TS. .’G1 is now referred as ‘Gibbs energy (or) Gibbs function.

“∆Gsys is negative” for spontaneous reactions (or) processes. Thus the spontaneity of a reaction and the algebraic signs of ∆H, ∆S, ∆G and the magnitude of T can be related as follows :
AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics 11

Question 32.
Is decrease in enthaply a crieterion for spontaneity ? Explain. [A.P. Mar. 16]
Answer:
Change in enthalpy “(∆H) = negative” may be a condition but not a necessary and sufficient condition for the spontaneous nature of a reaction. ;
Observe the following spontaneous reactions

  1. H2(g) + 1/2 O2(g) ⇌ H2O ∆ H0 = -285.8 kJ/mole
  2. C(graphite) + 2S(g) → CS2(g) ∆ H0 = 91.91 kJ/mole

In reaction (1) ∆H = -ve, in reaction (2) ∆H = +ve
∴ The decrease in enthalpy may be a condition but not a necessary and sufficient condition for spontaneity of reaction.

Question 33.
What is entropy ? Explain with examples.
OR
Explain “Entropy” concept.
Answer:
Entropy (S): Entropy is taken as a measure of disorder of molecules (or) randomness of the system. Greater the disorder of molecules in a system, the higher is the entropy. Entropy is a state function. It depends on the temperature, pressure of the state.
Entropy change, ∆S = \(\frac{\mathrm{q}_{\mathrm{rev}}}{\mathrm{T}}\) [qrev = heat absorbed by the system isothermally and reversibly at T]
For a spontaneous process in an isolated system the entropy change is positive.
(∆S = positive)
∆Stotal = ∆Ssystem + ∆Ssurroundings
AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics 12

AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics

Question 34.
Is increase in entropy a criterion for spontaneity ? Explain.
Answer:
Change in entropy “(∆S) = positive” is a condition but is not necessary and sufficient condition for the spontaneous nature of a reaction.
For a spontaneous process change in entropy is positive
∆S = +ve (spontaneous)
∆S = -ve (backward reaction is spontaneous)
∆S = 0 (Equilibrium reaction)
∴ ∆H = -ve, ∆S = +ve, ∆G = -ve then the reaction
Spontaneous at all T values (Reaction is irreversible)

Question 35.
Can All and ∆S discriminate between irreversible and reversible processes ? Explain.
Answer:

  • Change in internal energy (∆U)
  • In exothermic process and internal energy lowered (Reaction is Spontaneous)
  • In case of reversible reaction the reaction favoured in the direction where there is decrease in the internal energy.

Change in entropy “(∆S) = positive” is a condition but is not necessary and sufficient condition for the spontaneous nature of a reaction.
For a spontaneous process change in entropy is positive
∆S = +ve (spontaneous)
∆S = -ve (backward reaction is spontaneous)
∆S = 0 (Equilibrium reaction)
∴ ∆H = -ve, AS = +ve, ∆G = -ve then the reaction
Spontaneous at all T values (Reaction is irreversible)

Question 36.
In which of the following processes entropy increases ?
a) A liquid evaporates to vapour.
b) Temperature of a crystalline solid lowered from 115K to OK
c) CaCO3(s) → CaO(s) + CO2(g)
d) Cl2(g) → 2Cl(g)
Answer:
a) When a liquid evaporates to vapour randomness increases in the vapour state ∆S is +ve. Hence the process is spontaneous.

b) If the temperature of the crystalline solid is lowered to OK it becomes perfectly ordered. No disorder is noticed. Hence such a process is not spontaneous.

c) CaCO3(s) → CaO(s) + CO2(g)
∆S is very slightly positive. If temp is increased ΔS becomes more positive.

d) Cl2(g) → 2Cl(g)
It is an endothermic process there is no considerable increase in entropy. The process is not spontaneous.

AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics

Question 37.
For the oxidation of iron
4Fe(s) + 3O2(g) → 2Fe(2)O3(s), the entropy change is -549.45JK-1 mol-1 at 298K. Though it has negative entropy change the reaction is spontaneous. Why ?
(∆rHΘ = -1648 × 103 J mol-1)
Answer:
One decides the spontaneity of a reaction by considering
∆Stotal i.e< (∆Ssys + ∆Ssurr). For calculating ∆Ssurr, we have to consider the heat absorbed by the surroundings which is equal to -∆rHΘ. At temperature T, entropy change of the surroundings is
∆Ssurr = \(\frac{\Delta_{\mathrm{r}} \mathrm{H}^\theta}{\mathrm{T}}\) (at constant pressure)
= \(-\frac{\left(-1648 \times 10^3 \mathrm{Jmol}^{-1}\right.}{298 \mathrm{~K}}\) = 5530 JK-1 mol-1
Thus, total entropy change for this reaction
rStotal = 5530 JK-1 mol-1 + (-549.4 JK-1 mol-1) = 4980.6 JK-1 mol-1
This shows that the above reaction is spontaneous.

Question 38.
Which formalae in the following are correct ?
a) G = H – TS
b) ∆Gsys = ∆Hsys = ∆Hsys – ∆Ssys
c) ∆Ssurr = \(\frac{\Delta H_{\text {sur }}}{T}=\frac{-\Delta H_{\text {sys }}}{T}\)
d) ∆Ssurr = ∆Hsys + \(\frac{-\Delta \mathrm{H}_{\text {sys }}}{T}\)
e) TASTotal = T∆Hsys – ∆Hsys
Answer:
a) Correct
b) Correct at constant temperature
c) Correct
d) Correct
e) Not correct

Question 39.
Calculate ∆rGθ for conversion of oxygen to ozone ∆rGθ=163 kJ mol-1 O2(g) → O3(g) at 298K. Kp for the \(\frac{3}{2}\) the reaction is 2.43 × 10-29
Answer:
∆G0 = -2.303 RT log Kp
= -2.303 × 8.314 × 298 × log 2.43 × 10-29
= -16300 J/mole
= 163 kJ/mole

AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics

Question 40.
State the second law of thermodynamics and explain it.
Answer:
II law of thermodynamics :
Statements :

  • “It is impossible for a self acting machine unaided by any external agency to convert heat from a body at low temperature to a body at higher temperature”.
  • “All spontaneous processes are thermodynamically irreversible and entropy of the system increases in all spontaneous processes.”
  • “For any spontaneous process taking place in an isolated system, the change in entropy (∆S) is positive.”

Question 41.
State the third law of thermodynamics. What do you understand by it ?
Answer:
Third law of thermodynamics: “The entropy of a pure and perfectly crystalline substance is zero ; at .the absolute zero of a temperature (-273° C)”.
Slim T → 0 = 0
This law imposes a limitation on the value of entropy.
ST = \(\int_0^T \frac{C_p}{T} \cdot d T\)
Thus entropy (S) of a substance at any temperature is calculated if the temperature dependence of ‘Cp‘ is known in evaluating absolute entropy of any substance.
Importance of this law : One of the most important applications of third law is the calculation of entropy changes in chemical reactions.

Question 42.
Explain spontaneity of a process in terms of Gibbs energy.
Answer:
∆G = ∆H – T∆S
∆Gsys = negative for spontaneous reactions (∆G < 0)
∆Gsys = positive for non – spontaneous reactions (∆G > 0)
∆Gsys = zero for equilibrium reactions (∆G = 0)
∆G =-ve, ∆S =+ve, ∆H =-ve are the conditions for the sponaneity of reaction.

Question 43.
The sign and magnitude of Gibbs energy change of a chemical process tells about its spontaneity and useful work that could be extracted from it. Explain.
Answer:
To predict about the spontaneity of the reaction and
To redict the useful work that could be extract is the sign and magnitude of the quantities in Gibb’s equation is useful.
When a reaction is reversible it cannot proceed in both the direction with same case. In general the system would spontaneously change to configuration of lower free energy.
At equilibrium ∆G = 0
0 = ∆G° + RTLnK
rG° = -RTLnK = -2.303 RT log k
We know ∆rG° = ∆rH° – T∆rS° = -RT log k
For endothermic reaction ∆rH° may be large and positive
In such a case K is much less (K < 1)
In exothermic reactions ∆rH° is large and negative (K >1)
∴ For exothermic reaction
a) K must be large
b) ∆rG° is highly negative
c) ∆S° is highly positive
Hence with the sign and magnitude of the factors the yield of the product may be estimated.

AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics

Question 44.
In a process 701 J of heat is absorbed by a system and 394J of work is done by the system. What is the change in internal energy for the process ?
Answer:
q = 701 J
w = -394 J
∆U = q + w
∆U = 701 – 394
= 307 J

Question 45.
The reaction of cyanamide (s). with dioxygen was carried out in a bomb calorimeter and ∆U was found to be – 742, 7kJ mol-1 at 298K. Calculate the enthalpy change for the reaction at 298k.
NH2CN(g) + \(\frac{3}{2}\) O2(g) → N2(g) + H2O(1)
Answer:
NH2CN(g) + \(\frac{3}{2}\) O2(g) → N2(g) + H2O(1)
∆n = 2 – \(\frac{5}{2}\)
= 0.5
∆U = – 742 kJ/mol
∆H = ∆U + ∆n RT
∆H = – 742 + (-0.5) × 8.314 × 10-3 × 298
= -742.7 – 1.238
= – 743.94 kJ/mol4.

Question 46.
Calculate the number of kJ of heat necessary to rise the temperature of 60.0 g of aluminium from 35°C to 55°C. Molar heat capacity of aluminimum is 24 J mol-1 K-1.
Answer:
Given molar heat capacity of Al = 24 J/mole.k
= \(\frac{24}{27}\) J/gm. k
The heat necessary to raise the temp, of 60 g. of Al from 35°C to 55°C
= 60 × \(\frac{24}{27}\) × 55 – 35
= 60 × \(\frac{24}{27}\) × 20 = 1.067 kJ

Question 47.
Calculate the enthalpy change on freezing of 1.0 mol of water at 10.0°C to ice at 10.0°C. ∆fus H = 6.03 kJ mol-1 at 0°C.
Cp [H2 O(1)] = 75.3J mol-1 K-1
Cp [H2O(s)] = 36.8 J mol-1 K-1
Answer:
AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics 13
Hfus = 6.03 kJ
1 mole of water taken to -10°C from 0°C
H = Cp ∆T = 36.8 × -10 = – 3685
= 0.368 kJ.
Total enthalpy = 6.03 + 0.753 – 0.368
= 6.415 kJ/mole.

AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics

Question 48.
Enthalpy of combustion of carbon to CO2 is – 393.5 kJ mol-1. Calculate the heat released upon formation of 35.2 g of CO2 from carbon and dioxygen gas.
Answer:
C(s) + O2(g) → CO2(g) ∆H = -393.5 kJ/mole
44gms of CO2 → -393.5 kJ
35.2 gms of CO2 → ?
\(\frac{35.2 \times 393.5}{44}\) = -314.8 kJ

Question 49.
Enthalpies of formation of CO(g), CO2(g) and N2O4(g) are – 110, -393, 81 and 9.7 kJ mol-1 respectively. Find the value of ∆rH for the reaction :
N2O4(g) + 3CO(g) → N2O(g) + 3CO2(g)
Answer:
N2O4(g) + 3CO(g) → N2O(g) + 3CO2(g)
∆HCO = -110 KJ
∆HCO2 = -393 KJ
∆HN2O = -81 KJ
∆HN2O4 = -9.7 KJ
∆H = sum of enthalpies of products – sum of enthalpies of reactants
= -1017 – (-320.3)
∆H = -696.7kJ

Question 50.
Given N2(g) + 3H2(g) → 2NH3(g); ∆rH° = – 92.4 kJ mol-1 What is the standard enthalpy of formation of NH3 gas ?
Answer:
N2(g) + 3H2(g) → 2NH3(g)
2 moles of NH3 → -92.4kJ
For 1 mole of NH3 → \(\frac{-92.4}{2}\) = -46.2
∆H°(NH3) = -46.2 KJ l mole

AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics

Question 51.
Calculate the standard enthalpy of formation of CH3OH(l) from the following data :
CH3OH(l) + \(\frac{3}{2}\)O2(g) → CO3(g) + 2H3OΘ; ∆rH° = -726 kJ mol-1
C(graphite) + O2(g) → CO2(g); ∆cH° = – 393 kJ mol-1
H2(g) + O2(g) → H2O(l); ∆fH° = -286 kJ -mol-1.
Answer:
Given ∆fH° (CO2) = -393 KJ mol-1
fH° (H2O(1)) = -286 KJ mol-1
AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics 14
rH° = -726 KJ mol-1
rH°= ∆H°total total of the products – ∆H°total total of the reactants
-726 = [-393 + 2 (-286)] – [∆fH° CH3OH].
rH° (CH3OH) = 726 + [-393 + 2(-286)] KJ mol-1
= – 239 KJ mol-1

Question 52.
Calculate the enthalpy change for the process
CCl4(g) → C(g) + 4 Cl(g)
and calculate bond enthalpy of C – Cl in CCl4(g).
vap H°(CCl4) = 30.5 kJ mol-1.
fH°(CCl4) = – 135.5 kJ mol-1.
aH° (C) = 715.0 kJ mol-1, where ∆aH° is enthalpy of atomisation
aH° (Cl2) = 242 kJ mol-1
Answer:
CCl4(1) → CCl4(g) → C(g) + 4Cl(g)
vap H°(CCl4) = 30.5 kJ mol-1.
fH°(CCl4) = – 135.5 kJ mol-1.
∆H°(C) = 715.0 kJ mol-1
rH° (Cl2) = 242 kJ mol-1
AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics 15
rH° = 715 + 484 + 135.5 – 30.5
rH° = 1334.5 – 30.5 KJ mol-1 = 1304 KJ
C – Cl bond enthalpy = \(\frac{1304}{4}\) = 326 KJ mol-1

Question 53.
For an isolated system, ∆U = 0, what will be ∆S ?
Answer:
Given an Isolated system
For isolated system ∆U = 0
∴ ∆H = ∆U + ∆nRT
∆H = ∆nRT
∆H = \(\frac{\Delta \mathrm{H}}{\mathrm{T}}=\frac{\Delta \mathrm{nRT}}{\mathrm{T}}\)
∴ ∆S > 0 ∆S = + ve

AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics

Question 54.
For the reaction at 298 K.
2A + B → C
∆H = 400 kJ mol-1 and ∆S = 0.2 kJ K-1 mol-1
At what temperature will be reaction become spontaneous considering ∆H and ∆S to be constant over the temperature range.
Answer:
Formula ∆S = \(\frac{\Delta \mathrm{H}}{\mathrm{T}}\)
T = \(\frac{\Delta \mathrm{H}}{\Delta \mathrm{S}}=\frac{400}{0.2}\) = 2000K

Question 55.
For the reaction, 2 Cl(g) → Cl2(g), what are the signs of ∆H and ∆S ?
Answer:
2 Cl(g) → Cl2(g)
Here in the above reaction chlorine atoms (reactive)
Converted into stable Cl2 molecule It is an exothermic reaction
∆H = – ve, ∆S = + ve

Question 56.
For the reaction 2 A(g) + B(g) → 2D(g)
∆U° = – 10.5 kJ and ∆S° = – 44.1 JK-1.
Calculate ∆G° for the reaction and predict whether the reaction can occur spontaneously or not.
Answer:
2A(g) + B(g) → 2D(g)
ng = 2 – 3 = -1
∆H = ∆v + ng RT
∆H = – 10.5 + (-1) × \(\frac{8.314}{10^3}\) × 298
– 10.5 – 2.477 = – 12.977 KJ mol-1
∆S = -44.1 JK-1
∆G° = ∆H – T∆S
∆G° = – 12.977 – 298 (-44.1)
= – 12977 + 13141.8 = 1648 J or + 0.1648 J
∴ ∆G° is +ve the reaction is non spontaneous.

Question 57.
The equilibrium constant for a reaction is 10. What will be the value of ∆G° ? R = 8.314 JK-1 mol-1, T = 300 K.
Answer:
Formula ∆G° = – RT lnk
∆G° = – 2.303 RT log k
∆G° = – 2.303 × 8.314 × 300 × log 10
= – 5744 J/mole
∆G° = -5.744 KJ/mole

AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics

Question 58.
State the first law of thermodynamics. Explain its mathematical notation.
Answer:
First law of thermodynamics :
The law of conservation of energy is taken as the first law of thermodynamics.
Statements :

  1. “Energy can neither be created nor destroyed, although it can be transformed from one form to another”.
  2. “It is impossible to construct a perpetual motion machine of first kind (which produce work without consuming energy)”.
  3. “The total energy of the system and surroundings is constant”.

Mathematical form :
Let a system in state “A”, of internal energy EA absorb from the surroundings acertain amount of heat (Q) and undergo a change in its state to B. Let the internal energy in the state B is EB. If W is the work done by the system in the process, the increase in internal energy ∆E of the system is given by the equation, ∆E = EB – EA
AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics 16
From first law, net gain of energy (Q – W) must be equal to ∆E
∆E = (EB – EA) = (Q – W)
(or) Q = ∆E + W For infinitesimally small changes
∴ q = increase in internal energy + energy used in doing work.
Above equation is mathematical form of I law of thermodynamics.

Question 59.
State the 2nd law of thermodynamics in any two ways.
Answer:
II law of thermodynamics :
Statements :

  1. “It is impossible for a self acting machine unaided by any external agency to convert heat from a body at low temperature to a body at higher temperature”.
  2. “All spontaneous processes are thermodynamically irreversible and entropy of the system increases in all spontaneous processes.”
  3. “For any spontaneous pirocess taking place in an isolated system, the change in entropy (∆S) is positive.”

Question 60.
Explain Gibbs energy.
Answer:
Gibbs energy (∆G) : The amount of energy available from a system which can be put to useful work at constant temperature and pressure is called Gibb’s energy. For explaining the criteria for spontaneity of a reaction, Gibb’s introduced thermodynamic function which involved both enthalpy (H) and entropy (S) functions. This is known as free energy function (G). This is given by the equation,
G = H – TS
Where ‘G’ is known as Gibbs energy (or) Gibbs function.
∆G = ∆H – T ∆S
∆Gsys = negative for spontaneous reactions (∆G < 0)
∆Gsys = positive for non – spontaneous reactions (∆G >0)
∆Gsys = zero equilibrium reactions (∆G = 0)

AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics

Question 61.
Explain the sponteneity of a reaction interms of Gibbs energy.
Answer:
To predict about the spontaneity of the reaction and
To redict the useful work that could be extract is the sign and magnitude of the quantities in Gibb’s equation is useful.

When a reaction is reversible it cannot proceed in both the direction with same case. In general the system would spontaneously change to configuration of lower free energy.
At equilibrium ∆G = 0
0 = ∆G° + RTLnK
rG° = -RTLnK = -2.303 RT log k
We know ∆rG° = ∆rH° – T∆rS° = -RT log k
For endothermic reaction ∆rH° may be large and positive
In such a case K is much less (K < 1)
In exothermic reactions ∆rH° is large and negative (K >1)
∴ For exothermic reaction
a) K must be large
b) ∆rG° is highly negative
c) ∆S° is highly positive
Hence with the sign and magnitude of the factors the yield of the product my be estimated.
∆Gsys = negative for spontaneous reactions (∆G < 0)
∆Gsys = positive for non – spontaneous reactions (∆G > 0)
∆Gsys = zero equilibrium reactions (∆G = 0)

Long Answer Questions

Question 1.
State and explain Hess’s Law of constant heat summation. Give example.
Answer:
Hess’s Law : Hess’s law states that the total amount of heat evolved or absorbed in a chemical reaction is always same whether the reaction is carried out in one step (or) in several steps. Illustration : This means that the heat of reaction depends only on the initial and final stages and not on the intermediate stages through which the reaction is carried out. Let us consider a reaction in which A gives D. The reaction is brought out in one step and let the heat of reaction be ∆H.
AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics 17
A → D; ∆H
Suppose the same reaction is brought out in three stages as follows –
A → B : ∆H1
B → C : ∆H2
C → D : ∆H3
The net heat of reaction is ∆H1 +∆H2+∆H3.
According to Hess law ∆H = ∆H1 + ∆H2 + ∆H3.
Ex : Consider the formation of CO2. It can be prepared in two ways.
1) Direct method : By heating carbon in excess of O2.
C(s) + O2(g) → CO2(g); ∆H = – 393.5 kJ
2) Indirect method : Carbon can be converted into C02 in the following two steps.
C(s) + \(\frac{1}{2}\)O2(g) → CO2(g); ∆H1 = – 110.5 kJ
CO (g) + \(\frac{1}{2}\) O2(g) → CO2(g); ∆H2 = – 283.02 kJ
Total ∆H = -393.52 kJ
The two ∆H values are same.

AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics

Question 2.
Explain the experiment to determine the internal energy change of a chemical reaction.
Answer:
Internal energy (U): “The total energy stored in the substance at constant temperature and pressure is called its internal energy (U)”.
It is state function and is an extensive property.
Internal energy change (∆U) is considered as the difference the internal energies of products and reactants.
∆U = (Up – Uf)
In any path taken (Q – W) is equal to ∆U.
Where Q = heat and W = work.

  • The technique of measuring heats of reactions is called calorimetres. The apparatus used is calorimeter.
  • Calorimeter is a vessel properly insulated and contains water as the calorimetric liquids
  • A combustioble substance is burnt in oxygen.
  • Heat evolved during the reaction is measured from the rise in 5 temperature.
  • Here there is no worke done as ∆V = 0
  • By the raise of temperature of water bath and calorimeter the amount of heat evolved can be determined.
    AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics 18
    Internal Enegy All = Q. ∆t × \(\frac{\mathrm{M}}{\mathrm{m}}\)
    Q = heat capacity
    ∆t = t2 – t1
    M = molar mass of substance
    m = mass of the substance

Question 3.
Explain the experiment to determine the enthalpy change of chemical reaction.
Answer:
Enthalpy (H) : “Enthalpy is the amount of heat exchanged by a system with its surroundings at constant pressure and temperature”.
AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics 19
Mathematically, change in enthalpy ∆H = ∆E. P∆V
Enthalpy is state function. Thus, the magnitude of enthalpy change depends only upon the enthalpies in the initial and final states.
∆H = [Hproducts – Hreactants]
Heat change at constant pressure can be measured in a calorimeter at constant pressure ∆H = qp
∴ For Exothermic reaction ∆H = -ve
For Endothermic reaction ∆H = +ve

AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics

Question 4.
Explain the spontaneity of a reaction interms of enthalpy change, entropy change and Gibb’s energy change.
Answer:
Spontaneous reaction :
A process is said to be spontaneous if it occurs on its own without intervention of any external agency of any kind.
All natural processes are spontaneous.
Entropy increases in all spontaneous processes.
Change in entropy “(∆S) = positive” is a condition but is not necessary and sufficient condition for the spontaneous nature of a reaction.
Change in enthalpy “(∆H) = negative” may be a condition but not a necessary and sufficient condition for the spontaneous nature of a reaction.
At this juncture, “Gibbs” introduced another thermodynamic function which involves both enthalpy (H) and entropy (S) functions. This is known as free energy function (G). This is given by the equation, G = H – TS. ‘G’ is now referred as ‘Gibbs energy (or) Gibbs function.
“∆Gsys is negative” for spontaneous reactions (or) processes. Thus the spontaneity of a reaction and the algebraic signs of ∆H, ∆S, ∆G and the magnitude of ‘T’ can be related as follows :
AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics 20
Change in entropy “(∆S) = positive” is a condition but is not necessary and sufficient condition for the spontaneous nature of a reaction.
For a spontaneous process change in entropy is- positive
∆S = +ve (spontaneous)
∆S = -ve (backward reaction is spontaneous)
∆S = 0 (Equilibrium reaction)
∴ ∆H = -ve, ∆S = +ve, ∆G = -ve then the reaction
Spontaneous at all T values (Reaction is irreversible)

Solved Problems

Question 1.
Express the change in internal energy of a system when

  1. No heat is absorbed by the system from the surroundings, but work (w) is done on the system. What type of wall does the system have?
  2. No work is done on the system, but q amount of heat is taken out from the system and given to the surroundings. What type of wall does the system have?
  3. w amount of work is done by the system and q amount of heat is supplied to the system. What type of system would it be?

Solution:

  1. ∆U = wad wall is adiabatic
  2. ∆U = -q, thermally conducting walls
  3. ∆U = -q – w, closed system.

Question 2.
Two litres of an ideal gas at a pressure of 10 atm expand isothermally into a vacuum until its total volume is 10 litres. How much heat is absorbed and how much work is done in the expansion ?
Solution:
We have q = – w = pex (10 – 2) = 0(8) = 0
No work is done; no heat is absorbed.

Question 3.
Consider the same expansion, but this time against a constant external pressure of 1 atm.
Solution:
We have q = – w = pex (8) = (8) = 8 litre-atm

AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics

Question 4.
Consider the same expansion, to a final volume of 10 litres conducted reversibly.
Solution:
We have q = – w = 2.303 × 20 log \(\frac{10}{2}\)
= 32.2 litre-atm

Question 5.
If water vapour is assumed to be a perfect gas, molar enthalpy change for vapourisation of 1 mol of water at 1 bar and 100°C is 41 kJ mol-1. Calculate the internal energy change, when
i) 1 mol of water is vaporised at 1 bar pressure and 100°C.
ii) 1 mol of water is converted into ice.
Solution:
i) The change H2O (l) → H2O (g)
∆H = ∆U + ∆ngRT
or ∆U = ∆H – ∆ngRT, substituting the values, we get
∆U = 41.00 kJ mol-1 -1 × 8.3 J mol-1 K-1× 8.3 J mol-1 K-1 × 3731 |K|
= 41.00 kJ mol-1– 3.096 kJ mol-1
= 37.904 kJ mol-1

ii) The change H2O (l) → H2O (s)
There is negligible change in volume,
So, we can put p∆V = ∆ngRT ≈ 0 in this case,
∆H ≅ ∆U
so, ∆U = 41.00 kJ mol-1

Question 6.
1g of graphite is burnt in a bomb calorimeter in excess of oxygen at 298 K and 1 atmospheric pressure according to the equation
C (graphite) + O2 (g) → CO2 (g)
During the reaction, temperature rises from 298 K to 299 K. If the heat capacity of the bomb calorimeter is 20.7 kJ / K, what is the enthalpy change for the above reaction at 298 K and 1 atm ?
Solution:
Suppose q is the quantity of heat from the reaction mixture and Cv is the heat capacity of the calorimeter, then the quantity of heat absorbed by the calorimeter.
q = Cv × ∆T
Quantity of heat from the reaction will have the same magnitude but opposite sign because the heat lost by the system (reaction mixture) is equal to the heat gained by the calorimeter.
q = -Cv × ∆T = -20.7 kJ/K × (299 – 298)K
= -20.7 kJ
(Here, negative sign indicates the exothermic nature of the reaction)
Thus, ∆U for the combustion of the 1g of graphite = -20.7 kJ
For combustion of 1 mol of graphite,
= \(\frac{12.0 \mathrm{~g} \mathrm{~mol}^{-1} \times(-20.7 \mathrm{~kJ})}{1 \mathrm{~g}}\)
= -2.48 × 102 kJmol-1, Since ∆ng = 0,
∆ H = ∆ U = -2.48 × 102kJ mol-1

AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics

Question 7.
A swimmer coming out from a pool is covered with a film of water weighing about 18g. How much heat must be supplied to evaporate this water at 298 K ? Calculate the internal energy of vaporisation at 100°C.
vapH° for water at 373k = 40.66 kJ mol-1
Solution:
We can represent the process of evaporation as
AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics 21

Question 8.
The combustion of one mole of benzene takes place at 298 K and 1 atm. After combustion, CO2(g) and H2O (l) are produced and 3267.0 kJ of heat is liberated. Calculate the standard enthalpy of formation. ∆f H° of benzene. Standard enthalpies of formation of CO2(g) and H2O(l) are -393.5 kJ mol-1 and -285.83 kJ mol-1respectively.
Solution:
The formation reaction of benezene is given by:
6C (graphite) + 3H2 (g) → C6H6 (l);
fH° = ? ………………… (i)
The enthalpy of combustion of I mol of benzene is :
C6H6 (l) + \(\frac{15}{2}\) O2 (g) → 6CO2 (g) + 3H2O(l);
cH° =-3267 kJ mol-1 …………… (ii)
The enthalpy of formation of 1 mol of CO2 (g):
C(graphite) + O2 (g) → CO2 (g);
f H° = -393.5 kJ mol-1 ………………. (iii)
The enthalpy of formation of 1 mol of H2O(l) is:
H2 (g) + \(\frac{1}{2}\) O2 (g) → H2O(l) .
fH° = -285.83 kJ mol-1 ………………….. (iv)
Multiplying eqn. (iii) by 6 and eqn. (iv) by 3 we get:
6C (graphite) + 6O2 (g) → 6CO2 (g);
fH° = – 2361 kJ mol-1
3H2 (g) + \(\frac{3}{2}\) O2 (g) → 3H2O(l) .
fH° = – 857.49 kJ mol-1
Summing up the above two equations :
6C(graphite) + 3H2 (g) + \(\frac{15}{2}\) O2 (g) → 6CO2 (g) + 3H2O(l)
fH° = -3218.49 kJ mol-1 ………………. (v)
Reversing equation (ii);
6CO2 (g) + 3H2O(l) → C6H6 (l) + \(\frac{15}{2}\) O2
fH° = – 3267.0 kJ mol-1 ………………….. (vi)
adding equations (v) and (vi), we get
6C(graphite) + 3H2 (g) → C6H6 (l);
fH° = 48.51 kJ mol-1

AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics

Question 9.
Predict in which of the following, entropy increases/decreases :
i) A liquid crystallizes into a solid.
ii) Temperature of a crystalline solid is raised from 0 K to 115 K.
iii) 2NaHCO3 (s) → Na2 CO3 (s) + CO2 (g) + H2 O (g)
iv) H2 (g) → 2H(g)
Solution:
i) After freezing, the molecules attain an ordered state and therefore, entropy decreases.

ii) At 0 K, the contituent particles are static and entropy is minimum. If temperature is raised to 115 K, these begin to move and oscillate about their equilibrium positions in the lattice and system becomes more disordered, therefore entropy increases.

iii) Reactant, NaHCO3 is a solid and it has low entropy. Among products there are one solid and two gases. Therefore, the products represent a condition of higher entropy.

iv) Here one molecule gives two atoms i.e., number of particles increases leading to more disordered state. Two moles of H atoms have higher entropy than one mole of dihydrogen molecule.

Question 10.
For oxidation of iron,
4Fe(s) + 3O2 (g) → 2Fe2 O3 (s) entropy change is – 549.4 JK-1 mol-1 at 298 K. Inspite of negative entropy change of this reaction, why is the reaction spontaneous ?
(∆fH° for this reaction is -1648 × 103 J mol-1)
Solution:
One decides the spontaneity of a reaction by considering.
∆Stotal i.e., (∆Ssys + ∆Ssurr)- For calculating ∆Ssurr. we have to consider the heat absorbed by the surroundings which is equal to – ∆r ,H°. At temperature T, entropy change of the surroundings is ∆Ssurr = – \(\frac{\Delta_{\mathrm{r}} \mathrm{H}^{\circ}}{\mathrm{T}}\) (at constant pressure)
= –\(\frac{\left(-1648 \times 10^3 \mathrm{~J} \mathrm{~mol}^{-1}\right)}{298 \mathrm{~K}}\)
= 5530 JK-1 mol-1
Thus, total entropy change for this reaction.
rStotal = 5530 JK-1 mol-1 + (-549.4JK-1 mol-1)
= 4980.6 JK-1 mol-1
This shows that the above reaction is spontaneous.

Question 11.
Calculate ∆rG° for conversion of oxygen to ozone, 3/2 O2(g) → O3(g) at 298 K. if Kp for this conversion is 2.47 × 10-29.
Solution:
We know ∆rG° = – 2.303 RT log Kp and
R = 8.314 JK-1 mol-1
Therefore, ∆rG° = – 2.303 (8.314 J K-1 mol-1) × (298 K) (log 2.47 × 10-29)
= 163000 J mol-1
= 163 kJ mol-1

AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics

Question 12.
Find out the value of equilibrium constant for the following reaction at 298 K.
2NH3 (g) + CO2 (g) ⇌ NH2 CONH2 (aq) + H2O(l)
Standard Gibbs energy change, ∆rG° at the given temperature is – 13.6 kJ mol-1.
Solution:
AP Inter 1st Year Chemistry Study Material Chapter 6 Thermodynamics 22

Question 13.
At 60°C, dinitrogen tetroxide is fifty percent dissociated. Calculate the standard free energy change at this temperature and at one atmosphere.
Solution:
N2O4(g) ⇌ 2NO2(g)
If N2O4 is 50% dissociated, the mole fraction of both the substances is given by
xN2O4 = \(\frac{1-0.5}{1+0.5}\) ; xNO2 = \(\frac{2 \times 0.5}{1+0.5}\)
PN2O4 = \(\frac{0.5}{1.5}\) × 1 atm, PNO2 = \(\frac{1}{1.5}\) × 1 atm.
The equilibrium constant Kp is given by
Kp = \(\frac{\left(\mathrm{p}_{\mathrm{NO}_2}\right)^2}{\mathrm{p}_{\mathrm{N}_2 \mathrm{O}_4}}=\frac{1.5}{(1.5)^2(0.5)}\)
= 1.33 atm.
Since
rG° = – RT l n Kp
rG° = (-8.314 JK-1 mol-1) × (333 K) × (2.303) × (0.1239)
= – 763.8 kJ mol-1

Inter 2nd Year Maths 2A Binomial Theorem Solutions Ex 6(c)

Practicing the Intermediate 2nd Year Maths 2A Textbook Solutions Inter 2nd Year Maths 2A Binomial Theorem Solutions Exercise 6(c) will help students to clear their doubts quickly.

Intermediate 2nd Year Maths 2A Binomial Theorem Solutions Exercise 6(c)

Question 1.
Find an approximate value of the following corrected to 4 decimal places.
(i) \(\sqrt[5]{242}\)
Solution:
Inter 2nd Year Maths 2A Binomial Theorem Solutions Ex 6(c) Q1(i)

(ii) \(\sqrt[7]{127}\)
Solution:
Inter 2nd Year Maths 2A Binomial Theorem Solutions Ex 6(c) Q1(ii)

(iii) \(\sqrt[5]{32.16}\)
Solution:
Inter 2nd Year Maths 2A Binomial Theorem Solutions Ex 6(c) Q1(iii)

(iv) \(\sqrt{199}\)
Solution:
Inter 2nd Year Maths 2A Binomial Theorem Solutions Ex 6(c) Q1(iv)

Inter 2nd Year Maths 2A Binomial Theorem Solutions Ex 6(c)

(v) \(\sqrt[3]{1002}-\sqrt[3]{998}\)
Solution:
Inter 2nd Year Maths 2A Binomial Theorem Solutions Ex 6(c) Q1(v)

(vi) \((1.02)^{3 / 2}-(0.98)^{3 / 2}\)
Solution:
Inter 2nd Year Maths 2A Binomial Theorem Solutions Ex 6(c) Q1(vi)

Question 2.
If |x| is so small that x2 and higher powers of x may be neglected then find the approximate values of the following.
(i) \(\frac{(4+3 x)^{1 / 2}}{(3-2 x)^2}\)
Solution:
Inter 2nd Year Maths 2A Binomial Theorem Solutions Ex 6(c) Q2(i)

(ii) \(\frac{\left(1-\frac{2 x}{3}\right)^{3 / 2}(32+5 x)^{1 / 5}}{(3-x)^3}\)
Solution:
Inter 2nd Year Maths 2A Binomial Theorem Solutions Ex 6(c) Q2(ii)
Inter 2nd Year Maths 2A Binomial Theorem Solutions Ex 6(c) Q2(ii).1

(iii) \(\sqrt{4-x}\left(3-\frac{x}{2}\right)^{-1}\)
Solution:
Inter 2nd Year Maths 2A Binomial Theorem Solutions Ex 6(c) Q2(iii)

(iv) \(\frac{\sqrt{4+x}+\sqrt[3]{8+x}}{(1+2 x)+(1-2 x)^{-1 / 3}}\)
Solution:
Inter 2nd Year Maths 2A Binomial Theorem Solutions Ex 6(c) Q2(iv)
Inter 2nd Year Maths 2A Binomial Theorem Solutions Ex 6(c) Q2(iv).1

Inter 2nd Year Maths 2A Binomial Theorem Solutions Ex 6(c)

(v) \(\frac{(8+3 x)^{2 / 3}}{(2+3 x) \sqrt{4-5 x}}\)
Solution:
Inter 2nd Year Maths 2A Binomial Theorem Solutions Ex 6(c) Q2(v)
Inter 2nd Year Maths 2A Binomial Theorem Solutions Ex 6(c) Q2(v).1

Question 3.
Suppose s and t are positive and t is very small when compared to s. Then find an approximate value of \(\left(\frac{s}{s+t}\right)^{1 / 3}-\left(\frac{s}{s-t}\right)^{1 / 3}\)
Solution:
Since t is very small when compared with s, \(\frac{t}{s}\) is very very small.
Inter 2nd Year Maths 2A Binomial Theorem Solutions Ex 6(c) Q3

Question 4.
Suppose p, q are positive and p is very small when compared to q. Then find an approximate value of \(\left(\frac{q}{q+p}\right)^{1 / 2}+\left(\frac{q}{q-p}\right)^{1 / 2}\)
Solution:
Inter 2nd Year Maths 2A Binomial Theorem Solutions Ex 6(c) Q4
Inter 2nd Year Maths 2A Binomial Theorem Solutions Ex 6(c) Q4.1

Inter 2nd Year Maths 2A Binomial Theorem Solutions Ex 6(c)

Question 5.
By neglecting x4 and higher powers of x, find an approximate value of \(\sqrt[3]{x^2+64}-\sqrt[3]{x^2+27}\)
Solution:
Inter 2nd Year Maths 2A Binomial Theorem Solutions Ex 6(c) Q5

Question 6.
Expand 3√3 in increasing powers of \(\frac{2}{3}\).
Solution:
Inter 2nd Year Maths 2A Binomial Theorem Solutions Ex 6(c) Q6

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Andhra Pradesh BIEAP AP Inter 1st Year Chemistry Study Material 7th Lesson Chemical Equilibrium and Acids-Bases Textbook Questions and Answers.

AP Inter 1st Year Chemistry Study Material 7th Lesson Chemical Equilibrium and Acids-Bases

Very Short Answer Questions

Question 1.
State law of chemical equilibrium.
Answer:
The ratio of product of molar concentrations of the products to the product of molar concentrations of the reactants has a constrant value. This is known as the equilibrium law (or) law of chemical equilibrium.

Question 2.
Can equilibrium be achieved between water and its vapours in an open vessel. Explain.
Answer:
Equilibrium does not acheived between water and its vapour in an open vessel.
But the equilibrium achieved between water and its vapour in a closed vessel.
H2Ol) ⇌ H2O(g)

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 3.
Why the concentrations of pure liquids and pure solids are ignored from equilibrium constant expressions ?
Answer:
The concentrations of pure solids and pure liquids are ignored from equilibrium constant expres-sions because the concentrations of pure solids and pure liquids are taken as unity.

Question 4.
What is homogenous equilibrium ? Write two homogenous reactions.
Answer:
It the physical states of the participating substances are same they the equilibrium is homogeneous equilibrium
e.g.: H2(g) + I2(g) ⇌ 2HI(g))
N2(g) + 3H2(g) ⇌ 2NH3(g)

Question 5.
What is heterogeneous equilibrium? Write two heterogeneous reactions.
Answer:
If the physical states of all (or) some of the participating substances are same then the equilibrium is Heterogeneous equilibrium.
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 1

Question 6.
Write reaction quotient, Q, for each of the following reactions.
a) 3O2(g) ⇌ 2O3(g)
b) 4NH3(g) + 7O2(g) ⇌ 4NO2(g) + 6H2O(g)
Answer:
a) 3O2(g) ⇌ 2O3(g)
Q = \(\frac{\left(\mathrm{O}_3\right)_{\mathrm{t}}^2}{\left(\mathrm{O}_2\right)_{\mathrm{t}}^2}\)

b) 4NH3(g) + 7O2(g) ⇌ 4NO2(g) + 6H2O(s)
Q = \(\frac{\left[\mathrm{NO}_2\right]_t^4\left[\mathrm{H}_2 \mathrm{O}\right]_t^6}{\left[\mathrm{NH}_3\right]_t^4\left[\mathrm{O}_2\right]_t^7}\)

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 7.
Define equilibrium constant.
Answer:
The ratio of product of molar concentrations of products to the product of molar concentrations of reactants is called equilibrium constant (Kc)

Question 8.
The equilibrium constant expression for a gas reaction is Kc = \(\frac{\left[\mathrm{NH}_3\right]^4\left[\mathrm{O}_2\right]^5}{[\mathrm{NO}]^4\left[\mathrm{H}_2 \mathrm{O}\right]^6}\) Write the balanced chemical equation corresponding to this expression.
Answer:
Given Kc = \(\frac{\left[\mathrm{NH}_3\right]^4\left[\mathrm{O}_2\right]^5}{[\mathrm{NO}]^4\left[\mathrm{H}_2 \mathrm{O}\right]^6}\)
Balanced chemical equation is
4NO + 6H2O ⇌ 4NH3 + 5O2

Question 9.
Write the relation between Kp and Kc.
Answer:
Kp = Kc(RT)∆n
∆n = No. of gaseous products – No. of gaseous reactants
R = Universal gas constant

Question 10.
Under what conditions for a reaction Kp and Kc are numerically equal?
Answer:
If ∆n = 0 i.e. No. of gaseous products = No. of gaseous reactants
then Kp = Kc (RT)∆n
Kp = Kc(RT)0
then Kp = Kc

Question 11.
Give two chemical equilibrium reactions for which Kp = Kc
Answer:

  1. H2(g) + I2(g) ⇌ 2 HI(g)
    Here ∆n = 0
    ∴ Kp = Kc
  2. N2(g) + O2(g) ⇌ 2NO(g)
    Here ∆n = 0
    ∴ Kp = Kc

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 12.
Give two chemical equilibrium reactions for which Kp > Kc.
Answer:
1) 2NH3(g) ⇌ N2(g) + 3H2(g)
∆n = np – nR = (3 + 1) – 2 = 2
∆n = 2 (+ve)
∴ Kp > Kc

2) 2SO3(g) ⇌ 2SO2(g) + O2(g)
∆n = np – nR = (2 + 1) – 2 = 1
∆n = 1 (+ve)
∴ Kp > Kc

Question 13.
Give two chemical equilibrium reactions for which Kp < Kc
Answer:
1) N2(g) + 3H2(g) ⇌ 2NH3(g)
∆n = np – nR = 2 – (1 + 3) = -2
∆n = -2 (- ve)
∴ Kp < Kc

2) 2SO2(g) + O2(g) ⇌ 2SO3(g)
∆n = np – nR
= 2 – (1 + 2) = -1
∆n = -1 (-ve)
∴ Kp < Kc

Question 14.
Write the equations for the conversion of Kc to Kp for each of the following reactions.
a) CO(g) + H2O (g) ⇌ CO2 (g) + H2 (g)
b) C3H8(g) + 5O2(g) ⇌ 3CO2 (g) + 4H2O (g)
Answer:
a) CO(g) + H2O(g) ⇌ CO2(g) + H2(g)
Kp = \(\frac{\mathrm{P}_{\mathrm{CO}_2} \times \mathrm{P}_{\mathrm{H}_2}}{\mathrm{P}_{\mathrm{CO}} \times \mathrm{P}_{\mathrm{H}_2 \mathrm{O}}}\)
Kc = \(\frac{\left[\mathrm{CO}_2\right]\left[\mathrm{H}_2\right]}{[\mathrm{CO}]\left[\mathrm{H}_2 \mathrm{O}\right]}\)
Kc to Kp conversion

b) C3H8(g) + 5O2(g) ⇌ 3CO2 (g) + 4H2O (g)
Kp = \(\frac{\mathrm{P}_{\mathrm{CO}_2}^3 \times \mathrm{P}_{\mathrm{H}_2 \mathrm{O}^4}}{\mathrm{P}_{\mathrm{C}_3 \mathrm{H}_8} \times \mathrm{P}_{\mathrm{O}_2}^5}\)
Kc = \(\frac{\left[\mathrm{CO}_2\right]^3\left[\mathrm{H}_2 \mathrm{O}\right]^4}{\left[\mathrm{C}_3 \mathrm{H}_8\right]\left[\mathrm{O}_2\right]^5}\)
Kc to Kp conversion

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 15.
What are the factors which influence the chemical equilibrium?
Answer:
The factors influencing chemical equilibrium are

  1. Concentration of reactants and the products
  2. Temperature of reaction
  3. Pressure of reaction
  4. Inert gas addition etc.

Question 16.
What is the effect of pressure on a gaseous chemical equilibrium?
Answer:
Increase of external pressure of the reaction at equlibrium favours the reaction in the direction in which the volume (or) the No. of molecules decreases.
Decrease of external pressure of reaction at equilibrium favour the reaction in the direction in which the volume (or) no. of molecules increase.

Question 17.
What is the effect of increase in concentration of reactants of a chemical reaction at equilibrium?
Answer:
Increase of concentration of the reactants in the reaction mixture at equilibrium favours the forward reaction (Reactants → products)

Question 18.
Can catalyst disturb the state of equilibrium?
Answer:
Addition of catalyst to the reaction does not effect the equilibrium position. It only speedy up the attainment of the equilibrium.

Question 19.
On which factor, the equilibrium constant value changes?
Answer:
The numerical value of equilibrium constant depends on the standard state of products and reactants chosen.

Question 20.
The equilibrium constants of a reaction at 27° C and at 127°C are 1.6 × 10-3 and 7.6 × 10-2 respectively. Is the reaction exothermic or endothermic.
Answer:
Given equilibrium constants
At 27°C – 1.6 × 10-3
At 127°C -7.6 × 10-2
From the above values as the temperature raised euilibrium constant value increased. Here ∆H = +ve
∴ The reaction is endothermic.

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 21.
What is the effect of temperature on a system at equilibrium?
Answer:

  • Increase in temperature of the reaction mixture at equilibrium favours the endothermic reaction.
  • Decrease in temperature of the reaction mixture at equilibrium favours the exothermic reaction.

Question 22.
For an exothermic reaction, what happens to the equilibrium constant if temperature is raised?
Answer:
For an exothermic reaction when the temperature is raised equilibrium constant value decreased.

Question 23.
What kind of equilibrium constant can be calculated form ∆G° value for a reaction involving only gases?
Answer:
Mathematical expression in the thermodynamic view of equilibrium is
∆G = ∆G° + RTlnQ
At equilibrium ∆G = 0 and Q = K
∴ ∆G = ∆G° + RTlnK = 0
∆G° = -RTlnK
knK = \(\frac{-\Delta G^{\circ}}{R}\)
K = e-∆G°/RT
K = Equilibrium constant using the above equation the sponteinity of reaction can be predicted.

Question 24.
What is a Bronsted base ? Give one example.
Answer:
The substance which accepts a proton from the other substance is called Bronsted base
e.g.: NH3, H2O etc.

Qeustion 25.
What is Lewis acid ? Give one example.
Answer:
A substance which can accept an electron pair to from a co-ordinate covalent bond with donor is called Lewis acid.
e.g. : H+, BF3, SnCl2 etc.

Question 26.
What is meant by ionic product of water ?
Answer:
At a given temp, the product of the concentrations of H+ and OH ions in water is called ionic product
Ionic product Kw = [H+] [OH]
At 25°C Kw = 1.008 × 10-14 mole2/lit2

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 27.
What is the value of Kw? What are its units ?
Answer:
At 25°C Ionic product of water Kw = 1.008 × 10-14 mole2/lit2
Units : mole2/lit2

Question 28.
What is the effect of temperature on ionic product of water ?
Answer:
As the temperature increases ionisation of water increases and hence Ionic product alue (Kw) also increases.

Question 29.
H2O + H2O ⇌ H2O+ + OH
The ionic product of water is 1 × 10-14 at 25°C and 3.0 × 10-14 at 40°C.
Is the above process endothermic or exothermic ?
Answer:
Given reaction
H2O + H2O ⇌ H2O+ + OH
Kwat25°C – 1 × 10-14 mole2/lit2
Kwat40°C – 3 × 10-14 mole2/lit2
The above reaction is endo thermic reaction because Kw value increases with increase of temperature.

Question 30.
All Bronsted bases are Lewis bases. Explain.
Answer:
Bronsted base means proton acceptor and lewis base means electron pair donor.
But according to two theories base provides a lone pair.
∴ All bronsted bases are lewisbases

Question 31.
All Lewis acids are not Bronsted acids. Why ?
Answer:

  • Lewis acid means electron pair acceptor and bronsted acid means proton donar.
  • According to Lewis concept many acids donot have proton.
    So all Lewis acids are not Bronsted bases.

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 32.
What is degree of ionization?
Answer:

  • It is the expression used to find the extent of ionisation of electrolytes.
  • It is devoted by ‘α’.

Question 33.
What is the measure of strength of an acid and base ?
Answer:

  • Strength of an acid can be measured by the acid disociation constant (Ka)
  • Strength of a base can be measured by the base disaciating constant (kb)

Question 34.
Give two examples of salts whose aqueous solutions are basic.
Answer:
Sodium acetate (CH3COONa), Sodium carbonate (Na2CO3) salts aqueous solutions are basic in nature due to anionic hydrolysis.

Question 35.
Give two examples of salts whose aqueous solutions are acidic.
Answer:
Ammonium chloride (NH4Cl), Ammonium sulphate [(NH4)2 SO4] salts aqueous solutions are acidic in nature.

Question 36.
What equation is used for calculating the pH of an acid buffer ?
Answer:
Hendersen’s equation is used to calculate the PH of an acid buffer
PH = PKa + log \(\frac{\text { [Salt] }}{[\text { Acid] }}\)

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 37.
Phosphoric acid (H3PO4) have three ionization constants Ka1, Ka2 and Ka3. Among these ionization constants which has a lower value ? Give reason for it ? Answer:
Phosphoric acid (H3PO4) has
Ka1 – 7.5 × 10-3
Ka2 – 6.2 × 10-8
Ka3 – 4.2 × 10-13
Ka3 has lower value.
Reason :

  • It is more difficult to remove a proton from the negativeion due to electro static forces.
  • It is difficult to remove a proton from \(\mathrm{HPO}_4^{-2}\) anion.

Question 38.
Ice melts slowly at high altitudes. Explain. Why ?
Answer:
Ice melts slowly at high altitudes
Reason :

  • Ice and water are equilibrium only at a particular temperature and pressure conditions.
  • At high attitude temp and pressure conditions are changed.

Short Answer Questions

Qeustion 1.
Write expression for the equilibrium constant, Kc for each of the following reactions.
(i) 2NOCl(g) ⇌ 2NO(g) + Cl2(g)
(ii) 2Cu(NO3)2(s) ⇌ 2CUO(s) + 4NO2(g) + O2(g)
(iii) CH3COOC2H5(aq) + H2O(l) ⇌ CH3COOH(aq) + C2H5OH(aq)
(iv) Fe+3(aq) + 3OH(aq) ⇌ Fe(OH)3(s)
Answer:
i) Given equation
2NOCl(g) ⇌ 2NO(g) + Cl2(g)
Equilibrium constant Kc = \(\frac{\left[\mathrm{NO}^2\left[\mathrm{Cl}_2\right]\right.}{[\mathrm{NoCl}]}\)
ii) Given equation
2Cu(NO3)2(s) ⇌ 2CUO(s) + 4NO2(g) + O2
Kc = (NO2)4 (O2)

iii) CH3COOC2H5(aq) + H2O(l) ⇌ CH3COOH(aq) + C2H5OH(aq)
Kc = \(\frac{\left[\mathrm{CH}_3 \mathrm{COOH}\right]\left[\mathrm{C}_2 \mathrm{H}_5 \mathrm{OH}\right]}{\left[\mathrm{CH}_3 \mathrm{COO} \mathrm{C}_2 \mathrm{H}_5\right]}\)

iv) Fe+3(aq) + 3OH(aq) ⇌ Fe(OH)3(s)
Kc = \(\frac{1}{\left[\mathrm{Fe}^{+3}\right]\left[\mathrm{OH}^{-}\right]^3}\)

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 2.
Derive the relation between Kp and Kc for the equilibrium reaction
N2(g) + 3H2(g) ⇌ 2NH3(g) [Mar. 13]
Answer:
Given equation (Equilibrium reaction)
N2(g) + 3H2(g) ⇌ 2NH3(g)
We can make the following generalisation concerning the composition of equilibrium mixtures : If Kc > 103 products predominate over reactants, i.e., if Kc is very large, the reaction proceeds nearly to completion. Consider the following examples :
a) The reaction of H2 with O2 at 500 K has a very large equilibrium constant Kc = 2.4 × 1047
b) H2(g) + Cl2(g) ⇌ 2Hl(g) at 300 K has Kc = 4.0 × 1031
c) H2(g) + Br2(g) ⇌ 2HBr(g) at 300 K. Kc = 5.4 × 1018
If Kc < 10-3 reactants predominate over products i.e, if Kc is very small the reaction proceeds rarely.

Consider the following examples :
a) The decomposition of H2O into H2 and O2 at 500 K has a very small equilibrium constant.
Kc = 4.1 × 10-48

b) N2(g) + O2(g) ⇌ 2NO(g)
at 298 K has Kc = 4.8 × 10-31
If Kc is in the range of 10-3 to 103 appreciable concentrations of both reactants and products are present. Consider the following examples :
a) For reaction of H2 with I2 to give HI.
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 2
b) Also, gas phase decomposition of N2O4 to NO2 is another reaction with a value of Kc = 4.64 × 10-3 at 25°C which is neither too small nor too large. Hence equilibrium mixtures contain appreciable concentrations of both N2O4 and NO2.
These generarlisations are illustrated.

Question 3.
State law of chemical equilibrium ? What is Kc for the following equilibrium when the equilibrium concentration of each substance is [SO2] = 0.60 M,
[O2] = 0.82 M and [SO3] = 1.90 M
2SO2(g) + O2(g) ⇌ 2SO3(g)
Answer:
The ratio of product of molar concentrations of the products to the product of molar concentrations of the reactants has a constrant value. This is known as the equilibrium law (or) law of chemical equilibrium.
Given equation is
2SO2(g) + O2(g) ⇌ 2SO3(g)
Given [SO2] = 0.60M, [O2] = 0.82M, [SO3] – 1.90 M
Kc = \(\frac{\left[\mathrm{SO}_3\right]^2}{\left[\mathrm{SO}_2\right]^2\left[\mathrm{O}_2\right]}=\frac{(1.9)^2}{(0.6)^2(0.82)}\)
= \(\frac{3.61}{0.36 \times 0.86}=\frac{3.61}{0.2952}\)
= 12.22

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 4.
Why sealed soda water bottle on opening shows the evolution of gas with effervescence ?
Answer:

  • Sealed soda bottle on opening shows evolution of gas with effervescence
  • This is due to difference in solubility of CO2 at different pressures there is equilibrium between molecules of gaseous state and molecules of dissolved in the liquid state.
    CO2(g) ⇌ O2 (Jusolutcoiny)

Question 5.
Explain the significance of
a) a very large value of K,
b) a very small value of K and
c) a value of K of about 1.0
Answer:
a) A very large value of K indicates reaction proceeds almost to complete.
b) A very small value of K indicates reaction hardly proceeds.
c) A value of K about 1.0 indicates that reactants and products are at equilibrium.

Qeustion 6.
Why is it useful to compare Q with K ? What is the situation when
(a) Q = K
(b) Q < K (c) Q > K
Answer:
Q and K are compared to predict the direction of reaction
a) Q = K indicates that the reaction mixture is already at equilibrium
b) Q < K indicates that the reaction proceeds in the direction. c) Q > K indicates that the reaction proceeds in the direction of reactants (Reverse reaction)

Qeustion 7.
For the reaction Cl2(g) + F2(g) 2Cl F (g), Kc = 19.9
What will happen in a mixture originally containing [Cl2] = 0.4 mol L; [F2] = 0.2 mol L-1 and [Cl F] = 7.3 mol L?
Answer:
Given reaction
Cl2(g) + F2(g) ⇌ 2ClF(g), Kc = 19.9
Given [Cl2] = 0.4 mole/lit, [F2] = 0.2 mole/lit
[ClF] = 7.3 moles/lit
∴ \(\frac{[\mathrm{ClF}]^2}{\left[\mathrm{Cl}_2\right]\left[\mathrm{F}_2\right]}=\frac{[7.3]^2}{0.4 \times 0.2}=\frac{53.29}{0.8}\) = 66.6125
Qc > Kc
∴ The reaction proceeds in the backward direction (Reactants side, Reverse Reaction)

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 8.
Predict which of the following reaction will have appreciable concentration of reactants and products :
a) Cl2(g) 2Cl (g), Kc = 5 × 10-39
b) Cl2(g) + 2NO(g) ⇌ 2NOG (g), Kc = 3.7 × 108
c) Cl2(g) + 2NO2 (g) ⇌ 2N02a(g) Kc = 1.8
Answer:
a) Cl2(g) ⇌ 2ClF(g) Kc = 5 × 10-39
b) Cl2(g) +2NO(g) ⇌ 2NOCl(g) Kc = 3.7 × 108
c) Cl2(g) + 2 NO2(g) ⇌ 2NO2Cl(g) Kc = 1.8

  • This equation (c) have appreciable concentrations of reactants and products.
  • ‘Kc‘ value is neither too small nor too large. Hence equilibrium mixtures contain appreciable concentrations of reactants and products.

Question 9.
How to recognise the conditions under which changes in pressure would effect system in equilibrium.
Answer:
Effect of pressure charge : A pressure change obtained by changing the volume can affect the yield of products in case of a gaseous reaction where the total number of moles of gaseous reactants and total number of moles of ga seous products are different. In applying Le Chatelier’s principle to a heterogeneous equilibrium the effect of pressure changes on solids and liquids can be ignored because the volume (and concentration) of a solution/liquid is nearly independent of pressure.
Consider the reaction.
CO(g) + 3H2(g) ⇌ CH4(g) + H2O(g)
Here, 4 mol of gaseous reactants (CO + 3H2) become 2 mol of gaseous products (CH4 + H2O). Suppose equilibrium mixture (for above reaction) kept in a cylinder fitted with a piston at constant temperature is compressed to one half of its original volume. Then, total pressure will be doubled (according to pV = constant). The partial pressure and therefore, concentration of reactants and products have changed and the mixture is no longer at equilibrium. The direction in which the reaction goes to re-establish equilibrium can be predicted by applying the Le Chatelier’s principle.

Since pressure is doubled, the equilibrium now shifts in the forward direction, a direction in which the number of moles of the gas or pressure decreases (we know pressure is proportional to moles of the gas). The can also be understood by using reaction quotient, Qc. Let [CO], [H2], [CH4] and [H2O] be the molar concentrations at equilibrium for methanation reaction. When volume of the reaction mixture is halved, the partial pressure and the concentration are doubled. We obtain the reaction quotient by replacing each equilibrium concentration by double its value.
QC = \(\frac{\left[\mathrm{CH}_{4(\mathrm{~g})}\right]\left[\mathrm{H}_2 \mathrm{O}_{(\mathrm{g})}\right]}{\left[\mathrm{CO}_{(\mathrm{g})}\right]\left[\mathrm{H}_{2_{(\mathrm{g})}}\right]^3}\)
As QC < KC, the reaction proceeds in the forward direction.
In reaction C(s) + CO2(g) ⇌ 2CO(g), when pressure is increased, the reaction goes in the reverse direction because the number of moles of gas increases in the forward direction.

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 10.
What property of a reaction can be used to predict the effect of a change in temperature on the magnitude of an equilibrium constant ?
Answer:

  • Change in temperature leads to the charge in the value of equilibrium constant (Kc)
  • The temperature dependence of the equilibrium constant depends on the sign of ∆H for the reaction.
  • The equilibrium constant for an exothermic reactions (∆H = -ve) decreases as the temperature increases.
  • The equilibrium constant for an endothermic reaction (∆H = +ve) increases as the temperature increases.
  • Temp changes affect the equilibrium constant and rates of reactions.

Question 11.
Does the number of moles of reaction products increase, decrease or remains same when each of the following equilibria is subjected to a decrease in pressure by increasing the volume ?
i) PCl5(g) ⇌ PCl3 (g) + Cl2 (g)
ii) CaO(s) + CO3(g) ⇌ CaCO3 (s)
Answer:
i) PCl5(g) ⇌ PCl3(g) + Cl2(g)
Decrease in pressure by increasing volume leads to the reaction goes in forward direction be cause no. of moles of gas decrease in the forward direction.
∴ No. of moles of reaction products decreases.

ii) CaO(s) + CO2(g) ⇌ CaCO3(s)
Decrease in pressure by increasing volume the reaction does not move in any direction because only one gaseous product present i.e. CO2(g). So nochange in no. of moles of reaction products.

Question 12.
Which of the following reactions will get affected by increasing the pressure ? Also mention whether change will cause the reaction to go into forward or back ward direction.
i) COCl2(g) ⇌ CO(g) + Cl2 (g)
ii) CH4(g) + 2S2(g) ⇌ CS2 (g) + 2H2S(g)
iii) CO2(g) + C(s) ⇌ 2CO(g)
iv) 4NH3(g) + 5O2(g) ⇌ 4NO(g) + 6H2O(g)
Answer:
i) COCl2(g) ⇌ CO(g) + Cl2(g) (affected by pressaure)
Increase in pressure the reaction goes in reverse direction and No. of moles of gas increase in the forward direction. .

ii) CH4(g) + 2S2(g) ⇌ CS2(g) + 2H2(s)
No. of reactants and No. of products are same, pressure has no affect.

iii) CO2(g) + C(s) ⇌ 2 CO(g) (affected by pressure)

iv) 4NH3(g) + 5O2(g) ⇌ 4 NO(g) + 6H2O(g) (affected by pressure)
Increase in pressure, the reaction goes in reverse direction and No. of moles of gas increase in the forward direction.

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 13.
How will an increase in pressure affect each of the following equilibria ? An increase in temperature ?
i) 2NH3(g) ⇌ N2(g) + 3H2(g) ∆ H = 92 kJ
ii) N2(g) + O2(g) ⇌ 2NO(g) ∆ H = 181 kJ
iii) 2O3(g) ⇌ 3O2(g) ∆ H = – 285 kJ
iv) CaO(S) + CO2(g) ⇌ CaCO2(s) ∆ H = – 176 kJ
Answer:
i) 2NH3(g) ⇌ N2 + 3H2(g) ∆H = +92KJ

  • Increase in pressure favours backword reaction
  • Increase in temperatures favours to shift the equilibrium right.

ii) N2(g) + O2(s) ⇌ 2NO(s) ∆H = 181 KJ

  • Pressure has no effect. NP = NR
  • Increase in temperature favour to shift the equilibrium right

jjj) 2O3(g) 3O2(g) ⇌ ∆H = -285 KJ .

  • Increase in pressure favours backward reaction
  • Increase in temperative favours to shift the equilibrium left

iv) CaO(s) + CO2(g) ⇌ CaCO3(s) ∆H = -176 KJ

  • Pressure has no effect.
  • Increase in temperature favour to shieft the equilibrium left.

Question 14.
The dissociation of HI is independent of pressure, while the dissociation of PCl5 depends upon the pressure applied explain.
Answer:
Dissociation of HI
2HI(g) ⇌ H2(g) + I2(g)
Here np = nR
∴ Pressure has no effect on the above reaction.

Dissociation of PCl5
PCl5(g) ⇌ PCl3(g) + Cl2(g)
Here np ≠ nR
∴ This reaction has affected by pressure

Question 15.
Explain the terms :

  1. electrolyte
  2. Non-electrolyte
  3. strong and weak electrolytes
  4. ionic equilibrium

Answer:

  1. Electrolyte: The substances either in molten state (or) in solution conducts electricity and at the same time undergoes chemical decomposition, e.g.: HCl, HNO3 etc.
  2. Non-electrolytes: The substances which donot ionise (or) chemical decomposed when dis-solved in a solvent.
    e.g.: Sugar, urea etc.
  3. Strong Electrolytes: The electrolytes which readily under go chemical decomposition are called strong electrolytes.
    eg.: NaOH, HG, H2SO4 etc.,
    Weak electrolystes : The electrolytes which donot readtions undergo chemical decomposition are called weak electrolytes.
    eg. : CH3COOH, NH4OH etc.,
  4. Ionic equilibrium is a state of equilibrium set up between ionised and unionised molecule. In weak electrolytes, equilibrium is established between ions and the unionized molecules. This type of equilibrium involving ions in aq. solution is called ionic equilibrium.
    AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 3

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 16.
Explain the terms :
i) extent of ionization and on what factors it depends
ii) dissociation
iii) ionization
Answer:
Acids are generally represented as ‘HX’ and bases are represented as ‘BOH’.
Ionization of acids and bases : HX AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 4 H+ + X
BOH AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 4 B+ + OH
This process of splitting of acids and bases when dissolved in water is called ‘ionization’ or ‘dissociation1. The extent of ionization of acids and bases depends on the polarity of the molecule. For strong acids and strong bases extent of ionization is large. The extent of ionization is less for weak acids or bases.
This extent of ionization depends on the concentration of the solution of the substance. Process of ionization is an equilibrium process.
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 5
Hence, the equilibrium constants are expressed as Ka = \(\frac{\left[\mathrm{H}^{+}\right]\left[\mathrm{X}^{-}\right]}{[\mathrm{HX}]}\)
Ka = acid ionization constant (or) acid dissociation constant.
Kb = \(\frac{\left[\mathrm{B}^{+}\right]\left[\mathrm{OH}^{-}\right]}{[\mathrm{BOH}]}\)
Kb = base ionization constant (or) base dissociation constant.

Question 17.
Explain the Arrhenins concept of acids and bases.
Answer:
According to Arrhenius
Acid : Acids are the substances that dissociates in water to gave H+ ions.
e.g. : HCl, H2SO4 etc..
H x(aq) → H+(aq) + x(aq)
Base : Bases are the substances that produce hydrorylions (OH) ion dissociation in water
e.g. : NaOH, KOH MOH,
MOH(aq) → M+(aq) + OH(aq)

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 18.
What is a, conjugate acid – base pair ? Illustrate with examples. [A.P. Mar. 16]
Answer:
Conjugate acid – base pair:
A pair of a Bronsted acid and a base that differs by one proton (H+) is known as conjugate acid – base pair.
Consider the following reaction :
CH3COOH + H2O ⇌ H3O + CH3COO
It involves two conjugate pairs. The acid – base pairs are :
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 6

Question 19.
Acetic acid is a weak acid. List, in order of descending concentration, all of the ionic and molecular species present in 1 M aqueous solution of acetic acid.
Answer:
[H2O] > [CH3COOH] > [H3O+] [CH3COO] > [OH]

Question 20.
Show by suitable equations that each of the following species can act as a Brostead acid.
a) H3O+
b) HCl
c) NH3
d) HSO4
Answer:
a) H3O+ → H2O + H+
Proton donar Hence a Bronsted acid

b) HCl → H+ + Cl
Proton donar Hence a Bronsted acid

c) NH3 is bronsted base it’s not an acid
(Sometimes NH3 → NH2 + H+ as acid]

d) HSO4 → H+ + SO4-2
proton donar Hence Bronsted acid.

Question 21.
Show by suitable equations that each of the following spices can act as a Bronsted base.
a) H3O
b) OH
c) C2H5OH
d) HPO4-2
Answer:
a) H2O + H+ → H3O+
Proton acceptor Hence a Bronsted base
b) OH + H+ → H2O
Proton acceptor Hence a Bronsted base
c) C2H5OH is a proton donar
So it is not a bronsted base it is acid
d) HPO4-2 + H+ → H2PO4
proton acceptor Hence it is Bronsted base.

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 22.
The species H2O, HCO4, HSO4 and NH3 can act both as Bronsted acids and bases. Give the corresponding conjugate acid and base for each of them.
Answer:
The answer is given in the following Table :
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 7

Question 23.
Write equation that shows H2PO4 acting both as an acid and as a base.
Answer:
As an acid :
H2Po4 → H+ + HP2PO4-2
Proton donar
As a base:
H2Po4 + H+ → H3PO4
Proton acceptor

Question 24.
Write the conjugate acid and conjugate base of each of the following :
a) OH
b) H2O
c) HCO3
d) H2O2
Answer:
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 8

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 25.
Identity and label the Bronstead acid and its conjugate base, the Bronsted base and its conjugate acid in each of the following equations.
a) H2SO4 + Cl → HCl + HSO4
b) H2S + NH2 → HS + NH3
c) CN + H2O → HCN + OH
d) O-2 + H2O → 2OH
Answer:
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 9

Question 26.
Classify the species AlCl3, NH3, Mg+2 and H2O into Lewis acids and Lewis bases and justify your answer ?
Answer:

  • AlCl3, Mg+2 are lewis acids because These are electron pair acceptory
  • NH3, H2O are lewis bases becuse. These are electron pair donors.

Question 27.
What are the strengths of conjugate bases of a strong acid and a weak acid ?
Answer:
A strong acid has a weak conjugate base
Eg : HCl, Cl
A weak acid has a strong conjugate base
Eg : CH3COOH, CH3COO

Question 28.
What are the strength of conjugates acids of a strong base and weak base ?
Answer:
Strong base has a weak conjugate acid
Eg : HS, H2S
A weak base has a stong conjugatge acid
Eg : ClO4, HClO4

Question 29.
Define ionic product of water. What is its value at room temperature ?
Answer:
At a given temp, the product of the concentrations of H+ and OH ions in water is called ionic product
Ionic product Kw = [H] [OH]
At 25°C Kw = 1.008 × 10-14 mole2/lit2
At 25°C Ionic product of water Kw = 1.008 × 10-14 mole2/lit2
Units: mole2/lit2

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 30.
Define pH. pH cannot be calculated directly from the molar concentration of a weak acid or weak base. Why ? Derive an equation for the pH of a weak acid.
Answer:
The negative logarithm on value to the 10 of the hydrogen ion concentration expressed in moles/lit, in a solution is know as pH. of the solution.
pH : – log (H+) = – log10 \(\frac{1}{\left(\mathrm{H}^{+}\right)}\)
In case of a weak acid (or) weak base extent of ionisation is low so pH cannot be calculated directly from their molar concentrations.

Derivation of pH of a weak acid :
A general step wise approch can be adopted to evaluate the pH of the weak acid as follows.

Step 1. The species present before dissociation are identified as Bronsted – Lowry acids/ bases.
Step 2. Balanced equations for all possibel reactions i.e., with a species acting both as acid as well as base are written.
Step 3. The reaction with the higher Ka is identified as the primary reaction whilst the other is a subsidiary reaction.
Step 4. Enlist in a tabular form the following values for each of the species in the primary reaction.
a) Initial concentration. C.
b) Change in concentration on proceeding to equilibrium in terms of a, degree of ionization.
Step 5. Substitute equilibrium concentrations into equilibrium constrant equation for principal reaction and solve for α.
Step 6. Calculate the concentration of species in principal raction.
Step 7. Calculate pH = – log [H3O+]
The above mentioned methodology has been elucidated in the following examples.

Question 31.
Write equations to show the step wise ionization of the polyprotic acids H2SO4 and H2PO4.
Answer:
Step wise Ionisation H2SO4
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 10

Question 32.
Explain how acid strength changes among

  1. the hydrides of the group elements and
  2. the hydrides in the same row of the periodic table.

Answer:

  1. The hydrides of the group elements acid strength decreases from top to bottom.
  2. The hydrides in the same row of Periodic table i.e., in a period the acid strength increases from left to right.

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 33.
Justify the statement that water behaves like an acid and also like base on the basis of protonic concept.
Answer:

  • Water has the ability to acts as an acid and as well as a base i.e., it behaves as an amphoteric substance.
  • The following are the supporting reactions for the fact that water behaves as an amphoteric substance.
  • As an acid :
    H2O(l) + NH3(aq) ⇌ OH(aq) + NH4+(aq)
    H2O is proton do nor i.e., an acid according to Bronsted sense
  • As a base:
    H2O(l) + H2S(aq) ⇌ H3O+(aq) + HS(aq)
    H2O is proton acceptor i.e., a base according to Bronsted Sense

Question 34.
What is common ion effect ? Illustrate.
Answer:
Common ion effect : The decrease in the ionisation (dissociation) of a weak electrolyte, by the addition of a strong electrolyte having an ion common with the weak electrolyte, is known as common ion effect.
Example : The dissociation of NH4OH is diminished by the addition of NH4Cl due to the common ion, NH4+ ion
NH4OH ⇌ NH4 + OH
NH4OH AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 11 NH4+ + OH (addition of NH4Cl)
Importance of this effect In qualitative analysis :

  1. This principle is used in the systematic qualitative analysis of cations.
  2. The concentration of S2- in II group or OH ion in III group of chemical analysis is controlled by HCl and NH4OH on the basis of common ion effect.
    H+ ion is common to H2S (II group)
    NH4+ ion is common to NH4OH (IV group) .
  3. The common ion effect principle is also used in controlling the H+ ion concentration in buffer solutions.
  4. It is also used in the purification of common salts by passing dry HCl gas into impure salt solution.

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 35.
Define solubility product ? Write solubility product expressions for the following :
(i) Ag2Cr2O7
ii) Zr3(PO4)4
Answer:
Solubility product (KSP) : The product of the concentrations of the cation and the anion in a saturated solution of a salt at room temperature is called solubility product (Ksp).
Ksp = [Mn+][An-]
Significance of solubility product in chemical analysis : In any solution of a salt, if the product of [Mn+] and [An-] is greater than the solubility product (KSP) of the salt, then only the salt gets precipitated. For example,
i) If the product of [Cu 2+] & [S-2] in the II group (in presence of acid) of chemical analysis exceeds the value of KSP of CuS, then only CuS will be precipitated in the group. Similar argument applies to other metal ions in the group namely (Cd2+, Bi3+, Hg2+, Sb2+, etc).

ii) If the product of [Zn2+] and [S2-] in ammonical medium in IV group of chemical analysis exceeds the value of KSP of ZnS, it will be precipitated. Similar arguments apply to Mn2+, Ni2+ etc.
a) Solubility product expression for Ag2Cr2O7
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 12

Question 36.
Give the classification of salts. What type of salts undergo hydrolysis ?
Answer:
The following are the types of salts : –

  1. Salts of a strong acid and a strong base. Ex : NaCl
  2. Salts of a strong acid and a weak base. Ex : NH4Cl
  3. Salts of a weak acid and a strong base. Ex : CH3COONa
  4. Salts of a weak acid and a weak base. Ex : CH3COONH4
    Types 2, 3, 4 under go hydrolysis reactions.

Question 37.
What must be true of the value of ∆G° for a reaction if
a) K > 1
b) K = 1
c) K < 1
Answer:
a) If K > 1 then ∆ G° < 0
b) If K = 1 then ∆ G° = 0
c) If K < 1 then ∆ G° > 0

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 38.
Aqueous solution of NH4Cl is acidic. Explain.
Answer:
Given salt NH4Cl is obtained by the reaction of strong acid (HCl) and weak base (NH4OH).
The ionisation of NH4Cl as follows
\(\mathrm{NH}_4^{+}\) + Cl ⇌ NH4Cl
The above salt under goes cationic hydrolysis
\(\mathrm{NH}_4^{+}\) + H2O ⇌ NH4OH + H+
∴ The equation solution of salt is acidic due to greater H+ ion concentration than OH ion concentration.
∴ PH < 7

Question 39.
Aqueous solution of CH3COONa is basic explain.
Answer:
Given salt CH3COONa is obtained by the reaction of weak acid (CH3COOH) and strong base (NaOH).
The ionisation CH3COONa as follows
CH3COONa ⇌ CH3COO + Na+
The above salt under go anionic hydrolysis
CH3COO + H2O ⇌ CH3COOH + OH
∴ The equation solution of salt is basic due to greater OH ion concentration than H+ ion concentration.
∴ PH > 7

Question 40.
Give reason that acetic acid is less acidic in sodium acetate solution than in sodium chloride solution.
Answer:
Acetic acid in sodium acetate solution is less acidic.
Reasons:

  • Because of common ion effect ionisation process is decreased.
  • Sodium acetate is obtained by the reaction of weak acid and strong base:
    Acetic acid in NaG solution is more acidic.
    Here no common ion effect observed.
    NaCl is the salt obtained by strong acid and strong base.

Question 41.
AgCl is less soluble in AgNO3 solution than in pure water explain.
Answer:

  • AgCl is less soluble in AgNO3 because of common ion effect solubility is decreased.
  • AgCl is more soluble in pure water because common ion effect not observed here.

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 42.
Predict whether the following reaction will proceed from left to right to any measurable extent:
CH3COOH (aq) + Cl(aq) →
Answer:
Given CH3COOH(aq) + \(\mathrm{Cl}_{(\mathrm{aq})}^{-}\) →
The reaction does not proceed from left to right because Cl ion is obtained from strong acid do not under go hydrolysis and CH3COOH has low extent of ionisation.

Question 43.
Aqueous solution of H2S contains H2S, HS, S-2, H3O+, OH and H2O in varying concentrations. Which of these species can act only as a base ? Which can act only as an acid ? Which can act both an acid and as a base ?
Answer:

  • H23S, H3O+ acts as only acids
  • HS, OH, H2O acts as acids as well as bases
  • S-2 acts as only base

Long Answer Questions

Question 1.
What are equilibrium processes ? Explain equilibrium in Physical and Chemical pro-cesses with examples.
Answer:
Equilibrium State (or) Equilibrium Process : The state at which the velocity of forward reaction becomes equal to the velocity of backward reaction or reverse reaction is called ‘equilibrium state’. Consider a reversible reaction.
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 13
A + B ⇌ C + D taking place in a closed vessel. At the beginning we have only the reactants A and B. Their concentrations are maximum. As the reaction proceeds the reactants A and B change into the products C and D. The concentrations of the products increase gradually.
The rate of forward reaction diminishes while the reverse reaction sets in and proceeds with increasing speed. A state is soon reached where the speeds of forward and backward reactions become equal. If the rate of forward reaction Vf = The rate of reverse reaction (Vb).

The system is said to have attained a state of equilibrium. Once equilibrium is reached there is no further change in the composition of the system. The system appears to be stand still although it is dynamic. The products are formed by forward reaction just as fast as they change back into reactants by the reverse reaction.

The state at which the rate of forward reaction is equal to the rate of the reverse reaction in a reversible reaction is known as the equilibrium state or chemical equilibrium.

Examples :
i) Equilibrium physical processes :
a) Equilibrium between two phases (i.e.,) physical transformation processes.
Solid ⇌ Liquid (melting or fusion)
Solid ⇌ Vapour (sublimation)
b) Equilibrium between two different allotropic forms of same substance
α – Sulphur ⇌ β – Sulphur

ii) Equilibrium chemical processes :
a) N2 (g) + 3H2 (g) ⇌ 2NH3 (g)
b) 2SO2 (g) + O2 (g) ⇌ 2SO3 (g)
c) H2(g) + I2(g) ⇌ 2HI(g)

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 2.
What is meant by dynamic equilibrium ? Explain with suitable examples.
Answer:
Dynamic equilibrium : The forward and reverse reactions of a reversible reaction continue to take place with equal rates simultaneously at the equilibrium stage also. Hence the equilibrium is called Dynamic equilibrium.

Explanation : In order to understand the dynamic nature of the reaction, synthesis of ammonia is carried with exactly the same starting conditions but using D2 (deuterium) in place of H2. The reaction mixtures starting either with H2 or D2 reach equilibrium with the same composition, except that D2 and ND3 are present instead of H2 and NH3. After equilibrium is attained, these two mixtures (H2, N2, NH3 and D2, N2, ND3) are mixed together and left for a while. Later when this mixture analysed, it is found that the concentration of ammonia is just the same as before.

However, when this mixture is analysed by a moss spectrometer, it is found that ammonia and all deuterium containing forms of ammonia (NH3, NH2D, NHD2 and ND3) and dihydrogen and its deutrated forms (H2, HD and D2) are present. Thus one can conclude that scrambling of H and D atoms in the molecules must result from a continuation of the forward and reverse reactions in the mixture. If the reaction had simply stopped when they reached equilibrium, then there would have been no mixing of isotopes in this way.

Use of isotope (deuterium) in the formation of ammonia clearly indicates that chemical reaction react a state of dynamic equilibrium in which the rates of forward and reverse reactions are equal and there is no net change in composition.

Question 3.
Give the General characteristics of equilibria involving physical processes ?
Answer:
Characteristics of general equilibria :

  1. The rate of the forward reaction is equal to the rate of the reverse reaction (uf = ub).
  2. The concentrations of the reactants and the products remain unchanged with time.
  3. The observable properties such as pressure, concentration, density, colour also remain unchanged with time.
  4. The equilibrium is dynamic in nature. Both the forward reaction and the reverse reaction continue to take place simultaneously with equal rates.
  5. A catalyst does not alter the state of equilibrium or the composition of the chemical substances of the reaction at the equilibrium. It only speeds up the attainment of the equilibrium.
  6. Chemical equilibrium can be established from either side of the reversible reaction.

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 4.
What are the important features of equilibrium constant ? Discuss any two applications of equilibrium constant.
Answer:
The ratio of product of molar concentrations of products to the product of molar concentrations of reactants is called equilibrium constant (Kc)
Ex : aA + bB ⇌ cC + dD
KC = \(\frac{[C]^c[D]^d}{[A]^a[B]^b}\)
KC for reverse reaction is the inverse of the KC for the reaction in the forward direction.
Concentration equilibrium constant (KC) and pressure equilibrium constant (Kp) are related as follows
Kp = KC(RT)∆n
∆n = np – nR
While writing KC expression pure liquids, pure solids concentration are ignored Applications :

a) In predicting the extent of reaction :
The numerical value of the equilibrium constant for a reaction indicates the extent of the reac-tion. But it is important to note that an equilibrium constant does not give any information about the rate at which the equilibrium is reached the Magnitude of ke or Kp is directly proportional to the concentrations of products (as these appear in the numerator of equilibrium constant expression) and inversely proportional to the concentrations of the reactants (these appear in the denominator). This implies that a high value of K is suggestive of a high concentration of products and vice- versa.

We can make the following generalisation concerning the composition of equilibrium mixtures : If Kc > 103 products predominate over reactants, i.e., if Kc is very large, the reaction proceeds nearly to completion. Consider the following examples :

a) The reaction of H2 with O2 at 500 K has a very large equilibrium constant Kc = 2.4 × 1047
b) H2(g) + Cl2(g) ⇌ 2HCl(g) at 300 K has Kc = 4.0 × 1031
c) H2(g) + Br2(g) ⇌ 2HBr(g) at 300 K. Kc = 5.4 × 1018
If Kc = < 10-3 reactants predominate over products i.e, if Kc is very small the reaction proceeds rarely. Consider the following examples : –
a) The decomposition of H2O into H2 and O2 at 500 K has a very small equilibrium constant.
Kc = 4.1 × 10-48
b) N2(g) + O2(g) ⇌ 2NO(g)
at 298 K has Kc = 4.8 × 10-31
If Kc is in the range of 10-3 to 103 appreciable concentrations of both reactants and products are present. Consider the following examples :
a) For reaction of H2 with I2 to give HI.
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 14
b) Also, gas phase decomposition of N2O4 to NO2 is another reaction with a value of Kc = 4.64 × 10-3 at 25°C which is neither too small nor too large. Hence equilibrium mixtures contain appreciable concentrations of both N2O4 and NO2.
These generarlisations are illustrated.
b) In predicting the direction of reaction : A and K are compared to predict the direction of reaction
a) Q = K indicates that the reaction mixture is already at equilibrium
b) Q < K indicates that the reaction proceeds in the direction. c) Q > K indicates that the reaction proceeds in the direction of reactants (Reverse reaction)

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 5.
What is Lechatlier’s principle ? Discuss briefly the factors which can influence the equilibrium.
Answer:
Le Chatelier’s Principle – Statement: “If a system at equilibrium is subjected to the change of pressure, temperature (or) concentration, the system is shifted in such a way as to nullify the effect of change”.
Explanation:
1) Concentration: Increase of reactant concentrations pushes the equilibrium state to the products side and increase of the reactions concentrations pushes to the equilibrium to the reactants side. For example in the chemical equilibrium
H2(g) + I2(g) ⇌ 2HI(g)
Increase of H2 or I2 concentrations pushes the equilibrium in favours of HI and similarly the increase of HI concentrations pushes the equilibrium in favour of H2 and I2.

2) Pressure : Pressure will have no effect on the equilibrium reactions in which there is no change in the number of moles of the reactants and the products.
For example pressure will have no effect on
H2(g) + I2(g) ⇌ 2HI(g) equilibrium
(or) N2 (g) + O2 (g) ⇌ 2NO (g) equilibrium
But in the case of reactions in which there is a change in the number of moles of the reactants and the products increase of pressure pushes the reaction equilibrium in the direction in which there is a decrease in the number of moles.
For example in the reaction PCl5 (g) ⇌ PCl3 (g) + Cl2 (g)
Increase of pressure favours the backward reaction PCl3 (g) + Cl2(g) → PCl5 (g) and the decrease of pressure favour the forward reaction PCl5 (g) → PCl3 (g) + Cl2 (g)

3) Temperature : Increase of temperature favour the endothermic reaction and decrease of temperature favours exothermic reactions.
For example .
N2 (g) + O2 (g) ⇌ 2N0(g) – 180.8 kJ
Increase of temperature favours The forward reaction
N2 (g) + O2 (g) → 2NO (g) (since endothermic)
decrease of temperature favour the dissociation of NO into N2 and O2 since it is exothermic.
The backward reaction
2NO (g) → N2 (g) + O2 (g) (since exothermic)

Question 6.
Discuss the application of Lechaterlier’s principle for the industrial synthesis of Ammonia and sulphur trioxide.
Answer:
Applications of Le Chatelier’s principle to synthesis of Ammonia by Haber’s process :
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 15
Nitrogen and Hydrogen combine to form ammonia. The formation of ammonia is reversible and exothermic reaction. It is accompanied by decrease in volume.

Effect of Pressure: 1 volume of N2 combines with 3 volumes of H2 to form 2 volumes of NH3. There is decrease in volume in the forward reaction (4 volumes to 2 volumes). According to Le Chatelier’s principle increase of pressure favours the reaction where there is decrease in volume. So higher the pressure, greater the yield of ammonia. In practice 200 atmospheres are used in the manufacture of ammonia by Haber’s process.

Low pressures favour the reverse reaction i.e., decomposition of NH3 already formed.

Effect of Temperature : The formation of ammonia (forward reaction) is exothermic reaction
N2 + 3H2 → 2 NH3; ∆H = – 92 kJ.
Low temperatures favour the forward reaction. But at low temperatures the reaction is too slow. Therefore an optimum temperature (725 K- 775 K) is chosen in Haber’s process. To speed up the reaction, a catalyst, finely divided iron is used. To increase the activity of the catalyst molybdenum or a mixture of oxides of K after Al is used as promoter.

The reverse reaction (i.e.,) decomposition of NH3 is an endothermic reaction. High temperatures favour the decomposition of NH3. Therefore high temperatures are avoided in Haber’s process. Thus the optimum conditions are
Pressure : 200 atm
Temperature : 725-775 K
Catalyst : Fe (Powered)
Promoter : Mo (or) (K2O + Al2O3)
Application of Le Chatelier’s principle to the synthesis of SO3 :
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 16
The formation of SO3 is reversible and exothermic reaction. It is accompanied by decrease in volume.
1. Effect of Pressure : 2 volumes of SO2 and one volume of O2 combine to give 2 volumes of SO3. According to Le Chatelier’s Principle increase of pressure favours the reaction where there is decrease in volume. The formation of SO3 is accompanied by decrease in volume (3 volumes to 2 volumes). Higher the pressure greater is the yield of S03. But in contact process high pressures are not used because towers used in the manufacture are corroded by the acid at these high pressures.

Low pressures favour the decomposition of SO3 as there is increase in the volume of (2 volumes to 3 volumes). Therefore optimum pressures are used (1.5 to 1.7 atmosphere).

2. Effect of Temperature : The formation of SO3 is exothermic. 189 Id of heat is evolved. High temperatures favour the reverse reaction which is endothermic and do not favour the forward reaction which is exothermic. Low temperatures are favourable for the formation of SO3. At low temperature the reaction is too slow. Therefore an optimum temperature 673 K is used. To speed up the reaction V205 is used as catalyst.
Thus the optimum conditions are
Pressure : 1.5 to 1.7 atm
Temperature : 673 K
Catalyst : V2O5 (or) Platinised asbestos

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 7.
Dihydrogen gas is obtained from natural gas by partial oxidation with steam as per the following endothermic reaction. .
CH4(g) + H2O(g) ⇌ CO(g) + 3H2 (g)
a) Write an expression for Kp for the above reaction.
b) How will the values of Kp and composition of equilibrium mixture be affected by

  1. increasing the pressure
  2. increasing the temperature
  3. using a catalyst ?

Solution:
a) Given equation
CH4(g) + H2O(g) ⇌ CO(g) + 3H2(g)
Kp = \(\frac{\mathrm{P}_{\mathrm{CO}} \times \mathrm{P}_{\mathrm{H}_2}^3}{\mathrm{P}_{\mathrm{CH}_4} \times \mathrm{P}_{\mathrm{H}_2 \mathrm{O}}}\)

b)

  1. Increase of pressure the reaction proceeds in reverse direction because the no. of moles of gas increases in the forward direction.
    Increase pressure Kp value increases.
  2. Increase of temperature of the reaction mixture at equilibrium favours the reaction in the direction in which heat is absorbed.
    Increase of temperature Kp value increases.
  3. A catalyst does not disturbs (or) change the chemical reaction it just speed up the attaining of equilibrium.

Question 8.
Describe the effect of :
a) addition of H2
b) addition of CH3OH
c) removal of CO
d) removal of CH3OH on the equilibrium of the reaction
2H2(g) + CO(g) ⇌ CH3OH(g)
Solution:
a) Addition of H2 (Reactant):
Increasing the concentration of reactants in the reaction mixture at equilibrium favours the forward reaction.

b) Addition of CH3OH (Product):
Increasing the concentration of products in the reaction mixture at equilibrium favours the reverse reaction.

c) Removal of CO (Reactant):
Decreasing the concentration of reactants in the reaction mixture at equilibrium favours the reverse reaction.

d) Removal of CH3OH (Product):
Decreasing the concentration of products in the reaction mixture at equilibrium favours the forward reaction.

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 9.
At 473K, equilibrium constant Kc for the decomposition of phosphorous pentachloride, PCl5, is 8.3 × 10-3. If the decomposition is depicted as :
pCl5 (g) pCl3 (g) + Cl2 (g) ∆H = 124.0 kJ mol-1.
a) Write an expression of Kc for the reaction
b) What is the value of Kc for the reverse reaction at the same temperature ?
c) What would be effect on Kc if
(i) more PCl5 is added
(ii) pressure is increased
(iii) the temperature is increased.
Solution:
Given
PCl5(g) ⇌ PCl3(g) + Cl2(g) ∆H = 124.0 kJ/mole
KC = 8.3 × 10-3
a) KC = \(\frac{\left[\mathrm{PCl}_3\right]\left[\mathrm{Cl}_2\right]}{\left[\mathrm{PCl}_5\right]}\)

b) Given KC = 8.3 × 10-3
We have to calculate KC for reverse reaction
i.e., PCl3(g) + Cl2(g) ⇌ PCl5(g)
∴ Kc for the reverse reaction = \(\frac{1}{8.3 \times 10^{-3}}\)
= 0.12048 × 103
= 120.48

c) i) More PCl5 is added KC value is decreased.
ii) By the increase of pressure KC value is increased.
iii) Given reaction is Endothermic so increase of temperature increases KC value.

Question 10.
Explain the concept of Bronsted acids and Bronsted bases. Illustrate the answer with suitable examples.
Answer:
According to Bronsted theory o substance which can donate a proton to the other substance is known as acid. A substance which accept a proton from other substance is a base.
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 17
Here HCl donates a proton to water and behaves as Bronsted Lowry acid. Similarly NH3 accepts a proton from H2O and acts as Bronsted Lowry base.

Above reaction is a reversible reaction so that H3O+ can donate proton to act as acid. Cl can accept a proton to act as base. Thus each acid base reaction equilibrium involves two acids and two bases. Each pair differs by a proton, such acid base pair is called Conjugate acid base pain.
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 18
According to this theory strengths of acids and bases can be explained.

An acid which show great tendency to donate protons is a strong acid and an acid which shows less tendency to donate proton is a weak acid.

A base which shows great tendency to accept a proton is a strong base and a base which shows less tendency to accept a proton is a weak base.

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 11.
Explain Lewis acid-base theory with suitable example. Classify the following species into Lewis acids and Lewis bases and show these act as Lewis acid / base.
a) OH
b) F
c) H+
d) BCl3
Answer:
Lewis theory of acids and bases : G.N. Lewis proposed a more generalised theory of acids and bases. According to this theory.
Acid : A substance that can accept an electron pair to form a co-ordinate covalent bond is called an acid.
Types of Lewis acids :
Lewis acids are of 5 types.
1) All Cations : Simple cations Ag, CO+3, Cu+2, Fe+3, Al+3 can act as Lewis acids.

2) Compounds in which the central atom has an incomplete octet and possessing an empty orbital can act as Lewis acids.
Ex : BF3, BCl3, AlCl3, FeCl3.

3) Compounds in which the central atom has vacant d-orbitals and may expand its octet can act as lewis acids.
Ex : SiF4, SF4, TeF4, SnCl4, FeCl3.

4) Molecules having multiple bonds between atoms of dissimilar electronegativities can behave like Lewis acids.
Ex : CO2, SO2, SO3, NO2, Cl2O7, P4O10

5) Elements with six electrons in the valence shell or electron sextet can act as Lewis acids.
Ex : O, S
Base : According to Lewis theory a base is a substance which can donate an electron pair to form a co-ordinate covalent bond.

Types of Lewis bases : Lewis bases are divided into three types.
1) All anions
Ex: Cl, OH, CN, \(\mathrm{NH}_2^{\ominus}\), F, SCN All simple anions can act as Lewis bases.
2) Molecules with one or two lone pairs on the central atom can act as Lewis bases.
Ex : AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 19
a) Hydrenyl ion is a lewis base as it can donate an electron lone pair (:OH).
b) Flouride ion acts as a Lewis base as it can donate any one of its four electron lone pairs.
c) A Proton is a lewis acid as it can accept alone pair of electrons from bases like hydrenyl ion and flowride ion.
d) BCl3 acts as a lewis acid as it can accept a lane pair electrons from species like ammenia or amine molecules.

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 12.
What is degree of ionization in respect of weak acids and weak bases ? Derive the relationship between degree of ionization (α) and ionization constant (Ka) for the weak acid HX.
Answer:
Degree of ionization (α):
The extent of ionization (or) dissociation of a weak acid or weak base is termed as its Degree of ionization (α). Let the concentration of weak acid (HX) be ‘C moles/lit and the degree of ionization be ‘α’. Hence at equilibrium,
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 20
If ‘α’ is neglected compared to ‘one’ in the denominator.
We have Ka = Cα2 ⇒ α2 = \(\frac{K_a}{C} \Rightarrow \alpha=\sqrt{\frac{K_a}{C}}\)
Similar equation can be written for the ionization of weak base (BOH) in aq. solution.
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 21
Therefore strengths of two acids or two bases are compared generally by their Ka or Kb values. Higher Ka or Kb value, stronger is the acid or base.

Question 13.
Define pH. What is buffer solution ? Derive Henderson – Hasselbalch equation for calculating the pH of an acid buffer solution.
Answer:
pH : “The negative value of the logarithm to the base 10, of the hydrogen ion concentration, expressed in moles / lit, in a solution is known as the pH of the solution”.
Mathematically, pH = – log10 [H+]
Buffer Solution:
“A buffer solution is that solution which resists any change in its pH value on dilution (or) on addition of a small amount of a strong acid or a strong base
Ex : Acidic buffer: (CH3COOH + CH3COONa)
Basic buffer: (NH4OH + NH4Cl)

Preparation of buffer solutions :
1) Acid buffer solutions : An acid buffer consists of weak acid and its salt with strong base.
Ex : (CH3COOH + CH3COONa)
Acid buffer solutions are normally prepared by mixing either equal or different volumes of equimolar solutions of a weak acid and its salt.

2) Base buffer solutions : A basic buffer solution consists of a mixture of a weak base and its salt with a strong acid.
Ex : (NH4OH + NH4Cl)
Base buffer solutions are prepared generally, by mixing either equal or different volumes of equimolar solutions of a weak base and its salt.

Derivation of Hendersen’s equation for an acid buffer :
i) Consider the acid buffer HA + NaA
HA ⇌ H+ + A-
NaA ⇌ Na+ + A
Acid dissociation constant Ka = \(\frac{\left[\mathrm{A}^{+}\right]\left[\mathrm{A}^{-}\right]}{[\mathrm{HA}]}\)
Here due to common ion effect
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 22

Question 14.
Explain the term “Hydrolysis of salts” with examples. Discuss the pH of the following types of salt solutions.
(i) Salts of weak acid and strong base.
(ii) Salts of strong acid and weak base. [T.S. Mar. 16]
Solution:
It may be defined as a phenomenon in which the anion or cation or both of a salt react with water producing excess of OH ions or H ions or both in aq. solution”.
i) -4 Given salt CH3COONa is obtained by the reaction of weak acid (CH3COOH) and strong base (NaOH).
-4 The ionisation CH3COONa as follows
CH3COONa ⇌ CH3COO + Na+
The above salt undergo anionic hydrolysis
CH3COO + H2O ⇌ CH3COOH + OH
The aq solution of salt is basic due to greater OH ion concentration than H+ ion concentration.
∴ PH > 7

ii) -4 Given salt NH4;C3 is obtained by the reaction of strong acid (HCl) and weak base (NH4OH).
-4 The ionisation of NH4Cl as follows
\(\mathrm{NH}_4^{+}\) + Cl ⇌ NH4Cl
-4 The above salt under goes catioric hydrolysis
\(\mathrm{NH}_4^{+}\) + H2O ⇌ NH4OH + H+
∴ The equation solution of salt is acidic due to greater H+ ion concentration than OH ion concentration.
∴ PH < 7

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 15.
What is solubility product ? Explain the common ion effect on solubility of ionic salts.
Solution:
Solubility product (KSP) : The product of the concentrations of the cation and the anion in a saturated solution of a salt at room temperature is called solubility product (KSP).
KSP = [Mn+][An-]
Significance of solubility product in chemical analysis :
In any solution of a salt, if the product of [Mn+] and [An-] is greater than the solubility product (KSP) of the salt, then only the salt gets precipitated. For example,

i) If the product of [Cu2+] & [S-2] in the II group (in presence of acid) of chemical analysis exceeds the value of KSP of CuS, then only CuS will be precipitated in the group. Similar argument applies to other metal ions in the group namely (Cd2+, Bi3+, Hg2+, Sb2+, etc).

ii) If the product of [Zn2+] and [S2-] in ammonical medium in IV group of chemical analysis exceeds the value of Ksp of ZnS, it will be precipitated. Similar arguments apply to Mn2+, Ni2+ etc.
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 23

 

  • Given salt CH3COONa is obtained by the reaction of weak acid (CH3COOH) and strong base (NaOH).
  • The ionisation CH3COONa as follows
    CH3COONa ⇌ CH3COO + Na+
  • The above salt undergo anionic hydrolysis
    CH3COO + H2O ⇌ CH3COOH + OH

∴ The aq solution of salt is basic due to greater OH ion concentration than H+ ion concentration.
∴ PH > 7

Given salt NH4Cl is obtained by the reaction of strong acid (HCl) and weak base (NH4OH).
The ionisation of NH4Cl as follows
\(\mathrm{NH}_4^{+}\) + Cl ⇌ NH4Cl
The above salt under goes catioric hydrolysis
\(\mathrm{NH}_4^{+}\) + H2O ⇌ NH4OH + H+
∴ The equation solution of salt is acidic due to greater H+ ion concentration than OH ion concentration.
∴ PH < 7

Common ion effect on solubility of ionic salts
The common ion effect is used for almost complete precipitation of a particular ion as its sparingly soluble salt, with every low value of solubility product for gravimetric estimation. Thus we can precipitate silver ion as silver chloride, ferric ion as its hydroxide (or hydrated ferric oxide) and barium ion as its sulphate for quantitative estimations.

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 16.
Write notes on
(i) Common ion effect.
(ii) The relation between Ksp and solubility (S) of a spaningly soluble salt BaSO4.
Solution:
i) Common ion effect : The decrease in the ionisation (dissociation) of a weak electrolyte, by the addition of a strong electrolyte having an ion common with the weak electrolyte, is known as common ion effect.
Example : The dissociation of NH4OH is diminished by the addition of NH4Cl due to the common ion, NH4+ ion
NH4OH ⇌ NH4 + OH
NH4OH AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 11 NH4+ + OH (addition of NH4Cl)
Importance of this effect In qualitative analysis :

  1. This principle is used in the systematic qualitative analysis of cations.
  2. The concentration of S2- in II group or OH ion in III group of chemical analysis is controlled by HCl and NH4OH on the basis of common ion effect.
    H+ ion is common to H2S (II group)
    NH4+ ion is common to NH4OH (IV group) .
  3. The common ion effect principle is also used in controlling the H+ ion concentration in buffer solutions.
  4. It is also used in the purification of common salts by passing dry HCl gas into impure salt solution.

ii) Relation between KSP and ‘S’ of salt at BaSO4
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 24

Problems

Question 1.
1 Mole of PCl5 is heated in a closed vessel of 1 litre capacity. At equilibrium 0.4 moles of chlorine is found. Calculate the equilibrium Constant.
Solutioni:
PCl5(g) ⇌ PCl3(g) + Cl2(g)
Initial [PCl5] = 1 mole/lit
At equilibrium [Cl2] = 0.4 moles/lit
[PCl3] at equilibrium = 0.4 moles/lit
[PCl5] at equilibrium = 0.6 moles/lit
KC = \(\frac{\left[\mathrm{PCl}_3\right]\left[\mathrm{Cl}_2\right]}{\left[\mathrm{PCl}_5\right]}\)
= \(\frac{0.4 \times 0.4}{0.6}=\frac{0.16}{0.6}\)
= 0.2606

Question 2.
Nitrogen dioxide forms dinitrogen tetroxide according to the equation 2 NO2 (g) ⇌ N2O4 (g) when 0.1 mole of NO2 as added to a 1 litre flask at 25°C, the concentration changes so that at equilibrium [NO2] = 0.016M and [N2O4] = 0.042 M
a. What is the value of the reaction Quotient before any reaction occurs.
b. What is the value of the equilibrium constant for the reaction.
Solution:
Given equation
2NO2(g) ⇌ N2O4(g)
a) The value of reaction quotient before any reaction occurs = 0

b) KC = \(\frac{\left[\mathrm{N}_2 \mathrm{O}_4\right]}{\left[\mathrm{NO}_2\right]^2}\)
= \(\frac{0.042}{(0.016)^2}=\frac{0.042}{0.000256}\)
= 164.06
= 1.64 × 102

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 3.
The equilibrium constant for the reaction:
N2(g) + 3H2(g) ⇌ 2 NH3(g) at 725K is 6.0 × 10-2. At equilibrium, [H2] = 0.25 mol L-1 and [NO3] = 0.06 mol L-1. Calculate the equilibrium concentration of N2.
Solution:
Given
N2(g) + 3H2(g) ⇌ 2NH3(g)
KC = 6 × 10-2
[H2] = 0.25 moles/lit
[NH3] = 0.06 moles/lit
KC = \(\frac{\left[\mathrm{NH}_3\right]^2}{\left[\mathrm{~N}_2\right]\left[\mathrm{H}_2\right]^3}\)
6 × 10-2 = \(\frac{(0.06)^2}{\left(N_2\right)(0.25)^3}\)
[N2] = \(\frac{(0.06)^2}{6 \times 10^{-2} \times(0.25)^3}\)
= \(\frac{0.0036}{0.015625 \times 6 \times 10^{-2}}\)
= \(\frac{0.0036}{0.09375 \times 10^{-2}}\)
= 3.84 moles/lit

Question 4.
At certain temperature, Kc for the re action.
SO2 (g) + NO2 (g) ⇌ SO3 (g) + NO (g)
is 16. If intially one mole each of all the four gases are taken in one litre vessel. What are the equilibrium con centrations of NO and NO2.
Solution:
Given equations is
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 25
x = 8 – 4x
5x = 8
x = 1.6 moles/lit ∴ (NO] = 1.6 moles/lit
2 – x = 2 – 1.6 = 0.4 moles/lit
[NO2] = 0.4 moles/lit

Question 5.
Under certain conditions, the equilibrium constant for the decomposition of PCl5 (g) into PCl3 (g) and Cl2 (g) is 00211 mol L-1. What are the equilibrium concentrations of PCl5, PCl3 and Cl2, if the intial concentration of PCl5 was 1.00 M?
Solution:
PCl5(g) ⇌ PCl3 + Cl2(g)
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 26
0.0211 (1 – x) = x2
0.0211 – 0.0211 x = x2
x = 0.14 M
∴ 1 – x = 1 – 0.14
= 0.86 M
∴ 1 [PCl5] = 0.86 M
∴ [PCl3] = 0.86 M
[PCl3] = [Cl2] = 0.14 M

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 6.
For the reaction A + B ⇌ 3C at 25°C, a 3 litre vessel contains 1, 2, 4 mole of A, B and C respectively predict the direction of reaction if
a. Kc for the reaction is 10
b. Kc for the reaction is 15
c. Kc for the reaction is 10.66
Answer:
A + B ⇌ 3C at 25°C
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 27
a) Given Kc = 10
Kc< Qc
∴ Backward reaction takes place

b) Given Kc = 15
Kc > Qc
∴ Forward reaction takes place

c) Given Kc = 10.66 Q = Kc
∴ Represents equilibrium state

Question 7.
A mixture of H2, N2 and NH3 with molar concentrations 5.0 × 10-3 mol L-1, 4.0 × 10-3 mol L-1 and 2.0 × 10-3 mol L-1 respectively was prepared and heated to 500K. The value of Kc for the reaction : 3H2 (g) + N2 (g) ⇌ 2NH3 (g) at this temperature is 60. Predicts whether ammonia tends to form or decompose at this stage of concentration.
Answer:
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 28
Given Kc = 60
Kc < Qc
∴ Backward reaction i.e decomposition of NH3 takes place

Question 8.
At 500 K, Kp value for the reaction
2 SO2 (g) + O2 (g) ⇌ 2SO3 (g) is 2.5 × 1010.
Find the value of Kp for each of following reactions at the same temperature.
a. SO2(g) + 1/2O2(g) ⇌ SO3 (g)
b. SO3(g) ⇌ SO2(g) + 1/2O2(g)
c. 3SO2(g) + 3/2 O2(g) ⇌ 3SO3(g)
Solution:
2SO2(g) + O2(g) ⇌ 2SO3(g)
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 29
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 30

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 9.
Kc for the reaction N2O4 (g) ⇌ 2NO2 (g) is 4.63 × 10-3 at 25°C.
a. What is the value of K at this termperature
b. At 25°C, if the partial pressure of N2O4 (g) at equilibrium is 0.2 atm1 calculate equilibrium pressure of NO2(g).
Solution:
a) Given
Na2O4(g) ⇌ 2NO2(g)
kc = 4.63 × 10-3
kp = kc (RT)∆n
= 4.63 × 10-3 × 0.0821 × 298 [∆n = 1]
= 113.27 × 10-3
= 0.1132

b) kp = \(\frac{\mathrm{P}_{\mathrm{NO}_2{ }^2}}{\mathrm{P}_{\mathrm{N}_2 \mathrm{O}_4}}\)
0.1132 = \(\frac{\mathrm{P}_{\mathrm{NO}_2{ }^2}}{0.2}\)
PNO22 = 0.2264
PNO2 = 0.15 atm

Question 10.
At 27°C, Kp value for the reversible reaction PCl5(g) PCl5(g) + Cl2(g) is 0.65, calculate Kc
Solution:
PCl5(g) ⇌ PCl3(g) +Cl2(g)
Given kp = 0.65 .
kp = k<subC (RT)∆n
0.65 = k<subC 0.082 1 × 300 [T =27°C = 300K; ∆n = 1]
k<subC = \(\frac{0.65}{24.63}\) = 0.02639

Question 11.
Kc for the reaction, N2 (g) + 3H2 (g) 2NH3(g) is 0.5 at 400K, find Kp.
Solution:
N2(g) + 3 H2(g) ⇌ 2 NH3(g)
Given kC = 0.5
T = 400 K
kp = kc (RT)∆n
kp = 0.5 × (0.0821 × 400)-2 ∆n = -2
kp = 0.5 × (8.21 × 4)-2 = 0.5 × (32.84)-2
\(\frac{0.5}{1078.46}\) = 4.63 × 10-4

Question 12.
1 mole of A and 1 mole of B are taken in a 5 litre flask, 0.5 mole of C is formed in the equilibrium of
A + B ⇌ C + D.
What is molar concentration of each species if the reaction is carried with 2 mole of A 1 mole of B in a 5 litre flask at the same temperature.
Solution:
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 31
= 0.2 – 0.133 = 0.067 M
[C] = [D] = 0.133 M

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 13.
For the following reaction
PCl5(g) ⇌ PCl3(g) + Cl2(g) 0.4 mole of PCl5 0.2 mole of PCl3 and 0.6 mole of Cl2 are taken in a 1 litre flask. If Kc = 0.2, Predict the direction in which reaction proceeds.
Solution:
Given
PCl(g) ⇌ PCl3(g) + Cl2 (g)
QC = \(\frac{\left[\mathrm{PCl}_3\right]\left[\mathrm{Cl}_2\right]}{\left[\mathrm{PCl}_5\right]}\)
= \(\frac{0.2 \times 0.6}{0.4}\)
= 0.03
Given KC = 0.2
QC < KC
∴ Reaction proceds in forward direction.

Question 14.
In an equilibrium A + B C + D; A and B are mixed in a vessel at temperature T. The initial concentration of A was twice the initial concentration of B. After the attainment of equilibrium, concentration of C was thrice concentration of B, calculate Kc.
Solution:
Given
A + B ⇌ C + D
The initial concentrations of A was twice the initial concentration of B and after the attaining of equilibrium concentration of ‘C’ was thrice the concentration of ‘B’.
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 32

Question 15.
A mixture of SO2, SO3 and O2 gases are maintained at equilibrium in 10 litre flast at a temperature at which Kc for the reactin 2SO2(g) + O2(g) ⇌ 2SO3(g) is 100. At equilibrium.
a. If no. of moles of SO3 and SO2 inflask are same, how many moles of O2 are present.
b. If no. of moles of SO3 in flask is twice the no. of moles SO2, how many moles of O2 are present.
Solution:
a) Given
2 SO2(g) + O2(g) ⇌ 2 SO3(g)
Given KC = 100

a) Given no.of moles of SO3 = No. of moles of SO2
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 33
y = \(\frac{4}{10}\) = 0.4
∴ 0.4 moles of O2 present.

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 16.
For A + B ⇌ C, the equilibrium concentrations of A and B at a temperature are 15 mol L-1. When volume is doubled the reaction has euilibrium concentration of A as 10 mol L-1. Calculate.
a) Kc
b) Concentration of C in original equilibrium
Solution:
a) Given
A + B ⇌ C + D
a) KC
b) Concentration of C in original equilibrium

a) KC = \(\frac{[\mathrm{C}]}{[\mathrm{A}][\mathrm{B}]}\)
[A] = 10 moles/lit
[B] = 10 moles/lit
[CI = 20 moles/lit
= \(\frac{20}{10 \times 10}\)
= \(\frac{20}{100}\) = 0.2

b) KC = \(\frac{[\mathrm{C}]}{[\mathrm{A}][\mathrm{B}]}\)
[AI = [B] = 15 moles/lit
0.2 = \(\frac{[C]}{15 \times 15}\)
[C] = 225 × 0.2
= 45 M

Question 17.
A vessel at 100 K contains CO2 with a pressure of 0.5 atm. Some of the CO2 is converted into CO on addition of graphite. Calculate the value of K, if total pressure at equilibnum is 0.8 atm.
Solution:
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 34
0.5 – x + 2x = 0.8
x = 0.3
∴ PCO = 0.2 atm
PCO2 = 0.6 atm
K = \(\frac{\mathrm{P}_{\mathrm{CO}_2^2}}{\mathrm{P}_{\mathrm{CO}}}\)
= \(\frac{0.36}{0.2}\) = 1.8

Question 18.
The K1, va)ue for the reaction H2(g)+ I2(g) ⇌ 2 HI(g) at 460°C is 49. If the initial pressure of H2 and I2 are 0.5 atm respectively, determine the partial pressure of each gases a equilibrium.
Solution:
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 35
x = \(\frac{3.5}{9}\)
= 0.3888
∴ PHI = 2 × 0.3888
= 0.7788 atm
PH2 = 0.5 – 0.388
= 0.111 atm
PI2 = 0.5 – 0.3888
= 0.111 atm

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 19.
0.5 mol of H2 and 0.5 mole of I2 react in 10 litre flask at 448°C. The equilibrium constant Kc is 50 for
H2(g) + I2(g) ⇌ 2HI(g).
a) What is the value of Kp
b) Calculate mole of I2 at equilibrium.
Solution:
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 36
∴ No of moles of I2 = 0.5 – 0.3888
= 0.111 moles

Question 20.
How much PCl5 must be added to a one little vesel at 250°C in order to obtain a concentration of 0.1 mole of Cl2 at equilibrium. Kc for
PCl5(g) ⇌ PCl3(g) + Cl2(g) is 0.0414 M.
Solution:
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 37
∴ KC = \(\frac{\left[\mathrm{PCl}_3\right]\left[\mathrm{Cl}_2\right]}{\left[\mathrm{PCl}_5\right]}\)
0.0414 = \(\frac{0.1 \times 0.1}{x-0.0414}\)
0.0414 x – 0.00414 = 0.01
0.0414 x = 0.01 + 0.00414
x = 0.34154 moles

Question 21.
K for the reaction
N2(g) + 3H2(g) ⇌ 2NH3(g) at 400°C is 1.64 × 10-4.
a) Calculate Kc
b) Calculate ∆G° value using Kc value.
Solution:
Given equation N2(g) + 3H2(g) ⇌ 2NH3(g)
a) Kp = Kc (RT)∆n
∆n = 2 – (4)
= -2
1.64 × 10-4 = Kc (0.0821 × 673)-2
Kc = 1.64 × (55.2533)2 × 10-4
Kc = 0.5006

b) ∆G° = -2.303 RT log K
= -2.303 × 0.0821 × 673 × log 0.5005
= 3874 J

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 22.
Calculate the pH of
a) 10-3 M HCl
b) 10-3 M H2SO4
c) 10-6 M HNO3
d) 0.02 M H2SO4
Solution:
a) 10-3 M HCl
pH = -log (H+)
= – log 10-3
= 3

b) 10-3 M H2SO4
pH = – log10 [H+]
= 0.001 × 2
= 0.002
pH = – log 0.002
= – log2 × 10-3
= – log2 – log 10-3
= 3 – log2
= 3 – 0.3010
= 2.699

c) 10-6 M HNO3
0.0008 M Ba(OH)2)
pOH = – log [OH]
[OH] = 0.0008 × 2
= 0.0016 N
∴ pOH = – log 0.0016
= – log 16 × 10-4
= – log 16 + 4 log 10
= – log24 + 4
= – 4 log 2 + 4 = 2.796
pH + pOH = 14
pH = 14 – 2.796
pH = 11.204

d) 0.02 M H2SO4
pH = -log10(H+)
[H+] = 0.02 × 2
= 0.04N
pH = – log 0.04
= -log 4 × 102
= -log4 – log 10-2
= 2 – log4
= 1.3010

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 23.
Calculate of pH for
a) 0.001 M NaOH
b) 0.01 M Ca (OH)2
c) 0.0008 M Ba(OH)2
d) 0.004M NaOH
Solution:
a) 0.001 M NaOH
pOH = -log[OH]
pOH = – log (0.001)
= – log 10-3
= 3
pH = 14 – pOH = 14 – 3 = 11

b) 0.01 M Ca (OH)2
pOH = -log [OH]
[OH] = 0.01 × 2
= 0.02 N
pOH = – log 0.02
pOH = -log 2 × 10-2
pOH = – log2 + 2 log 10
= 2 – 0.3010
= 1.699
∴ pH = 14 – pOH
= 14 – 1.699 = 12.301

c) 00008M Ba(OH)2
pOH = -log [OW]
[OH] = 0.0008 × 2
= 0.0016N
∴ pOH = – log 0.0016
= -log 16 × 10
= -log 16 + 4log 10
= – log 24 + 4
= -4 log 2 + 4 = 2.796
pH + pOH = 14
pH = 14 – 2.796
pH = 11.204
10-6M HNO3
PH = -log [H+]
= -log 10-6
= 6 log 10
= 6

d) 0.004M NaOH
pOH = – log[OH]
= – log 0.004
= – log 4 × 10-3
= 3 – log22
= 3 – 0.6020
= 2.398
pH = 14 – POH
= 14 – 2.398
= 11.602

Question 24.
The pH of a solution is 3.6 Calculate H3O+ ion concentration.
Solution:
pH = – log [H+]
log [H+] = – 3.6 (or) \(\overline{4} .4000\)
[H+] = anti log of \(\overline{4} .4\)
= 2.512 × 10-4.

Question 25.
The pH of a solution is 8.6 calculate the OH ion concentration.
Solution:
pH = 8.6
∴ pOH = 14 – pH
= 14 – 8.6
= 5.4
∴ pOH = – log [OH]
log [OH] = -5.4 (or) \(\overline{6} .6000\)
[OH] = anti log of \(\overline{6} .6000\)
= 3.981 × 10-6

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 26.
What is [H+] for a solution in which
a) pH = 3
b) pH = 4.75
c) pH = 4.4
Solution:
a) Given
pH = 3
∴ pH = – log [H+]
log [H+] = -3
∴ [H+] = 10-3 M

b) Given
pH = 4.75
PH = – log [H+]
log [H++] = – 4.75 (or) \(\overline{5} .2500\)
[H+] = anti log of 5.2500
= 1.77 × 10-5 M

c) Given
pH = 4.4
PH = – log [H+]
log [H+] = – 4.4 (or) \(\overline{5} .6\)
[H+] = anti log of \(\overline{5} .6\)
= 2.512 × 10-6 M

Question 27.
A solution of 0.005 M H2SO4 is diluted 100 times. Calculate the pH of diluted solution.
Solution:
0.005 M H2SO4
[H+] = 0.005 × 2 = 0.01
It is diluted to 100 times
[H+] = \(\frac{0.01}{100}\) = 0.0001
pH = – log [H+]
= – log 0.0001
= – log 10-4 = 4

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 28.
A solution of HCl has a pH = 3. If one ml of it is diluted to 1 litre, what will be the pH of the resulting solution ?
Solution:
Given
pH = 3 for HCZ solution
∴ [H+] = 10-3 M
It is diluted to litre
∴ [H+] = \(\frac{10^{-3}}{10^3}\) = 10-6
∴ pH = – log [H+]
= – log 10-6 = 6

Question 29.
What is the pH of 10-8 M HCl ?
Solution:
The given acid solution is very dilute. Hence the H+ obtained from the acid and water must be taken into consideration.
H+ ion concentration of acid is 10-8 M
H+ ion concentration of water is 10-7 M
Total [H+] = 10-8 + 10-7
= 1.1 × 10-7 or 1.1 × 10-7 M
∴ pH = – log 1.1 × 10-7
= 7 – log 1.1 = 7 – 0.0414 = 6.995

Question 30.
Calculate the pH of the following basic solutions.
a. [OH] = 0.05 M
b. [OH] = 2 × 10-4 M
Solution:
a) [OH] = 0.05 M
pOH = – log (0.05)
= – log 5 × 10-2
= – log 5 + 2 log 10
= 2 – log 5
= 1.3010
pH = 14 – pOH
pH = 14 – 1.3010
= 12.699

b) [OH] = 2 × 10-4 M
pOH = – log 2 × 10-4
= – log 2 + log 10
= 4 – log 2
= 4 – 0.3010
= 3.699
pH = 14 – 3.699 = 10.301

Question 31.
2g of NaOH is dissolved in water to give 1 litre solution. What is the pH of the solution ?
Solution:
N = \(\frac{w t}{G \in w} \times \frac{1}{1}\)
= \(\frac{2}{40}\) = \(\frac{1}{20}\) = 0.05 N = [OH]
∴ pOH = – log 0.05
= – log 5 × 10-2
= 2 – log 5
= 1.3010
pH = 14 – pOH
= 13 – 1.3010
= 12.699

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 32.
Calculate the pH of the following solutions.
a. 0.37 g of Ca(OH)2 dissolved in water to give 500 ml solution
b. 0.3 g of NaOH dissolved in water to give 200 ml solution
c. 0.1825% HCl aqueous solution
d. 1ml of 13.6 M HCl is diluted with water to give 1 litre solution.
Solution:
a) 0.37 gms of Ca(OH)2 in 500 ml
N = \(\frac{w t}{G \times w} \times \frac{1}{1}\)
N = \(\frac{0.37}{37} \times \frac{1000}{500}\)
= 0.01 × 2
= 0.02 = pOH
pOH = – log 0.02
= 1.699
pH = 14 – pOH = 12.301

b) 0.3 gms of NaOH in 200 mZ solution
N = \(\frac{0.3}{40} \times \frac{1000}{200}\)
= \(\frac{0.3 \times 5}{40}=\frac{1.5}{40}\)
= 0.375 = [OH]
pOH = – log [OH]
= – log 0.375
= 1.426
pH = 14 – pOH
= 14 – 1.426
= 12.574

c) 0.1825% HCl solution means 100 ml
solution contain 0.1825 g. of HCl
Normality (N) = \(\frac{W}{E W} \times \frac{1000}{V}\)
= \(\frac{0.1825}{36.5} \times \frac{1000}{100}\) = 0.05 N
∴ [H+] = 0.05 N
∴ pH = – log[H+]
= – log [0.051
= – log [5 × 10-2]
= 2 – log 5
= 2 – 0.692
∴ pH = 1.308

d) ∴ [H+] = 13.6
Diluted with water to 1 litre solution
∴ [H+] = \(\frac{13.6}{1000}\) = 0.016
pH = – log [H+]
= – log 0.0136 = 1.87

Question 33.
How many grams of NaOH are present in 100 ml solution if pH of the solution is 10?
Solution:
Given
pH = 10
pOH = 14 – 10 = 4
∴ Normality = 10-4 N
N = \(\frac{w t}{G M w} \times \frac{1000}{V(m l)}\)
10-4 = \(\frac{w t}{40} \times \frac{1000}{100}\)
∴ wt = 4 × 10-4 gms

Question 34.
The value of Kw is 9.55 × 10-14 at certain temperature. Calculate the pH of water at this temperature.
Solution:
Since Kw = [H+] [OH]
Kw = 9.55 × 10-14 mol2/lit2
[H+] [OH] = 9.55 × 10-14 mol2/lit2
∴ [H+] = \(\sqrt{9.55 \times 10^{-14}}\)
= 3.09 × 10-7 mol/lit
pH = – log[H+]
= – log [3.09 × 10-7]
= – [log 3.09 + log 10-7]
= – [0.49 – 7]
pH = 7 – 0.49 = 6.51

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 35.
Calculate the pH of 10-8 M NaOH.
Solution:
The given box solution is very dilute. Hence OH obtained from the box and water must be taken into consideration.
∴ [OH] = 10-8 + 10-7 = 1.1 × 10-7
pOH = -log 1.1 × 10-7
pOH = 6.995
∴ pOH = 14 – 6.995 = 7.005

Question 36.
150 ml of 0.5 M HCl and 100 ml of 0.2 M HCl are mixed. Find the pH of the resulting solution.
Solution:
N = \(\frac{N_1 V_1+N_2 V_2}{V_1+V_2}\)
= \(\frac{150 \times 0.5+100 \times 0.2}{250}\)
= \(\frac{75+20}{250}=\frac{95}{250}\) = 0.38
pH = – log [H+] = – log 1 (0.38) = 0.42

Question 37.
Calculate the pH of solution obtained by mixing 10ml of 0.1 M HCl and 40ml of 0.2M H2SO4.
Solution:
M = \(\frac{10 \times 0.1+40 \times 0.4}{50}\)
= \(\frac{1+16}{50}=\frac{17}{50}\) = 0.34
pH = – log [H+]
= – log (0.34) = 0.4685

Question 38.
100 ml of pH = 4 solution is mixed with 100 ml of pH = 6 solution. What is the pH of resulting solution.
Solution:
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 38

Question 39.
Equal volumes of 0.5 M NaOH and 0.3 M KOH are mixed in an experiment. Find the pOH and pH of the resulting solution.
Solution:
N = \(\frac{N_1 V_1+N_2 V_2}{V_1+V_2}\)
N = \(\frac{x(0.5)+x(0.3)}{x+x}\) Given value equal
= \(\frac{0.8}{2}\) = 0.4 = [OH]
pOH = – log [OH]
= – log (0.4)
= 0.3979
pH = 14 – 0.3979
= 13.6021

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 40.
60 m2 of 1 M HCl is mixed with 40 ml of 1 M NaOH. What is the pH of resulutant solution ?
Solution:
VANA > VBNB ⇒ N = \(\frac{V_A N_A-V_B N_B}{V_A+V_B}\)
∴ N = \(\frac{60-40}{60+40}\)
= \(\frac{20}{100}\) = \(\frac{1}{5}\) = 0.2
pH = – log [H+] = – log 0.2
= 0.699

Question 41.
Calculate the pH of a solution which contains 100 m2 of 0.1 M HCl and 9.9 m2 of 1.0M NaOH.
Solution:
Here VANA > VBNB
N = \(\frac{V_A N_A-V_B N_B}{V_A+V_B}\)
= \(\frac{10-9.9}{109.9}\)
= \(\frac{0.1}{109.9}\)
= 0.000909.
pH = – log [H+]
= – log (0.000909)
= 3.0416

Question 42.
What will be the resultant pH when 200 ml of an aqueons solution of HCl having pH = 2 is mixed with 300 m2 of an aqueous of NaOH having pH = 12?
Solution:
VA = 200 ml
NA = 10-2
[∵ pH = 2]

VB = 300 ml
NB = 10-2
[∵ pOH = 2]
∴ VBNB > VANA
∴ N = \(\frac{300 \times 10^{-2}-200 \times 10^{-2}}{300+200}\)
= \(\frac{1}{500}\) = 0.002 = [OH]
pOH = – log [OH]
= – log [0.002]
= -log 2 × 10-3
= – log 2 + 3 log 10
= 3 – log 2
= 2.699
∴ pH = 14 – pOH
= 14 – 2.699
= 11.3010

Question 43.
50 ml of 0.2 M HCl is added to 30 ml of 0.1 M KOH solution. Find pH of the solution.
Solution:
VA = 50 ml ; VB = 30 ml
NA = 0.2 N; NB = 0.1 ml
VANA > VBNB
∴ N = \(\frac{V_A N_A-V_B N_B}{V_A+V_B}\)
= \(\frac{50 \times 0.2-30 \times 0.1}{50+30}\)
= \(\frac{1-0.3}{80}=\frac{0.7}{80}\)
= 0.00875
∴ pH = – log [H+]
= – log (0.00875)
= 1.058.

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 44.
40 ml of 0.2 M HNO3 when reacted with 60 ml of 0.3 M NaOH, gave a mixed solution. What is the pH ofo the resulting solution.
Solution:
VBNB > VANA
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 39

Question 45.
50 ml of 0.1 M H2SO4 were added to 100 ml of 0.2 M HNO3. Then the solution is diluted to 300 ml. What is the pH of the solution.
Solution:
V1 = 50 ml ; V2 = 100 ml
N1 = 0.1 × 2 = 0.2N ; N2 = 02N
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 40

Question 46.
What is the Kw value in an aqueons solution of pKw = 13.725
Solution:
Given pKw = 13.725
pKw = – log Kw
Kw = antilog of 13.725
= 1.884 × 10-14

Question 47.
The ionic product of water at 80°C is 2.44 × 10-13. What are the concentrations of hydroniumion and the hydroxide in pure water at 80°C?
Solution:
Kw = [H+] [OH]
Kw = 2.44 × 10-13
For pure water [H+] = [OH]
∴ [H+] = Kw
[H+] = \(\sqrt{2.44 \times 10^{-13}}\)
= 4.94 × 10-7 moles / lit
∴ [OH] = 4.94 × 10-7 mole/lit

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 48.
The ionization constant for water is 2.9 × 10-14 at 40°C. Calculate [H3O+], [OH]. pH and pOH for pure water at 40C.
Solution:
Kw = [H-1] [OH]
2.9 × 10-14 = [H+] [OH]
For pure water [H+] = [OH]
∴ [H+]2 = 2.9 × 10-14
[H+] = 1.7 × 10-7 = [H3O+]
∴ [OH] = 1.7 × 10-7
pH = – log [H+]
= -log [1.7 × 10-7]
= 7 – log 1.7 = 6.7689
pOH = – log [OH]
= – log [1.7 × 10-7]
= 7 – log 1.7 = 6.7689

Question 49.
Calculate the pH of
a. 0.002 M acetic acid having 2.3% dissociation.
b. 0.002 M NH4OH having 2.3% dissociation.
Solution:
a) [H+] = Cα
= 0.002 × 2.3 = 0.0046
pH = – log [H+]
= – log 0.0046
= 4.3372

b) [OH] = Cα
= 0.002 × 2.3
= 0.0046
pOH = – log [OH]
= – log 0.0046
= 4.3372
pH = 14 – pOH
= 14 – 4.3372
= 9.6628

Question 50.
Calculate Ka of acetic acid from equilibrium concentration given below
[H3O+] = [CH3COO] = 1.34 × 10-3M, [CH3COOH] = 9.866 × 10-2 M
Solution:
Given
[H3O+] = [CH3COO] = = 1.34 × 10-3 M
[CH3COOH] = 9.866 × 10-2 M
Ka = \(\frac{\left[\mathrm{H}_3 \mathrm{O}^{+}\right]\left[\mathrm{CH}_3 \mathrm{COO}^{-}\right]}{\left[\mathrm{CH}_3 \mathrm{COOH}\right]}\)
= \(\frac{1.34 \times 10^{-3} \times 1.34 \times 10^{-3}}{9.866 \times 10^{-2}}\)
= 1.819 × 10-5

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 51.
Calculate pH of 0.1 M acetic acid having Ka =1.8 × 10-5.
Solution:
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 41

Question 52.
The pH of 0.1 M solution of weak mono protic acid is 4.0 calculate its [H+] and Ka.
Solution:
Given pH = 4.0
∴ [H+] = ?
pH = – log[H+]
[H+] = 10-pH
= 10-4
∴ [H+] = C × α
α = \(\frac{10^{-4}}{0.1}\)
= 10-3
Ka = C × α2
= 0.1 × (10-3)2
= 0.1 × 10-6
= 10-7.

Question 53.
Ka of 0.02 M CH3 COOH is 1.8 × 10-5. Calculate
a. [H3O+]
b. % ionization
c. pH
Solution:
(a), (b)
d2 = \(\frac{K_a}{C}\)
= \(\frac{1.8 \times 10^{-5}}{0.02}\)
α2 = 9 × 10-4
α = 3 × 10-2
% of ionisation
[H3O+] = C × α
= 0.02 × 3 × 10-2 = 6 × 10-4

c) pH = -log[H3O+]
= -log(6 × 10-4) = 3.24.

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 54.
Calculate the pH of 0.01 M solution of CH3 COOH. Ka for CHa OOH at 298K is 1.8 × 10-5.
Solution:
α2 = \(\frac{K_a}{C}\)
= \(\frac{1.8 \times 10^{-5}}{0.01}\)
= 180 × 10-5
= 18 × 10-4
α = 4.1 × 10-2
∴ [H3O+] = 0.01 × 4.1 × 10-2
= 4.1 × 10-4
pH = – log [H3O+]
= -log (4.1 × 10-4) = 3.38.

Question 55.
The pH of 0.1 M solution of an organic acid is 4.0 calculate the dissociation constant of the acid.
Solution:
Given pH = 4.0
∴ [H+] = ?
pH = – log[H+]
∴ [H+] = 10-pH = 10-4
∴ [H+] = C × α
α = \(\frac{10^{-4}}{0.1}\) = 10-3
Ka = C × α2
= 0.1 × (10-3)2 = 0.1 × 10-6 = 10-7

Question 56.
Three ionization constants for HF, HCOOH and HCN at 298 K are 6.8 × 10-4, 1.8 × 10-4 and 4.8 × 10-9 respectively. Calculate the ionization constants of the corresponding conjugate base.
Solution:
i) Consider the case of HF
Ka = 6.8 × 10-4
Kb = \(\frac{\mathrm{K}_{\mathrm{W}}}{\mathrm{K}_{\mathrm{a}}}=\frac{1 \times 10^{-14}}{6.8 \times 10^{-4}}\) = 0.147 × 10-10
ii) Consider the case of HCOOH
Ka = 1.8 × 10-4
Kb = \(\frac{K_w}{K_a}=\frac{1 \times 10^{-14}}{1.8 \times 10^{-4}}\) = 0.56 × 10-10

iii) Consider the case of HCN
Ka = 4.8 × 10-9
Kb = \(\frac{\mathrm{K}_{\mathrm{w}}}{\mathrm{K}_{\mathrm{a}}}=\frac{1 \times 10^{-14}}{4.8 \times 10^{-9}}\) = 0.208 × 10-5

Question 57.
Find the concentration of hydroxide ion in a 0.25 M solution of trimethylamine, a weak base.
(CH3)3N + H2O ⇌ (CH3)3N+H+OH; Kb = 7.4 × 10-5.
Solution:
σα = \(\sqrt{\frac{K_b}{c}}=\sqrt{\frac{7.4 \times 10^{-5}}{0.25}}\)
= 1.74 × 10-3
[OH] = c × α
= 0.25 × 1.74 × 10-3
= 4.32 × 10-3 M

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 58.
The 0.005 M monobasic acid has a pH of 5. What is the extent of ionization ?
Solution:
pH = 5
∴ [H+] = 10-5
Concentration (C) = 0.005 M = 5 × 10-3 M
[H+] = Cα
α = degree of dissociation per mole = ?
∴ α = \(\frac{\left[\mathrm{H}^{+}\right]}{\mathrm{C}}=\frac{\left[\mathrm{H}^{+}\right]}{5 \times 10^{-3}}=\frac{10^{-5}}{5 \times 10^{-3}}\)
= \(\frac{1}{5}\) × 10-2 = 0.2 × 10-2
Percentage of degree of dissociates
= 0.02 × 10-2 × 100 = 0.2 %

Question 59.
50 ml of 0.1 M NH4OH, 25m1 of 2 M NH4Cl were used to make a buffer. What is the pH if PKa is 4.8.
Solution:
The given buffer is a base buffer.
For a base buffer pOH = pKb + log \(\frac{[\text { Salt }]}{[\text { Base }]}\)
∴ pOH = 4.8 + \(\log \left\{\frac{\frac{(25 \times 2)}{75}}{\frac{(50 \times 0.1)}{75}}\right\}\)
= 4.8 + log(10) = 4.8 + 1 = 5.8
∴ pH = (14.0 – pOH) = (14.0 – 5.8) = 8.2
∵ pH + pOH = 14.0 at room temperature.

Question 60.
The pH of a buffer prepared by mixing 50 ml of 0.2 M CH3 COOH and 25 ml of CH3 COONa is 4.8. If the PKa is 4.8. What is the strength of CH3 COONa.
Solution:
The given solution is an acid buffer.
∴ The pHof the buffer is given by,
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 42

Question 61.
50 ml of 0.1 M sodium acetate, 25ml of 0.2 M acetic acid were added together to form the buffer solution. PKa of CH3 COOH is 4.8 Find the pH of the solution.
Solution:
The buffer solution is an acid buffer.
Hendersen’s equation for an acid buffer is
pH = pKa + log
\(\frac{[\text { Salt }]}{[\text { Base }]}=4.8+\left\{\log \frac{\frac{(0.1 \times 50)}{75}}{\frac{(0.2 \times 25)}{75}}\right\}\)
= 4.8 + log (1)
∴ pH = 4.8.

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 62.
When 20 ml of 0.1 M NH4OH are added to 20ml of 1M NH4Cl solution, the pH of the buffer formed is 8.2. What is the pKa of NH4OH?
Solution:
For base buffer, pH is given by
pH = 14.0 – pOH and pOH
= pKb + \(\log \frac{[\text { Salt }]}{[\text { Base }]}\)
The pH of the buffer = 8.2;
∴ pOH =(14 – pH) = 5.8 then
5.8 = pKb + \(\log \frac{[\text { Salt }]}{[\text { Base }]}\)
= pKb + \(\log \left[\frac{\frac{20 \times 1}{40}}{\frac{20 \times 0.1}{40}}\right]\)
= pKb + log (10) = (pKb + 1.0)
∴ pKb = (5.8 – 1) = 4.8

Question 63.
One litre of buffer solution contains 0.1 mole of aceticacid add 1 mole of sodium acetate. Find its pH if PKa of CH3COOH is 4.8.
Solution:
For an acid buffer pH = PKa + \(\log \frac{[\text { Salt }]}{[\text { Base }]}\)
pH = 4.8 + log \(\frac{1}{0.1}\) = 4.8 + log 10
pH = 4.8 + 1 ⇒ pH = 5.8

Question 64.
50 ml of 1 M CH3 COOH solution, when added to 50 ml of 0.5 M NaOH gives a solution with a pH ‘X’. Find the value of ‘X’. PKa of acetic acid is 4.8.
Solution:
Acetic acid reacts with sodium hydroxide to form sodium acetate.
CH3COOH + NaOH → CH3COONa + H2O
No.of millimoles of CH3COOH = 50 × 1 = 50
No.of millimoles of NaOH = 50 × 0.5 = 25
No.of millimoles of CH3. COOH remained unreacted = 50 – 25 = 25
Total Vol. of solution = 50 + 50 = 100 ml
No. of millimoles of CH3COONa formed = 25
∴ Molarity of CH3COONa = \(\frac{25}{100}\) = .25 M
Molarity of CH3COOH = \(\frac{25}{100}\) = .25 M
pH = pKa + \(\log \frac{[\text { Salt }]}{[\text { Base }]}\)
pH = 4.8 + log \(\frac{1}{0.1}\) = 4.8 + log 10
= 4.8 + log \(\frac{.25}{.24}\) = 4.8 + log(1)
∴ pH of the solution = 4.8 + 0 = 4.8

Question 65.
The solubility product of AgCl is 1.6 × 10-10 mol2/L2. What is solubility?
Solution:
KSP = S2
1.6 × 10-10 = S2
∴ S = \(\sqrt{10^{-10} \times 1.6}\) = 1.3 × 10-5 moles /litre.

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 66.
The solutbility product of Zr (OH)2 is 4.5 × 10-17 mol3L-3 what is solubility?
Solution:
Zn(OH)2 ⇌ Zn++ + 2OH
∴ KSP = S × 4S2 = 4S3
∴ 4S3 = 4.5 × 10-17
S3 = 4.5 × 10-18/4 = 11.7 × 10-18
S = \(\sqrt[3]{11.7} \times \sqrt[3]{10^{-18}}\)
S = 2.6 × 10-6

Question 67.
The solubility of Ag2 CrO4 is 1.3 × 10-4 mol L-1. What is the solubility product ?
Solution:
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 43
KSP = 4S2 × S = 4S3
∴ KSP = 4 × S3
= 4 × (1.3 × 10-4)3
KSP = 9 × 10-12

Question 68.
The solubility of A2B = 2 × 10-3 mol L-1. What is solubility product?
Solution:
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 44
KSP = [2s]2[S]
S = 3\(\sqrt{\frac{K_{S P}}{4}}\)
= 3\(\sqrt{\frac{2 \times 10^{-3}}{4}}\) = 3.2 × 10-10

Question 69.
The sotubility product of a salt AB = 10-10 mol2 L-2. What is solubility?
Solution:
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 45
KSP = s2
S = \(\sqrt{K_{s p}}=\sqrt{10^{-10}}\) = 10-5 moles/lit

Question 70.
PQ and RS2 are two sparingly soluble salts. Their solubility products are equal and each equal to 4.0 × 10-18. Which salt is more soluble ?
Solution:
Given KSP of PQ and RS2 salts = 4 × 10-18
S = \(\sqrt{\frac{K_{\mathrm{SP}}}{4}}\) for RS2
S = \(\sqrt{K_{s p}}\) for PQ
∴ Solubility of RS2 is more.

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 71.
In a 0.1 M solution, acetic acid is 1.34% ionized. Calculate [H+], [CH3COO] and [CH3COOH] in the solution and Calculate Ka of acetic acid.
Solution:
Given 1.34 % ionised
[H+] = C × α
= 0.1 × 1.34 × 10-2 = 1.34 × 10-3 M
∴ [CH3COOH] =1.34 × 10-3 M
[H+] = Cα = 1.34 × 10-3]
α = \(\frac{1.34 \times 10^{-3}}{0.1}\) = 1.34 × 10-2]
[CH3COO] (1 – α) = 0.1 (1 – 0.0134)
= 0.09866 M
Ka = α2 × C
= (1.34 × 10-2) × 0.1 = 1.79 × 10-5

Solved Problems

Question 1.
The following concentrations were obtained for the formation of NH3 from N2 and H2 at equilibrium at 500 K. [N2] = 1.5 × 10-2 M. [H2] = 3.0 × 10-2 M and [NH3] = 1.2 × 10-2 M. Calculate equilibrium constant.
Solution:
The equilibrium constant for the reaction. N2(g) + 3 H2 (g) ⇌ NH3(g) can be written as,
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 46

Question 2.
At equilibrium, the concentrations of N2 = 3.0 × 10-3 M, O2 = 4.2 × 10-3 M and NO = 2.8 × 10-3 M in sealed vessel at 800K. What will be Kc for the reaction
N2(g) + O2 (g) ⇌ 2 NO(g)
Solution:
For the reaction equilibrium constant, KC can be written as,
KC = \(\frac{[\mathrm{NO}]^2}{\left[\mathrm{~N}_2\right]\left[\mathrm{O}_2\right]}\)
= \(\frac{\left(2.8 \times 10^{-3} \mathrm{M}\right)^2}{\left(3.0 \times 10^{-3} \mathrm{M}\right)\left(4.2 \times 10^{-3} \mathrm{M}\right)}\)
= 0.622

Question 3.
PCl5, PCl3 and Care at equilibrium at 500 K and having concentration 1.59 M PCl3, 159M Cl2 and 1.41 M PCl5, Calculate K for the reaction
PCl5 ⇌ PCl3 + Cl2
Solution:
The equilibrium constant K. for the above reaction can be written as,
KC = \(\frac{\left[\mathrm{PCl}_3\right]\left[\mathrm{Cl}_2\right]}{\left[\mathrm{PCl}_5\right]}=\frac{(1.59)^2}{(1.41)}\) = 1.79

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 4.
The value of KC = 4.24 of 800 K for the reaction
CO (g) + H2O (g) ⇌ CO2 (g) + H2 (g)
Calculate equilibrium concentrations of CO2, H2, CO and H2O at 800 K, if only CO and H2O are present initially at concentrations of 0.10 M each.
Solution:
For the reaction,
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 47
where x is the amount of CO2 and H2 at equilibrium.
Hence, equilibrium constant can be written as,
KC = x2/(0.1 – x)2 = 4.24
x2 = 4.24(0.01 + x2 – 0.2 x)
x2 = 0.0424 + 4.24 x2 – 0.848 x
3.24 x2 – 0.848 x + 0.0424 = 0
a = 3.24, b = – 0.848, c = 0.0424
(for quadratic equation ax2 + bx + c = 0
x = \(\frac{\left(-b \pm \sqrt{b^2-4 a c}\right)}{2 a}\)
x = 0.848 ± \(\sqrt{(0.848)^2}\) – 4(3.24) (0.04224) / (3.24 × 2)
x = (0.848 ± 0.4118)/6.48
x1 = (0.848 – 0.4118)] 6.48 = 0.067
x2 = (0.848 + 0.4118)/6.48 = 0.194
the value 0.194 should be neglected because it will give concentration of the reactant which is more than initial concentration. Hence the equilibrium concentrations are,
[CO2] [H2] = x = 0.067 M
[CO] = [H2O] = 0.1 – 0.067 = 0.033 M

Question 5.
For the equilibrium, 3 NOCl (g) ⇌ 2 NO(g) + C2 (g) the value of the equilibrium constant, KC is 3.75 × 10-6 at 1069 K. Calculate the Kp for the reaction at this temperature ?
Solution:
We know that,
KP = KC(RT)∆n
For the above reaction,
∆n = (2 + 1) – 2 = 1
KP = 3.75 × 10-6 (0.0831 × 1069)
KP = 0.033

Question 6.
The value of KP for the reaction.
CO2 (g) + C(s) ⇌ 2CO (g) is 3.0 at 1000 K. If initially PCO2 = 0.48 bar and PCO = 0 bar and pure graphite is present, calculate the equilibrium partial pressures of CO and CO2.
Solution:
For the reaction,
let ‘x’ be the decrease in pressure of CO2, then
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 48
= (-3 ± 5.66)/8
= (-3 + 5.66)/8 (as value of x cannot be negative hence we neglect that value)
x = 2.66/8 = 0.33
The equilibrium partial pressures are,
PCO = 2x = 2 × 0.33 = 0.66 bar
PCO2 = 0.48 – x = 0.48 – 0.33 = 0.15 bar

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 7.
The value of Kc for the reaction
2A ⇌ B + C is 2 × 10-3. At a given time, the composition of reaction mixture is [A] = [B] = [C] = 3 × 10-4 M. In which direction the reaction will proceed ?
Solution:
For the reaction the reaction quotient QC is given by.
QC = [B] [C] / [A]2
as [A] = [B] = [C] = 3 × 10-4 M
QC = (3 × 10-4) (3 × 10-4)/ (3 × 10-4)2 = 1
as QC > KC so the reaction will proceed in the reverse direction.

Question 8.
13.8 g of N2O4 was placed in a 1L reaction vessel at 400 K and allowed to attain equilibrium
N2O4(g) ⇌ 2 NO2(g)
The total pressure at equilbrium was found to be 9.15 bar. Calculate KC, KP and partial pressure at equilibrium.
Solution:
We know pV = nRT
Total volume (V) = 1 L
Molecular mass of N2O4 = 92 g
Number of moles = 13.8 g / 92 g = 0.15 of the gas (n)
Gas constant (R) = 0.083 bar L mol-1K-1
Temperature (T) = 400 K
pV = nRT
p × 1L = 0.15 mol × 0.083 bar L mol-1K-1 × 400 K
p = 4.98 bar
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 49
Hence,
ptotal at equilibrium = pN2O4 + pNO2
9.15 = (4.98 – x) + 2x
9.15 = 4.98+ x
x = 9.15 – 4.98 = 4.17 bar
Partial pressures at equilibrium are,
pN2O4 = 4.98 – 4.17 = 0.81 bar
pNO2 = 2x = 2 × 4.17 = 8.34 bar
Kp = (pNO2)2 / pN2O4
= (8.34)2 / 0.81 = 85.87
Kp = KC (RT)∆n
85.87 = KC(0.083 × 400)1
KC = 2.586 = 2.6

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 9.
300 mol of PCl5 kept in 1L closed reaction vessel was allowed to attain, equilibrium at 380K. Calculate composition of the mixture at equilibrium. Kc = 1.80
Solution:
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 50

Question 10.
The value of ∆G° for the phosphorylation of glucose in glycolysis is 13.8 kJ/mol. Find the value of Kc at 298 K.
Solution:
∆G° = 13.8 ki/ mol = 13.8 × 103J/mol
Also, ∆G° = – RT Inkc
Hence, In Kc = – 13.8 × 103 J/mol
(8.314 Jmol-1K-1 × 298 K)
In Kc = 5.569
Kc = e-5 569
Kc = 3.81 × 10-3

Question 11.
Hydrolysis of sucrose gives.
Sucrose + H2O ⇌ Glucose + Fructose Equilibrium constatnt Kc for the reaction is 2 × 1013 at 300K. Calculate ∆GΘ at 300K.
Solution:
∆G° = -RT InKc
∆G° = -8.314J mol-1K-1 × 300K × In (2 × 1013)
∆G° = – 7.64 × 104J mo-1

Question 12.
What will be the conjugate bases for the following Bronsted acides: HF, H2SO4 and \(\mathrm{HCO}_3^{-}\) ?
Solution:
The conjugate bases should have one proton less in each case and therefore the corresponding conjugate bases are: F, \(\mathrm{HSO}_3^{-}\) and CO2-3 respectively.

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 13.
Write the conjugate acids for the following Bronsted bases: NH2, NH3 and HCOO.
Solution:
The conjugate acid should have one extra proton in each case and therefore the corresponding conjugate acids are : NH3, NHO+4 and HCOOH respectively.

Question 14.
The species; H2O, \(\mathrm{HCO}_{3^{\prime}}^{-}\), \(\mathrm{HCO}_{4^{\prime}}^{-}\) and NH3 can act both as Bronsted acids andbases. For each case give the corresponding conjugate acid and conjugate base.
Solution:
The answer is given in the following Table:
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 51

Question 15.
Classify the following species into Lewis acids and Lewis bases and show how these act as such :
(a) HO
(b) F
(c) H+
(d) BCl3
Solution:
(a) Hydroxyl ion is a Lewis base as it can donate an electron lone pair (: OH).
(b) Flouride ion acts as a Lewis base as it can donate any one of its four electron lone pairs.
(c) A proton is a Lewis acid as it can accept a lone pair of electrons from bases like hydroxyl ion and fluoride ion.
(d) BCl2 acts a Lewis acid as it can accept a lone pair of electrons from species like ammonia or amine molecules.

Question 16.
The concentration of hydrogen ion in a sample of soft drink is 3.8 × 10-3M. What is its pH ?
Solution:
pH = – Log[3.8 × 10-3]
= -{log[3.8] + log[10-3}]
= – {(0.58) + (-3.0)} = -2.42} = 2.42
Therefore, the pH of the soft drink is 2.42 and it can be inferred that it is acidic.

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 17.
Calculate pH of a 1.0 -8 M solution of HCl.
Solution:
2H2O(1) H3O+ (aq) + OH(aq)
Kw = [OH][H3O+]
= 10-14
Let, x = [OH] = [H3O+] from H2O. The H3O+ concentration is generated (i) from the ionization of HCl dissolved i.e., HCl (aq) + H2O (1) ⇌ H3O+ (aq) + Cl(aq), and (ii) from ionization of H2O. Inthese very dilute solutions. both sources of H3O+ must be considered:
[H3O+] = 10-8 + X
Kw = (10-8 + x)(x) = 10-14
or x2 + 10-8 x – 10-14 = 0
[OH] = x = 9.5 × 10-8
So, pOH = 7.02 and pH = 6.98

Question 18.
The ionization costant of HF is 3.2 × 10-4. Calculate the degree of dissociation of HF in its 0.02 M solution. Calculate the concentration of all species present (H3O+.F and HF) in the solution and its pH.
Solution:
The following proton transfer reactions are possible :
1) HF + H2O ⇌ H3O+FKa = 3.2 × 10-4
2) H2O + H2O H3O+ + OH Kw = 1.0 × 10-14
As ka >> Kw, [1] is the principle reaction.
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 52
Substituting equilibrium concentrations in the equailibrium reaction for principal reaction gives.
Ka (0.02α)2/ (0.02 – 0.02α)
= 0.02α2/(1 – α) = 3.2 × 10-4
We obtain the following quadratic equation:
α2 + 1.6 × 10-2α – 1.6 × 10-2 = 0
The quadratic equation in α can be solved and the two values of the roots are :
α = + 0.12 and – 0.12
The negative root is not acceptable and hence.
a = 0.12
This means that the degree of ionization α = 0.12, then equilibrium concentrations of other species viz., HF, F and H3O+ are given by:
[H3O+] = [F] = cα = 0.02 × 0.12
= 2.4 × 10-3M
[HF] = c(1 – α) = 0.02(1 – 0.12)
= 17.6 × 10-3M
pH = -log[H+] = -log (2.4 × 10-3) = 2.62

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 19.
The pH of 0.1 M monobasic acid is 4.50. Calculate the concentration of species H+, A and HA at equilibrium. Also, determine the value of Ka and pKa of the monobasic acid.
Solution:
pH = – log [H+]
Therefore [H+] = 10-pH = 10-4.50
= 3.16 × 10-5
[H+] = [A] = 3.16 × 10-5
Thus, Ka = [H+] [A] / [HA]
[HA]eqlbm = 0.1 – (3.16 × 10-5) ≅ 0.1
Ka = (3.16 × 10-5)2/0.1 = 1.0 × 10-8
pKa = – log (10-8) = 8
Alternatively, “Percent dissociation” is another useful method for measure of strength of a weak acid and is given as : Percent dissociation
= [HA]dissociated × 100/[HA]initial (7.32)

Question 20.
Calculate the pH of 0.08M solution of hypochlorous acid. HOCI. The ionization Constant of the acid is 2.5 × 10-5, Determine the percent dissociation of HOCl
Solution:
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 53
[H+] = 1.41 × 10-3 M Therefore,
Percent dissociation
= {[HOCI]dissociated × 100/[HOCI]initial}
= 1.41 × 10-3 /0.08 = 1.76%,
pH = -log (1.41 × 10-3) = 2.85

Question 21.
The pH of 0.004 M hydrazine solution is 9.7. Calculate is ionization constant Kb and pKb.
Solution:
NH2NH2 + H2O ⇌ \(\mathrm{NH}_2 \mathrm{NH}_3^{+}\) +OH
From the pH we can calculate the hydrogen ion concentration. Knowing hydrogen ion concentration and the ionic product of water we can calculate the concentration of hydroxyl ions. Thus we have :
[H+] = antilog (-pH)
= antilog (-9.7) = 1.67 × 10-10
[OH] = Kw/[H+] = 1 × 10-14/ 1.67 × 10-10
= 5.98 × 10-5
The concentration fo the corresponding hydrazinium ion is also the same as that of hydroxyl ion. The concentrations of both these ions are very small so the concentration of the undissociated base can be taken equal to 0.004 M.
Thus,
Kb = [NH2NH3+] [OH] / [NH2NH2]
= (5.98 × 10-5)2/0.004 = 8.96 × 10-7
pKb = -logKb = -log(8.96 × 10-7 ) = 6.04

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 22.
Calculate the pH of the solution in which 0.2M NH4Cl and 0.1 M NH3 are present. The pKb of ammonia solution is 4.75.
Solution:
NH3 + H2O ⇌ \(\mathrm{NH}_4^{+}\) + OH
The ionization constant of NH3,
Kb= antilog (-pKb) i.e.
Kb = 10-4.75 = 1.77 × 10-5 M
3 + H2O ⇌ \(\mathrm{NH}_4^{+}\) + OH Initial concentration (M)
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 54
Change in concentration to reach equilibrium (M)
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 55
Kb = [latex]\mathrm{NH}_4^{+}[/latex][OH-1] / [NH3]
= (0.20 × x) (x) / (0.1 – x) = 1.77 × 10-5
As Kb is small, we can neglect x in comparison to 0.1 M and 0.2 M. Thus
[OH] = x = 0.88 × 10-5
Therefore, [H+] = 1.12 × 10-9
pH = -log [H+] = 8.95 .

Question 23.
Determine the degree of ionization and pH of a 0.05M of ammonia solution. The ionization constant of ammonia can be taken from Table 7.7. Also, calculate the ionization constant of the conjugate acid of ammonia.
Solution:
The ionization of NH3 in water is represented by equation :
NH3 + H2O ⇌ \(\mathrm{NH}_4^{+}\) + OH
We use equation (7.33) to calculate hydroxyl ion concentration,
[OH] c α = 0.05 α
Kb = 0.05 α2/(1 – α)
The value of α is small, therefore the quadratic equation can be simplified by neglecting a in comparison to 1 in the denominator on right hand side of the equation,
Thus,
Kb = C α2 or α = \(\sqrt{\left(1.77 \times 10^{-5} / 0.05\right)}\)
= 0.018.
[OH] = C α = 0.05 × 0.018 = 9.4 × 10-4 M.
[H+] = Kw/[OH] = 10-14 / (9.4 × 10-4)
= 1.06 × 10-11
pH = -log (1.06 × 10-11) = 10.97
Now, using the relation for conjugate acid-base pair,
Ka × Kb = Kw
using the value of Kb of NH3 from Tabel 7.7.
We can determine the concentration of con-jugate acid \(\mathrm{NH}_4^{+}\)
Ka × Kw/ Kb = 10-14/ 1.77 × 10-5
= 5.64 × 10-10.

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 24.
Calculate the pH of a 0.10M ammonia solution. Calculate the pH after 50.0 mL of this solution is treated with 25.0 mL of 0.10M HCl. The dissociation constant of ammonia, Kb = 1.77 × 10-5.
Solution:
NH3 + H2O ⇌ \(\mathrm{NH}_4^{+}\) + OH
Kb = [latex]\mathrm{NH}_4^{+}[/latex] [OH] / [NH3] = 1.77 × 10-5
Before neutralization,
[latex]\mathrm{NH}_4^{+}[/latex] [OH] = x
[NH3] = 0.10 – x – 0.10, x2 / 0.10 = 1.77 × 10-5
Thus, x= 1.33 × 10-3 = [OH]
Therefore, [H+] = Kw / [OH] = 10-14 / (1.33 × 10-3)= 7.51 × 10-12
pH = -log (7.5 × 10-12 ) = 11.12
On addition of 25 mL of 0.1 M HCl solution (i.e., 2.5 mmol of HCl) to 50 mL of 0.1 M ammonia solution (i.e., 5 mmol of NH3), 2.5 mmol of ammonia molecules are neutralized. The resulting 75 mL solution contains the remaining unneutralized 2.5 mmol of NH3 molecules and 2.5 mmol of \(\mathrm{NH}_4^{+}\)
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 56
The resulting 75 mL of solution contains 2.5 mmol of \(\mathrm{NH}_4^{+}\) ions (i.e., 0.033M) and 2.5 mmol (i.e., 0.033M) of uneutralised NH3 molecules. This NH3 exists as NH4OH in the following equilibrium :
AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases 57
The final 75 mL solution after neutralisation already contains 2.5 m mol [latex]\mathrm{NH}_4^{+}[/latex] ions (i.e. 0.033M), thus total concentration of [latex]\mathrm{NH}_4^{+}[/latex] ions is given as : [NHJ] = 0.033 + y
As y is small, [NH4OH] ≃ 0.033 M and [latex]\mathrm{NH}_4^{+}[/latex] ≃ 0.033M.
We know,
Kb = [latex]\mathrm{NH}_4^{+}[/latex] [OH] / [NH4OH]
= y(0.033)/(0.033) 1.77 × 10-5 M
Thus, y = 1.77 × 10-5 = [OH]
[H+] = 10-14/ 1.77 × 10-5 = 0.56 × 10-9
Hence, pH = 9.24

Question 25.
The pKa of acetic acid and pKa of ammonium hydroxide are 4.76 and 4.75 respectively. Calculate the pH of am-monium acetate solution.
Solution:
pH = 7 + 1/2 [pKa – pKb]
= 7 + 1/2 [4.76 – 4.75]
= 7 + 1/2 [0.01] = 7 + 0.005 = 7.005

Question 26.
Calculate the solubility of A2X3 in pure water, assuming that neither kind of ion reacts with water. The solubility product of A2X3, Ksp = 1.1 × 10-23.
Solution:
A2X3 ⇌ 2A3+ + 3X2-
Ksp = [A3+]2 [X2-]3 = 1.1 × 10-23
If S = solubility of A2X3, then
[A3+] = 2S; [X2-] = 3S
therefore, Ksp = (2S)2 (3S)3 = 108S5
= 1.1 × 10-23
thus, S5 = 1 × 10-23
S = 1.0 × 10-5 mol/L.

AP Inter 1st Year Chemistry Study Material Chapter 7 Chemical Equilibrium and Acids-Bases

Question 27.
The values of Ksp of two sparingly soluble salts Ni(OH)2 and AgCN are 2.0 × 10-15 and 6 × 10-17 respectively. Which salt is more soluble ? Explain.
Solution:
AgCN ⇌ Ag+ + CN
Ksp = [Ag+] [CN] = 6 × 10-17
Ni(OH)2 ⇌ Ni2 + 2OH
Ksp = [Ni2+] [OH]2 = 2 × 10-15
Let [Ag+] = S1 then [CN] = S1
Let [Ni2+] = S2 then [OH] = 2S2
S12 = 6 × 10-17 = S1 = 7.8 × 10-9
(S2) (2S2)2 = 2 × 10-15, S2 = 0.58 × 10-4
Ni(OH)2 is more soluble than AgCN.

Question 28.
Calculate the molar solubility of Ni(OH)2 in 0.10 M NaOH. The ionic product of Ni(OH)2 is 2.0 × 10-15.
Solution:
Let the solubility of Ni(OH)2 be equal to S. Dissolution of S mol/L of Ni(OH)2 provides S mol/L of Ni2+ and 2S mol/L of OH, but the total concentration of OH = (0.10 + 2S) mol/L because the solution already contains 0.10 mol/L of OH from NaOH.
Ksp = 2.0 × 10-15 = [Ni2+] [OH]2
(S) (0.10 + 2S)2 As Ksp is small, 2S < < 0.10,
thus, (0.10 + 2S) ≈ 0.10
Hence,
2.0 × 10-15 S(0.10)2
S = 2.0 × 10-13 M = [Ni2+]