Faraday's First Law

Q.

If $0.224\mathrm{L}$ ${\mathrm{H}}_2$ gas is formed at the cathode, the volume of ${\mathrm{O}}_2$ gas formed at the anode under identical conditions is?

Q.

One mole of ${\mathrm{N}}_2{\mathrm{O}}_4$ in a $1\mathrm{L}$ flask decomposes to attain the equilibrium ${\mathrm{N}}_2{\mathrm{O}}_4\left(\mathrm{g}\right)\rightleftharpoons 2{\mathrm{NO}}_2\left(\mathrm{g}\right)$. At the equilibrium the mole fraction of ${\mathrm{N}}_2{\mathrm{O}}_4$ is $1/2$ . then ${\mathrm{K}}_{\mathrm{C}}$ will be:

Q. Electrochemical equivalent (Z) is same as equivalent weight of the deposited substance

1. True
2. False

Q. When electricity is passed through a solution of $AlC{l}_3$, 13.5 g of Al is deposited. The number of Faradays (charge) must be:

1. 1
2. 1.5
3. 0.5
4. 2

Q. $2Faraday$ of electricity is passed through a solution of $CuS{O}_4$ . The mass of copper deposited at the cathode is:

1. $0g$
2. $63.5g$
3. $2.0g$
4. $127.0g$

Q. Electrolysis of dilute aqueous $NaCl$ solution was carried out by passing $10mA$ current. The time required to liberate $0.01mole$ of $H_2$ gas at the cathode is ($1F=96500Cmo{l}^{-1})$:

1. $9.65\times {10}^{4}s$
2. $19.3\times {10}^{4}s$
3. $28.95\times {10}^{4}s$
4. $38.6\times {10}^{4}s$

Q. Calculate the volume of gas liberated at anode at NTP from the electrolysis of CuSO4 solution by a current of 2 ampere passed for 10 minutes.

1. 0.0696 liters

2. 0.864 liters

3. 1.68 liters

4. 3.45 liters

Q. A certain amount of current liberates $0.5g$ of hydrogen in 2 hours. How many grams of copper can be produced by the same current flowing for the same time in copper sulphate solution?

Q. In a fuel cell, $H_2$ and $O_2$ react to produce electricity. In the process, $H_2$ gas is oxidized at the anode and $O_2$ gas is reduced at the cathode. If $67.2\text{litres}$ of $H_2$ at $STP$ react in 15 minutes and the entire current is used for electrode deposition of $Cu$ from $Cu\left(II\right)$ solution, how many grams of $Cu$ will be deposited and what is the average current produced?
Anode : $H_2+2O{H}^{-}\to 2H_{2}O+2e^{-}$
Cathode : $O_2+2H_{2}O+4e^{-}\to 4O{H}^{-}$

1. $\text{190.5 g of Cu}$
2. $\text{381 g of Cu}$
3. $\text{643.33 A}$
4. $\text{321.6 A}$

Q. During the discharge of a lead storage battery the density of sulphuric acid fell from $1.294$ to $1.139{\text{g mL}}^{-1}$ . $H_{2}S{O}_4$ of density $1.294{\text{g mL}}^{-1}$ is 39% and that of density $1.139{\text{g mL}}^{-1}$ is 20% by weight. The battery holds $3.5\text{L}$ of acid and the volume practically remains constant during the discharge. Calculate the number of Ampere hours for which the battery must have been used. The discharging reactions are:
$Pb+S{O}_4^{2-}\to PbS{O}_4+2e^{-}\left(\text{anode}\right)$
$Pb{O}_2+4H^{+}+S{O}_4^{2-}+2e^{-}\to PbS{O}_4+2H_{2}O\left(\text{cathode}\right)$

1. $190$
2. $530$
3. $265$
4. $132.5$

Q. In a hydrogen oxygen fuel cell, electricity is produced. In this process $H_2\left(\text{g}\right)$ is oxidised at anode and $O_2\left(\text{g}\right)$ reduced at cathode. Given : Cathode: $O_2\left(g\right)+2H_{2}O\left(l\right)+4e^{-}\to 4O{H}^{-}\left(aq\right)$ Anode : $H_2\left(g\right)+2O{H}^{-}\left(aq\right)\to 2H_{2}O\left(l\right)+2e^{-}$ $4.48\text{L}$ of $H_2$ at $1\text{atm}$ and $273\text{K}$ is oxidised in $9650\text{sec}$.
If current produced in fuel cell, use for the deposition of $C{u}^{2+}$ in $1\text{L}$, $2\text{M}CuS{O}_4\left(\text{aq}\right)$ solution for $241.25\text{sec}$ using $Pt$ electrode. The $pH$ of solution after electrolysis is:

1. $1$
2. $2$
3. $3$
4. $4$

Q.

When aq. AgNO₃ is electrolysed using Pt electrodes if 3.2 g of O₂ is liberated at the anode, the amount Ag that will be deposited at the cathode in grams is (at wt. of Ag = 108)

1. 32 g

2. 1.6 g

3. 4.32 g

4. 43.2 g

Q. Two beakers that have the same solution have different masses of metals deposited on their electrodes when the current passed .

1. is the same
2. is different
3. is zero

Q.

In an electroplating experiment m g of silver is deposited, when 4 amperes

of current flows for 2 minutes. The amount (in gms) of silver deposited by 6

amperes of current flowing for 40 seconds will be

[MNR 1991]

1. 4 m

2. m/2

3. m/4

4. 2m

Q.

During the electrolysis of an electrolyte, the number of ions produced, is directly proportional to the

[AFMC 2002]

1. Time consumed

2. Electro chemical equivalent of electrolysis

3. Quantity of electricity passed

4. Mass of electrons

Q.

What is the mass of silver deposited on the passage of 2 A current during the electrolysis of silver?
Electrochemical equivalent of silver = $1.18\times {10}^{-3}gm{C}^{-1}$

1. $2.28\times {10}^{-3}gm$
2. $2.36\times {10}^{-3}gm$
3. $2\times {10}^{-3}gm$
4. $Cannotbedetermined$

Q. The amount of electricity required to deposit 0.9gof aluminum the electrode reaction is
$A{l}^{3+}+3e^{-}\to Al$
(atomic mass of Al = 27)

1. $9.65\times {10}^{3}C$
2. $1.93\times {10}^{4}C$
3. $9.65\times {10}^{4}C$
4. $4.34\times {10}^{5}C$

Q.

In an electroplating experiment m g of silver is deposited, when 4 amperes

of current flows for 2 minutes. The amount (in gms) of silver deposited by 6

amperes of current flowing for 40 seconds will be

[MNR 1991]

1. 4 m

2. m/2

3. m/4

4. 2m

Q. In a fuel cell, $H_2$ and $O_2$ react to produce electricity. In the process, $H_2$ gas is oxidized at the anode and $O_2$ gas is reduced at the cathode. If $67.2\text{litres}$ of $H_2$ at $STP$ react in 15 minutes and the entire current is used for electrode deposition of $Cu$ from $Cu\left(II\right)$ solution, how many grams of $Cu$ will be deposited and what is the average current produced?
Anode : $H_2+2O{H}^{-}\to 2H_{2}O+2e^{-}$
Cathode : $O_2+2H_{2}O+4e^{-}\to 4O{H}^{-}$

1. $\text{190.5 g of Cu}$
2. $\text{381 g of Cu}$
3. $\text{643.33 A}$
4. $\text{321.6 A}$

Q. A solution of $Ni(N{O}_3{)}_2$ is electrolyzed between platinum electrodes using $0.1Faraday$ electricity. Moles of $Ni$ that will be deposited at the cathode:

1. $0.20$
2. $0.10$
3. $0.15$
4. $0.05$

Q. The amount of electricity required to deposit 0.9gof aluminum the electrode reaction is
$A{l}^{3+}+3e^{-}\to Al$
(atomic mass of Al = 27)

1. $9.65\times {10}^{3}C$
2. $1.93\times {10}^{4}C$
3. $9.65\times {10}^{4}C$
4. $4.34\times {10}^{5}C$