The electricity required in coulombs for the oxidation of 1 mole of FeO to \(Fe_{2} O_{3 }\) is:

The amount of charge required for the reduction of 1 mol of ${\mathrm{MnO}}_4^{-}$ to ${\mathrm{Mn}}^{2+}$ is:

$1.$ $4.8$ $\times$ ${10}^5$ $\mathrm{C}$
$2.$ $3.2$ $\times$ ${10}^6$ $\mathrm{C}$
$3.$ $1.8$ $\times$ ${10}^5$ $\mathrm{C}$
$4.$ $4.1$ $\times$ ${10}^4$ $\mathrm{C}$

The number of electrons delivered at the cathode during electrolysis by a current of 1 ampere in 60 seconds is:

(Given: Charge on electron = 1.60 × 10^–19 C)

1.  $6\times {10}^{23}$

2.  $6\times {10}^{20}$

3.  $3.75\times {10}^{20}$

4.  $7.48\times {10}^{20}$

When 0.1 mol MnO₄^2– is oxidized, the quantity of electricity required to completely oxidise MnO₄^2– to MnO₄^– is:

The number of Faradays required to produce 20.0 g of Ca from molten CaCl₂ is-

1. 2F

2. 1F

3. 4F

4. 3F

A steady current of 1.5 A flows through a copper voltmeter for 10 min. If the electrochemical equivalent of copper is 30 × 10^-5 g C^-1, the mass of copper deposited on the electrode will be:

1. 0.40 g

2. 0.50 g

3. 0.67 g

4. 0.27 g

For the cell, Ti/Ti⁺(0.001M)||Cu²⁺(0.1M)|Cu, ${\mathrm{E}}_{\mathrm{cell}}^{\mathrm{o}}$ $$at

25 $°$C is 0.83 V. E_cell can be increased :

1. By increasing [Cu²⁺]

2. By increasing [Ti⁺]

3. By decreasing [Cu²⁺]

4. None of the above.

The  number of electrons involved in the deposition of 63.5 g of Cu from a solution of $CuS{O}_4$ is:

1. $6.022$ $\times$ ${10}^{23}$ $$

2. $3.011$ $\times$ ${10}^{23}$ $$

3. $12.044$ $\times$ ${10}^{23}$ $$

4. $6.022$ $\times$ ${10}^{22}$