A solution of copper sulphate is electrolysed for 10 minutes with a current of 1.5 amperes. The mass of copper deposited at cathode is :

(Given : Molar mass of $\mathrm{Cu}=63 \mathrm{~g} \mathrm{~mol}^{-1}$;

$$ \left.1 \mathrm{~F}=96487 \mathrm{C} \mathrm{~mol}^{-1}\right) $$

Calculate emf of the half cell given below :

$$ \begin{aligned} & \mathrm{Pt}(\mathrm{~s})\left|\mathrm{H}_2(\mathrm{~g}, 2 \mathrm{~atm})\right| \mathrm{HCl}(\mathrm{aq}, 0.02 \mathrm{M}) \\ & \mathrm{E}_{\mathrm{H}_2 / \mathrm{H}^{+}}^{\circ}=0 \mathrm{~V} \end{aligned} $$

(Given: $\frac{2.303 R T}{F}=0.059, \log 2=0.3010$ )

From the following select the one which is not an example of corrosion.

The standard cell potential of the following cell $$\mathrm{Zn}\left|\mathrm{Zn}^{2+}(\mathrm{aq})\right| \mathrm{Fe}^{2+}(\mathrm{aq}) \mid \mathrm{Fe}$$ is $$0.32 \mathrm{~V}$$. Calculate the standard Gibbs energy change for the reaction:

$$\mathrm{Zn}(\mathrm{s})+\mathrm{Fe}^{2+}(\mathrm{aq}) \rightarrow \mathrm{Zn}^{2+}(\mathrm{aq})+\mathrm{Fe}(\mathrm{s})$$

(Given : $$1 \mathrm{~F}=96487 \mathrm{C}$$)