Dr. S.N. DHAWAN, Dr. S.C.KHETERPAL and, P.N.KAPIL Solutions for Chapter: Redox Reactions, Exercise 1: PROBLEMS FOR PRACTICE

The e.m.f. $\left(\mathrm{E}°\right)$ of the following cells are
$\mathrm{Ag}\hspace{0.17em}\left|{\mathrm{Ag}}^{+}\right(1 \mathrm{M})\Vert {\mathrm{Cu}}^{2+}(1 \mathrm{M}\left)\right|\mathrm{Cu}$:  $\mathrm{E}= –0.46 \mathrm{V}$, $\mathrm{Zn}\hspace{0.17em}\left|{\mathrm{Zn}}^{2+}\right(1 \mathrm{M})\Vert {\mathrm{Cu}}^{2+}(1 \mathrm{M}\left)\right|\mathrm{Cu}$, $\mathrm{E}= +1.10 \mathrm{V}$
Calculate the e.m.f. of the cell $\mathrm{Zn}\hspace{0.17em}\left|{\mathrm{Zn}}^{2+}\right(1 \mathrm{M})\Vert {\mathrm{Ag}}^{+}(1 \mathrm{M}\left)\right|\mathrm{Ag}$

Two half cells are ${\mathrm{Al}}^{3+}\left(\mathrm{aq}\right)/\mathrm{Al}$ and ${\mathrm{Mg}}^{2+}\left(\mathrm{aq}\right)/\mathrm{Mg}$
The reduction potentials these half cells are $–1.66 \mathrm{V}$ and $–2.36 \mathrm{V}$ respectively. Calculate the cell potential. Write the cell reaction also.

Calculate E° for the cell: $\mathrm{Al}\hspace{0.17em}|{\mathrm{Al}}^{3+}\left(1\hspace{0.17em}\mathrm{M}\right)\Vert {\mathrm{Cu}}^{2+}\left(1\hspace{0.17em}\mathrm{M}\right)|\mathrm{Cu}$
Given: ${{\mathrm{E}}^{\circ }}_{{\mathrm{Al}}^{3+}/\mathrm{Al}}$ and ${{\mathrm{E}}^{\circ }}_{{\mathrm{Cu}}^{2+}/\mathrm{Cu}}$ as $–1.66 \mathrm{V}$ and $0.34 \mathrm{V}$ respectively.

Predict reaction of $1 \mathrm{N}$ sulphuric acid with the following metals: (i) Copper (ii) lead (iii) iron
Given : ${{\mathrm{E}}^{\circ }}_{{\mathrm{Cu}}^{2+},\hspace{0.17em}\mathrm{Cu}}$$=+0.34 \mathrm{volt}$ ; ${{\mathrm{E}}^{\circ }}_{{\mathrm{Pb}}^{2+},\hspace{0.17em}\mathrm{Pb}}$$= –0.13 \mathrm{volt}$ ; ${{\mathrm{E}}^{\circ }}_{{\mathrm{Fe}}^{2+},\hspace{0.17em}\mathrm{Fe}}$$= –0.44 \mathrm{volt}$

Can a solution of $1 \mathrm{M}$ ${\mathrm{ZnSO}}_4$ be stored in a vessel made of copper? Given that
${{\mathrm{E}}^{\circ }}_{\mathrm{Zn},{\mathrm{Zn}}^{2+}}=0.76\hspace{0.17em}\mathrm{volt}$ and ${{\mathrm{E}}^{\circ }}_{{\mathrm{Cu}}^{2+},\hspace{0.17em}\mathrm{Cu}}=0.34\hspace{0.17em}\mathrm{volt}$

Is it safe to stir $1 \mathrm{M}$ ${\mathrm{AgNO}}_3$ solution with a copper spoon? Given: ${{\mathrm{E}}^{\circ }}_{{\mathrm{Ag}}^{+},\hspace{0.17em}\mathrm{Ag}}=0.80\hspace{0.17em}\mathrm{volt}$ and ${{\mathrm{E}}^{\circ }}_{\mathrm{Cu},\hspace{0.17em}\hspace{0.17em}{\mathrm{Cu}}^{2+}}=–0.34\hspace{0.17em}\mathrm{volt}.$ Explain.

Can we use a copper vessel to store $1 \mathrm{M}$ ${\mathrm{AgNO}}_3$ solution? Given that, ${{\mathrm{E}}^{\circ }}_{{\mathrm{Cu}}^{2+}/\hspace{0.17em}\mathrm{Cu}}=+\hspace{0.17em}0.34\hspace{0.17em}\mathrm{volt}$ and ${{\mathrm{E}}^{\circ }}_{{\mathrm{Ag}}^{+}/\hspace{0.17em}\mathrm{Ag}}=+\hspace{0.17em}0.80\hspace{0.17em}\mathrm{volt}.$

Why blue colour of copper sulphate solution gets discharged when zinc rod is dipped in it? Given: ${{\mathrm{E}}^{\circ }}_{{\mathrm{Cu}}^{2+}/\hspace{0.17em}\mathrm{Cu}}=\hspace{0.17em}0.34\hspace{0.17em}\mathrm{V}$ and ${{\mathrm{E}}^{\circ }}_{\mathrm{Zn}/\hspace{0.17em}{\mathrm{Zn}}^{2+}}=\hspace{0.17em}0.76\hspace{0.17em}\mathrm{V}$.