Nernst Equation

$\left(\mathrm{i}\right)$ The  ${\mathrm{E}}_{\mathrm{cell}}^{°}$ of the electrochemical cell representing the reaction is

$2\mathrm{Cr}\left(\mathrm{s}\right)+3{\mathrm{Fe}}^{2+}\left(\mathrm{aq}\right)\to 2{\mathrm{Cr}}^{3+}\left(\mathrm{aq}\right)+3\mathrm{Fe}\left(\mathrm{s}\right)$

$\left(\mathrm{ii}\right)$ The ${\mathrm{E}}_{\mathrm{cell}}$ at  ${25}^{°}\mathrm{C}$ when $\left[{\mathrm{Cr}}^{+3}\right]=0.1 \mathrm{M}$ and $\left[{\mathrm{Fe}}^{2+}\right]=0.01 \mathrm{M}$ will be?

$$\begin{gathered} {\mathrm{E}}_{\raisebox{1ex}{${\mathrm{Cr}}^{+3}$}\!\left/ \!\raisebox{-1ex}{$\mathrm{Cr}$}\right.}^{°}=-0.74\mathrm{\hspace{0.17em}}\mathrm{V} \\ {\mathrm{E}}_{\raisebox{1ex}{${\mathrm{Fe}}^{+2}$}\!\left/ \!\raisebox{-1ex}{$\mathrm{Fe}$}\right.}^{°}=-0.44\mathrm{\hspace{0.17em}}\mathrm{V} \end{gathered}$$

The EMF of a cell corresponding to the reaction:

$\mathrm{Zn}\left(\mathrm{s}\right)+2{\mathrm{H}}^{+}\left(\mathrm{aq}\right)\to {\mathrm{Zn}}^{2+}\left(\mathrm{aq}\right)(0.1\hspace{0.17em}\hspace{0.17em}\mathrm{M})+{\mathrm{H}}_2\left(\mathrm{g}\right)\hspace{0.17em}\hspace{0.17em}(1\hspace{0.17em}\hspace{0.17em}\mathrm{atm}.)\hspace{0.17em}\hspace{0.17em}\mathrm{is}\hspace{0.17em}\hspace{0.17em}0.28\hspace{0.17em}\hspace{0.17em}\mathrm{volt}\hspace{0.17em}\hspace{0.17em}\mathrm{at}\hspace{0.17em}\hspace{0.17em}25°\mathrm{C}.$

The pH of the solution at the hydrogen electrode?
${\text{E}}_{\text{Z}{\text{n}}^{2+}/\text{Z}\text{n}}^{\text{o}}=-0.76\text{v}\text{o}\text{l}\text{t};{\text{E}}_{{\text{H}}^{+}/{\text{H}}_2}^{\text{o}}=0$

If $6.539\times {10}^{-2} \mathrm{g}$ of metallic zinc is added to $100 \mathrm{mL}$ saturated solution of $\mathrm{AgCl}$, then $\mathrm{Ag} \left(\mathrm{s}\right)$ precipitates in the above reaction.

$\left[{\mathrm{K}}_{\mathrm{sp}} \left(\mathrm{AgCl}\right)={10}^{-10}\right]$

(Atomic mass of $\mathrm{Zn}=65.39$)
What is the value of $\log \frac{\left[{\mathrm{Zn}}^{2+}\right]}{{\left[{\mathrm{Ag}}^{+}\right]}^2}$?

$$\begin{gathered} {{\mathrm{E}}^{°}}_{\mathrm{Ag}}=0\mathrm{.}80 \mathrm{V} \\ {{\mathrm{E}}^{°}}_{\mathrm{Zn}}=-0\mathrm{.}76 \mathrm{V} \end{gathered}$$

Find the number of moles of $\mathrm{Ag}$ formed.

A cell, $Ag|A{g}^{+}\parallel C{u}^{2+}|Cu,$ initially contains 1 M $A{g}^{+}$  and 1  $MC{u}^{2+}$ ions.

Calculate the change in the cell potential after the passage of 9.65 A of current for 1 h.

For the reduction of $N{O}_3^{-}$ ion in an aqueous solution, $E°$ is $+0.96V$. Values $E°$ for some metal ions are given below

$\begin{array}{l}{V}^{2+}(aq)+2e^{-}\to VE°=-1.19V\\ F{e}^{3+}(aq)+3e^{-}\to FeE°=-0.04V\\ A{u}^{3+}(aq)+3e^{-}\to AuE°=+1.40V\\ H{g}^{2+}(aq)+2e^{-}\to HgE°=+0.86V\end{array}$

The pair(s) of metals that is (are) oxidized by ${\mathrm{NO}}_3^{-}$ in aqueous solution is/are

The standard reduction potential values of three metallic cations, X, Y and Z are $0.52,-3.03 \mathrm{and}-1.18\mathrm{V}$ respectively. The order of reducing power of the corresponding metals is:

Using the Nernst equation, calculate emf of the following cell at $298 \mathrm{K}$:

$\mathrm{Cr}\left(\mathrm{s}\right) \left| {\mathrm{Cr}}^{3+}\left(0.1 \mathrm{M}\right) \right|\left| {\mathrm{Fe}}^{2+}\left(0.01\mathrm{M}\right) \right|\mathrm{Fe}\left(\mathrm{s}\right)$

Given that,

$$\begin{gathered} {\mathrm{E}}_{{\mathrm{Cr}}^{3+}/\mathrm{Cr}}^{\circ }=-0.75 \mathrm{V} \\ {\mathrm{E}}_{{\mathrm{Fe}}^{2+}/\mathrm{Fe}}^{\circ }=-0.45 \mathrm{V} \end{gathered}$$

$\frac{2.303\mathrm{RT}}{\mathrm{F}}=0.0591\mathrm{V}$

Higher reduction potential increases the reactivity of non-metals.

What is the affect of reduction potential on the reactivity of non-metals?

Which halogen will be more reactive?

Reduction potential of bromine- $+1.07\mathrm{V}$ or reduction potential of chlorine- $1.36\mathrm{V}$.

Give an example of non-metal displacing a non-metal in a chemical reaction.