276100 Given, $\mathrm{E}_{\mathrm{Cr}^{3+} / \mathrm{Cr}}^{0}=0.72 \mathrm{~V}$,
$\mathrm{E}_{\mathrm{Fe}^{2+} / \mathrm{Fe}}^{0}=-0.42 \mathrm{~V}$
The potential for the cell $\mathrm{Cr} \mid \mathrm{Cr}^{3+}(0.1 \mathrm{M}) \ \vert \mathrm{Fe}^{2+}$ $(0.01 \mathrm{M}) \mid \mathrm{Fe}$ is

276101 The standard electrode potential $E^{\circ}$ and its temperature coefficient $\left(\frac{d E^{0}}{d T}\right)$ for a cell are $2 V$ and $-5 \times 10^{-4} \mathrm{VK}^{-1}$ at $300 \mathrm{~K}$ respectively. The cell reaction is $\mathrm{Zn}(\mathrm{s})+\mathrm{Cu}^{2+}(\mathrm{aq}) \rightarrow \mathrm{Zn}^{2+}(\mathrm{aq})+$ $\mathrm{Cu}(\mathrm{s})$. The standard reaction enthalpy $\left(\Delta_{\mathrm{r}} \mathrm{H}^{\ominus}\right)$ at $300 \mathrm{~K}$ in kJ mol ${ }^{-1}$ is, [Use, $R=8 \mathrm{JK}^{-1} \mathrm{~mol}^{-1}$ and $F=96,000 \mathrm{C} \mathrm{mol}^{-1}$ ]

276102 Calculate the standard cell potential (in V) of the cell in which following reaction takes place
$\mathrm{Fe}^{2+}(\mathrm{aq})+\mathrm{Ag}^{+}(\mathrm{aq}) \rightarrow \mathrm{Fe}^{3+} \text { (aq) }+\mathrm{Ag}(\mathrm{s})$
Given that, $\mathbf{E}_{{ }_{\mathrm{Ag}+/ \mathrm{Ag}}}^{\mathbf{0}}=\mathrm{x}$ V,
$\mathbf{E}_{\mathrm{Fe}^{2+} / \mathrm{Fe}}^{\mathbf{0}}=\mathbf{y} \mathbf{V}$
$\mathbf{E}_{\mathrm{Fe}^{3+} / \mathrm{Fe}}^{\mathbf{0}}=\mathbf{z} \mathbf{V}$

276104 For the given cell;
$\mathrm{Cu}(\mathrm{s})\left \vert\mathrm{Cu}^{2+}\left(\mathrm{C}_{1} \mathrm{M}\right) \ \vert \mathrm{Cu}^{2+}\left(\mathrm{C}_{2} \mathrm{M}\right)\right \vert \mathrm{Cu}(\mathrm{s})$
change in Gibbs energy $(\Delta G)$ is negative, if

276100 Given, $\mathrm{E}_{\mathrm{Cr}^{3+} / \mathrm{Cr}}^{0}=0.72 \mathrm{~V}$,
$\mathrm{E}_{\mathrm{Fe}^{2+} / \mathrm{Fe}}^{0}=-0.42 \mathrm{~V}$
The potential for the cell $\mathrm{Cr} \mid \mathrm{Cr}^{3+}(0.1 \mathrm{M}) \ \vert \mathrm{Fe}^{2+}$ $(0.01 \mathrm{M}) \mid \mathrm{Fe}$ is

276101 The standard electrode potential $E^{\circ}$ and its temperature coefficient $\left(\frac{d E^{0}}{d T}\right)$ for a cell are $2 V$ and $-5 \times 10^{-4} \mathrm{VK}^{-1}$ at $300 \mathrm{~K}$ respectively. The cell reaction is $\mathrm{Zn}(\mathrm{s})+\mathrm{Cu}^{2+}(\mathrm{aq}) \rightarrow \mathrm{Zn}^{2+}(\mathrm{aq})+$ $\mathrm{Cu}(\mathrm{s})$. The standard reaction enthalpy $\left(\Delta_{\mathrm{r}} \mathrm{H}^{\ominus}\right)$ at $300 \mathrm{~K}$ in kJ mol ${ }^{-1}$ is, [Use, $R=8 \mathrm{JK}^{-1} \mathrm{~mol}^{-1}$ and $F=96,000 \mathrm{C} \mathrm{mol}^{-1}$ ]

276102 Calculate the standard cell potential (in V) of the cell in which following reaction takes place
$\mathrm{Fe}^{2+}(\mathrm{aq})+\mathrm{Ag}^{+}(\mathrm{aq}) \rightarrow \mathrm{Fe}^{3+} \text { (aq) }+\mathrm{Ag}(\mathrm{s})$
Given that, $\mathbf{E}_{{ }_{\mathrm{Ag}+/ \mathrm{Ag}}}^{\mathbf{0}}=\mathrm{x}$ V,
$\mathbf{E}_{\mathrm{Fe}^{2+} / \mathrm{Fe}}^{\mathbf{0}}=\mathbf{y} \mathbf{V}$
$\mathbf{E}_{\mathrm{Fe}^{3+} / \mathrm{Fe}}^{\mathbf{0}}=\mathbf{z} \mathbf{V}$

276104 For the given cell;
$\mathrm{Cu}(\mathrm{s})\left \vert\mathrm{Cu}^{2+}\left(\mathrm{C}_{1} \mathrm{M}\right) \ \vert \mathrm{Cu}^{2+}\left(\mathrm{C}_{2} \mathrm{M}\right)\right \vert \mathrm{Cu}(\mathrm{s})$
change in Gibbs energy $(\Delta G)$ is negative, if

276100 Given, $\mathrm{E}_{\mathrm{Cr}^{3+} / \mathrm{Cr}}^{0}=0.72 \mathrm{~V}$,
$\mathrm{E}_{\mathrm{Fe}^{2+} / \mathrm{Fe}}^{0}=-0.42 \mathrm{~V}$
The potential for the cell $\mathrm{Cr} \mid \mathrm{Cr}^{3+}(0.1 \mathrm{M}) \ \vert \mathrm{Fe}^{2+}$ $(0.01 \mathrm{M}) \mid \mathrm{Fe}$ is

276101 The standard electrode potential $E^{\circ}$ and its temperature coefficient $\left(\frac{d E^{0}}{d T}\right)$ for a cell are $2 V$ and $-5 \times 10^{-4} \mathrm{VK}^{-1}$ at $300 \mathrm{~K}$ respectively. The cell reaction is $\mathrm{Zn}(\mathrm{s})+\mathrm{Cu}^{2+}(\mathrm{aq}) \rightarrow \mathrm{Zn}^{2+}(\mathrm{aq})+$ $\mathrm{Cu}(\mathrm{s})$. The standard reaction enthalpy $\left(\Delta_{\mathrm{r}} \mathrm{H}^{\ominus}\right)$ at $300 \mathrm{~K}$ in kJ mol ${ }^{-1}$ is, [Use, $R=8 \mathrm{JK}^{-1} \mathrm{~mol}^{-1}$ and $F=96,000 \mathrm{C} \mathrm{mol}^{-1}$ ]

276102 Calculate the standard cell potential (in V) of the cell in which following reaction takes place
$\mathrm{Fe}^{2+}(\mathrm{aq})+\mathrm{Ag}^{+}(\mathrm{aq}) \rightarrow \mathrm{Fe}^{3+} \text { (aq) }+\mathrm{Ag}(\mathrm{s})$
Given that, $\mathbf{E}_{{ }_{\mathrm{Ag}+/ \mathrm{Ag}}}^{\mathbf{0}}=\mathrm{x}$ V,
$\mathbf{E}_{\mathrm{Fe}^{2+} / \mathrm{Fe}}^{\mathbf{0}}=\mathbf{y} \mathbf{V}$
$\mathbf{E}_{\mathrm{Fe}^{3+} / \mathrm{Fe}}^{\mathbf{0}}=\mathbf{z} \mathbf{V}$

276104 For the given cell;
$\mathrm{Cu}(\mathrm{s})\left \vert\mathrm{Cu}^{2+}\left(\mathrm{C}_{1} \mathrm{M}\right) \ \vert \mathrm{Cu}^{2+}\left(\mathrm{C}_{2} \mathrm{M}\right)\right \vert \mathrm{Cu}(\mathrm{s})$
change in Gibbs energy $(\Delta G)$ is negative, if

276100 Given, $\mathrm{E}_{\mathrm{Cr}^{3+} / \mathrm{Cr}}^{0}=0.72 \mathrm{~V}$,
$\mathrm{E}_{\mathrm{Fe}^{2+} / \mathrm{Fe}}^{0}=-0.42 \mathrm{~V}$
The potential for the cell $\mathrm{Cr} \mid \mathrm{Cr}^{3+}(0.1 \mathrm{M}) \ \vert \mathrm{Fe}^{2+}$ $(0.01 \mathrm{M}) \mid \mathrm{Fe}$ is

276101 The standard electrode potential $E^{\circ}$ and its temperature coefficient $\left(\frac{d E^{0}}{d T}\right)$ for a cell are $2 V$ and $-5 \times 10^{-4} \mathrm{VK}^{-1}$ at $300 \mathrm{~K}$ respectively. The cell reaction is $\mathrm{Zn}(\mathrm{s})+\mathrm{Cu}^{2+}(\mathrm{aq}) \rightarrow \mathrm{Zn}^{2+}(\mathrm{aq})+$ $\mathrm{Cu}(\mathrm{s})$. The standard reaction enthalpy $\left(\Delta_{\mathrm{r}} \mathrm{H}^{\ominus}\right)$ at $300 \mathrm{~K}$ in kJ mol ${ }^{-1}$ is, [Use, $R=8 \mathrm{JK}^{-1} \mathrm{~mol}^{-1}$ and $F=96,000 \mathrm{C} \mathrm{mol}^{-1}$ ]

276102 Calculate the standard cell potential (in V) of the cell in which following reaction takes place
$\mathrm{Fe}^{2+}(\mathrm{aq})+\mathrm{Ag}^{+}(\mathrm{aq}) \rightarrow \mathrm{Fe}^{3+} \text { (aq) }+\mathrm{Ag}(\mathrm{s})$
Given that, $\mathbf{E}_{{ }_{\mathrm{Ag}+/ \mathrm{Ag}}}^{\mathbf{0}}=\mathrm{x}$ V,
$\mathbf{E}_{\mathrm{Fe}^{2+} / \mathrm{Fe}}^{\mathbf{0}}=\mathbf{y} \mathbf{V}$
$\mathbf{E}_{\mathrm{Fe}^{3+} / \mathrm{Fe}}^{\mathbf{0}}=\mathbf{z} \mathbf{V}$

276104 For the given cell;
$\mathrm{Cu}(\mathrm{s})\left \vert\mathrm{Cu}^{2+}\left(\mathrm{C}_{1} \mathrm{M}\right) \ \vert \mathrm{Cu}^{2+}\left(\mathrm{C}_{2} \mathrm{M}\right)\right \vert \mathrm{Cu}(\mathrm{s})$
change in Gibbs energy $(\Delta G)$ is negative, if