Chapter 17

For the disproportionation reaction $$ 2 \mathrm{Cu}^{+}(\mathrm{aq}) \rightleftharpoons \mathrm{Cu}(\mathrm{s})+\mathrm{Cu}^{2+}(\mathrm{aq}) \text { at } 298 \mathrm{~K}, \ln \mathrm{K} $$ (where \(\mathrm{K}\) is the equilibrium constant) is \(\times 10^{-1}\). Given [Main Sep. 02, 2020 (II)] $$ \left(\mathrm{E}_{\mathrm{Cu}^{2+} / \mathrm{Cu}^{+}}^{0}=0.16 \mathrm{~V} ; \mathrm{E}_{\mathrm{Cu}^{+} / \mathrm{Cu}}^{0}=0.52 \mathrm{~V} ; \frac{\mathrm{RT}}{\mathrm{F}}=0.025\right) $$

An aqueous solution of \(X\) is added slowly to an aqueous solution of \(\mathrm{Y}\) as shown in List \(\mathrm{I}\). The variation in conductivity of these reactions is given in List II. Match list I with List II and select the correct answer using the code given below the lists: [Adv. 2013] List I List II P. \(\left(\mathrm{C}_{2} \mathrm{H}_{5}\right)_{3} \mathrm{~N}+\mathrm{CH}_{3} \mathrm{COOH}\) 1\. Conductivity \(\mathrm{X} \quad \mathrm{Y}\) decreases and then increases Q. KI \((0.1 \mathrm{M})+\mathrm{AgNO}_{3}(0.01 \mathrm{M}) \quad 2\) Conductivity. \(\begin{array}{lll}\mathrm{X} & \mathrm{Y} & \text { decreases and then }\end{array}\) does not change much R. \(\quad \mathrm{CH}_{3} \mathrm{COOH}+\mathrm{KOH}\) 3\. Conductivity X \(\mathrm{Y} \quad\) increases and then does not change much S. \(\mathrm{NaOH}+\mathrm{HI}\) 4\. Conductivity does not X \(Y \quad\) change much and then increases Codes : \(\mathrm{P}\) \(\mathrm{Q}\) \(\mathrm{R} \quad \mathrm{S}\) \(\mathbf{Q}^{\prime}\) 4 3 1 \(\begin{array}{llll}\text { (a) } & 3 & 4 & 2 \\ \text { (b) } & 4 & 3 & 2 \\\ \text { (c) } & 2 & 3 & 2\end{array}\) 1 3 3 4 4 1 (d) 1 \(3 \quad 2\)

Match the following, choosing one item from column \(\mathbf{X}\) and one from column Y. [Multiple Concepts, 1982 - 2 Marks] \(\mathbf{X} \quad \mathbf{Y}\) (i) neutrons (p) Kohlrausch (ii) molecular speed (q) van der Waals (iii) intermolecular forces (r) Maxwell (iv) conductance of ions (s) Chadwick

Consider an electrochemical cell: \(A(\mathrm{~s})\left|A^{\mathrm{n}+}(\mathrm{aq}, 2 \mathrm{M})\right| B^{2 \mathrm{n}+}(\mathrm{aq}, 1 \mathrm{M}) \mid B(\mathrm{~s})\) The value of \(\Delta H^{\circ}\) for the cell reaction is twice that of \(\Delta G^{\circ}\) at \(300 \mathrm{~K}\). If the emf of the cell is zero, the \(\Delta S^{\circ}\) (in \(\mathrm{J} \mathrm{K}^{-1} \mathrm{~mol}^{-1}\) ) of the cell reaction per mole of \(B\) formed at \(300 \mathrm{~K}\) is (Given: \(\ln (2)=0.7, R\) (universal gas constant) \(=8.3 \mathrm{~J} \mathrm{~K}^{-1} \mathrm{~mol}^{-1} . H, S\) and \(G\) are enthalpy, entropy and Gibbs energy, respectively.)

Two students use same stock solution of \(\mathrm{ZnSO}_{4}\) and a solution of \(\mathrm{CuSO}_{4}\). The emf of one cell is \(0.03 \mathrm{~V}\) higher than the other. The conc. of \(\mathrm{CuSO}_{4}\) in the cell with higher emf value is \(0.5 \mathrm{M}\). Find out the conc. of \(\mathrm{CuSO}_{4}\) in the other cell \((2.203 \mathrm{RT} / \mathrm{F}=0.06\) ).

An excess of liquid mercury is added to an acidified solution of \(1.0 \times\) \(10^{-3} \mathrm{M} \mathrm{Fe}^{3+}\). It is found that \(5 \%\) of \(\mathrm{Fe}^{3+}\) remains at equilibrium at \(25^{\circ} \mathrm{C}\). Calculate \(E_{\mathrm{Hg}_{2}^{2+} \mid \mathrm{Hg}}\), assuming that the only reaction that occurs is \(2 \mathrm{Hg}+2 \mathrm{Fe}^{3+} \longrightarrow \mathrm{Hg}_{2}^{2+}+2 \mathrm{Fe}^{2+}\) (Given \(\left.E_{\mathrm{Fe}^{3+} \mid \mathrm{Fe}^{2+}}^{\circ}=0.77 \mathrm{~V} .\right)\)

How many grams of silver could be plated out on a serving tray by electrolysis of a solution containing silver in \(+1\) oxidation state for a period of \(8.0\) hours at a current of \(8.46\) amperes? What is the area of the tray if the thickness of the silver plating is \(0.00254 \mathrm{~cm}\) ? Density of silver is \(10.5 \mathrm{~g} / \mathrm{cm}^{3}\).

In a galvanic cell, the salt bridge (a) Does not participate chemically in the cell reaction (b) Stops the diffusion of ions from one electrode to another (c) Is necessary for the occurrence of the cell reaction (d) Ensures mixing of the two electrolytic solutions

Chromium metal can be plated out from an acidic solution containing \(\mathrm{CrO}_{3}\) according to the following equation. \(\mathrm{CrO}_{3}(\mathrm{aq})+6 \mathrm{H}^{+}(\mathrm{aq})+6 \mathrm{e}^{-} \rightarrow \mathrm{Cr}(\mathrm{s})+3 \mathrm{H}_{2} \mathrm{O}\) Calculate (i) how many grams of chromium will be plated out by 24,000 coulombs and (ii) how long will it take to plate out \(1.5 \mathrm{~g}\) of chromium by using \(12.5\) amp current.

For the reduction of \(\mathrm{NO}_{3}^{-}\)ion in an aqueous solution, \(E^{\circ}\) is \(+0.96 \mathrm{~V}\). Values of \(E^{\circ}\) for some metal ions are given below \(\mathrm{V}^{2+}(\mathrm{aq})+2 \mathrm{e}^{-} \rightarrow \mathrm{V}\) \(E^{\circ}=-1.19 \mathrm{~V}\) \(\mathrm{Fe}^{3+}(\mathrm{aq})+3 \mathrm{e}^{-} \rightarrow \mathrm{Fe} \quad E^{\circ}=-0.04 \mathrm{~V}\) \(\mathrm{Au}^{3+}(\mathrm{aq})+3 \mathrm{e}^{-} \rightarrow \mathrm{Au}\) \(E^{\circ}=+1.40 \mathrm{~V}\) \(\mathrm{Hg}^{2+}(\mathrm{aq})+2 \mathrm{e}^{-} \rightarrow \mathrm{Hg} \quad E^{\circ}=+0.86 \mathrm{~V}\) The pair(s) of metals that is (are) oxidized by \(\mathrm{NQ}^{-}\)in aqueous solution is (are) [2009] (a) \(\mathrm{V}\) and \(\mathrm{Hg}\) (b) \(\mathrm{Hg}\) and \(\mathrm{Fe}\) (c) \(\mathrm{Fe}\) and \(\mathrm{Au}\) (d) Fe and \(\mathrm{V}\)

The standard reduction potential values of three metallic cations, \(X, Y\) and \(Z\) are \(0.52,-3.03\) and \(-1.18 \mathrm{~V}\) respectively. The order of reducing power of the corresponding metals is [1998 - 2 Marks] (a) \(Y>Z>X\) (b) \(X>Y>Z\) (c) \(Z>Y>X\) (d) \(Z>X>Y\)

The standard reduction potential data at \(25^{\circ} \mathrm{C}\) is given below : [Adv. 2013] \(E^{\circ}\left(\mathrm{Fe}^{3+}, \mathrm{Fe}^{2+}\right)=+0.77 \mathrm{~V} ; E^{\circ}\left(\mathrm{Fe}^{2+}, \mathrm{Fe}\right)=-0.44 \mathrm{~V} ; E^{\circ}\left(\mathrm{Cu}^{2+}, \mathrm{Cu}\right)=+\) \(0.34 \mathrm{~V} ; E^{\circ}\left(\mathrm{Cu}^{+}, \mathrm{Cu}\right)=+0.52 \mathrm{~V}\) \(E^{\circ}\left[\mathrm{O}_{2}(\mathrm{~g})+4 \mathrm{H}^{+}+4 \mathrm{e}^{-} \rightarrow 2 \mathrm{H}_{2} \mathrm{O}\right]=+1.23 \mathrm{~V} ; E^{\circ}\left[\mathrm{O}_{2}(\mathrm{~g})+2 \mathrm{H}_{2} \mathrm{O}+4 \mathrm{e}^{-} \rightarrow\right.\) \(\left.4 \mathrm{OH}^{-}\right]=+0.40 \mathrm{~V}\) \(E^{\circ}\left(\mathrm{Cr}^{3+}, \mathrm{Cr}\right)=-0.74 \mathrm{~V} ; E^{\circ}\left(\mathrm{Cr}^{2+}, \mathrm{Cr}\right)=-0.91 \mathrm{~V}\) Match \(E^{\circ}\) of the redox pair in List I with the values given in List II and select the correct answer using the code given below the lists: List I \(\quad\) List II P. \(E^{\circ}\left(\mathrm{Fe}^{3+}, \mathrm{Fe}\right)\) 1\. \(-0.18 \mathrm{~V}\) Q. \(\quad E^{\circ}\left(4 \mathrm{H}_{2} \mathrm{O} \rightleftharpoons 4 \mathrm{H}^{+}+4 \mathrm{OH}^{-}\right) \quad\) 2. \(-0.4 \mathrm{~V}\) R. \(\quad E^{\circ}\left(\mathrm{Cu}^{2+}+\mathrm{Cu} \rightarrow 2 \mathrm{Cu}^{+}\right)\) 3\. \(-0.04 \mathrm{~V}\) S. \(E^{\circ}\left(\mathrm{Cr}^{3+}, \mathrm{Cr}^{2+}\right)\) 4\. \(-0.83 \mathrm{~V}\) Codes : \(\begin{array}{lllll} & \mathrm{P} & \mathrm{Q} & \mathrm{R} & \mathrm{S} \\\ \text { (a) } & 4 & 1 & 2 & 3\end{array}\) (b) \(2 \quad 3 \quad 4\) (c) \(1 \quad 2 \quad 3 \quad 4\) (d) 3 4 1 2

In a fuel cell, hydrogen and oxygen react to produce electricity. In the process, hydrogen gas is oxidised at the anode and oxygen at the cathode. If \(67.2\) litre of \(\mathrm{H}_{2}\) at STP react in 15 minutes, what is the average current produced? If the entire current is used for electro deposition of copper from copper (II) solution, how many grams of copper will be deposited? [1988 - 4 Marks] Anode reaction : \(\mathrm{H}_{2}+2 \mathrm{OH}^{-} \rightarrow 2 \mathrm{H}_{2} \mathrm{O}+2 \mathrm{e}^{-}\) Cathode reaction : \(\frac{1}{2} \mathrm{O}_{2}+\mathrm{H}_{2} \mathrm{O}+2 \mathrm{e}^{-} \rightarrow 2 \mathrm{OH}^{-}\).

During the discharge of a lead storage battery, the density of sulphuric acid fell from \(1.294\) to \(1.139 \mathrm{~g} / \mathrm{mL}\). Sulphuric acid of density \(1.294\) \(\mathrm{g} / \mathrm{mL}\) is \(39 \%\) by weight and that of \(1.139 \mathrm{~g} / \mathrm{mL}\) is \(20 \% \mathrm{H}_{2} \mathrm{SO}_{4}\) by weight. The battery holds \(3.5\) litres of the acid and the volume remained practically constant during the discharge. Calculate the number of ampere- hours for which the battery must have been used. The charging and discharging reactions are : [1986-5 Marks] Anode : \(\mathrm{Pb}+\mathrm{SO}_{4}^{2-}=\mathrm{PbSO}_{4}+2 \mathrm{e}^{-}\)(discharging) Cathode : \(\mathrm{PbO}_{2}+4 \mathrm{H}^{+}+\mathrm{SO}_{4}^{2-}+2 \mathrm{e}^{-}=\mathrm{PbSO}_{4}+2 \mathrm{H}_{2} \mathrm{O}\) (discharging) Note : Both the reactions take place at the anode and cathode respectively during discharge. Both reaction get reverse during charging.

In an electrolysis experiment, current, was passed for 5 hours through two cells connected in series. The first cell contains a solution of gold and the second contains copper sulphate solution. \(9.85 \mathrm{~g}\) of gold was deposited in the first cell. If the oxidation number of gold is \(+3\), find the amount of copper deposited on the cathode of the second cell. Also calculate the magnitude of the current in amperes. (1 faraday \(=96,500\) coulombs \()\)

Among the following, identify the correct statement. (a) Chloride ion is oxidised by \(\mathrm{O}_{2}\) (b) \(\mathrm{Fe}^{2+}\) is oxidised by iodine (c) Iodide ion is oxidised by chlorine (d) \(\mathrm{Mn}^{2+}\) is oxidised by chlorine

While \(\mathrm{Fe}^{3+}\) is stable, \(\mathrm{Mn}^{3+}\) is not stable in acid solution because (a) \(\mathrm{O}_{2}\) oxideses \(\mathrm{Mn}^{2+}\) to \(\mathrm{Mn}^{3+}\) (b) \(\mathrm{O}_{2}\) oxideses both \(\mathrm{Mn}^{2+}\) to \(\mathrm{Mn}^{3+}\) and \(\mathrm{Fe}^{2+}\) to \(\mathrm{Fe}^{3+}\) (c) \(\mathrm{Fe}^{3+}\) oxideses \(\mathrm{H}_{2} \mathrm{O}\) to \(\mathrm{O}_{2}\) (d) \(\mathrm{Mn}^{3+}\) oxideses \(\mathrm{H}_{2} \mathrm{O}\) to \(\mathrm{O}_{2}\)

Sodium fusion extract, obtained from aniline, on treatment with iron (II) sulphate and \(\mathrm{H}_{2} \mathrm{SO}_{4}\) in presence of air gives a Prussian blue precipitate. The blue colour is due to the formation of (a) \(\mathrm{Fe}_{4}\left[\mathrm{Fe}(\mathrm{CN})_{6}\right]_{3}\) (b) \(\mathrm{Fe}_{3}\left[\mathrm{Fe}(\mathrm{CN})_{6}\right]_{2}\) (c) \(\mathrm{Fe}_{4}\left[\mathrm{Fe}(\mathrm{CN})_{6}\right]_{2}\) (d) \(\mathrm{Fe}_{3}\left[\mathrm{Fe}(\mathrm{CN})_{6}\right]_{3}\)

(a) For the reaction \(\mathrm{Ag}^{+}(\mathrm{aq})+\mathrm{Cl}^{-}(\mathrm{aq}) \rightleftharpoons \mathrm{AgCl}(\mathrm{s})\) Given: \begin{tabular}{|c|c|} \hline Species & \(\Delta G_{f}^{\circ}(\mathrm{kJ} / \mathrm{mol})\) \\ \hline \(\mathrm{Ag}^{+}(\mathrm{aq})\) & \(+77\) \\ \hline \(\mathrm{Cl}^{-}(\mathrm{aq})\) & \(-129\) \\ \hline \(\mathrm{AgCl}(\mathrm{s})\) & \(-109\) \\ \hline \end{tabular} Write the cell representation of above reaction and calculate \(E_{\text {cell }}^{0}\) at \(298 \mathrm{~K}\). Also find the solubility product of \(\mathrm{AgCl}\). (b) If \(6.539 \times 10^{-2} \mathrm{~g}\) of metallic zinc is added to \(100 \mathrm{~mL}\) saturated solution of \(\mathrm{AgCl}\). Find the value of \(\log _{0} \frac{\left[\mathrm{Zn}^{2+}\right]}{\left[\mathrm{Ag}^{+}\right]^{2}}\) How many moles of \(\mathrm{Ag}\) will be precipitated in the above reaction. Given that [2005 - 6 Marks] \(\mathrm{Ag}^{+}+\mathrm{e}^{-} \square \rightarrow \mathrm{Ag} ; E^{\circ}=0.80 \mathrm{~V} ;\) \(\mathrm{Zn}^{2+}+2 \mathrm{e}^{-} \square \rightarrow \mathrm{Zn} ; E^{\circ}=-0.76 \mathrm{~V}\) (It was given that Atomic mass of \(\mathrm{Zn}=65.39\) )

The following electrochemical cell has been set up. \(\mathrm{Pt}(1)\left|\mathrm{Fe}^{3+}, \mathrm{Fe}^{2+}(a=1)\right| \mathrm{Ce}^{4+}, \mathrm{Ce}^{3+}(a=1) \mid \mathrm{Pt}(2)\) \(E^{\circ}\left(\mathrm{Fe}^{3+}, \mathrm{Fe}^{2+}\right)=0.77 \mathrm{~V}: E^{\circ}\left(\mathrm{Ce}^{4+} / \mathrm{Ce}^{3+}\right)=1.61 \mathrm{~V}\) If an ammeter is connected between the two platinum electrodes, predict the direction of flow of current. Will the current increase or decrease with time?

Calculate the equilibrium constant for the reaction \(\mathrm{Fe}^{2+}+\mathrm{Ce}^{4+} \rightleftharpoons \mathrm{Fe}^{3+}+\mathrm{Ce}^{3+}\) [1997 - 2 Marks] (given \(E_{\mathrm{Ce}^{4+} / \mathrm{Ce}^{3+}}^{\circ}=1.44 \mathrm{~V} ; E_{\mathrm{Fe}^{3+} / \mathrm{Fe}^{2+}}^{\circ}=0.68 \mathrm{~V} ;\) )

Although aluminium is above hydrogen in the electrochemical series, it is stable in air and water. Explain.

The Edison storage cells is represented as \(\mathrm{Fe}(\mathrm{s})|\mathrm{FeO}(\mathrm{s})| \mathrm{KOH}(\mathrm{aq})\left|\mathrm{Ni}_{2} \mathrm{O}_{3}(\mathrm{~s})\right| \mathrm{Ni}(\mathrm{s})\) The half-cell reactions are : \(\mathrm{Ni}_{2} \mathrm{O}_{3}(\mathrm{~s})+\mathrm{H}_{2} \mathrm{O}(\mathrm{l})+\rightleftharpoons 2 \mathrm{e}^{-} 2 \mathrm{NiO}_{(s)}+2 \mathrm{OH}^{-}\) \(E^{\circ}=+0.40 \mathrm{~V}\) \(\mathrm{FeO}(\mathrm{s})+\mathrm{H}_{2} \mathrm{O}(1)+2 \mathrm{e}^{-} \rightleftharpoons \mathrm{Fe}(\mathrm{s})+2 \mathrm{OH}^{-} ; E^{0}=-0.87 \mathrm{~V}\) (i) What is the cell reaction? (ii) What is the cell e.m.f ? How does it depend on the concentration of \(\mathrm{KOH}\) ? (iii) What is the maximum amount of electrical energy that can be obtained from one mole of \(\mathrm{Ni}_{2} \mathrm{O}_{3} ?\)

The standard reduction potential of the \(\mathrm{Ag}^{+} / \mathrm{Ag}\) electrode at \(298 \mathrm{~K}\) is \(0.799 \mathrm{~V}\). Given that for AgI, \(K_{s p}=8.7 \times 10^{-17}\), evaluate the potential of the \(\mathrm{Ag}^{+} / \mathrm{Ag}\) electrode in a saturated solution of AgI. Also calculate the standard reduction potential of the \(\mathrm{I}^{-} / \mathrm{AgI} / \mathrm{Ag}\) electrode.

The standard reduction potential for the half-cell \(\mathrm{NO}_{3}^{-}(\mathrm{aq})+2 \mathrm{H}^{+}(\mathrm{aq})+\mathrm{e} \rightarrow \mathrm{NO}_{2}(\mathrm{~g})+\mathrm{H}_{2} \mathrm{O}\) is \(0.78 \mathrm{~V}\) (i) Calculate the reduction potential in \(8 \mathrm{M} \mathrm{H}^{+}\) (ii) What will be the reduction potential of the half-cell in a neutral solution? Assume all the other species to be at unit concentration.

An aqueous solution of \(\mathrm{NaCl}\) on electrolysis gives \(\mathrm{H}_{2}(\mathrm{~g}), \mathrm{Cl}_{2}(\mathrm{~g})\) and \(\mathrm{NaOH}\) according to the reaction : \(2 \mathrm{Cl}^{-}(\mathrm{aq})+2 \mathrm{H}_{2} \mathrm{O}=2 \mathrm{OH}^{-}(\mathrm{aq})+\mathrm{H}_{2}(\mathrm{~g})+\mathrm{Cl}_{2}(\mathrm{~g})\) A direct current of 25 amperes with a current efficiency of \(62 \%\) is passed through 20 litres of \(\mathrm{NaCl}\) solution \((20 \%\) by weight). Write down the reactions taking place at the anode and the cathode. How long will it take to produce \(1 \mathrm{~kg}\) of \(\mathrm{Cl}_{2}\) ? What will be the molarity of the solution with respect to hydroxide ion? (Assume no loss due to evaporation.)

For the galvanic cell. [1992 - 4 Marks] \(\mathrm{Ag}|\mathrm{AgCl}(\mathrm{s}), \mathrm{KCl}(0.2 \mathrm{M}) \| \mathrm{KBr}(0.001 \mathrm{M}), \mathrm{AgBr}(\mathrm{s})| \mathrm{Ag}\) Calculate the EMF generated and assign correct polarity to each electrode for a spontaneous process after taking into account the cell reaction at \(25^{\circ} \mathrm{C}\). \(\left[K_{s p}(\mathrm{AgCl})=2.8 \times 10^{-10} ; K_{s p}(\mathrm{AgBr})=3.3 \times 10^{-13}\right]\)

Zinc granules are added in excess to a \(500 \mathrm{~mL}\). of \(1.0 \mathrm{M}\) nickel nitrate solution at \(25^{\circ} \mathrm{C}\) until the equilibrium is reached. If the standard reduction potential of \(\mathrm{Zn}^{2+} \mid \mathrm{Zn}\) and \(\mathrm{Ni}^{2+} \mid \mathrm{Ni}\) are \(-0.75 \mathrm{~V}\) and \(-0.24 \mathrm{~V}\) respectively, find out the concentration of \(\mathrm{Ni}^{2+}\) in solution at equilibrium.

The standard reduction potential of \(\mathrm{Cu}^{++} / \mathrm{Cu}\) and \(\mathrm{Ag}^{+} / \mathrm{Ag}\) electrodes are \(0.337\) and \(0.799\) volt respectively. Construct a galvanic cell using these electrodes so that its standard e.m.f. is positive. For what concentration of \(\mathrm{Ag}^{+}\)will the e.m.f. of the cell, at \(25^{\circ} \mathrm{C}\), be zero if the concentration of \(\mathrm{Cu}^{++}\)is \(0.01 \mathrm{M}\) ?

A cell contains two hydrogen electrodes. The negative electrode is in contact with a solution of \(10^{-6} \mathrm{M}\) hydrogen ions. The EMF of the cell is \(0.118 \mathrm{~V}\) at \(25^{\circ} \mathrm{C}\). Calculate the concentration of hydrogen ions at the positive electrode.

The EMF of a cell corresponding to the reaction: \(\mathrm{Zn}(\mathrm{s})+2 \mathrm{H}^{+}(\mathrm{aq}) \rightarrow \mathrm{Zn}^{2+}+(0.1 \mathrm{M})+\mathrm{H}_{2}(\mathrm{~g})(1 \mathrm{~atm})\) is \(0.28\) volt at \(25^{\circ} \mathrm{C}\). Write the half-cell reactions and calculate the \(\mathrm{pH}\) of the solution at the hydrogen electrode. \(E_{\mathrm{Zn}^{2+} / \mathrm{Zn}}^{\circ}=-0.76 \mathrm{volt} ; E_{\mathrm{H}^{+} / \mathrm{H}_{2}}^{0}=0\)

Consider the cell [1982 - 2 Marks] \(\mathrm{Zn} \mid \mathrm{Zn}^{2+}\) (aq) \((1.0 \mathrm{M}) \| \mathrm{Cu}^{2+}\) (aq) \((1.0 \mathrm{M}) \mid \mathrm{Cu}\) The standard reduction potentials are : \(+0.350\) volts for \(2 \mathrm{e}^{-}+\mathrm{Cu}^{2+}\) (aq) \(\rightarrow \mathrm{Cu}\) and \(-0.763\) volts for \(2 \mathrm{e}^{-}+\mathrm{Zn}^{2+}(\mathrm{aq})\) \(\rightarrow \mathrm{Zn}\) (i) Write down the cell reaction. (ii) Calculate the emf of the cell. (iii) Is the cell reaction spontaneous or not?