Chapter 17

For the given cell; \(\mathrm{Cu}(\mathrm{s})\left|\mathrm{Cu}^{2+}\left(\mathrm{C}_{1} \mathrm{M}\right) \| \mathrm{Cu}^{2+}\left(\mathrm{C}_{2} \mathrm{M}\right)\right| \mathrm{Cu}(\mathrm{s})\) change in Gibbs energy \((\Delta \mathrm{G})\) is negative, if : (a) \(\mathrm{C}_{1}=\mathrm{C}_{2}\) (b) \(\mathrm{C}_{2}=\frac{\mathrm{C}_{1}}{\sqrt{2}}\) (c) \(\mathrm{C}_{1}=2 \mathrm{C}_{2}\) (d) \(\mathrm{C}_{2}=\sqrt{2} \mathrm{C}_{1}\)

Given that the standard potentials \(\left(\mathrm{E}^{0}\right)\) of \(\mathrm{Cu}^{2+} / \mathrm{Cu}\) and \(\mathrm{Cu}^{+} / \mathrm{Cu}\) are \(0.34\) \(\mathrm{V}\) and \(0.522 \mathrm{~V}\) respectively, the \(\mathrm{E}^{0}\) of \(\mathrm{Cu}^{2+} / \mathrm{Cu}^{+}\)is: (a) \(+0.182 \mathrm{~V}\) (b) \(+0.158 \mathrm{~V}\) (c) \(-0.182 \mathrm{~V}\) (d) \(-0.158 \mathrm{~V}\)

The decreasing order of electrical conductivity of the following aqueous solutions is: [Main April 12, 2019 (II)] \(0.1 \mathrm{M}\) Formic acid (A), \(0.1 \mathrm{M}\) Acetic acid (B), \(0.1 \mathrm{M}\) Benzoic acid (C). (a) \(\mathrm{A}>\mathrm{C}>\mathrm{B}\) (b) \(\mathrm{C}>\mathrm{B}>\mathrm{A}\) (c) \(\mathrm{A}>\mathrm{B}>\mathrm{C}\) (d) \(\mathrm{C}>\mathrm{A}>\mathrm{B}\)

The standard Gibbs energy for the given cell reaction in \(\mathrm{kJ} \mathrm{mol}^{-1}\) at \(298 \mathrm{~K}\) is: [Main April 9, 2019 (I)] \(\mathrm{Zn}(\mathrm{s})+\mathrm{Cu}^{2+}(\mathrm{aq}) \rightarrow \mathrm{Zn}^{2+}(\mathrm{aq})+\mathrm{Cu}(\mathrm{s})\) \(\mathrm{E}^{\circ}=2 \mathrm{~V}\) at \(298 \mathrm{~K}\) (Faraday's constant, \(\left.\mathrm{F}=96000 \mathrm{C} \mathrm{mol}^{-1}\right)\) (a) \(-384\) (b) 384 (c) 192 (d) \(-192\)

Consider the statements S1 and S2: S1 : Conductivity always increases with decrease in the concentration of electrolyte. S2 : Molar conductivity always increases with decrease in the concentration of electrolyte. The correct option among the following is : [Main April 10, 2019 (I)] (a) Both S1 and \(\mathrm{S} 2\) are wrong (b) \(\mathrm{S} 1\) is wrong and \(\mathrm{S} 2\) is correct (c) Both S1 and S2 are correct (d) \(\mathrm{S} 1\) is correct and \(\mathrm{S} 2\) is wrong

A solution of \(\mathrm{Ni}\left(\mathrm{NO}_{3}\right)_{2}\) is electrolysed between platinum electrodes using \(0.1\) Faraday electricity. How many mole of Ni will be deposited at the cathode? (a) \(0.05\) (b) \(0.20\) (c) \(0.15\) (d) \(0.10\)

For the cell \(\mathrm{Zn}(\mathrm{s}) \mid \mathrm{Zn}^{2+}(\mathrm{aq}) \| \mathrm{M}^{\mathrm{x}+}\) (aq) \(\mid \mathrm{M}(\mathrm{s})\), different half cells and their standard electrode potentials are given below: (aq)/ (aq)/ \(\mathrm{Au}^{3+}(\mathrm{aq}) / \mathrm{Ag}^{+}(\mathrm{aq}) / \mathrm{Fe}^{3+}(\mathrm{aq}) / \mathrm{Fe}^{2+}(\mathrm{a}\) \(\mathrm{Au}(\mathrm{s}) \quad \mathrm{Ag}(\mathrm{s}) \quad \mathrm{Fe}^{2+}(\mathrm{aq}) \mathrm{Fe}(\mathrm{s})\) \(\frac{\mathrm{M}^{\mathrm{x}+}(\mathrm{aq}) /}{\mathrm{M}(\mathrm{s})}\) (aq) 1 \(\mathrm{E}_{\mathrm{M}^{\mathrm{x}+} / \mathrm{M}}^{\circ} /(\mathrm{V}) \overline{1.40}\) 1 0 \(\begin{array}{lll}0.80 & 0.77 & -0.44\end{array}\) If \(\mathrm{E}_{\mathrm{Zn}^{2+} / \mathrm{Zn}}^{\circ}=-0.76 \mathrm{~V}\), which cathode will give a maximum value of \(\mathrm{E}_{\mathrm{cell}}^{\circ}\) per electron transferred? [Main Jan. \(11,2019(\mathrm{I})]\) (a) \(\mathrm{Ag}^{+} / \mathrm{Ag}\) (b) \(\mathrm{Fe}^{3+} / \mathrm{Fe}^{2+}\) (c) \(\mathrm{Au}^{3+} / \mathrm{Au}\) (d) \(\mathrm{Fe}^{2+} / \mathrm{Fe}\)

The anodic half-cell of lead-acid battery is recharged using electricity of \(0.05\) Faraday. The amount of \(\mathrm{PbSO}_{4}\) electrolyzed in g during the process is: (Molar mass of \(\mathrm{PbSO}_{4}=303 \mathrm{~g} \mathrm{~mol}^{-1}\) ) [Main Jan. 9, 2019 (I)] (a) \(22.8\) (b) \(15.2\) (c) \(7.6\) (d) \(11.4\)

If the standard electrode potential for a cell is \(2 \mathrm{~V}\) at \(300 \mathrm{~K}\), the equilibrium constant \((\mathrm{K})\) for the reaction \(\mathrm{Zn}(\mathrm{s})+\mathrm{Cu}^{2+}(\mathrm{aq}) \rightleftharpoons \mathrm{Zn}^{2+}(\mathrm{aq})+\mathrm{Cu}(\mathrm{s})\) at \(300 \mathrm{~K}\) is approximately \(\left(\mathrm{R}=8 \mathrm{JK}^{-1} \mathrm{~mol}^{-1}, \mathrm{~F}=96000 \mathrm{C} \mathrm{mol}^{-1}\right)\) [Main Jan. 9, 2019 (II)] (a) \(\mathrm{e}^{-80}\) (b) \(\mathrm{e}^{-160}\) (c) \(\mathrm{e}^{320}\) (d) \(\mathrm{e}^{160}\)

What is the standard reduction potential \(\left(\mathrm{E}^{\circ}\right)\) for \(\mathrm{Fe}^{3+} \rightarrow \mathrm{Fe}\) ? Given that : [Main Online April 8, 2017] \(\mathrm{Fe}^{2+}+2 \mathrm{e}^{-} \rightarrow \mathrm{Fe} ; \mathrm{E}_{\mathrm{Fe}^{2+} / \mathrm{Fe}}^{\circ}=-0.47 \mathrm{~V}\) \(\mathrm{Fe}^{3+}+\mathrm{e}^{-} \rightarrow \mathrm{Fe}^{2+} ; \mathrm{E}_{\mathrm{Fe}^{3+} / \mathrm{Fe}^{2+}}=+0.77 \mathrm{~V}\) (a) \(-0.057 \mathrm{~V}\) (b) \(+0.057 \mathrm{~V}\) (c) \(+0.30 \mathrm{~V}\) (d) \(-0.30 \mathrm{~V}\)

How long (approximate) should water be electrolysed by passing through 100 amperes current so that the oxygen released can completely burn \(27.66 \mathrm{~g}\) of diborane? (Atomic weight of \(\mathrm{B}=10.8 \mathrm{u}\) ) [Main 2018] (a) \(6.4\) hours (b) \(0.8\) hours (c) \(3.2\) hours (d) \(1.6\) hours

For the following cell, \(\mathrm{Zn}(\mathrm{s})\left|\mathrm{ZnSO}_{4}(\mathrm{aq}) \| \mathrm{CuSO}_{4}(\mathrm{aq})\right| \mathrm{Cu}(\mathrm{s})\) when the concentration of \(\mathrm{Zn}^{2+}\) is 10 times the concentration of \(\mathrm{Cu}^{2+}\), the expression for \(\Delta G\) (in \(\mathrm{J} \mathrm{mol}^{-1}\) ) is \([F\) is Faraday constant; \(R\) is gas constant; \(T\) is temperature; \(E^{0}\) (cell) \(\left.=1.1 \mathrm{~V}\right]\) [Adv. 2017] (a) \(1.1 \mathrm{~F}\) (b) \(2.303 R T-2.2 F\) (c) \(2.303 R T+1.1 F\) (d) \(-2.2 F\)

When an electric current is passes through acidified water, \(112 \mathrm{~mL}\) of hydrogen gas at N.T.P. was collected at the cathode in 965 seconds. The current passed, in ampere, is : [Main Online April 15, 2018 (I)] (a) \(2.0\) (b) \(0.1\) (c) \(0.5\) (d) \(1.0\)

Galvanization is applying a coating of: (a) \(\mathrm{Cu}\) (b) \(\mathrm{Zn}\) (c) \(\mathrm{Pb}\) (d) \(\mathrm{Cr}\)

Given \(\mathrm{E}_{\mathrm{Cl}_{2} / \mathrm{Cl}^{-}}=1.36 \mathrm{~V}, \mathrm{E}_{\mathrm{Cr}^{3+} / \mathrm{Cr}}^{\circ}=-0.74 \mathrm{~V}\) \(\mathrm{E}_{\mathrm{Cr}_{2} / \mathrm{O}_{7}^{2-} / \mathrm{Cr}^{3+}}=1.33 \mathrm{~V}, \mathrm{E}_{\mathrm{MnO}_{4}^{\circ} / \mathrm{Mn}^{2+}}=1.51 \mathrm{~V}\) Among the following, the strongest reducing agent is (a) \(\mathrm{Cr}\) (b) \(\mathrm{Mn}^{2+}\) (c) \(\mathrm{Cr}^{3+}\) (d) \(\mathrm{Cl}^{-}\)

Identify the correct statement : [Main Online April 10, 2016] (a) Corrosion of iron can be minimized by forming a contact with another metal with a higher reduction potential (b) Iron corrodes in oxygen free water (c) Corrosion of iron can be minimized by forming an impermeable barrier at its surface (d) Iron corrodes more rapidly in salt water because its electrochemical potential is higher

What will occur if a block of copper metal is dropped into a beaker containing a solution of \(1 \mathrm{M} \mathrm{ZnSO}_{4}\) ? [Main Online April 9, 2016] (a) The copper metal will dissolve with evolution of oxygen gas (b) The copper metal will dissolve with evolution of hydrogen gas (c) No reaction will occur (d) The copper metal will dissolve and zinc metal will be deposited.

For the following electrochemical cell at \(298 \mathrm{~K}\), \(\operatorname{Pt}(\mathrm{s})\left|\mathrm{H}_{2}(\mathrm{~g}, 1 \mathrm{bar})\right| \mathrm{H}^{+}(\mathrm{aq}, 1 \mathrm{M}) \| M^{4+}(\mathrm{aq}), M^{2+}(\mathrm{aq}) \mid \mathrm{Pt}(\mathrm{s})\) \(E_{\mathrm{ccll}}=0.092 \mathrm{~V}\) when \(\frac{\left[M^{2+}(\mathrm{aq})\right]}{\left[M^{4+}(\mathrm{aq})\right]}=10^{x} .\) [Adv. 2016] Given : \(E_{M^{4+} / M^{2+}}^{\circ}=0.151 \mathrm{~V} ; 2.303 \frac{R T}{F}=0.059 \mathrm{~V}\) The value of \(x\) is (a) \(-2\) (b) \(-1\) (c) 1 (d) 2

Two Faraday of electricity is passed through a solution of \(\mathrm{CuSO}_{4}\). The mass of copper deposited at the cathode is (at. mass of \(\mathrm{Cu}=63.5 \mathrm{amu}\) ) [Main 2015] (a) \(2 \mathrm{~g}\) (b) \(127 \mathrm{~g}\) (c) \(0 \mathrm{~g}\) (d) \(63.5 \mathrm{~g}\)

Given below are the half-cell reactions: $$ \begin{aligned} &\mathrm{Mn}^{2+}+2 \mathrm{e}^{-} \rightarrow \mathrm{Mn} ; \mathrm{E}^{\circ}=-1.18 \mathrm{~V} \\ &2\left(\mathrm{Mn}^{3+}+\mathrm{e}^{-} \rightarrow \mathrm{Mn}^{2+}\right) ; \mathrm{E}^{\circ}=+1.51 \mathrm{~V} \end{aligned} $$ The \(\mathrm{E}^{\circ}\) for \(3 \mathrm{Mn}^{2+} \rightarrow \mathrm{Mn}+2 \mathrm{Mn}^{3+}\) will be: (a) \(-2.69 \mathrm{~V} ;\) the reaction will not occur (b) \(-2.69 \mathrm{~V}\); the reaction will occur (c) \(-0.33 \mathrm{~V}\); the reaction will not occur (d) \(-0.33 \mathrm{~V}\); the reaction will occur

A variable, opposite external potential \(\left(\mathrm{E}_{\text {ex }}\right.\) ) is applied to the cell \(\mathrm{Zn} \mathrm{Zn}^{2+}(1 \mathrm{M}) \| \mathrm{Cu}^{2+}(1 \mathrm{M}) \mid \mathrm{Cu}\), of potential \(1.1 \mathrm{~V}\). When \(\mathrm{E}_{\mathrm{ext}}<1.1 \mathrm{~V}\) and \(\mathrm{E}_{\text {ext }}>1.1 \mathrm{~V}\), respectively electrons flow from : [Main Online April 10, 2015] (a) anode to cathode in both cases (b) cathode to anode and anode to cathode (c) anode to cathode and cathode to anode (d) cathode to anode in both cases

Given : \(E_{\mathrm{Cr}^{3+} / \mathrm{Cr}}^{\circ}=-0.74 \mathrm{~V} ; \mathrm{E}_{\mathrm{MnO}_{4}^{-} / \mathrm{Mn}^{2+}}^{\circ}=1.51 \mathrm{~V}\) \(\mathrm{E}_{\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-} / \mathrm{Cr}^{3+}}=1.33 \mathrm{~V} ; \mathrm{E}_{\mathrm{Cl} / \mathrm{Cl}^{-}}^{\circ}=1.36 \mathrm{~V}\) Based on the data given above, strongest oxidising agent will be : [Main 2013] (a) \(\mathrm{Cl}\) (b) \(\mathrm{Cr}^{3+}\) (c) \(\mathrm{Mn}^{2+}\) (d) \(\mathrm{Mn} \mathrm{O}_{4}\)

Resistance of \(0.2 \mathrm{M}\) solution of an electrolyte is \(50 \Omega\). The specific conductance of the solution is \(1.4 \mathrm{~S} \mathrm{~m}^{-1}\). The resistance of \(0.5 \mathrm{M}\) solution of the same electrolyte is \(280 \Omega\). The molar conductivity of \(0.5 \mathrm{M}\) solution of the electrolyte in \(\mathrm{S} \mathrm{m}^{2} \mathrm{~mol}^{-1}\) is: [Main 2014] (a) \(5 \times 10^{-4}\) (b) \(5 \times 10^{-3}\) (c) \(5 \times 10^{3}\) (d) \(5 \times 10^{2}\)

Consider the following cell reaction: [2011] \(2 \mathrm{Fe}(\mathrm{s})+\mathrm{O}_{2}(\mathrm{~g})+4 \mathrm{H}^{+}(\mathrm{aq}) \rightarrow 2 \mathrm{Fe}^{2+}(\mathrm{aq})+2 \mathrm{H}_{2} \mathrm{O}(\mathrm{l}) ; E^{\circ}=1.67 \mathrm{~V}\) At \(\left[\mathrm{Fe}^{2+}\right]=10^{-3} \mathrm{M}, \mathrm{P}\left(\mathrm{O}_{2}\right)=0.1\) atm and \(\mathrm{pH}=3\), the cell potential at \(25^{\circ} \mathrm{C}\) is (a) \(1.47 \mathrm{~V}\) (b) \(1.77 \mathrm{~V}\) (c) \(1.87 \mathrm{~V}\) (d) \(1.57 \mathrm{~V}\)

The rusting of iron takes place as follows \(2 \mathrm{H}^{+}+2 \mathrm{e}^{-}+1 / 2 \mathrm{O}_{2} \longrightarrow \mathrm{H}_{2} \mathrm{O}(1) ; E^{\circ}=+1.23 \mathrm{~V}\) \(\mathrm{Fe}^{2+}+2 \mathrm{e}^{-} \longrightarrow \mathrm{Fe}(\mathrm{s}) ; E^{\circ}=-0.44 \mathrm{~V}\) Calculate \(\Delta G^{\circ}\) for the net process (a) \(-322 \mathrm{~kJ} \mathrm{~mol}^{-1}\) (b) \(-161 \mathrm{~kJ} \mathrm{~mol}^{-1}\) (c) \(-152 \mathrm{~kJ} \mathrm{~mol}^{-1}\) (d) \(-76 \mathrm{~kJ} \mathrm{~mol}^{-1}\)

The standard electrode potentials \(\left(\mathrm{E}^{\circ} \mathrm{M}^{+} / \mathrm{M}\right)\) of four metals \(\mathrm{A}, \mathrm{B}, \mathrm{C}\) and \(\mathrm{D}\) are \(-1.2 \mathrm{~V}, 0.6 \mathrm{~V}, 0.85 \mathrm{~V}\) and \(-0.76 \mathrm{~V}\), respectively. The sequence of deposition of metals on applying potential is: [Main Online April 9, 2014] (a) \(A, C, B, D\) (b) \(\mathrm{B}, \mathrm{D}, \mathrm{C}, \mathrm{A}\) (c) \(\mathrm{C}, \mathrm{B}, \mathrm{D}, \mathrm{A}\) (d) \(\mathrm{D}, \mathrm{A}, \mathrm{B}, \mathrm{C}\)

Given \(\mathrm{Fe}^{3+}(\mathrm{aq})+\mathrm{e}^{-} \rightarrow \mathrm{Fe}^{2+}(\mathrm{aq}) ; \mathrm{E}^{\circ}=+0.77 \mathrm{~V}\) \(\mathrm{Al}^{3+}(\mathrm{aq})+3 \mathrm{e}^{-} \rightarrow \mathrm{Al}(\mathrm{s}) ; \mathrm{E}^{\circ}=-1.66 \mathrm{~V}\) \(\mathrm{Br}_{2}(\mathrm{aq})+2 \mathrm{e}^{-} \rightarrow 2 \mathrm{Br} ; \mathrm{E}^{\circ}=+1.09 \mathrm{~V}\) Considering the electrode potentials, which of the following represents the correct order of reducing power? [Main Online April 11, 2014] (a) \(\mathrm{Fe}^{2+}<\mathrm{Al}<\mathrm{Br}\) (b) \(\mathrm{Br}^{-<\mathrm{Fe}^{2+}<\mathrm{Al}}\) (c) \(\mathrm{Al}<\mathrm{Br}^{-}<\mathrm{Fe}^{2+}\) (d) \(\mathrm{Al}<\mathrm{Fe}^{2+}<\mathrm{Br}\)

In the electrolytic cell, flow of electrons is from (a) Cathode to anode in solution (b) Cathode to anode through external supply (c) Cathode to anode through internal supply (d) Anode to cathode through internal supply

How many electrons would be required to deposit \(6.35 \mathrm{~g}\) of copper at the cathode during the electrolysis of an aqueous solution of copper sulphate? (Atomic mass of copper \(=63.5 \mathrm{u}, \mathrm{N}_{\mathrm{A}}=\) Avogadro's constant): (a) \(\frac{\mathrm{N}_{\mathrm{A}}}{20}\) (b) \(\frac{\mathrm{N}_{\mathrm{A}}}{10}\) (c) \(\frac{\mathrm{N}_{\mathrm{A}}}{5}\) (d) \(\frac{\mathrm{N}_{\mathrm{A}}}{2}\)

A solution of copper sulphate \(\left(\mathrm{CuSO}_{4}\right)\) is electrolysed for 10 minutes with a current of \(1.5\) amperes. The mass of copper deposited at the cathode (at. mass of \(\mathrm{Cu}=63 \mathrm{u}\) ) is: [Main Online April 25, 2013] (a) \(0.3892 \mathrm{~g}\) (b) \(0.2938 \mathrm{~g}\) (c) \(0.2398 \mathrm{~g}\) (d) \(0.3928 \mathrm{~g}\)

For the electrochemical cell, \(M\left|M^{+} \| X^{-}\right| X, E^{\circ}\left(M^{+} / M\right)=0.44 \mathrm{~V}\) and \(E^{\circ}\left(X / X^{-}\right)=0.33 \mathrm{~V}\) From this data one can deduce that [2000S] (a) \(M+X \rightarrow M^{+}+X^{-}\)is the spontaneous reaction (b) \(M^{+}+X^{-} \rightarrow M+X\) is the spontaneous reaction (c) \(\mathrm{E}_{\text {cell }}=0.77 \mathrm{~V}\) (d) \(\mathrm{E}_{\mathrm{ccll}}=-0.77 \mathrm{~V}\)

Electrolysis of dilute aqueous \(\mathrm{NaCl}\) solution was carried out by passing 10 milli ampere current. The time required to liberate \(0.01\) mol of \(\mathrm{H}_{2}\) gas at the cathode is \(\left(1\right.\) Faraday \(\left.=96500 \mathrm{C} \mathrm{mol}^{-1}\right)\) [2008S] (a) \(9.65 \times 10^{4} \mathrm{sec}\) (b) \(19.3 \times 10^{4} \mathrm{sec}\) (c) \(28.95 \times 10^{4} \mathrm{sec}\) (d) \(38.6 \times 10^{4} \mathrm{sec}\)

A gas \(X\) at 1 atm is bubbled through a solution containing a mixture of \(1 \mathrm{M} Y^{-}\)and \(1 \mathrm{M} Z^{-}\)at \(25^{\circ} \mathrm{C}\). If the reduction potential of \(Z>Y>X\) then, [1999 - 2 Marks] (a) \(Y\) will oxidize \(X\) and \(\operatorname{not} Z\) (b) \(Y\) will oxidize \(Z\) and \(\operatorname{not} X\) (c) \(Y\) will oxidize both \(X\) and \(Z\) (d) \(Y\) will reduce both \(X\) and \(Z\)

The correct order of equivalent conductance at infinite dilution of \(\mathrm{LiCl}, \mathrm{NaCl}\) and \(\mathrm{KCl}\) is [2001S] (a) \(\mathrm{LiCl}>\mathrm{NaCl}>\mathrm{KCl}\) (b) \(\mathrm{KCl}>\mathrm{NaCl}>\mathrm{LiCl}\) (c) \(\mathrm{NaCl}>\mathrm{KCl}>\mathrm{LiCl}\) (d) \(\mathrm{LiCl}>\mathrm{KCl}>\mathrm{NaCl}\)

The standard reduction potentials of \(\mathrm{Cu}^{2+} \mid \mathrm{Cu}\) and \(\mathrm{Cu}^{2+} \mid \mathrm{Cu}^{+}\)are \(0.337\) \(\mathrm{V}\) and \(0.153\) respectively. The standard electrode potential of \(\mathrm{Cu}^{+} \mid \mathrm{Cu}\) half cell is [1997-1 Mark] (a) \(0.184 \mathrm{~V}\) (b) \(0.827 \mathrm{~V}\) (c) \(0.521 \mathrm{~V}\) (d) \(0.490 \mathrm{~V}\)

The electric charge for electrode deposition of one gram equivalent of a substance is : [1984-1 Mark] (a) one ampere per second. (b) 96,500 coloumbs per second. (c) one ampere for one hour. (d) charge on one mole of electrons.

A dilute aqueous solution of \(\mathrm{Na}_{2} \mathrm{SO}_{4}\) is electrolyzed using platinum electrodes. The products at the anode and cathode are: [1996-1 Mark] (a) \(\mathrm{O}_{2}, \mathrm{H}_{2}\) (b) \(\mathrm{S}_{2} \mathrm{O}_{8}^{2-}, \mathrm{Na}\) (c) \(\mathrm{O}_{2}, \mathrm{Na}\) (d) \(\mathrm{S}_{2} \mathrm{O}_{8}^{2-}, \mathrm{H}_{2}\)

Faraday's laws of electrolysis are related to the (a) atomic number of the reactants. (b) atomic number of the anion. (c) equivalent weight of the electrolyte. (d) speed of the cation.

A solution of sodium sulphate in water is electrolysed using inert electrodes. The products at the cathode and anode are respectively [1987 - 1 Mark] (a) \(\mathrm{H}_{2}, \mathrm{O}_{2}\) (b) \(\mathrm{O}_{2}, \mathrm{H}_{2}\) (c) \(\mathrm{O}_{2}, \mathrm{Na}\) (d) \(\mathrm{O}_{2}, \mathrm{SO}_{2}\)

The molar conductivity of a solution of a weak acid \(\mathrm{H} X(0.01 \mathrm{M})\) is 10 times smaller than the molar conductivity of a solution of a weak acid \(\mathrm{H} Y(0.10 \mathrm{M})\). If \(\lambda_{\mathrm{X}^{-}}^{0} \approx \lambda_{\mathrm{Y}}^{0}\) the difference in their \(\mathrm{p} K_{a}\) values, \(\mathrm{p} K_{a}(\mathrm{H} X)\) \(-\mathrm{p} K_{a}(\mathrm{H} Y)\), is (consider degree of ionization of both acids to be \(<<1\) )

The reaction : \(1 / 2 \mathrm{H}_{2}(\mathrm{~g})+\mathrm{AgCl}(\mathrm{s}) \rightarrow \mathrm{H}^{+}(\mathrm{aq})+\mathrm{Cl}^{-}(\mathrm{aq})+\mathrm{Ag}(\mathrm{s})\) occurs in the galvanic cell (a) \(\mathrm{Ag}|\mathrm{AgCl}(\mathrm{s})| \mathrm{KCl}(\mathrm{soln})\left|\mathrm{AgNO}_{3}(\mathrm{soln})\right| \mathrm{Ag}\) (b) \(\mathrm{Pt}\left|\mathrm{H}_{2}(\mathrm{~g})\right| \mathrm{HCl}(\mathrm{soln})\left|\mathrm{AgNO}_{3}(\mathrm{soln})\right| \mathrm{Ag}\) (c) \(\mathrm{Pt}\left|\mathrm{H}_{2}(\mathrm{~g})\right| \mathrm{HCl}(\mathrm{soln})|\mathrm{AgCl}(\mathrm{s})| \mathrm{Ag}\) (d) \(\mathrm{Pt}\left|\mathrm{H}_{2}(\mathrm{~g})\right| \mathrm{KCl}(\mathrm{soln})|\mathrm{AgCl}(\mathrm{s})| \mathrm{Ag}\)

Potassium chlorate is prepared by the electrolysis of \(\mathrm{KCl}\) in basic solution \(6 \mathrm{OH}^{-}+\mathrm{Cl}^{-} \rightarrow \mathrm{ClO}_{3}+3 \mathrm{H}_{2} \mathrm{O}+6 \mathrm{e}^{-}\) If only \(60 \%\) of the current is utilized in the reaction, the time (rounded to the nearest hour) required to produce \(10 \mathrm{~g}\) of \(\mathrm{KClO}_{3}\) using a current of \(2 \mathrm{~A}\) is (Given : \(\mathrm{F}=\overline{96,500 \mathrm{C}} \mathrm{mol}^{-1} ;\) molar mass of \(\mathrm{KClO}_{3}=122 \mathrm{~g} \mathrm{~mol}^{-1}\) )

A solution containing one mole per litre of each \(\mathrm{Cu}\left(\mathrm{NO}_{3}\right)_{2} ; \mathrm{AgNO}_{3}\); \(\mathrm{Hg}_{2}\left(\mathrm{NO}_{3}\right)_{2}\); is being electrolysed by using inert electrodes. The values of standard electrode potentials in volts (reduction potentials) are : [1984-1 Mark] \(\mathrm{Ag} / \mathrm{Ag}^{+}=+0.80,2 \mathrm{Hg} / \mathrm{Hg}_{2}^{++}=+0.79\) \(\mathrm{Cu} / \mathrm{Cu}^{++}=+0.34, \mathrm{Mg} / \mathrm{Mg}^{++}=-2.37\) With increasing voltage, the sequence of deposition of metals on the cathode will be : (a) \(\mathrm{Ag}, \mathrm{Hg}, \mathrm{Cu}, \mathrm{Mg}\) (b) \(\mathrm{Mg}, \mathrm{Cu}, \mathrm{Hg}, \mathrm{Ag}\) (c) \(\mathrm{Ag}, \mathrm{Hg}, \mathrm{Cu}\) (d) \(\mathrm{Cu}, \mathrm{Hg}, \mathrm{Ag}\)

The standard reduction potentials at \(298 \mathrm{~K}\) for the following half reactions are given against each [1981-1 Mark] \(\mathrm{Zn}^{2+}(\mathrm{aq})+2 \mathrm{e} \rightleftharpoons \mathrm{Zn}(\mathrm{s}) \quad-0.762\) \(\mathrm{Cr}^{3+}(\mathrm{aq})+2 \mathrm{e} \rightleftharpoons \mathrm{Cr}(\mathrm{s}) \quad-0.740\) \(2 \mathrm{H}^{+}(\mathrm{aq})+2 \mathrm{e} \rightleftharpoons \mathrm{H}_{2}(\mathrm{~g}) \quad 0.000\) \(\mathrm{Fe}^{3+}(\mathrm{aq})+2 \mathrm{e} \rightleftharpoons \mathrm{Fe}^{2+}(\) aq \() \quad 0.770\) which is the strongest reducing agent? (a) \(\mathrm{Zn}(\mathrm{s})\) (b) \(\mathrm{Cr}(\mathrm{s})\) (c) \(\mathrm{H}_{2}(\mathrm{~g})\) (d) \(\mathrm{Fe}^{2+}(\mathrm{aq})\)

Consider a \(70 \%\) efficient hydrogen-oxygen fuel cell working under standard conditions at 1 bar and \(298 \mathrm{~K}\). Its cell reaction is $$ \mathrm{H}_{2}(g)+\frac{1}{2} \mathrm{O}_{2}(g) \rightarrow \mathrm{H}_{2} \mathrm{O}(l) $$ The work derived from the cell on the consumption of \(1.0 \times 10^{-3} \mathrm{~mol}\) of \(\mathrm{H}_{2}(g)\) is used to compress \(1.00 \mathrm{~mol}\) of a monoatomic ideal gas in a thermally insulated container. What is the change in the temperature (in \(\mathrm{K}\) ) of the ideal gas? The standard reduction potentials for the two half-cells are given below. $$ \begin{aligned} &\mathrm{O}_{2}(g)+4 \mathrm{H}^{+}(a q)+4 e^{-} \rightarrow 2 \mathrm{H}_{2} \mathrm{O}(l), E^{0}=1.23 \mathrm{~V}, \\ &2 \mathrm{H}^{+}(a q)+2 e^{-} \rightarrow \mathrm{H}_{2}(g), E^{0}=0.00 \mathrm{~V} \end{aligned} $$ Use \(F=96500 \mathrm{C} \mathrm{mol}^{-1}, R=8.314 \mathrm{~J} \mathrm{~mol}^{-1} \mathrm{~K}^{-1}\).

For the electrochemical cell, [Adv. 2018] \(\mathrm{Mg}(\mathrm{s})\left|\mathrm{Mg}^{2+}(\mathrm{aq}, 1 \mathrm{M}) \| \mathrm{Cu}^{2+}(\mathrm{aq}, 1 \mathrm{M})\right| \mathrm{Cu}(\mathrm{s})\) the standard emf of the cell is \(2.70 \mathrm{~V}\) at \(300 \mathrm{~K}\). When the concentration of \(\mathrm{Mg}^{2+}\) is changed to \(x \mathrm{M}\), the cell potential changes to \(2.67 \mathrm{~V}\) at \(300 \mathrm{~K}\). The value of \(x\) is (given, \(\frac{F}{R}=11500 \mathrm{~K} \mathrm{~V}^{-1}\), where \(F\) is the Faraday constant and \(R\) is the gas constant, \(\ln 10=2.30\) )

A current of \(1.70 \mathrm{~A}\) is passed through \(300.0 \mathrm{~mL}\) of \(0.160 \mathrm{M}\) solution of a \(\mathrm{ZnSO}_{4}\) for \(230 \mathrm{sec}\). with a current efficiency of \(90 \%\). Find out the molarity of \(\mathrm{Zn}^{2+}\) after the deposition of \(\mathrm{Zn}\). Assume the volume of the solution to remain constant during the electrolysis.

All the energy released from the reaction \(X \rightarrow Y, \Delta_{1} G^{\circ}=-193 \mathrm{~kJ} \mathrm{~mol}^{-1}\) is used for oxidizing \(M^{*}\) as \(M^{+} \rightarrow M^{3+}+2 \mathrm{e}^{-}, E^{\circ}=-0.25 \mathrm{~V}\) Under standard conditions, the number of moles of \(M^{+}\)oxidized when one mole of \(X\) is converted to \(Y\) is \(\left[\mathrm{F}=96500 \mathrm{C} \mathrm{mol}^{-1}\right]\)

An oxidation-reduction reaction in which 3 electrons are transferred has a \(\Delta \mathrm{G}^{0}\) of \(17.37 \mathrm{~kJ} \mathrm{~mol}^{-1}\) at \(25^{\circ} \mathrm{C}\). The value of \(\mathrm{E}_{\text {cell }}^{0}\) (in \(\mathrm{V}\) ) is \(\times 10^{-2}\) \(\left(1 \mathrm{~F}=96,500 \mathrm{C} \mathrm{mol}^{-1}\right)\)

The Gibbs energy change (in \(\mathrm{J}\) ) for the given reaction at \(\left[\mathrm{Cu}^{2+}\right]=\) \(\left[\mathrm{Sn}^{2+}\right]=1 \mathrm{M}\) and \(298 \mathrm{~K}\) is : \(\mathrm{Cu}(\mathrm{s})+\mathrm{Sn}^{2+}(\mathrm{aq}) \rightarrow \mathrm{Cu}^{2+}(\mathrm{aq})+\mathrm{Sn}(\mathrm{s})\) \(\left(\mathrm{E}_{\mathrm{Sa}^{2+} \mid \mathrm{Su}}^{0}=-0.16 \mathrm{~V}, \mathrm{E}_{\mathrm{Cu}^{2+} \mid \mathrm{Cu}}^{0}=0.34 \mathrm{~V}\right.\) Take \(\mathrm{F}=96500 \mathrm{C} \mathrm{mol}^{-1}\) )