Chapter 9

Electrolysis of dilute aqueous \(\mathrm{NaCl}\) solution was carried out by passing 10 mili ampere current. The time required to librate \(0.01\) mole of \(\mathrm{H}_{2}\) gas at the cathode? (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) \(38.6 \times 10^{4} \mathrm{Sec}\)

The standard reduction potentials at \(298 \mathrm{~K}\) for the following half- reactions are given against each \(\mathrm{Zn}^{21}(\mathrm{aq})+2 \mathrm{e}=\mathrm{Zn}(\mathrm{s})-0.762\) \(\mathrm{Cr}^{3+}(\mathrm{aq})+2 \mathrm{e} \cdots \mathrm{Cr}(\mathrm{s}) \quad-0.740\) \(2 \mathrm{H}^{+}(\mathrm{aq})+2 \mathrm{e} \rightleftharpoons \mathrm{H}_{2}(\mathrm{~g}) \quad 0.000\) \(\mathrm{Fe}^{3+}(\mathrm{aq})+2 \mathrm{e}=\mathrm{Fe}^{2+}\) (aq) \(0.770\) Which is the strongest reducing agent? (a) \(\mathrm{H}_{2}(\mathrm{~g})\) (b) \(\mathrm{Cr}(\mathrm{s})\) (c) \(\mathrm{Zn}(\mathrm{s})\) (d) \(\mathrm{Fe}^{2+}(\mathrm{aq})\)

\(\mathrm{Ag}\left|\mathrm{Ag}^{+}(\mathrm{IM}) \| \mathrm{Ag}^{+}(2 \mathrm{M})\right| \mathrm{Ag}\) 1 L solution 1 L solution \(0.5 \mathrm{~F}\) of electricity in the LHS (anode) the \(1 F\) electricity in the RHS (cathode) is first passed making them independent electro cells at \(298 \mathrm{~K}\). The emf of the cell after electrolysis will (a) increase (b) decrease (c) not change (d) time is also required

One litre of \(1 \mathrm{M} \mathrm{CuSO}_{4}\) solution is electrolysed. After passing \(2 \mathrm{~F}\) of electricity, molarity of \(\mathrm{CuSO}_{4}\) solution will be (a) \(\mathrm{M} / 2\) (b) \(\mathrm{M} / 4\) (c) \(\mathrm{M}\) (d) 0

In acidic medium \(\mathrm{MnO}_{4}^{-}\)is an oxidizing agent \(\mathrm{MnO}_{4}^{-}+8 \mathrm{H}^{+}+5 \mathrm{e}^{-} \longrightarrow \mathrm{Mn}^{2+}+4 \mathrm{H}_{2} \mathrm{O} .\) If \(\mathrm{H}^{+}\)ion concentration is doubled, electrode potential of the half cell \(\mathrm{MnO}_{4}^{-}, \mathrm{Mn}^{2+} / \mathrm{Pt}\) will (a) increase by \(28.46 \mathrm{mV}\) (b) decrease by \(28.46 \mathrm{mV}\) (c) increase by \(14.23 \mathrm{mV}\) (d) decrease by \(142.30 \mathrm{mV}\)

Calculate the weight of copper that will be deposited at the cathode in the electrolysis of a \(0.2 \mathrm{M}\) solution of copper sulphate, when quantity of electricity, equal to the required to liberate \(2.24 \mathrm{~L}\) of hydrogen at STP from a \(0.1 \mathrm{M}\) aqueous sulphuric acid, is passed (Atomic mass of \(\mathrm{Cu}=63.5\) ) (a) \(6.35 \mathrm{~g}\) (b) \(3.17 \mathrm{~g}\) (c) \(12.71 \mathrm{~g}\) (d) \(63.5 \mathrm{~g}\)

Given that \(E^{\circ} \quad\left(\mathrm{Zn}^{2} / \mathrm{Zn}\right)=-0.763 \mathrm{~V}\) and \(\mathrm{E}^{\circ}\left(\mathrm{Cd}^{2+} / \mathrm{Cd}\right)=-0.403 \mathrm{~V}\), the emf of the following cell \(\mathrm{Zn}\left|\mathrm{Zn}^{21}(\mathrm{a}=0.04) \| \mathrm{Cd}^{2+}(\mathrm{a}=0.2)\right| \mathrm{Cd}\) is given by (a) \(E=+0.36+[0.059 / 2][\log (0.2 / 0.004)]\) (b) \(E=-0.36+[0.059 / 2][\log (0.2 / 0.004)]\) (c) \(E=+0.36+[0.059 / 2][\log (0.004 / 0.2)]\) (d) \(E=-0.36+[0.059 / 2][\log (0.004 / 0.2)]\)

If the pressure of hydrogen gas is increased from 1 arm to \(100 \mathrm{~atm}\), keeping the hydrogen ion concentration constant at \(1 \mathrm{M}\), the voltage of the hydrogen half cell at \(25^{\circ} \mathrm{C}\) will be (a) \(-0.059 \mathrm{~V}\) (b) \(+0.059 \mathrm{~V}\) (c) \(5.09 \mathrm{~V}\) (d) \(0.259 \mathrm{~V}\)

The conductivity of \(0.01 \mathrm{~mol} / \mathrm{dm}^{3}\) aqueous acetic acid at \(300 \mathrm{~K}\) is \(19.5 \times 10^{-5} \mathrm{ohm}^{-1} \mathrm{~cm}^{-1}\) and limiting molar conductivity of acetic acid at the same temperature is \(390 \mathrm{ohm}^{-1} \mathrm{~cm}^{2} \mathrm{~mol}^{-1} .\) The degree of dissociation of acetic acid is (a) \(0.05\) (b) \(0.5 \times 10^{-2}\) (c) \(5 \times 10^{-7}\) (d) \(5 \times 10^{-3}\)

The hydrogen electrode is dipped in a solution of \(\mathrm{pH}=\) \(3.0\) at \(25^{\circ} \mathrm{C}\). The potential of hydrogen electrode would be (a) \(-0.177 \mathrm{~V}\) (b) \(0.177 \mathrm{~V}\) (c) \(1.77 \mathrm{~V}\) (d) \(0.277 \mathrm{~V}\)

Three faraday of electricity is passed through aqueous solutions of \(\mathrm{AgNO}_{3}, \mathrm{NiSO}_{4}\) and \(\mathrm{CrCl}_{3}\) kept in three vessels using inert electrodes. The ratio in moles in which the metals \(\mathrm{Ag}, \mathrm{Ni}\) and \(\mathrm{Cr}\) will be deposited is (a) \(1: 2: 3\) (b) \(2: 3: 6\) (c) \(6: 3: 2\) (d) \(3: 2: 6\)

When an electric current is passed through acidulated water, \(112 \mathrm{~mL}\) of hydrogen gas at NTP collects at the cathode in 965 seconds. The current passed, in ampere is (a) \(0.1\) (b) \(0.5\) (c) \(1.0\) (d) \(2.0\)

For the electrochemical cell, \(\mathrm{M}\left|\mathrm{M}^{+} \| \mathrm{X}^{-}\right| \mathrm{X}\) \(E^{\circ} \mathrm{M}^{+} / \mathrm{M}=0.44 \mathrm{~V}\) and \(E^{\circ} \mathrm{X} / \mathrm{X}^{-}=0.33 \mathrm{~V}\) From these data, one can deduce that (a) \(\mathrm{M}+\mathrm{X} \longrightarrow \mathrm{M}^{\prime}+\mathrm{X}^{-}\)is the spontaneous reaction (b) \(\mathrm{M}^{+}+\mathrm{X}^{-} \longrightarrow \mathrm{M}+\mathrm{X}\) is spontaneous reaction (c) \(E_{\text {cell }}=0.77 \mathrm{~V}\) (d) \(E_{\text {cell }}=-0.77 \mathrm{~V}\)

When a quantity of electricity is passed through \(\mathrm{CuSO}_{4}\) solution, \(0.16 \mathrm{~g}\) of copper gets deposited. If the same quantity of electricity is passed through acidulated water, then the volume of \(\mathrm{H}_{2}\) liberated at STP will be (At. wt of \(\mathrm{Cu}=64\) ) (a) \(4.0 \mathrm{~cm}^{3}\) (b) \(56 \mathrm{~cm}^{3}\) (c) \(604 \mathrm{~cm}^{3}\) (d) \(8.0 \mathrm{~cm}^{3}\)

\(4.5 \mathrm{~g}\) of aluminium (at. mass \(27 \mathrm{amu}\) ) is deposited at cathode from \(\mathrm{Al}^{3+}\) solution by a certain quantity of electric charge. The volume of hydrogen produced at STP from \(\mathrm{H}^{+}\)ions is solution by the same quantity of electric charge will be (a) \(44.8 \mathrm{~L}\) (b) \(22.4 \mathrm{~L}\) (c) \(11.2 \mathrm{~L}\) (d) \(5.6 \mathrm{~L}\)

The half cell reaction for the corrosion \(2 \mathrm{H}^{+}+1 / 2 \mathrm{O}_{2}+2 \mathrm{e}^{-} \longrightarrow \mathrm{H}_{2} \mathrm{O}, E^{\circ}=1.23 \mathrm{~V}\) \(\mathrm{Fe}^{2+}+2 \mathrm{e}^{-} \longrightarrow \mathrm{Fe}(\mathrm{s}) ; E^{\circ}=-0.44 \mathrm{~V}\) Find the \(\Delta \mathrm{G}^{\circ}\) (in \(\mathrm{kJ}\) ) for the overall reaction. (a) \(-76\) (b) \(-322\) (c) \(-161\) (d) \(-152\)

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

Equal quantities of electricity are passed through three voltameters containing \(\mathrm{FeSO}_{4}, \mathrm{Fe}_{2}\left(\mathrm{SO}_{4}\right)_{3}\), and \(\mathrm{Fe}\left(\mathrm{NO}_{3}\right)_{3}\) Consider the following statements in this regard (1) the amount of iron deposited in \(\mathrm{FeSO}_{4}\) and \(\mathrm{Fe}_{2}\left(\mathrm{SO}_{4}\right)_{3}\) is equal (2) the amount of iron deposited in \(\mathrm{Fe}\left(\mathrm{NO}_{3}\right)_{3}\) is two thirds of the amount of iron deposited in \(\mathrm{FeSO}_{4}\) (3) the amount of iron deposited in \(\mathrm{Fe}_{2}\left(\mathrm{SO}_{4}\right)_{3} \mathrm{Fe}\left(\mathrm{NO}_{3}\right)_{3}\) is equal Of these statements (a) 1 and 2 are correct (b) 2 and 3 are correct (c) 1 and 3 are correct (d) 1,2 and 3 are correct

The reversible reduction potential of pure water is \(-0.413 \mathrm{~V}\) under \(1.00 \mathrm{~atm} \mathrm{H}_{2}\) pressure. If the reduction is considered to be \(2 \mathrm{H}^{+}+2 \mathrm{e}^{2} \longrightarrow \mathrm{H}_{2}\), calculate \(\mathrm{pH}\) of pure water. (a) 6 (b) 7 (c) 3 (d) 5

For a \(\mathrm{Ag}-\mathrm{Zn}\) button cell, net reaction is \(\mathrm{Zn}(\mathrm{s})+\mathrm{Ag}_{2} \mathrm{O}(\mathrm{s}) \longrightarrow \mathrm{ZnO}(\mathrm{s})+2 \mathrm{Ag}(\mathrm{s})\) \(\Delta \mathrm{G}_{\mathrm{f}}^{\circ}\left(\mathrm{Ag}_{2} \mathrm{O}\right)=-11.21 \mathrm{~kJ} \mathrm{~mol}^{-1}\) \(\Delta \mathrm{G}_{f}^{\circ}(\mathrm{ZnO})=-318.3 \mathrm{~kJ} \mathrm{~mol}^{-1}\) Hence \(E_{\text {cell }}^{\circ}\) of the button cell is (a) \(3.591 \mathrm{~V}\) (b) \(2.591 \mathrm{~V}\) (c) \(-1.591 \mathrm{~V}\) (d) \(1.591 \mathrm{~V}\)

A current of 15 amp is employed to plate Nickel in a \(\mathrm{NiSO}_{4}\) bath. Both \(\mathrm{Ni}\) and \(\mathrm{H}_{2}\) are formed at the cathode. If \(9.9 \mathrm{~g}\) of \(\mathrm{Ni}\) are deposited with the simultaneous liberation of \(2.51\) litres of \(\mathrm{H}_{2}\) measured at STP, what is the current efficiency for the deposition of Ni? (Atomic weight of \(\mathrm{Ni}=58.7\) ) (a) \(60 \%\) (b) \(70 \%\) (c) \(80 \%\) (d) \(56 \%\)

Four elements \(\mathrm{A}, \mathrm{B}, \mathrm{C}\) and \(\mathrm{D}\) can form diatomic molecules and monoatomic anions with \(-1\) charge. Consider the following reactions about these. \(2 \mathrm{~B}+\mathrm{C}_{2} \longrightarrow 2 \mathrm{C}^{-}+\mathrm{B}_{2}\) \(\mathrm{B}_{2}+2 \mathrm{D} \longrightarrow{\longrightarrow}{\longrightarrow} \mathrm{B}^{-}+\mathrm{D}_{2}\) \(2 \mathrm{~A}^{-}+\mathrm{C}_{2}\) no reaction Select correct statement about these. (1) \(\mathrm{A}_{2}\) is strongest oxidizing agent while \(\mathrm{D}\) is strongest reducing agent (2) \(\mathrm{D}_{2}\) is strongest oxidizing agent while \(\mathrm{A}\) is strongest reducing agent (3) \(\mathrm{C}_{2}\) will oxidize \(\mathrm{B}^{-}\)and also \(\mathrm{D}^{-}\)to form \(\mathrm{B}_{2}\) and \(\mathrm{D}_{2}\) (4) \(\mathrm{E}^{\mathrm{O}} \mathrm{A}_{2} / \mathrm{A}^{-}\)is the lowest (a) 2 and 3 (b) 1 and 3 (c) 2 and 4 (d) 1,2 and 3

In which of the following aqueous solutions during electrolysis, \(\mathrm{H}_{2}\) and \(\mathrm{Cl}_{2}\) are liberated? (a) \(\mathrm{CuCl}_{2}(\mathrm{aq})\) (b) \(\mathrm{KCl}(\mathrm{aq})\) (c) \(\mathrm{MgCl}_{2}(\mathrm{aq})\) (d) \(\mathrm{NaCl}(\mathrm{aq})\)

For the electrolysis of \(\mathrm{CuSO}_{4}\) solution which is/are correct? (a) Cathode reaction: \(\mathrm{Cu}^{2+}+2 \mathrm{e}^{-} \rightarrow \mathrm{Cu}\) using \(\mathrm{Cu}\) electrode (b) Anode reaction: \(\mathrm{Cu} \rightarrow \mathrm{Cu}^{2}++2 \mathrm{e}^{-}\)using \(\mathrm{Cu}\) electrode (c) Cathode reaction: \(2 \mathrm{H}^{+}+2 \mathrm{e}^{-} \rightarrow \mathrm{H}_{2}\) using \(\mathrm{Pt}\) electrode (d) Anode reaction: \(\mathrm{Cu} \rightarrow \mathrm{Cu}^{2+{ }^{2}}+2 \mathrm{e}^{-}\)using \(\mathrm{Pt}\) electrode

Given that \(\mathrm{E}_{\mathrm{N}^{2+} / \mathrm{N}}^{0}=-0.25 \mathrm{~V} ; \mathrm{E}_{\mathrm{Cu}^{2+} / \mathrm{Cu}}^{0}=+0.34 \mathrm{~V}\) \(\mathrm{E}_{\mathrm{Ag}^{*} / \Lambda_{8}}^{0}=+0.80 \mathrm{~V} ; \mathrm{E}_{\mathrm{Zn}^{2+} / Z_{\mathrm{n}}}^{0}=-0.76 \mathrm{~V}\) Which of the following redox processes will not take place in specified direction? (a) \(\mathrm{Zn}(\mathrm{s})+2 \mathrm{H}^{+}(\mathrm{aq}) \rightarrow \mathrm{Zn}^{2+}(\mathrm{aq})+\mathrm{H}_{2}(\mathrm{~g})\) (b) \(\mathrm{Cu}(\mathrm{s})+2 \mathrm{H}^{+}\)(aq) \(\rightarrow \mathrm{Cu}^{2+}(\mathrm{aq})+\mathrm{H}_{2}(\mathrm{~g})\) (c) \(\mathrm{Cu}(\mathrm{s})+2 \mathrm{Ag}^{+}(\mathrm{aq}) \rightarrow \mathrm{Cu}^{2+}(\mathrm{aq})+2 \mathrm{Ag}(\mathrm{s})\) (d) \(\mathrm{Ni}^{2}\) (aq) \(+\mathrm{Cu}(\mathrm{s}) \rightarrow \mathrm{Ni}(\mathrm{s})+\mathrm{Cu}^{2+}(\mathrm{aq})\)

Which is/are correct statement about salt bridge? (a) Ions of salt bridge discharge at electrode (b) Ions of salt bridge do not discharge at electrode (c) Velocity of ions of salt bridge are almost equal (d) Salt bridge complete the electric circuit.

Which of the following statements are correct? (a) \(\mathrm{KMnO}_{4}\) is a powerful oxidising agent. (b) \(\mathrm{KMnO}_{4}\) is a weaker oxidising agent than \(\mathrm{K}_{2} \mathrm{Cr}_{2} \mathrm{O}_{7}\) in acid medium. (c) \(\mathrm{KMnO}_{4}\) is a stronger oxidising agent than \(\mathrm{K}_{2} \mathrm{Cr}_{2} \mathrm{O}_{7}\) in acid medium. (d) \(\mathrm{K}_{2} \mathrm{Cr}_{2} \mathrm{O}_{7}\) oxidises a secondary alcohol to a ketone.

When \(\mathrm{Cl}_{2}\) is passed through hot \(\mathrm{NaOH}\) solution, oxidation number of chlorine changes from (a) 0 to \(-1\) (b) 0 to \(+5\) (c) 0 to \(+7\) (d) \(-1\) to 0

In which of the following compounds the oxidation state of oxygen is other than \(-2 ?\) (a) \(\mathrm{H}_{2} \mathrm{O}_{2}\) (b) \(\mathrm{O}_{2}\) (c) \(\mathrm{O}_{2} \mathrm{~F}_{2}\) (d) \(\mathrm{H}_{2} \mathrm{O}\)

Identify the compounds in which the sulphur atoms are in different oxidation states? (a) \(\mathrm{K}_{2} \mathrm{~S}_{2} \mathrm{O}_{7}\) (b) \(\mathrm{Na}_{2} \mathrm{~S}_{2} \mathrm{O}_{3}\) (c) \(\mathrm{Na}_{2} \mathrm{~S}_{4} \mathrm{O}_{6}\) (d) \(\mathrm{K}_{2} \mathrm{~S}_{2} \mathrm{O}_{\mathrm{y}}\)

In an electrolytic cell, electrolysis is carried out. Electrical energy is converted into chemical energy. In an electrochemical cell, chemical reaction, i.e., redox reaction occurs and electricity is generated. So chemical energy is converted into electrical energy. Electrolysis is governed by Faraday's laws. The potential difference between the electrodes which is called electromotive force is responsible for the generation of electric energy in the electrochemical cells. The standard reduction potential values of three metallic cations \(\mathrm{X}, \mathrm{Y}\) and \(\mathrm{Z}\) are \(0.50 \mathrm{~V},-3.03 \mathrm{~V}\) and \(-1.2 \mathrm{~V}\) respectively. The order of reducing power of the corresponding metals is (a) \(\mathrm{X}>\mathrm{Y}>\mathrm{Z}\) (b) \(\mathrm{Z}>\mathrm{Y}>\mathrm{X}\) (c) \(\mathrm{Y}>\mathrm{Z}>\mathrm{X}\) (d) \(\mathrm{X}>\mathrm{Z}>\mathrm{Y}\)

In an electrolytic cell, electrolysis is carried out. Electrical energy is converted into chemical energy. In an electrochemical cell, chemical reaction, i.e., redox reaction occurs and electricity is generated. So chemical energy is converted into electrical energy. Electrolysis is governed by Faraday's laws. The potential difference between the electrodes which is called electromotive force is responsible for the generation of electric energy in the electrochemical cells. Two electrolytic cells, one containing acidified \(\mathrm{FeCl}_{2}\) and another acidified \(\mathrm{FeCl}_{3}\) are connected in series. The ratio of iron deposited at the cathodes in the tow cells will be (a) \(3: 1\) (b) \(2: 1\) (c) \(2: 3\) (d) \(3: 2\)

The electrochemical series is the arrangement of various electrode systems in the increasing order of their standard reduction potentials. It has several important features. On moving from the top to the bottom in the series, tendency to gain electrons, i.e., to get reduced increases. The electrode systems having negative values of standard reduction potentials act as anode when connected to a standard hydrogen electrode, while those having positive values act as cathode. HCl cannot be stored in an aluminium vessel because (a) \(\mathrm{Al}\) is a highly reactive metal. (b) \(\mathrm{HCl}\) is an oxidizing acid (c) \(\mathrm{E}_{\mathrm{A}^{3+} / \mathrm{Al}}^{0}\) is much smaller than \(E_{\mathrm{H}}^{0}+\mathrm{H}_{2}\) (d) All of these

The electrochemical series is the arrangement of various electrode systems in the increasing order of their standard reduction potentials. It has several important features. On moving from the top to the bottom in the series, tendency to gain electrons, i.e., to get reduced increases. The electrode systems having negative values of standard reduction potentials act as anode when connected to a standard hydrogen electrode, while those having positive values act as cathode. If \(E_{\mathrm{cu}^{2}}^{0}{\underline{\phantom{xx}}}_{\mathrm{Cu}}=0.34 \mathrm{~V}\) and \(E_{\mathrm{Ag}^{+} / \Lambda g}^{0}=0.8 \mathrm{~V}\), predict whether the reaction given below is feasible or not ? $$ \mathrm{Cu}^{2}+(\mathrm{aq})+2 \mathrm{Ag}(\mathrm{s}) \longrightarrow \mathrm{Cu}(\mathrm{s})+2 \mathrm{Ag}^{+}(\mathrm{aq}) $$ (a) Not feasible (b) Feasible (c) Feasible on increasing the conc. of \(\mathrm{Ag}^{+}(\mathrm{aq}) .\) (d) easible at high temp.

The electrochemical series is the arrangement of various electrode systems in the increasing order of their standard reduction potentials. It has several important features. On moving from the top to the bottom in the series, tendency to gain electrons, i.e., to get reduced increases. The electrode systems having negative values of standard reduction potentials act as anode when connected to a standard hydrogen electrode, while those having positive values act as cathode. If \(\mathrm{E}_{\mathrm{Fe}^{2+} / \mathrm{re}}^{0}=-0.44 \mathrm{~V}\) and \(\left.\mathrm{E}_{\mathrm{Mg}^{+2} \mathrm{Mg}}^{0}\right)=-2.37 \mathrm{~V}, \mathrm{E}_{\mathrm{Cu}^{2+}}^{0}\) \(=+0.34 \mathrm{~V}\) and \(\mathrm{E}_{\mathrm{Ag}^{+} / \mathrm{A}_{\mathrm{E}}}^{\mathrm{n}}=+0.80 \mathrm{~V}\), the correct order in which the metals displace each other is (a) \(\mathrm{Fe}>\mathrm{Cu}>\mathrm{Ag}>\mathrm{Mg}\) (b) \(\mathrm{Ag}>\mathrm{Cu}>\mathrm{Fe}>\mathrm{Mg}\) (c) \(\mathrm{Mg}>\mathrm{Fe}>\mathrm{Cu}>\mathrm{Ag}\) (d) \(\mathrm{Fe}>\mathrm{Ag}>\mathrm{Cu}>\mathrm{Mg}\)

Match the following $$ \begin{array}{ll} \hline \text { Column-I } & \text { Column-II } \\ \hline \begin{array}{l} \text { (a) Charge on one mole of } \\ \text { electron. } \end{array} & \text { (p) } 1 \text { Faraday } \\ \text { (b) } 108 \mathrm{~g} \text { of silver deposited } & \text { (q) } 96500 \text { coulomb } \\ \text { at electrode. } & \end{array} $$ $$ \begin{array}{ll} \ \text { (c) } 22.4 \text { L of hydrogen at } & \text { (r) 2 Faraday } \\ \text { STP collected. } & \\ \text { (d) } 8 \text { g of oxygen collected. } & \text { (t) } 5.6 \mathrm{~L} \text { at STP } \\ \hline \end{array} $$

Match the following $$ \begin{array}{ll} \hline \text { Column-I } & \text { Column-II } \\ \hline \begin{array}{l} \text { (a) } 50 \% \text { solution of } \mathrm{H}_{2} \mathrm{SO}_{4} \\ \text { using Pt electrodes } \end{array} & \text { (p) } \mathrm{H}_{2} \text { is evolved at } \\ \text { cathode } \\ \begin{array}{ll} \text { (b) } \text { Dilute solution } \mathrm{NaCl} & \text { (q) } \mathrm{O}_{2} \text { is evolved at } \\ \text { using Pt electrodes } & \text { anode } \end{array} \\ \begin{array}{ll} \text { (c) Dilute solution of } \mathrm{H}_{2} \mathrm{SO}_{4} \\ \text { using Cu electrodes } \end{array} & \text { (r) } \mathrm{Cl}_{2} \text { is evolved at } \\ \text { (d) Concentrated solution of } & \text { anode } \\ \mathrm{LiCl} \text { using Pt electrodes. } & \text { (s) } \mathrm{H}_{2} \mathrm{~S}_{2} \mathrm{O}_{8} \text { is } \\ \text { formed at anode } \\ & \text { (t) non-spontaneous } \\ \text { process } \end{array} $$

The current strength in ampere required to deposit \(8.0 \mathrm{~g}\) of silver in one hour is [At No. of \(\mathrm{Ag}=108\) ] approximately

The standard electrode potential of \(\mathrm{Cu}^{2} / \mathrm{Cu}=0.34 \mathrm{~V}\). The electrode potential will be zero, when the conc. of \(\mathrm{Cu}^{2+}\) is as \(\mathrm{x} \times 10^{-12} \mathrm{M}\). the value of \(\mathrm{x}\) is \([\) lig \(2=0.3010, \log 3=0.4771\) and \(\log 3.4=0.5315]\)

In the electrolysis of \(\mathrm{KI}, \mathrm{I}_{2}\) is formed at the anode by the reaction; \(2 \mathrm{I} \longrightarrow \mathrm{I}_{2}+2 \mathrm{e}^{-}\) After the passage of current of \(0.5\) ampere for 9650 seconds, \(I_{2}\) is formed which required \(40 \mathrm{ml}\) of \(0.1 \mathrm{M}\) \(\mathrm{Na}_{2} \mathrm{~S}_{2} \mathrm{O}_{3} \cdot 5 \mathrm{H}_{2} \mathrm{O}\) solution in the following reaction; \(\mathrm{I}_{2}+2 \mathrm{~S}_{2} \mathrm{O}_{3}^{2-} \longrightarrow \mathrm{S}_{4} \mathrm{O}_{6}^{2-}+2 \mathrm{I}\) What is the current efficiency?

Which of the following reaction is possible at anode? [2002] (a) \(\mathrm{F}_{2}+2 \mathrm{e}^{-} \longrightarrow 2 \mathrm{~F}^{-}\) (b) \(2 \mathrm{H}^{+}+{ }^{1} / 2 \mathrm{O}_{2}+2 \mathrm{e}^{-} \longrightarrow \mathrm{H}_{2} \mathrm{O}\) (c) \(2 \mathrm{Cr}_{2}{\underline{\phantom{xx}}}^{3+}+7 \mathrm{H}_{2} \mathrm{O} \longrightarrow \mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}+14 \mathrm{H}^{+}+6 \mathrm{e}^{-}\) (d) \(\mathrm{Fe}^{2+} \longrightarrow \mathrm{Fe}^{3+}+\mathrm{e}^{-}\)

Consider the following reaction at \(1100^{\circ} \mathrm{C}\) [2002] (I) \(2 \mathrm{C}+\mathrm{O}_{2} \longrightarrow 2 \mathrm{CO} \Delta \mathrm{G}^{\circ}=-460 \mathrm{~kJ} \mathrm{~mol}^{-1}\) (II) \(2 \mathrm{Zn}+\mathrm{O}_{2} \longrightarrow 2 \mathrm{ZnO} \Delta \mathrm{G}^{\circ}=-360 \mathrm{~kJ} \mathrm{~mol}^{-1}\) based on these, select correct alternate (a) zinc can be oxidized by \(\mathrm{CO}\) (b) zinc oxide can be reduced by carbon (c) both are correct (d) none is correct

Conductivity (Seimens S) is directly proportional to area of the vessel and the concentration of the solution in it and is inversely proportional to the length of the vessel, then constant of proportionality is expressed in [2002] (a) \(\mathrm{S} \mathrm{m} \mathrm{mol}^{-1}\) (b) \(\mathrm{S}^{2} \mathrm{~m}^{2} \mathrm{~mol}^{-2}\) (c) \(\mathrm{S} \mathrm{m}^{2} \mathrm{~mol}^{-1}\) (d) \(\mathrm{S}^{2} \mathrm{~m}^{2} \mathrm{~mol}\)

For the redox reaction \(\mathrm{Zn}(\mathrm{s})+\mathrm{Cu}^{2+}(0.1 \mathrm{M}) \longrightarrow \mathrm{Zn}^{2+}(1 \mathrm{M})+\mathrm{Cu}(\mathrm{s})\) Taking place in a cell, \(E_{\text {call }}^{\circ}\) is \(1.10\) volt. \(E_{\text {cell }}\) for the cell will be \((2.303 \mathrm{RT} / F=0.0591)\) [2003] (a) \(2.14 \mathrm{~V}\) (b) \(1.80 \mathrm{~V}\) (c) \(1.07 \mathrm{~V}\) (d) \(0.82 \mathrm{~V}\)

For a cell reaction involving two electrons, the standard emf of the cell is found to be \(0.295 \mathrm{~V}\) at \(25^{\circ} \mathrm{C}\). The equilibrium constant of the reaction at \(25^{\circ} \mathrm{C}\) will be [2003] (a) \(1 \times 10^{-10}\) (b) \(29.5 \times 10^{-2}\) (c) 10 (d) \(1 \times 10^{10}\)

During electrolysis of a solution of \(\mathrm{AgNO}_{3}, 9650\) coulombs of charge pass through the electroplating bath, the mass of silver deposited on the cathode will be [2003] (a) \(1.08 \mathrm{~g}\) (b) \(10.8 \mathrm{~g}\) (c) \(21.6 \mathrm{~g}\) (d) \(108 \mathrm{~g}\)

Standard reduction electrode potentials of three metals \(\mathrm{A}, \mathrm{B}\) and \(\mathrm{C}\) are \(+0.5 \mathrm{~V},-3.0 \mathrm{~V}\) and \(-1.2 \mathrm{~V}\) respectively. The reducing power of these metals are [2003] (a) \(\mathrm{B}>\mathrm{C}>\mathrm{A}\) (b) \(\mathrm{A}>\mathrm{B}>\mathrm{C}\) (c) \(\mathrm{C}>\mathrm{B}>\mathrm{A}\) (d) \(\mathrm{A}>\mathrm{C}>\mathrm{B}\)

In a hydrogen-oxygen fuel cell, combustion of hydrogen occurs to [2004] (a) produce high purity water (b) generate heat (c) remove adsorbed oxygen from electrode surfaces (d) create potential difference between the two electrodes

Consider the following \(E^{\circ}\) values \(E\left(\mathrm{Fe}^{3+} / \mathrm{Fe}^{2+}\right)=+0.77 \mathrm{~V}\) \(E\left(\mathrm{Sn}^{2+} / \mathrm{Sn}\right)=-0.14 \mathrm{~V}\) Under standard conditions, the potential for the reaction \(\mathrm{Sn}(\mathrm{s})+2 \mathrm{Fe}^{3+}(\mathrm{aq}) \longrightarrow 2 \mathrm{Fe}^{2+}(\mathrm{aq})+\mathrm{Sn}^{2+}(\mathrm{aq})\) is (a) \(1.68 \mathrm{~V}\) (b) \(0.91 \mathrm{~V}\) (c) \(0.63 \mathrm{~V}\) (d) \(1.46 \mathrm{~V}\)

The standard emf of a cell, involving one electron change is found to be \(0.591 \mathrm{~V}\) at \(25^{\circ} \mathrm{C}\). The equilibrium constant of the reaction is \(\left(F=96500 \mathrm{C} \mathrm{mol}^{-1}, \mathrm{R}\right.\) \(\left.=8.314 \mathrm{JK}^{-1} \mathrm{~mol}^{-1}\right) \quad[\mathbf{2 0 0 4}]\) (a) \(1.0 \times 10^{30}\) (b) \(1.0 \times 10^{1}\) (c) \(1.0 \times 10^{5}\) (d) \(1.0 \times 10^{10}\)