Chemistry · Textbook solutions

Electrochemistry

Every solved example, exercise, and miscellaneous question — in the order the textbook teaches them. · 56 questions

5. Electrochemistry — worked examples

12 q

Solved Examples

Worked · 12
  1. Solved Ex.5.1
    The molar conductivity of 0.05 M BaCl2\mathrm{BaCl_2} solution at 25∘C25^\circ\mathrm{C} is 223 Ω−1 cm2 mol−1\Omega^{-1}\ \mathrm{cm^2\ mol^{-1}}. What is its conductivity ?
  2. Solved Ex.5.2
    Calculate the molar conductivity of AgI\mathrm{AgI} at zero concentration if the molar conductivities of NaI\mathrm{NaI}, AgNO3\mathrm{AgNO_3} and NaNO3\mathrm{NaNO_3} at zero concentration are respectively, 126.9, 133.4 and 121.5 Ω−1 cm2 mol−1\Omega^{-1}\ \mathrm{cm^2\ mol^{-1}}.
  3. Solved Ex.5.3
    Calculate molar conductivities at zero concentration for CaCl2\mathrm{CaCl_2} and Na2SO4\mathrm{Na_2SO_4}. Given : molar ionic conductivitis of Ca2⊕\mathrm{Ca^{2\oplus}}, Cl⊖\mathrm{Cl^\ominus}, Na⊕\mathrm{Na^\oplus} and SO42⊖\mathrm{SO_4^{2\ominus}} ions are respectively, 104, 76.4, 50.1 and 159.6 Ω−1 cm2 mol−1\Omega^{-1}\ \mathrm{cm^2\ mol^{-1}}.
  4. Solved Ex.5.4
    The molar conductivity of 0.01M acetic acid at 25∘C25^\circ\mathrm{C} is 16.5 Ω−1 cm2 mol−1\Omega^{-1}\ \mathrm{cm^2\ mol^{-1}}. Calculate its degree of dissociation in 0.01 M solution and dissociation constant if molar conductivity of acetic acid at zero concentration is 390.7 Ω−1 cm2 mol−1\Omega^{-1}\ \mathrm{cm^2\ mol^{-1}}.
  5. Solved Ex.5.5
    A conductivity cell containing 0.01M KCl\mathrm{KCl} gives at 25∘C25^\circ\mathrm{C} the resistance of 604 ohms. The same cell containing 0.001M AgNO3\mathrm{AgNO_3} gives resistance of 6530 ohms. Calculate the molar conductivity of 0.001M AgNO3\mathrm{AgNO_3}. [Conductivity of 0.01M KCl\mathrm{KCl} at 25∘C25^\circ\mathrm{C} is 0.00141 Ω−1 cm−1\Omega^{-1}\ \mathrm{cm^{-1}}]
  6. Solved Ex.5.6
    What is the mass of Cu\mathrm{Cu} metal produced at the cathode during the passage of 5 ampere current through CuSO4\mathrm{CuSO_4} solution for 100 minutes. Molar mass of Cu\mathrm{Cu} is 63.5 g mol−1\mathrm{mol^{-1}}.
  7. Solved Ex.5.7
    How long will it take to produce 2.415 g of Ag\mathrm{Ag} metal from its salt solution by passing a current of 3 ampere ? Molar mass of Ag\mathrm{Ag} is 107.9 g mol−1\mathrm{mol^{-1}}.
  8. Solved Ex.5.8
    How many moles of electrons are required for reduction of 3 moles of Zn2⊕\mathrm{Zn^{2\oplus}} to Zn\mathrm{Zn} ? How many Faradays of electricity will be required ?
  9. Solved Ex.5.9
    In a certain electrolysis experiment 4.36 g of Zn\mathrm{Zn} are deposited in one cell containing ZnSO4\mathrm{ZnSO_4} solution. Calculate the mass of Al\mathrm{Al} deposited in another cell containing AlCl3\mathrm{AlCl_3} solution connected in series with ZnSO4\mathrm{ZnSO_4} cell. Molar masses of Zn\mathrm{Zn} and Al\mathrm{Al} are 65.4 g mol−1\mathrm{mol^{-1}} and 27 g mol−1\mathrm{mol^{-1}}, respectively.
  10. Solved Ex.5.10
    Calculate the voltage of the cell, Sn(s) ∣ Sn2⊕ (0.02M) ∣∣ Ag⊕ (0.01M) ∣ Ag(s)\mathrm{Sn(s)}\,|\,\mathrm{Sn^{2\oplus}}\,(0.02\mathrm{M})\,||\,\mathrm{Ag^\oplus}\,(0.01\mathrm{M})\,|\,\mathrm{Ag(s)} at 25∘C25^\circ\mathrm{C}. ESn0=−0.136E^0_{Sn} = -0.136 V, EAg0=0.800E^0_{Ag} = 0.800 V.
  11. Solved Ex.5.11
    The standard potential of the electrode, Zn2⊕ (0.02 M) ∣ Zn(s)\mathrm{Zn^{2\oplus}}\,(0.02\ \mathrm{M})\,|\,\mathrm{Zn(s)} is −0.76-0.76 V. Calculate its potential.
  12. Solved Ex.5.12
    Calculate standard Gibbs energy change and equilibrium constant at 25∘C25^\circ\mathrm{C} for the cell reaction, Cd(s)+Sn2⊕(aq)⟶Cd2⊕(aq)+Sn(s)\mathrm{Cd(s)} + \mathrm{Sn^{2\oplus}(aq)} \longrightarrow \mathrm{Cd^{2\oplus}(aq)} + \mathrm{Sn(s)} Given : ECd0=−0.403E^0_{Cd} = -0.403 V and ESn0=−0.136E^0_{Sn} = -0.136 V. Write formula of the cell.

Exercises

44 q

1. Choose the most correct option.

Practice · 10
  1. Choose the most correct option.
    Ex Q.1 (i)
    Two solutions have the ratio of their concentrations 0.4 and ratio of their conductivities 0.216. The ratio of their molar conductivities will be
    1. A.
      0.54
    2. B.
      11.574
    3. C.
      0.0864
    4. D.
      1.852
  2. Ex Q.1 (ii)
    On diluting the solution of an electrolyte
    1. A.
      both Λ\Lambda and kk increase
    2. B.
      both Λ\Lambda and kk decrease
    3. C.
      Λ\Lambda increases and kk decreases
    4. D.
      Λ\Lambda decreases and kk increases
  3. Ex Q.1 (iii)
    1 S m2 mol−1\mathrm{m^2\ mol^{-1}} is eual to
    1. A.
      10−410^{-4} S m2 mol−1\mathrm{m^2\ mol^{-1}}
    2. B.
      10410^{4} Ω−1 cm2 mol−1\Omega^{-1}\ \mathrm{cm^2\ mol^{-1}}
    3. C.
      10−210^{-2} S cm2 mol−1\mathrm{cm^2\ mol^{-1}}
    4. D.
      10210^{2} Ω−1 cm2 mol−1\Omega^{-1}\ \mathrm{cm^2\ mol^{-1}}
  4. Ex Q.1 (iv)
    The standard potential of the cell in which the following reaction occurs H2 (g, 1 atm)+Cu2⊕(1M)⟶2H⊕(1M)+Cu(s)\mathrm{H_2}\,(\mathrm{g},\ 1\,\mathrm{atm}) + \mathrm{Cu^{2\oplus}}(1\mathrm{M}) \longrightarrow 2\mathrm{H^\oplus}(1\mathrm{M}) + \mathrm{Cu(s)}, (ECu0=0.34 V)(E^0_{Cu} = 0.34\,\mathrm{V}) is
    1. A.
      −0.34-0.34 V
    2. B.
      0.34 V
    3. C.
      0.17 V
    4. D.
      −0.17-0.17 V
  5. Ex Q.1 (v)
    For the cell, Pb(s) ∣ Pb2⊕(1M) ∣∣ Ag⊕(1M) ∣ Ag(s)\mathrm{Pb(s)}\,|\,\mathrm{Pb^{2\oplus}}(1\mathrm{M})\,||\,\mathrm{Ag^\oplus}(1\mathrm{M})\,|\,\mathrm{Ag(s)}, if concentraton of an ion in the anode compartment is increased by a factor of 10, the emf of the cell will
    1. A.
      increase by 10 V
    2. B.
      increase by 0.0296 V
    3. C.
      decrease by 10 V
    4. D.
      decrease by 0.0296 V
  6. Ex Q.1 (vi)
    Consider the half reactions with standard potentials i. Ag⊕(aq)+e⊖⟶Ag(s)\mathrm{Ag^\oplus(aq)} + \mathrm{e^\ominus} \longrightarrow \mathrm{Ag(s)} E0=0.8VE^0 = 0.8\mathrm{V} ii. I2(s)+2e⊖⟶2I⊖(aq)\mathrm{I_2(s)} + 2\mathrm{e^\ominus} \longrightarrow 2\mathrm{I^\ominus(aq)} E0=0.53VE^0 = 0.53\mathrm{V} iii. Pb2⊕(aq)+2e⊖⟶Pb(s)\mathrm{Pb^{2\oplus}(aq)} + 2\mathrm{e^\ominus} \longrightarrow \mathrm{Pb(s)} E0=−0.13VE^0 = -0.13\mathrm{V} iv. Fe2⊕(aq)+2e⊖⟶Fe(s)\mathrm{Fe^{2\oplus}(aq)} + 2\mathrm{e^\ominus} \longrightarrow \mathrm{Fe(s)} E0=−0.44VE^0 = -0.44\mathrm{V} The strongest oxidising and reducing agents respectively are
    1. A.
      Ag\mathrm{Ag} and Fe2⊕\mathrm{Fe^{2\oplus}}
    2. B.
      Ag⊕\mathrm{Ag^\oplus} and Fe\mathrm{Fe}
    3. C.
      Pb2⊕\mathrm{Pb^{2\oplus}} and I⊖\mathrm{I^\ominus}
    4. D.
      I2\mathrm{I_2} and Fe2⊕\mathrm{Fe^{2\oplus}}
  7. Ex Q.1 (vii)
    For the reaction Ni(s)+Cu2⊕(1M)⟶Ni2⊕(1M)+Cu(s)\mathrm{Ni(s)} + \mathrm{Cu^{2\oplus}}(1\mathrm{M}) \longrightarrow \mathrm{Ni^{2\oplus}}(1\mathrm{M}) + \mathrm{Cu(s)}, Ecell0=0.57VE^0_{cell} = 0.57\mathrm{V} ΔG0\Delta G^0 of the reaction is
    1. A.
      110 kJ
    2. B.
      −110-110 kJ
    3. C.
      55 kJ
    4. D.
      −55-55 kJ
  8. Ex Q.1 (viii)
    Which of the following is not correct?
    1. A.
      Gibbs energy is an extensive property
    2. B.
      Electrode potential or cell potential is an intensive property.
    3. C.
      Electrical work =−ΔG= -\Delta G
    4. D.
      If half reaction is multiplied by a numerical factor, the corresponding E0E^0 value is also multiplied by the same factor.
  9. Ex Q.1 (ix)
    The oxidation reaction that takes place in lead storage battery during discharge is
    1. A.
      Pb2⊕(aq)+SO42⊖(aq)⟶PbSO4(s)\mathrm{Pb^{2\oplus}(aq)} + \mathrm{SO_4^{2\ominus}(aq)} \longrightarrow \mathrm{PbSO_4(s)}
    2. B.
      PbSO4(s)+2H2O(l)⟶PbO2(s)+4H⊕(aq)+SO42⊖(aq)+2e⊖\mathrm{PbSO_4(s)} + 2\mathrm{H_2O}(l) \longrightarrow \mathrm{PbO_2(s)} + 4\mathrm{H^\oplus(aq)} + \mathrm{SO_4^{2\ominus}(aq)} + 2\mathrm{e^\ominus}
    3. C.
      Pb(s)+SO42⊖(aq)⟶PbSO4(s)+2e⊖\mathrm{Pb(s)} + \mathrm{SO_4^{2\ominus}(aq)} \longrightarrow \mathrm{PbSO_4(s)} + 2\mathrm{e^\ominus}
    4. D.
      PbSO4(s)+2e⊖⟶Pb(s)+SO42⊖(aq)\mathrm{PbSO_4(s)} + 2\mathrm{e^\ominus} \longrightarrow \mathrm{Pb(s)} + \mathrm{SO_4^{2\ominus}(aq)}
  10. Ex Q.1 (x)
    Which of the following expressions represent molar conductivity of Al2(SO4)3\mathrm{Al_2(SO_4)_3}?
    1. A.
      3 λAl3⊕0+2 λSO42⊖03\,\lambda^0_{\mathrm{Al^{3\oplus}}} + 2\,\lambda^0_{\mathrm{SO_4^{2\ominus}}}
    2. B.
      2 λAl3⊕0+3 λSO42⊖02\,\lambda^0_{\mathrm{Al^{3\oplus}}} + 3\,\lambda^0_{\mathrm{SO_4^{2\ominus}}}
    3. C.
      13 λAl3⊕0+12 λSO42⊖0\tfrac{1}{3}\,\lambda^0_{\mathrm{Al^{3\oplus}}} + \tfrac{1}{2}\,\lambda^0_{\mathrm{SO_4^{2\ominus}}}
    4. D.
      λAl3⊕0+λSO42⊖0\lambda^0_{\mathrm{Al^{3\oplus}}} + \lambda^0_{\mathrm{SO_4^{2\ominus}}}

2. Answer the following in one or two sentences.

Practice · 10
  1. Answer the following in one or two sentences.
    Ex Q.2 (i)
    What is a cell constant ?
  2. Ex Q.2 (ii)
    Write the relationship between conductivity and molar conductivity and hence unit of molar conductivity.
  3. Ex Q.2 (iii)
    Write the electrode reactions during electrolysis of molten KCl\mathrm{KCl}.
  4. Ex Q.2 (iv)
    Write any two functions of salt bridge.
  5. Ex Q.2 (v)
    What is standard cell potential for the reaction 2Al(s)+3Ni2⊕(1M)⟶2Al3⊕(1M)+3Ni(s)2\mathrm{Al(s)} + 3\mathrm{Ni^{2\oplus}}(1\mathrm{M}) \longrightarrow 2\mathrm{Al^{3\oplus}}(1\mathrm{M}) + 3\mathrm{Ni(s)} if ENi0=−0.25E^0_{Ni} = -0.25 V and EAl0=−1.66E^0_{Al} = -1.66V ?
  6. Ex Q.2 (vi)
    Write Nerst equation. What part of it represents the correction factor for nonstandard state conditions ?
  7. Ex Q.2 (vii)
    Under what conditions the cell potential is called standard cell potential ?
  8. Ex Q.2 (viii)
    Formulate a cell from the following electrode reactions : Au3⊕(aq)+3e⊖⟶Au(s)\mathrm{Au^{3\oplus}(aq)} + 3\mathrm{e^\ominus} \longrightarrow \mathrm{Au(s)} Mg(s)⟶Mg2⊕(aq)+2e⊖\mathrm{Mg(s)} \longrightarrow \mathrm{Mg^{2\oplus}(aq)} + 2\mathrm{e^\ominus}
  9. Ex Q.2 (ix)
    How many electrons would have a total charge of 1 coulomb ?
  10. Ex Q.2 (x)
    What is the significance of the single vertical line and double vertical line in the formulation galvanic cell.

3. Answer the following in brief

Practice · 10
  1. Answer the following in brief
    Ex Q.3 (i)
    Explain the effect of dilution of solution on conductivity ?
  2. Ex Q.3 (ii)
    What is a salt bridge ?
  3. Ex Q.3 (iii)
    Write electrode reactions for the electrolysis of aqueous NaCl\mathrm{NaCl}.
  4. Ex Q.3 (iv)
    How many moles of electrons are passed when 0.8 ampere current is passed for 1 hour through molten CaCl2\mathrm{CaCl_2} ?
  5. Ex Q.3 (v)
    Construct a galvanic cell from the electrodes Co3⊕ ∣ Co\mathrm{Co^{3\oplus}}\,|\,\mathrm{Co} and Mn2⊕ ∣ Mn\mathrm{Mn^{2\oplus}}\,|\,\mathrm{Mn}. ECo0=1.82E^0_{Co} = 1.82 V, EMn0=−1.18E^0_{Mn} = -1.18V. Calculate Ecell0E^0_{cell}.
  6. Ex Q.3 (vi)
    Using the relationsip between ΔG0\Delta G^0 of cell reaction and the standard potential associated with it, how will you show that the electrical potential is an intensive property ?
  7. Ex Q.3 (viii)
    Derive the relationship between standard cell potential and equilibrium constant of cell reaction.
  8. Ex Q.3 (ix)
    It is impossible to measure the potential of a single electrode. Comment.
  9. Ex Q.3 (x)
    Why do the cell potential of lead accumulators decrease when it generates electricity ? How the cell potential can be increased ?
  10. Ex Q.3 (xi)
    Write the electrode reactions and net cell reaction in NICAD battery.

4. Answer the following :

Practice · 14
  1. Answer the following :
    Ex Q.4 (i)
    What is Kohrausch law of independent migration of ions? How is it useful in obtaining molar conductivity at zero concentration of a weak electrolyte ? Explain with an example.
  2. Ex Q.4 (ii)
    Explain electrolysis of molten NaCl\mathrm{NaCl}.
  3. Ex Q.4 (iii)
    What current strength in amperes will be required to produce 2.4 g of Cu\mathrm{Cu} from CuSO4\mathrm{CuSO_4} solution in 1 hour ? Molar mass of Cu\mathrm{Cu} = 63.5 g mol−1\mathrm{mol^{-1}}.
  4. Ex Q.4 (iv)
    Equilibrium constant of the reaction, 2Cu⊕(aq)⟶Cu2⊕(aq)+Cu(s)2\mathrm{Cu^\oplus(aq)} \longrightarrow \mathrm{Cu^{2\oplus}(aq)} + \mathrm{Cu(s)} is 1.2×1061.2 \times 10^{6}. What is the standard potential of the cell in which the reaction takes place ?
  5. Ex Q.4 (v)
    Calculate emf of the cell Zn(s) ∣ Zn2⊕(0.2M) ∣∣ H⊕(1.6M) ∣ H2(g, 1.8 atm) ∣ Pt\mathrm{Zn(s)}\,|\,\mathrm{Zn^{2\oplus}}(0.2\mathrm{M})\,||\,\mathrm{H^\oplus}(1.6\mathrm{M})\,|\,\mathrm{H_2}(\mathrm{g},\ 1.8\ \mathrm{atm})\,|\,\mathrm{Pt} at 25∘C25^\circ\mathrm{C}.
  6. Ex Q.4 (vi)
    Calculate emf of the following cell at 25∘C25^\circ\mathrm{C}. Zn(s) ∣ Zn2⊕(0.08M) ∣∣ Cr3⊕(0.1M) ∣ Cr\mathrm{Zn(s)}\,|\,\mathrm{Zn^{2\oplus}}(0.08\mathrm{M})\,||\,\mathrm{Cr^{3\oplus}}(0.1\mathrm{M})\,|\,\mathrm{Cr} EZn0=−0.76E^0_{Zn} = -0.76 V, ECr0=−0.74E^0_{Cr} = -0.74 V
  7. Ex Q.4 (vii)
    What is a cell constant ? What are its units? How is it determined experimentally?
  8. Ex Q.4 (viii)
    How will you calculate the moles of electrons passed and mass of the substance produced during electrolysis of a salt solution using reaction stoichiometry.
  9. Ex Q.4 (ix)
    Write the electrode reactions when lead storage cell generates electricity. What are the anode and cathode and the electrode reactions during its recharging?
  10. Ex Q.4 (x)
    What are anode and cathode of H2\mathrm{H_2}-O2\mathrm{O_2} fuel cell ? Name the electrolyte used in it. Write electrode reactions and net cell reaction taking place in the fuel cell.
  11. Ex Q.4 (xi)
    What are anode and cathode for Leclanche' dry cell ? Write electrode reactions and overall cell reaction when it generates electricity.
  12. Ex Q.4 (xii)
    Identify oxidising agents and arrange them in order of increasing strength under standard state conditions. The standard potentials are given in parenthesis. Al\mathrm{Al} (−1.66V)(-1.66\mathrm{V}), Al3⊕(−1.66V)\mathrm{Al^{3\oplus}}(-1.66\mathrm{V}), Cl2\mathrm{Cl_2} (1.36V)(1.36\mathrm{V}), Cd2⊕(−0.4V)\mathrm{Cd^{2\oplus}}(-0.4\mathrm{V}), Fe(−0.44V)\mathrm{Fe}(-0.44\mathrm{V}), I2(0.54V)\mathrm{I_2}(0.54\mathrm{V}), Br⊖(1.09V)\mathrm{Br^\ominus}(1.09\mathrm{V}).
  13. Ex Q.4 (xiii)
    Which of the following species are reducing agents? Arrange them in order of increasing strength under standard state conditions. The standard potentials are given in parenthesis. K\mathrm{K} (−2.93V)(-2.93\mathrm{V}), Br2(1.09V)\mathrm{Br_2}(1.09\mathrm{V}), Mg(−2.36V)\mathrm{Mg}(-2.36\mathrm{V}), Ce3⊕(1.61V)\mathrm{Ce^{3\oplus}}(1.61\mathrm{V}), Ti2⊕(−0.37V)\mathrm{Ti^{2\oplus}}(-0.37\mathrm{V}), Ag⊕(0.8 V)\mathrm{Ag^\oplus}(0.8\ \mathrm{V}), Ni\mathrm{Ni} (−0.23V)(-0.23\mathrm{V}).
  14. Answer the following : Standard aqueous electrode potentials at 298 K (from Table 5.1 of the chapter) :
    Half reactionE0E^0 / V
    Br2+2e⊖⟶2Br⊖\mathrm{Br_2} + 2\mathrm{e^\ominus} \longrightarrow 2\mathrm{Br^\ominus}+1.080+1.080
    Ag⊕+e⊖⟶Ag\mathrm{Ag^\oplus} + \mathrm{e^\ominus} \longrightarrow \mathrm{Ag}+0.799+0.799
    Sn2⊕+2e⊖⟶Sn\mathrm{Sn^{2\oplus}} + 2\mathrm{e^\ominus} \longrightarrow \mathrm{Sn}−0.136-0.136
    Ni2⊕+2e⊖⟶Ni\mathrm{Ni^{2\oplus}} + 2\mathrm{e^\ominus} \longrightarrow \mathrm{Ni}−0.257-0.257
    Cd2⊕+2e⊖⟶Cd\mathrm{Cd^{2\oplus}} + 2\mathrm{e^\ominus} \longrightarrow \mathrm{Cd}−0.403-0.403
    Ca2⊕+2e⊖⟶Ca\mathrm{Ca^{2\oplus}} + 2\mathrm{e^\ominus} \longrightarrow \mathrm{Ca}−2.866-2.866
    Ex Q.4 (xiv)
    Predict whether the following reactions would occur spontaneously under standard state conditions. a. Ca(s)+Cd2⊕(aq)⟶Ca2⊕(aq)+Cd(s)\mathrm{Ca(s)} + \mathrm{Cd^{2\oplus}(aq)} \longrightarrow \mathrm{Ca^{2\oplus}(aq)} + \mathrm{Cd(s)} b. 2 Br⊖(s)+Sn2⊕(aq)⟶Br2(l)+Sn(s)2\,\mathrm{Br^\ominus(s)} + \mathrm{Sn^{2\oplus}(aq)} \longrightarrow \mathrm{Br_2}(l) + \mathrm{Sn(s)} c. 2Ag(s)+Ni2⊕(aq)⟶2 Ag⊕(aq)+Ni(s)2\mathrm{Ag(s)} + \mathrm{Ni^{2\oplus}(aq)} \longrightarrow 2\,\mathrm{Ag^\oplus(aq)} + \mathrm{Ni(s)} (use information of Table 5.1)