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Chemistry · Textbook solutions

Electrochemistry

Every solved example, exercise, and miscellaneous question — in the order the textbook teaches them. · 43 questions

Worked Examples

10 q

Solved Examples

Worked · 10
  1. Eg 2.1
    Represent the cell in which the following reaction takes place Mg(s)+2Ag+(0.0001 M)→Mg2+(0.130 M)+2Ag(s)\text{Mg}(s) + 2\text{Ag}^{+}(0.0001\,\text{M}) \rightarrow \text{Mg}^{2+}(0.130\,\text{M}) + 2\text{Ag}(s) Calculate its E(cell)E_{(cell)} if E(cell)∘=3.17 VE^{\circ}_{(cell)} = 3.17\,\text{V}.
  2. Eg 2.2
    Calculate the equilibrium constant of the reaction: Cu(s)+2Ag+(aq)→Cu2+(aq)+2Ag(s)\text{Cu}(s) + 2\text{Ag}^{+}(aq) \rightarrow \text{Cu}^{2+}(aq) + 2\text{Ag}(s) E(cell)∘=0.46 VE^{\circ}_{(cell)} = 0.46\,\text{V}
  3. Eg 2.3
    The standard electrode potential for Daniell cell is 1.1 V1.1\,\text{V}. Calculate the standard Gibbs energy for the reaction: Zn(s)+Cu2+(aq)→Zn2+(aq)+Cu(s)\text{Zn}(s) + \text{Cu}^{2+}(aq) \rightarrow \text{Zn}^{2+}(aq) + \text{Cu}(s)
  4. Eg 2.4
    Resistance of a conductivity cell filled with 0.1 mol L−10.1\,\text{mol L}^{-1} KCl\text{KCl} solution is 100 Ω100\,\Omega. If the resistance of the same cell when filled with 0.02 mol L−10.02\,\text{mol L}^{-1} KCl\text{KCl} solution is 520 Ω520\,\Omega, calculate the conductivity and molar conductivity of 0.02 mol L−10.02\,\text{mol L}^{-1} KCl\text{KCl} solution. The conductivity of 0.1 mol L−10.1\,\text{mol L}^{-1} KCl\text{KCl} solution is 1.29 S/m1.29\,\text{S/m}.
  5. Eg 2.5
    The electrical resistance of a column of 0.05 mol L−10.05\,\text{mol L}^{-1} NaOH\text{NaOH} solution of diameter 1 cm1\,\text{cm} and length 50 cm50\,\text{cm} is 5.55×103 ohm5.55 \times 10^{3}\,\text{ohm}. Calculate its resistivity, conductivity and molar conductivity.
  6. Eg 2.6
    The molar conductivity of KCl\text{KCl} solutions at different concentrations at 298 K298\,\text{K} are given below:
    c/mol L−1c/\text{mol L}^{-1}Λm/S cm2 mol−1\Lambda_m/\text{S cm}^{2}\,\text{mol}^{-1}
    0.000198148.61
    0.000309148.29
    0.000521147.81
    0.000989147.09
    Show that a plot between Λm\Lambda_m and c1/2c^{1/2} is a straight line. Determine the values of Λm∘\Lambda_m^{\circ} and AA for KCl\text{KCl}.
  7. Eg 2.7
    Calculate Λm∘\Lambda_m^{\circ} for CaCl2\text{CaCl}_2 and MgSO4\text{MgSO}_4 from the data given in Table 3.4.
  8. Eg 2.8
    Λm∘\Lambda_m^{\circ} for NaCl\text{NaCl}, HCl\text{HCl} and NaAc\text{NaAc} are 126.4126.4, 425.9425.9 and 91.0 S cm2 mol−191.0\,\text{S cm}^{2}\,\text{mol}^{-1} respectively. Calculate Λ∘\Lambda^{\circ} for HAc\text{HAc}.
  9. Eg 2.9
    The conductivity of 0.001028 mol L−10.001028\,\text{mol L}^{-1} acetic acid is 4.95×10−5 S cm−14.95 \times 10^{-5}\,\text{S cm}^{-1}. Calculate its dissociation constant if Λm∘\Lambda_m^{\circ} for acetic acid is 390.5 S cm2 mol−1390.5\,\text{S cm}^{2}\,\text{mol}^{-1}.
  10. Eg 2.10
    A solution of CuSO4\text{CuSO}_4 is electrolysed for 10 minutes with a current of 1.5 amperes. What is the mass of copper deposited at the cathode?

Intext Questions

15 q
  1. Intext 2.1
    How would you determine the standard electrode potential of the system Mg2+ ∣ Mg\text{Mg}^{2+}\,|\,\text{Mg}?
  2. Intext 2.2
    Can you store copper sulphate solutions in a zinc pot?
  3. Intext 2.3
    Consult the table of standard electrode potentials and suggest three substances that can oxidise ferrous ions under suitable conditions.
  4. Intext 2.4
    Calculate the potential of hydrogen electrode in contact with a solution whose pH is 10.
  5. Intext 2.5
    Calculate the emf of the cell in which the following reaction takes place: Ni(s)+2Ag+(0.002 M)→Ni2+(0.160 M)+2Ag(s)\text{Ni}(s) + 2\text{Ag}^{+}(0.002\,\text{M}) \rightarrow \text{Ni}^{2+}(0.160\,\text{M}) + 2\text{Ag}(s) Given that Ecell∘=1.05 VE^{\circ}_{cell} = 1.05\,\text{V}
  6. Intext 2.6
    The cell in which the following reaction occurs: 2Fe3+(aq)+2I−(aq)→2Fe2+(aq)+I2(s)2\text{Fe}^{3+}(aq) + 2\text{I}^{-}(aq) \rightarrow 2\text{Fe}^{2+}(aq) + \text{I}_2(s) has Ecell∘=0.236 VE^{\circ}_{cell} = 0.236\,\text{V} at 298 K298\,\text{K}. Calculate the standard Gibbs energy and the equilibrium constant of the cell reaction.
  7. Intext 2.7
    Why does the conductivity of a solution decrease with dilution?
  8. Intext 2.8
    Suggest a way to determine the Λm∘\Lambda_m^{\circ} value of water.
  9. Intext 2.9
    The molar conductivity of 0.025 mol L−10.025\,\text{mol L}^{-1} methanoic acid is 46.1 S cm2 mol−146.1\,\text{S cm}^{2}\,\text{mol}^{-1}. Calculate its degree of dissociation and dissociation constant. Given λ0(H+)=349.6 S cm2 mol−1\lambda^{0}(\text{H}^{+}) = 349.6\,\text{S cm}^{2}\,\text{mol}^{-1} and λ0(HCOO−)=54.6 S cm2 mol−1\lambda^{0}(\text{HCOO}^{-}) = 54.6\,\text{S cm}^{2}\,\text{mol}^{-1}.
  10. Intext 2.10
    If a current of 0.5 ampere flows through a metallic wire for 2 hours, then how many electrons would flow through the wire?
  11. Intext 2.11
    Suggest a list of metals that are extracted electrolytically.
  12. Intext 2.12
    Consider the reaction: Cr2O72−+14H++6e−→2Cr3++7H2O\text{Cr}_2\text{O}_7^{2-} + 14\text{H}^{+} + 6e^{-} \rightarrow 2\text{Cr}^{3+} + 7\text{H}_2\text{O} What is the quantity of electricity in coulombs needed to reduce 1 mol of Cr2O72−\text{Cr}_2\text{O}_7^{2-}?
  13. Intext 2.13
    Write the chemistry of recharging the lead storage battery, highlighting all the materials that are involved during recharging.
  14. Intext 2.14
    Suggest two materials other than hydrogen that can be used as fuels in fuel cells.
  15. Intext 2.15
    Explain how rusting of iron is envisaged as setting up of an electrochemical cell.

Exercises

18 q
  1. Ex 2.1
    Arrange the following metals in the order in which they displace each other from the solution of their salts. Al\text{Al}, Cu\text{Cu}, Fe\text{Fe}, Mg\text{Mg} and Zn\text{Zn}.
  2. Ex 2.2
    Given the standard electrode potentials, K+/K=−2.93 V\text{K}^{+}/\text{K} = -2.93\,\text{V}, Ag+/Ag=0.80 V\text{Ag}^{+}/\text{Ag} = 0.80\,\text{V}, Hg2+/Hg=0.79 V\text{Hg}^{2+}/\text{Hg} = 0.79\,\text{V} Mg2+/Mg=−2.37 V\text{Mg}^{2+}/\text{Mg} = -2.37\,\text{V}, Cr3+/Cr=−0.74 V\text{Cr}^{3+}/\text{Cr} = -0.74\,\text{V} Arrange these metals in their increasing order of reducing power.
  3. Ex 2.3
    Depict the galvanic cell in which the reaction Zn(s)+2Ag+(aq)→Zn2+(aq)+2Ag(s)\text{Zn}(s) + 2\text{Ag}^{+}(aq) \rightarrow \text{Zn}^{2+}(aq) + 2\text{Ag}(s) takes place. Further show: (i) Which of the electrode is negatively charged? (ii) The carriers of the current in the cell. (iii) Individual reaction at each electrode.
  4. Ex 2.4
    Calculate the standard cell potentials of galvanic cell in which the following reactions take place: (i) 2Cr(s)+3Cd2+(aq)→2Cr3+(aq)+3Cd2\text{Cr}(s) + 3\text{Cd}^{2+}(aq) \rightarrow 2\text{Cr}^{3+}(aq) + 3\text{Cd} (ii) Fe2+(aq)+Ag+(aq)→Fe3+(aq)+Ag(s)\text{Fe}^{2+}(aq) + \text{Ag}^{+}(aq) \rightarrow \text{Fe}^{3+}(aq) + \text{Ag}(s) Calculate the ΔrG∘\Delta_r G^{\circ} and equilibrium constant of the reactions.
  5. Ex 2.5
    Write the Nernst equation and emf of the following cells at 298 K: (i) Mg(s) ∣ Mg2+(0.001 M) ∣∣ Cu2+(0.0001 M) ∣ Cu(s)\text{Mg}(s)\,|\,\text{Mg}^{2+}(0.001\,\text{M})\,||\,\text{Cu}^{2+}(0.0001\,\text{M})\,|\,\text{Cu}(s) (ii) Fe(s) ∣ Fe2+(0.001 M) ∣∣ H+(1 M) ∣ H2(g)(1 bar) ∣ Pt(s)\text{Fe}(s)\,|\,\text{Fe}^{2+}(0.001\,\text{M})\,||\,\text{H}^{+}(1\,\text{M})\,|\,\text{H}_2(g)(1\,\text{bar})\,|\,\text{Pt}(s) (iii) Sn(s) ∣ Sn2+(0.050 M) ∣∣ H+(0.020 M) ∣ H2(g) (1 bar) ∣ Pt(s)\text{Sn}(s)\,|\,\text{Sn}^{2+}(0.050\,\text{M})\,||\,\text{H}^{+}(0.020\,\text{M})\,|\,\text{H}_2(g)\ (1\,\text{bar})\,|\,\text{Pt}(s) (iv) Pt(s) ∣ Br−(0.010 M) ∣ Br2(l) ∣∣ H+(0.030 M) ∣ H2(g) (1 bar) ∣ Pt(s)\text{Pt}(s)\,|\,\text{Br}^{-}(0.010\,\text{M})\,|\,\text{Br}_2(l)\,||\,\text{H}^{+}(0.030\,\text{M})\,|\,\text{H}_2(g)\ (1\,\text{bar})\,|\,\text{Pt}(s).
  6. Ex 2.6
    In the button cells widely used in watches and other devices the following reaction takes place: Zn(s)+Ag2O(s)+H2O(l)→Zn2+(aq)+2Ag(s)+2OH−(aq)\text{Zn}(s) + \text{Ag}_2\text{O}(s) + \text{H}_2\text{O}(l) \rightarrow \text{Zn}^{2+}(aq) + 2\text{Ag}(s) + 2\text{OH}^{-}(aq) Determine ΔrG∘\Delta_r G^{\circ} and E∘E^{\circ} for the reaction.
  7. Ex 2.7
    Define conductivity and molar conductivity for the solution of an electrolyte. Discuss their variation with concentration.
  8. Ex 2.8
    The conductivity of 0.20 M solution of KCl\text{KCl} at 298 K is 0.0248 S cm−10.0248\,\text{S cm}^{-1}. Calculate its molar conductivity.
  9. Ex 2.9
    The resistance of a conductivity cell containing 0.001M KCl\text{KCl} solution at 298 K is 1500 Ω1500\,\Omega. What is the cell constant if conductivity of 0.001M KCl\text{KCl} solution at 298 K is 0.146×10−3 S cm−10.146 \times 10^{-3}\,\text{S cm}^{-1}.
  10. Ex 2.10
    The conductivity of sodium chloride at 298 K has been determined at different concentrations and the results are given below:
    Concentration/M0.0010.0100.0200.0500.100
    102×κ/S m−110^{2} \times \kappa/\text{S m}^{-1}1.23711.8523.1555.53106.74
    Calculate Λm\Lambda_m for all concentrations and draw a plot between Λm\Lambda_m and c1/2c^{1/2}. Find the value of Λm0\Lambda_m^{0}.
  11. Ex 2.11
    Conductivity of 0.00241 M acetic acid is 7.896×10−5 S cm−17.896 \times 10^{-5}\,\text{S cm}^{-1}. Calculate its molar conductivity. If Λm0\Lambda_m^{0} for acetic acid is 390.5 S cm2 mol−1390.5\,\text{S cm}^{2}\,\text{mol}^{-1}, what is its dissociation constant?
  12. Ex 2.12
    How much charge is required for the following reductions: (i) 1 mol of Al3+\text{Al}^{3+} to Al\text{Al}? (ii) 1 mol of Cu2+\text{Cu}^{2+} to Cu\text{Cu}? (iii) 1 mol of MnO4−\text{MnO}_4^{-} to Mn2+\text{Mn}^{2+}?
  13. Ex 2.13
    How much electricity in terms of Faraday is required to produce (i) 20.0 g of Ca\text{Ca} from molten CaCl2\text{CaCl}_2? (ii) 40.0 g of Al\text{Al} from molten Al2O3\text{Al}_2\text{O}_3?
  14. Ex 2.14
    How much electricity is required in coulomb for the oxidation of (i) 1 mol of H2O\text{H}_2\text{O} to O2\text{O}_2? (ii) 1 mol of FeO\text{FeO} to Fe2O3\text{Fe}_2\text{O}_3?
  15. Ex 2.15
    A solution of Ni(NO3)2\text{Ni(NO}_3)_2 is electrolysed between platinum electrodes using a current of 5 amperes for 20 minutes. What mass of Ni\text{Ni} is deposited at the cathode?
  16. Ex 2.16
    Three electrolytic cells A, B, C containing solutions of ZnSO4\text{ZnSO}_4, AgNO3\text{AgNO}_3 and CuSO4\text{CuSO}_4, respectively are connected in series. A steady current of 1.5 amperes was passed through them until 1.45 g of silver deposited at the cathode of cell B. How long did the current flow? What mass of copper and zinc were deposited?
  17. Ex 2.17
    Using the standard electrode potentials given in Table 3.1, predict if the reaction between the following is feasible: (i) Fe3+(aq)\text{Fe}^{3+}(aq) and I−(aq)\text{I}^{-}(aq) (ii) Ag+(aq)\text{Ag}^{+}(aq) and Cu(s)\text{Cu}(s) (iii) Fe3+(aq)\text{Fe}^{3+}(aq) and Br−(aq)\text{Br}^{-}(aq) (iv) Ag(s)\text{Ag}(s) and Fe3+(aq)\text{Fe}^{3+}(aq) (v) Br2(aq)\text{Br}_2(aq) and Fe2+(aq)\text{Fe}^{2+}(aq).
  18. Ex 2.18
    Predict the products of electrolysis in each of the following: (i) An aqueous solution of AgNO3\text{AgNO}_3 with silver electrodes. (ii) An aqueous solution of AgNO3\text{AgNO}_3 with platinum electrodes. (iii) A dilute solution of H2SO4\text{H}_2\text{SO}_4 with platinum electrodes. (iv) An aqueous solution of CuCl2\text{CuCl}_2 with platinum electrodes.