Electrochemistry

Electrochemistry studies the production of electricity from the energy released in spontaneous chemical reactions and the use of electrical energy to bring about non-spontaneous chemical changes. In this chapter you will study galvanic cells such as the Daniell cell, Zn(s) / Zn2+(aq) // Cu2+(aq) / Cu(s), with an emf of 1.10 V under standard conditions, the standard hydrogen electrode and standard electrode potentials. You will apply the Nernst equation, E = E0 - (0.059/n) log Q at 298 K, and relate the standard emf to the equilibrium constant and to the Gibbs energy change through Delta G0 = -nFE0. The chapter then explains conductance of electrolytic solutions: resistance, conductivity, cell constant and molar conductivity, how they vary with concentration for strong and weak electrolytes, and Kohlrausch's law of independent migration of ions and its applications. You will also study electrolytic cells, Faraday's laws and the products of electrolysis, primary and secondary batteries such as the dry cell, mercury cell and lead storage battery, the hydrogen-oxygen fuel cell, and corrosion as an electrochemical process.

Detailed Worksheet: Electrochemistry Section A - Definitions (10 marks) 1. Distinguish between a galvanic cell and an electrolytic cell on two bases. (2 marks) 2. What is the standard hydrogen electrode? Write its cell representation and the value of its potential. (2 marks) 3. Define conductivity and molar conductivity. Write the relation between them with units. (2 marks) 4. State Kohlrausch's law of independent migration of ions. (2 marks) 5. State Faraday's first and second laws of electrolysis. (2 marks) Section B - Calculations and Applications (15 marks) 6. Calculate the emf of the cell Zn(s) / Zn2+(0.001 M) // Cu2+(0.1 M) / Cu(s) at 298 K, given E0cell = 1.10 V. Also calculate Delta G0 for the cell reaction (F = 96,500 C mol^-1). (3 marks) 7. Calculate the equilibrium constant for the reaction Cu(s) + 2Ag+(aq) -> Cu2+(aq) + 2Ag(s), given E0cell = 0.46 V at 298 K (antilog 0.59 = 3.92). (3 marks) 8. The resistance of a 0.01 M KCl solution in a conductivity cell is 1,500 ohm, and the cell constant is 1.5 cm^-1. Calculate the conductivity and the molar conductivity of the solution. (3 marks) 9. Given the limiting molar conductivities of CH3COONa = 91.0, HCl = 425.9 and NaCl = 126.4 S cm2 mol^-1, calculate the limiting molar conductivity of acetic acid. If the molar conductivity of 0.001028 M acetic acid is 48.15 S cm2 mol^-1, calculate its degree of dissociation. (3 marks) 10. A current of 2 A is passed through a solution of AgNO3 for 30 minutes. Calculate the mass of silver deposited at the cathode (Ag = 108 g/mol, F = 96,500 C mol^-1). If the same charge is passed through CuSO4 solution, calculate the mass of copper deposited (Cu = 63.5 g/mol). (3 marks) Section C - Diagrams (10 marks) 11. Draw a labelled diagram of the Daniell cell showing the zinc and copper electrodes, the solutions, the salt bridge, the direction of electron flow and the direction of current. Write the half-cell reactions. (4 marks) 12. Draw a labelled diagram of the standard hydrogen electrode, showing the platinum foil coated with platinum black, hydrogen gas at 1 bar and the acid solution. (3 marks) 13. Draw a graph of molar conductivity versus the square root of concentration for a strong electrolyte (KCl) and a weak electrolyte (CH3COOH), and explain how the limiting molar conductivity of each is obtained. (3 marks) Section D - Analysis and Higher-order Thinking (15 marks) 14. Explain the products of electrolysis of (i) aqueous NaCl (ii) molten NaCl (iii) aqueous CuSO4 with copper electrodes (iv) aqueous CuSO4 with platinum electrodes. Explain why chlorine, rather than oxygen, is liberated at the anode in the electrolysis of aqueous NaCl, despite its higher reduction potential. (5 marks) 15. Describe the construction and working of the lead storage battery, with the reactions at the anode and cathode during discharge and charging. Analyse why it is called a secondary cell and why the density of the electrolyte falls during discharge. (5 marks) 16. Explain the electrochemical mechanism of rusting of iron with the half reactions. Analyse why rusting is faster in saline water and in the presence of moisture and carbon dioxide, and explain how galvanisation protects iron even when the zinc coating is scratched. (5 marks) Instructions: Time allowed 2 hours. Attempt all sections. Show all calculations with formulae and units. Use log and antilog values given in questions.