19.1 Electrochemical cells

Voltaic Cells

  • Electromotive Force (EMF): The energy supplied by a source divided by the electric charge transported through the source
  • In voltaic cell, a cell potential is generated, resulting in movement of electrons from anode to cathode via external circuit.
  • Cell potential: the potential difference between the cathode and the anode when the cell is operating
  • Under standard conditions, cell potential is called Standard Cell potential
    • Eᶱcell = Eᶱcat – Eᶱan
  • In order to calculate Eᶱcell for a spontaneous cell, the cathode is taken as the more positive value from the two electrodes
  • The more positive one is also the strongest oxidizing agent

Standard Hydrogen Electrode (SHE)

  • Consists of an inert platinum electrode in contact with 1 mol dm-3 hydrogen ions and hydrogen gas at 100 kPa and 298 K. This is an example of a gas electrode
  • Standard electrode potential of a single half-cell cannot be measured on its own. Has to be relative to another cell
  • Standard electrode potentials are measured relative to SHE
  • SHE has Eᶱcell of 0V
  • The reduction half equation corresponding to the SHE cell is
    •  Hydrogen gas

Cell potential and Gibbs free energy

  • Spontaneous:
    • Eᶱcell is positive, ΔG is negative
  • Non-Spontaneous
    • Eᶱcell is negative, ΔG is positive
  • When ΔG is 0, Eᶱcell is 0
  • Both are related by following equation:
    • Gibbs & Cell potential
    • Where:
      • n= amount, in mol, of electrons
      • F = Faraday’s constant = 96500 C mol-1

Electrolytic Cells

  • Convert electrical to chemical energy
  • In SL, we looked at electrolysis of molten salt, now we will look at types of electrolysis
  • The higher the reduction potential, higher the tendency to react
  1. Electrolysis aqueous NaCl
    1. Concentrated
      • You have to take into account water as well
      • At cathode, water is reduced to create hydrogen gas
      • At anode, Cl– is oxidized to create Cl2 gas
    2. Diluted
      • At cathode, hydrogen ions are reduced to create hydrogen gas
      • At anode, water is oxidized to produce oxygen gas
      • This is equivalent to electrolysis of water
  2. Electrolysis of CuSO4
    1. Inert graphite (carbon) electrodes
      • Electrodes don’t take part in reactions
      • At cathode, copper ions are reduced to create copper deposits
      • At anode, water is oxidized to produce oxygen gas
    2. Active copper electrodes
      • Electrodes take part in reaction
      • At cathode, copper ions are reduced to create copper deposits
      • At anode, sludge of impurities is found
      • Process known as electrorefining in which the impurities in copper are separated from copper itself
      • Also the basis of electroplating in which a thin layer of metal is deposited onto cathode of another
  3. Electrolysis of water
    • Water is poor conductor of electricity
    • Electrolysis of water is done in dilute solutions of sulfuric acid or sodium hydroxide using inert Pt electrodes
    • At cathode, hydrogen ions are reduced to create hydrogen gas
    • At anode, water is oxidized to produce oxygen gas

Factors affecting amount of product formed

  1. Current
    • Higher the current, greater yield
    • Q=It
  2. Duration of electrolysis
    • Longer the time, greater yield
  3. Charge on the ion
    • Na+ required 1 mol of electrons however Pb2+ requires 2 mols of electrons

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