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Electrolysis is the process that produces a non-spontaneous redox reaction by the passage of electrical energy from a power supply through a conducting liquid thereby converting it into chemical energy. It is essentially the opposite of a galvanic cell. The applications of electrolysis include the recharging of a secondary cell, production of reactive metals, purifying metals and electroplating. It still occurs between the strongest reductant and the strongest oxidant, but now it simply has an uphill gradient, so the smaller the distance the better and it is important to note that molten electrolytes are used in order to remove water so that it can’t cause side-reactions. However, energy must be expended to maintain the electrolyte in a molten state

Electrolytic and Galvanic Cells

In order to recharge a secondary cell i.e. a cell that can be recharged because the products are still in contact with the electrode, we must supply a voltage higher than the discharging voltage. For example, if a car battery produced 12V, then we must supply a voltage of around 13V in order to recharge it.

Galvanic Cells Electrolytic Cells
Produces Electricity Consumes Electricity
Spontaneous Reactions Non-spontaneous Reactions
Chemical to Electrical Energy Electrical to Chemical Energy
Anode is negative; Cathode is positive Anode is positive; Cathode is negative



Electroplating is a process that uses electrolysis to deposit a layer of metal on the surface of another material. The object to be plated is connected by a wire to the negative terminal of the power supply, becoming the negative electrode in the cell. When the cell is in operation, the power supply acts as an ‘electron pump’, pushing electrons onto the negative electrode and removing electrons from the positive electrode. Therefore, in the case of electroplating \(Sn\) onto a can, the can would be submerged within a \(Sn(NO_3)_2\) solution, and the \(Sn^{2+}\) ions in the solution would receive the electrons pumped over at the cathode to coat a layer of \(Sn\) on the can.

Faraday’s Laws

The mass of a substance produced or consumed at an electrode in an electrochemical cell is directly proportional to the electrical charge passed through the cell.

  • \(Q=I\times t\)
  • \(Q=\text{Charge (Coulombs)}\)
  • \(I=\text{Current in Amps}\)
  • \(t=\text{time in seconds}\)
In order to produce one mole of a substance at an electrode in an electrochemical cell, one, two, three or another whole number of moles of electrons must be consumed.
  • \(Q=n(e^-)\times F)\)
  • \(Q=\text{Charge (Coulombs)}\)
  • \(n(e^-)=\text{mols of electrons}\)
  • \(F=96500\text{ (Faraday’s Constant)}\)