DPST1031 Chap.12 Redox, Galvanic Cells and Electrolysis
Redox, Galvanic Cells and Electrolysis
Electrochemistry brings the course together: electron transfer, equilibrium and Gibbs energy in one system. This chapter assigns oxidation numbers, identifies what is oxidised and reduced and which species are the oxidising and reducing agents, and balances redox equations with half-equations in acidic and basic solution.
It then builds galvanic cells, writes and reads cell diagrams, calculates standard cell potentials from reduction potentials, and links E° to ΔG° and K. The Nernst equation corrects potentials for non-standard concentrations, including concentration cells, and the chapter closes with corrosion, batteries and electrolysis calculations with the Faraday equation.
What this chapter covers
- 01
Oxidation numbers and their rules
- 02
Oxidising and reducing agents
- 03
Balancing half-equations in acid and base
- 04
Galvanic cells, anode and cathode
- 05
Cell diagrams and the salt bridge
- 06
Standard reduction potentials and E°cell
- 07
E°, ΔG° and K
- 08
Nernst equation, corrosion, batteries and electrolysis
Worked example · free
Reading a zinc and copper galvanic cell
- +1Copper has the higher reduction potential, so Cu²⁺ is reduced at the cathode; zinc is oxidised at the anode.
- +1Cell diagram: Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s).
- +1E°cell = E°cathode − E°anode = 0.34 − (−0.76) = 1.10 V.
- +1Electrons flow through the external wire from the zinc anode to the copper cathode.
Key terms
- Oxidation number
- A bookkeeping charge assigned to an atom to track electron transfer in a reaction.
- Oxidising agent
- The species that is reduced and so removes electrons from another species.
- Half-equation
- An equation showing only the oxidation or only the reduction part of a redox reaction, with electrons.
- Galvanic cell
- A cell in which a spontaneous redox reaction produces an electric current.
- Salt bridge
- An electrolyte connection that lets ions move between half-cells to keep them neutral.
- Standard hydrogen electrode
- The reference half-cell assigned a potential of exactly zero volts.
- Electrolytic cell
- A cell in which an external power supply drives a non-spontaneous redox reaction.
Redox, Galvanic Cells and Electrolysis FAQ
How do I know which electrode is the anode?
In a galvanic cell, the half-cell with the lower standard reduction potential is reversed and becomes the anode, where oxidation happens. The half-cell with the higher potential is the cathode, where reduction happens. Oxidation is always at the anode in every type of cell.
Why do I not multiply E° by the coefficients?
A cell potential is energy per unit charge, an intensive quantity. Doubling a half-equation doubles the electrons and the energy together, so their ratio, the potential, is unchanged. Gibbs energy does scale, which is why ΔG° = −nFE° contains n.
What does the Nernst equation add?
It corrects the standard potential for actual concentrations through the reaction quotient Q. As a galvanic cell runs, products build up, Q rises and the cell potential falls, reaching zero when Q equals K and the cell is at equilibrium, which is when a battery is flat.
How do I calculate the mass deposited in electrolysis?
Find the charge with Q = It, using time in seconds. Divide by the Faraday constant to get moles of electrons, use the half-equation to convert to moles of metal, then multiply by the molar mass. The electron ratio in the half-equation is the step most often missed.
Why is the anode negative in a galvanic cell but positive in electrolysis?
Oxidation always happens at the anode. In a galvanic cell the spontaneous oxidation releases electrons there, so it is the negative terminal. In electrolysis an external supply pulls electrons away from the anode to force oxidation, so it is connected to the positive terminal of the supply.
Exam move
Start with oxidation numbers until you can assign them to any atom in an ion in seconds. Then balance ten redox reactions with half-equations, five in acid and five in base, checking atoms and charge every time. For cells, practise moving between a balanced equation, a cell diagram and a labelled sketch with anode, cathode, electron flow and salt bridge.
Next calculate E°cell, ΔG° and K for several pairs, and apply the Nernst equation, including to concentration cells. Finish with Faraday calculations in both directions, finding mass from current and time, and time from a required mass.
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