Electrochemical Cells and Cell Potentials

Full Notes

The background theory of electrochemical cells and half-cells for Edexcel A-level chemistry is covered here.

The standard cell potential is the overall voltage produced when two half-cells are connected under standard conditions.

It’s calculated by:

Note:

• The cathode is the half-cell where reduction happens.
• The anode is the half-cell where oxidation happens.

Meaning you can also write this as:

Cell Notation

In electrochemistry, cells are written using a shorthand:

• A single vertical line (|) separates different phases (solid, liquid, aqueous).
• A double line (||) represents the salt bridge.
• The anode (oxidation) is written on the left, and the cathode (reduction) on the right.

For example Write the conventional cell notation for an electrochemical cell made from the following:

Zn²⁺(aq) + 2e⁻ ⇌ Zn(s)  E° = −0.76 V
Cu²⁺(aq) + 2e⁻ ⇌ Cu(s)  E° = +0.34 V

We’ve already established (see above) Cu²⁺/Cu is the cathode, where reduction happens. Cu²⁺ will be reduced to Cu. Equally, Zn²⁺/Zn is the anode, where oxidation happens, Zn will be oxidised to Zn²⁺.

Anode is written on the left with the Zn(s) and Zn²⁺(aq) separated by a vertical line as they are in different phases. Cathode is written on the right with the Cu²⁺(aq) and Cu(s) again separated by a vertical line.

Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s)

Predicting Thermodynamic Feasibility

The link between feasibility, the equilibrium constant K and electrochemical cells has been outlined in detail here. However, on this page is just what you need to know for Edexcel A-level, don’t worry too much about the extra background - just learn the equations and practice using them :)

You can predict whether a redox reaction is feasible (able to happen) using the calculated standard cell potential:

If E°_cell is positive, the reaction is thermodynamically feasible under standard conditions.

However, a feasible reaction may not happen due to kinetic barriers like high activation energy.

Link to Entropy and Equilibrium Constant

Gibbs free energy change relates to the cell potential:

Where:

• ΔG° = Gibbs free energy change (J mol⁻¹)
• n = number of electrons transferred
• F = Faraday constant (96500 C mol⁻¹)
• E°_cell = standard cell potential (V)

If ΔG° is negative, the reaction is spontaneous.

So if E°_cell is large and positive, ΔG is large and negative.

This means the forward direction (production of products) of the reaction is favoured, giving a large equilibrium constant, K.

ΔG° and K (Equilibrium Constant) are linked by:

• To summarise – if E°_cell is positive, then ΔG is negative and K is large meaning products favoured at equilibrium.

Limitations of E° Predictions

Even if E°_cell suggests a reaction is feasible, the following factors may prevent it:

• Kinetic Factors:
  • A reaction may be too slow (e.g. due to a high activation energy barrier).
  • Without a catalyst, the reaction may not occur at a useful rate.
• Departure from Standard Conditions:
  • E° values are only valid at 298 K, 100 kPa, and 1.00 mol dm⁻³.
  • Changes in concentration, temperature or pressure can shift the position of equilibrium, altering the cell potential.
  • For example If concentrations change, then the actual potentials of each half-cell will change.
    • Increasing the concentration of oxidised species in the half-cell makes E more positive.
    • Increasing the concentration of reduced species in the half-cell makes E more negative.
  • This affects how easily redox reactions occur under non-standard conditions.

So, a reaction predicted as feasible might not occur in practice, or might proceed only very slowly.

Electrochemical Series

The electrochemical series is a list of half-cells arranged in order of increasing E° values:

• Top (most negative): strong reducing agents (e.g. metals like Na, Zn)
• Bottom (most positive): strong oxidising agents (e.g. halogens)

The series helps us predict which species will be oxidised or reduced in a reaction.

Disproportionation Reactions

In a disproportionation reaction, one species is both oxidised and reduced.

These reactions are likely when:

• E° for reduction is positive
• E° for oxidation (reverse reaction) is less positive or negative

Summary

• E°_cell = E°(cathode) minus E°(anode) under standard conditions.
• Electrons flow from anode to cathode and positive E°_cell indicates feasibility.
• Higher E° means stronger oxidising agent and lower E° means stronger reducing agent.
• ΔG° = −nFE°_cell and ΔG° = −RT ln K link cell potential to spontaneity and equilibrium.
• Predictions can fail in practice due to kinetics or non-standard conditions affecting potentials.