Describe the construction and operation of a galvanic cell, including the role of the salt bridge, the conventions of anode and cathode, and the calculation of standard cell potentials from the standard reduction potential table

What this dot point is asking

QCAA wants you to describe a galvanic cell at the level of components (electrodes, electrolytes, salt bridge, external circuit), identify the anode and cathode with their conventional signs, write the half-equations and overall ionic equation, calculate the standard cell potential from the reduction potential table, and explain the role of the salt bridge. This dot point underpins IA2 student experiments (constructing and measuring cell potentials) and feeds the EA short and extended-response questions.

What a galvanic cell is

A galvanic (voltaic) cell converts the chemical energy of a spontaneous redox reaction into electrical energy. The defining trick is to physically separate the oxidation and reduction half-reactions into two half-cells, forcing the electrons to travel through an external wire rather than transferring directly. The current in the wire can do useful work.

A galvanic cell is the textbook battery. The Zn/Cu Daniell cell is the canonical example.

Components and their roles

Anode and cathode

The two electrodes are named after the half-reaction occurring at them, regardless of cell type.

• Anode. Site of oxidation. In a galvanic cell, the anode is the negative terminal (it releases electrons to the external wire).
• Cathode. Site of reduction. In a galvanic cell, the cathode is the positive terminal (it accepts electrons from the external wire).

Memory aid: AnOx, RedCat. Anode oxidation, Reduction at the Cathode. Note that in an electrolytic cell the polarity is reversed (anode positive, cathode negative), but the oxidation-reduction assignment is the same.

Electron direction in the external wire: anode -> cathode.

Ion direction in the salt bridge: anions migrate toward the anode (to balance positive build-up); cations migrate toward the cathode (to replace positive charge being consumed).

Using the standard reduction potential table

The table lists half-reactions written as reductions, with associated E0 values. Highly positive E0 means the species is a strong oxidising agent (readily reduced). Highly negative E0 means the species is a strong reducing agent (readily oxidised; the reverse direction is favoured).

To predict spontaneity:

1. Look up the standard reduction potential for each half-reaction.
2. The half-reaction with the more positive E0 occurs as written (reduction, at the cathode).
3. The half-reaction with the less positive (more negative) E0 occurs in reverse (oxidation, at the anode).
4. Calculate the cell potential:

Both E0 values are taken from the table as reductions (sign as listed). If E0 cell is positive, the reaction is spontaneous as written. If negative, the reaction is non-spontaneous; the reverse direction would be spontaneous.

Standard conditions: 25 degrees C, 1.0 mol/L aqueous concentrations, 1 atm gas pressures, 1 mol/L hydrogen ion. Non-standard concentrations affect the actual potential (Nernst equation, treated in Unit 4 contexts; not required by name in Unit 3).

Worked cell potential calculations

Daniell cell (Zn/Cu). E0(Cu2+/Cu) = +0.34 V; E0(Zn2+/Zn) = -0.76 V.

• More positive: Cu2+. Reduced at cathode.
• More negative: Zn2+. Reversed; Zn oxidised at anode.
• E0 cell = +0.34 - (-0.76) = +1.10 V. Spontaneous.

Ag/Cu cell. E0(Ag+/Ag) = +0.80 V; E0(Cu2+/Cu) = +0.34 V.

• More positive: Ag+. Reduced at cathode.
• Cu2+ reversed; Cu oxidised at anode.
• E0 cell = +0.80 - 0.34 = +0.46 V.

Sn/Ag cell. E0(Ag+/Ag) = +0.80 V; E0(Sn2+/Sn) = -0.14 V.

• E0 cell = +0.80 - (-0.14) = +0.94 V.

Cell diagram notation

The conventional shorthand for a galvanic cell, used in QCAA EA and IA1:

• Single vertical bar: phase boundary between solid electrode and its solution.
• Double vertical bar: salt bridge.
• Anode on the left, cathode on the right.
• Concentrations in brackets where specified.

Daniell cell:

Cells with inert electrodes (Pt for solution-only redox, such as Fe2+/Fe3+) use Pt as the conducting surface but Pt itself is not involved.

(Solution-only species separated by commas, not vertical bars.)

Role of the salt bridge in detail

Without a salt bridge:

1. Zn -> Zn2+ + 2 e- proceeds initially.
2. Zn2+ accumulates in the anode compartment (positive charge build-up).
3. Cu2+ + 2 e- -> Cu proceeds initially.
4. SO4^2- (from CuSO4) is left without Cu2+ partners in the cathode compartment (net negative charge).
5. Charge separation builds an electric field opposing further electron flow.
6. The current drops to zero within fractions of a second.

The salt bridge prevents this by supplying mobile ions. NO3- migrates into the anode compartment to balance Zn2+. K+ migrates into the cathode compartment to balance SO4^2-. Charge neutrality is preserved, so the redox reaction continues and the current is sustained.

The salt bridge ions are chosen so they do not participate in either half-reaction. KNO3 is the QCAA default; KCl is avoided if Ag+ is in the cathode compartment (would form AgCl precipitate).

Predicting spontaneity from the table

The reduction potential table is a kind of "redox priority list". The species higher in the table (more positive E0) wins as the oxidising agent; the species lower (more negative E0) loses electrons.

Examples of predicting:

• Mix Cu metal with FeSO4 solution. E0(Cu2+/Cu) = +0.34 V; E0(Fe2+/Fe) = -0.44 V. Cu is below Fe2+ as a reducing agent; the reaction Cu + Fe2+ -> Cu2+ + Fe has E0 = -0.44 - 0.34 = -0.78 V. Non-spontaneous. Nothing happens.
• Mix Zn metal with CuSO4 solution. E0 cell = +0.34 - (-0.76) = +1.10 V. Spontaneous. Zn dissolves, Cu plates out.

This is exactly the kind of reasoning QCAA tests in IA1 short response when given a stimulus of two metals and two ion solutions and asked which reactions occur.

Connecting back to equilibrium

A galvanic cell delivers current until equilibrium is reached. The reaction proceeds in the spontaneous direction, but the consequent change in ion concentrations (Zn2+ rises, Cu2+ falls) shifts the actual cell potential toward zero. When the cell is fully discharged, the half-cells are at electrochemical equilibrium and the voltage is zero. This is the chemistry of a flat battery.

This continuity between equilibrium (Topic 1) and redox (Topic 2) is a frequent cross-topic EA Paper 2 question.

Worked example: an IA2 design

A common IA2: design and construct a galvanic cell, measure E0 with a high-impedance voltmeter, compare to the theoretical value from the table, and account for any discrepancy.

Strong report features:

• Clear cell schematic with the salt bridge labelled.
• Theoretical E0 calculated from the table.
• Measured E0 reported with uncertainty (voltmeter precision).
• Discussion of why measured may differ: non-standard concentrations (Nernst), surface contamination of electrodes, salt bridge electrolyte choice affecting junction potential, internal resistance.
• Conclusion linking observed potential to the spontaneity criterion.

The IA2 criteria reward the design justification, the data, and the evaluation, in that order. Knowing the dot-point chemistry cold is the foundation that lets the higher-cognitive criteria succeed.

Common traps

Calling the anode positive in a galvanic cell. It is negative. The convention flips in electrolytic cells.

Computing E0 cell as E0(cathode) + E0(anode). Subtract, not add. Both E0 values are taken from the table as reductions.

Forgetting to use the table values as-is. Do not flip the sign of E0(anode) before subtracting; the subtraction handles the reversal.

Confusing the salt bridge with the wire. Wire carries electrons (in the external circuit). Salt bridge carries ions (between the half-cells). Different particles, different paths.

Using a salt bridge ion that participates. KCl with Ag+ forms AgCl precipitate; KNO3 is the safe default.

Treating a negative E0 cell as "doesn't happen". It means the forward direction is non-spontaneous; the reverse is spontaneous. State which direction will actually occur.

In one sentence

A galvanic cell drives a spontaneous redox reaction by separating the oxidation half (at the negative anode) from the reduction half (at the positive cathode), connecting them by an external wire that carries electrons from anode to cathode and a salt bridge that carries ions to keep both half-cells charge-neutral, with the standard cell potential calculated as E0(cathode) - E0(anode) using the reduction potential table and a positive result confirming spontaneity.