Galvanic cells and standard electrode potentials

Separating the half-reactions

Drop zinc into copper(II) sulfate and the electrons pass directly from zinc atoms to copper ions; the energy just warms the solution. In a galvanic (voltaic) cell, the two half-reactions happen in separate containers, called half-cells, so the electrons have to flow through an external wire, where they can do work. You built simple cells in Year 10 (Science lesson 4C); here we add the reasons and the numbers.

• Anode: the electrode where oxidation happens. In a galvanic cell it is the negative electrode, because electrons are released there. Zn(s) → Zn²⁺(aq) + 2e⁻
• Cathode: the electrode where reduction happens. It is positive in a galvanic cell. Cu²⁺(aq) + 2e⁻ → Cu(s)
• External circuit: electrons flow through the wire from anode to cathode. A voltmeter measures the potential difference.
• Salt bridge: filter paper or a tube soaked in an unreactive electrolyte such as KNO₃. Anions (NO₃⁻) move towards the anode and cations (K⁺) towards the cathode, balancing the charge that builds up in each beaker. Without it, the current stops almost at once.

A memory aid: An Ox and a Red Cat: anode, oxidation; reduction, cathode. Over time the zinc electrode loses mass and the copper electrode gains mass, while the blue colour in the copper half-cell fades.

Inert electrodes

Some half-reactions don't involve a solid metal at all: Fe³⁺/Fe²⁺, Cl₂/Cl⁻ or I₂/I⁻, for example. These half-cells use an inert electrode, usually platinum or graphite. It conducts electrons to or from the solution but doesn't react itself.

The standard hydrogen electrode and E°

We can't measure the potential of one half-cell on its own: a voltmeter needs two connections. So every half-cell is measured against the same reference, the standard hydrogen electrode (SHE), which is defined as 0.00 V.

The standard reduction potential, E°, of a half-cell is the voltage it gives when connected to the SHE under standard conditions: 25 °C, 100 kPa for gases, and 1.0 mol L⁻¹ for all solutions. The HSC data sheet lists these as reduction half-equations:

• A large positive E° (F₂, Cl₂, MnO₄⁻) means the species on the left is easily reduced: a strong oxidising agent.
• A large negative E° (K⁺, Mg²⁺) means the reduction hardly happens, so the metal on the right is easily oxidised: a strong reducing agent.

Compare the bottom of the table with the activity series from lesson 3C: K, Ca, Na, Mg, Al, Zn, Fe, Pb, then H, then Cu and Ag. The orders match closely. The table is the activity series with numbers attached, measured under controlled conditions. (Lithium, at −3.04 V, is the most negative of all, even though it reacts more gently than sodium; E° measures energy, not speed.)

Calculating the cell potential

The half-cell with the more positive E° is reduced (the cathode); the other is oxidised (the anode). The standard cell potential is

$$E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}$$

using both values exactly as written in the reduction table. Don't multiply E° values when you multiply a half-equation: a potential is energy per unit of charge, so it doesn't depend on how many electrons are in the equation.

Predicting whether a reaction will happen

The same arithmetic predicts whether any redox reaction is spontaneous. Write the reaction as it would happen, decide which species is reduced and which is oxidised, and calculate E° = E°(reduced species' couple) − E°(oxidised species' couple). If E° is positive, the reaction is spontaneous; if it is negative, it doesn't happen (the reverse reaction does).

A positive E° tells you a reaction can happen, not how fast. Aluminium should react strongly with water (it is far below hydrogen), but its oxide layer makes the rate negligible. Rates are the subject of the next two lessons.