Electrochemical cells · Year 11
An electrochemical cell turns the energy of a redox reaction into electrical energy. Two half-cells, each a metal (or platinum) dipping into a solution of ions, are joined by a wire and a salt bridge. The half-cell that releases electrons more easily becomes the negative electrode. Standard electrode potentials, E°, measured against the standard hydrogen electrode, tell you which way the electrons flow and the e.m.f. of the cell.
Key facts
- A half-cell is written as a reduction: Zn²⁺(aq) + 2e⁻ ⇌ Zn(s). Its standard electrode potential, E°, is measured against the standard hydrogen electrode (E° = 0.00 V) at 298 K, 100 kPa and 1.00 mol dm⁻³.
- The more negative E° is, the more easily the metal loses electrons (a stronger reducing agent). The more positive E° is, the more easily the oxidised form gains electrons (a stronger oxidising agent).
- In a cell the half-cell with the more positive E° is the positive electrode, where reduction happens; the other is the negative electrode, where oxidation happens. Electrons flow through the wire from the negative to the positive electrode.
- E°cell = E°(positive electrode) − E°(negative electrode). In cell notation, Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s), the right-hand electrode is positive: E°cell = E°right − E°left.
- The salt bridge (filter paper soaked in potassium nitrate solution) lets ions move between the half-cells to balance the charge, completing the circuit.
- A reaction is feasible if E°cell is positive: an oxidising agent on the left of one half-equation reacts with a reducing agent on the right of a half-equation with a less positive E°.
- Non-rechargeable cells stop when a reactant is used up; rechargeable cells such as lithium-ion cells are reversed by charging; fuel cells are fed with fuel, and a hydrogen–oxygen fuel cell makes only water.