Electrochemistry definitions

Electrode potentials and the standard hydrogen electrode

Half-cell

An element in contact with a solution of its ions, or an inert electrode in a solution of an oxidised and a reduced form, at which a redox equilibrium is set up.

Electrode potential

The potential difference between an electrode and the solution it is in; it can be measured only against a second electrode.

Standard hydrogen electrode

A half-cell in which hydrogen gas at 100 kPa is bubbled over a platinum electrode in a solution of H⁺ ions at 1.00 mol dm⁻³ at 298 K; its electrode potential is defined as 0.00 V.

Standard electrode potential

The EMF of a cell made from the half-cell and the standard hydrogen electrode, measured under standard conditions.

Standard conditions

A temperature of 298 K, a pressure of 100 kPa for any gas, and a concentration of 1.00 mol dm⁻³ for every ion in the half-equation.

Inert electrode

An electrode, usually platinum, that conducts electrons to and from a redox equilibrium in the solution without reacting or setting up an equilibrium of its own; used when neither species of the couple is a metal.

Conventional representation

A one-line description of a cell: the electrodes at the ends, single vertical lines for phase boundaries, a double line for the salt bridge, and the oxidised form of each couple next to the salt bridge.

Standard cell potential

The EMF of a cell in which both half-cells are under standard conditions, E°cell = E°right − E°left.

Salt bridge

A connection between two half-cells, usually filter paper soaked in saturated potassium nitrate solution, that allows ions to move between them to complete the circuit without the solutions mixing.

EMF

The potential difference between the two electrodes of a cell when no current flows.

The electrochemical series and predicting feasibility

Electrochemical series

A list of half-equations written as reductions, in order of their standard electrode potentials.

Oxidising agent

A species that accepts electrons from another species and is itself reduced.

Reducing agent

A species that donates electrons to another species and is itself oxidised.

Feasible

Able to take place; a redox reaction is feasible under standard conditions when E°cell, calculated with the half-cell in which reduction occurs on the right, is positive.

Disproportionation

A redox reaction in which the same element in one species is simultaneously oxidised and reduced.

Kinetically stable

Describes a mixture that does not react at a measurable rate although the reaction is feasible, because the activation energy is high.

Nernst equation

E = E° + (0.059/z) log₁₀([oxidised form] / [reduced form]) at 298 K, giving the electrode potential of a half-cell at concentrations that are not standard.

Faraday constant

The charge on one mole of electrons, 96 500 C mol⁻¹.

Commercial cells, rechargeable cells and fuel cells

Non-rechargeable cell

A cell in which the reaction cannot be reversed, so the cell is discarded when its reactants are used up.

Rechargeable cell

A cell in which the reaction can be reversed by applying an external voltage greater than its EMF, regenerating the reactants.

Fuel cell

A cell that generates an electric current from a fuel and oxygen supplied continuously from outside, and does not need to be electrically recharged.

Electrolyte

The ionic solution, melt or solid between the electrodes of a cell, through which ions move to complete the circuit.

Anode

The electrode at which oxidation takes place; in a cell delivering a current it is the negative electrode.

Cathode

The electrode at which reduction takes place; in a cell delivering a current it is the positive electrode.

Electrolysis

The decomposition of an ionic compound, molten or in solution, by an electric current from an external supply.