Concepts: redox reactions and electrode potential

4.1 From electron transfer to electricity

The reaction between zinc and copper ions can be run directly, in which case the chemical energy appears as heat, or in a cell where the oxidation and reduction are separated and the electrons flow through an external wire. In the Daniell cell, a zinc rod dips into zinc sulphate solution and a copper rod dips into copper sulphate solution, and the two solutions are joined by a salt bridge.

Zinc, being more reactive, loses electrons and becomes zinc ions, so the zinc electrode is the anode. The electrons travel through the external wire to the copper electrode, where copper ions gain electrons and deposit as copper metal, so the copper electrode is the cathode. The salt bridge completes the circuit by allowing ions to migrate and keeps the solutions electrically neutral.

The figure shows the cell with the direction of electron flow and the direction of conventional current marked. Current flows opposite to the electrons by convention, and this is a favourite point in questions. The cell converts the free energy of the redox reaction into electrical work.

4.2 Standard electrode potentials

The tendency of a substance to gain electrons is measured by its electrode potential. A single electrode has no absolute potential, so the standard hydrogen electrode is assigned zero and all other potentials are measured against it under standard conditions of one molar concentration, one bar pressure and 298 K. A positive standard reduction potential means that the substance is reduced more readily than hydrogen ions.

The standard reduction potentials give a quantitative version of the activity series. Fluorine has a very large positive value and is the strongest oxidising agent, while lithium has a very negative value and is a strong reducing agent. A species with a more positive reduction potential is reduced by a species with a more negative value, and the reaction proceeds in that direction.

For a cell built from two couples, the cell potential is the reduction potential of the cathode minus the reduction potential of the anode. A positive cell potential means the cell reaction as written is spontaneous, and the free energy change is negative. The cell potential is an intensive property, so it does not change when the equation is multiplied by a factor, even though the free energy does.

The standard hydrogen electrode is a platinum plate coated with platinum black and dipped in one molar hydrogen ion solution, with hydrogen gas at one bar bubbled over it. It is assigned a potential of zero at all temperatures, and every other standard potential is quoted as a reduction potential against it. The salt bridge in a practical cell is usually a tube of potassium chloride or potassium nitrate in agar, and its role is to let ions move while preventing the two solutions from mixing quickly.

The Nernst equation extends the cell potential to non-standard concentrations by adding a term that depends on the reaction quotient and temperature. At equilibrium the cell potential becomes zero and the quotient equals the equilibrium constant, which links the standard potential to the equilibrium constant of the redox reaction. For school numericals the standard potential and the sign test are usually enough, but the qualitative statement that the potential falls as the reaction approaches equilibrium is worth remembering.

4.3 Feasibility and the redox series

A redox reaction is feasible when its cell potential is positive. This test explains the order of displacement reactions: zinc displaces copper because the zinc couple has a more negative potential, and copper displaces silver for the same reason. It also explains why some reactions need an acidic medium, because the hydrogen ion concentration affects the potential of couples such as permanganate and dichromate.

Comparing potentials also orders oxidising and reducing strength. The strongest oxidising agents are the species with the most positive reduction potentials, such as fluorine, permanganate in acid and dichromate in acid. The strongest reducing agents are the species with the most negative values, such as lithium and potassium. A reaction between a strong oxidising agent and a strong reducing agent has a large positive cell potential and goes almost to completion.

Electrode potential is a thermodynamic quantity, so it says nothing about how fast the reaction goes. A reaction with a positive cell potential may be very slow at room temperature, and this is the basis of kinetic control in many redox systems. The chapter's emphasis, however, is the thermodynamic feasibility, which is decided by the sign of the cell potential. The table of standard reduction potentials in the textbook is worth learning in its order, because most feasibility questions can then be answered by a glance at the relative positions of the two couples. Read the potentials carefully before deciding which species is oxidised and which is reduced.