Redox and Fuel Cells

A redox reaction is any reaction in which electrons are transferred between species. The name comes from combining reduction and oxidation, the two halves of the process that always happen together. One substance loses electrons while another gains them, so oxidation and reduction can never occur on their own.

The easiest way to remember the definitions is the mnemonic OIL RIG:

• Oxidation Is Loss of electrons
• Reduction Is Gain of electrons

For example, when zinc metal reacts with copper(II) ions, zinc atoms lose electrons and copper ions gain them:

\( Zn \)→\( Zn^{2+} + 2e^- \) (oxidation)

\( Cu^{2+} + 2e^- \)→\( Cu \) (reduction)

Oxidation states (oxidation numbers) let you follow electron transfer even in reactions that do not involve simple ions. A species is oxidised if its oxidation state increases, and reduced if its oxidation state decreases.

Key rules to assign oxidation states:

• An uncombined element has an oxidation state of 0.
• A simple ion has an oxidation state equal to its charge.
• Oxygen is usually \(-2\), and hydrogen is usually \(+1\).
• The oxidation states in a neutral compound sum to 0, and in an ion they sum to the ion's charge.

Some redox reactions fall into recognisable categories, such as combustion, displacement, and disproportionation, where the same element is simultaneously oxidised and reduced. Being able to identify the type of redox reaction quickly helps you decide how to balance it.

Every redox reaction can be split into two half-equations, one for oxidation and one for reduction. Balancing a redox reaction means writing each half-equation so that atoms and charge both balance, then combining them so the electrons cancel exactly.

Steps for balancing a redox reaction in acidic solution:

1. Balance the atoms of the element being oxidised or reduced.
2. Balance oxygen atoms by adding \( H_2O \).
3. Balance hydrogen atoms by adding \( H^+ \).
4. Balance charge by adding electrons, \( e^- \).
5. Multiply each half-equation so the number of electrons matches, then add them together and cancel any species appearing on both sides.

For example, combining the oxidation of iron(II) with the reduction of manganate(VII) ions gives:

\( Fe^{2+} \)→\( Fe^{3+} + e^- \)

\( MnO_4^- + 8H^+ + 5e^- \)→\( Mn^{2+} + 4H_2O \)

Multiplying the first half-equation by 5 and adding gives the overall equation:

\( MnO_4^- + 8H^+ + 5Fe^{2+} \)→\( Mn^{2+} + 4H_2O + 5Fe^{3+} \)

This exact combination of half-equations is also the basis of redox titrations, where a known concentration of one reagent is used to find the concentration of the other by tracking the colour change as electrons are transferred.

A hydrogen fuel cell converts the chemical energy of a redox reaction directly into electrical energy, without burning the fuel. Hydrogen gas is oxidised at one electrode and oxygen gas is reduced at the other, with the overall effect being the formation of water and a flow of electrons through an external circuit.

In an acidic (proton exchange membrane) fuel cell, the electrode half-equations are:

At the negative electrode (oxidation): \( 2H_2 \)→\( 4H^+ + 4e^- \)

At the positive electrode (reduction): \( O_2 + 4H^+ + 4e^- \)→\( 2H_2O \)

Adding these gives the overall cell reaction: \( 2H_2 + O_2 \)→\( 2H_2O \)

In an alkaline fuel cell, hydroxide ions carry the charge through the electrolyte instead of protons, so the half-equations are written slightly differently, but the overall reaction and the net transfer of electrons are the same.

Because a fuel cell keeps oxidation and reduction physically separated at two electrodes, the electron transfer is forced to travel through an external circuit rather than happening directly between particles, which is exactly what allows the redox reaction to do useful electrical work. This is the same principle used in the chemical cells you meet elsewhere in electrochemistry, and the voltage a cell produces can be predicted using standard electrode potentials, covered in introduction to cell potential.

Compared with burning hydrogen directly, a fuel cell produces electrical energy with only water as a product and no flame, which is why hydrogen and fuel cells are widely discussed as a cleaner alternative for vehicles and power generation. The main practical challenges are storing hydrogen safely and producing it efficiently in the first place.

When you meet a new equation, ask which species has changed oxidation state. If one element's oxidation state goes up and another's goes down, it is a redox reaction. If no oxidation states change (as in most acid-base reactions), it is not. Practising this check on a range of examples, including displacement reactions, combustion, and fuel cell electrode reactions, is the fastest way to become confident with redox chemistry.

Related lessons

• Uk A Level Chemistry