Understand Electron Transfer Reactions — Understand Reactivity 3.2—Electron transfer reactions

Humans have observed electron transfer reactions for thousands of years without understanding the underlying chemistry. Ancient metalworkers extracted copper and iron from ores using charcoal fires, unknowingly driving redox reactions — reactions in which electrons are transferred from one species to another. The rusting of iron, the tarnishing of silver, and the bleaching of dyes all share the same fundamental mechanism: the movement of electrons between chemical species.

It was not until the eighteenth and nineteenth centuries that scientists began to piece together the theory behind these transformations. The development of the electrochemical cell, the concept of oxidation states, and eventually the discovery of the electron itself all contributed to our modern understanding. Today, electron transfer reactions underpin technologies from smartphone batteries to water purification, and they are essential to biological processes such as cellular respiration and photosynthesis.

The central question this topic addresses is straightforward yet profound: what happens at the particle level when electrons move from one chemical species to another, and how can we predict and harness those changes?

Electron transfer reactions are built on a handful of interconnected ideas. Understanding each one is essential before you can analyse or predict redox behaviour in the lab or on an exam. Let's define the key concepts.

A clear diagram makes redox chemistry far easier to grasp. Below, you can see the classic reaction between zinc metal and copper(II) sulfate solution. Zinc atoms donate electrons to copper ions, causing zinc to dissolve and copper metal to deposit. This is the reaction that takes place when you place a piece of zinc into blue CuSO₄ solution and watch it turn colourless as copper metal appears.

Notice that the number of electrons lost by zinc equals the number gained by copper. This is a fundamental rule: electrons are conserved in redox reactions. You can never create or destroy electrons; they simply move from one species to another. Writing separate half-equations for each process and then combining them ensures that electron bookkeeping is correct.

Half-equations are the toolbox of redox chemistry. Each half-equation isolates either the oxidation or the reduction step, showing the electrons explicitly. When you combine two balanced half-equations so that the electrons cancel, you obtain the overall balanced redox equation.

Oxidation Half-Equation

Reduction Half-Equation

Combining Half-Equations

When the number of electrons in the two half-equations is not the same, you need to multiply one or both half-equations by appropriate integers so that the electrons cancel. For example, if one half-equation shows 2e⁻ and the other shows 3e⁻, you would multiply the first by 3 and the second by 2 to give 6e⁻ in each.

Balancing in Acidic or Neutral Aqueous Solutions

1. Step 1: Assign oxidation states to identify what is oxidised and what is reduced.
2. Step 2: Write separate, unbalanced half-equations for oxidation and reduction.
3. Step 3: Balance atoms other than O and H first, then balance O by adding H₂O, and H by adding H⁺ (in acidic solution).
4. Step 4: Balance charge by adding electrons to the more positive side.
5. Step 5: Equalise electron count, add the two half-equations, and cancel species that appear on both sides.

Not every metal reacts with every metal ion. Whether an electron transfer reaction proceeds spontaneously depends on the relative reactivity of the metals involved. The activity series (also called the reactivity series) ranks metals from most reactive to least reactive. A metal higher in the series will displace ions of any metal lower in the series from solution.

Using the series, you can predict that zinc will displace copper from a CuSO₄ solution because zinc is above copper. However, copper will not displace zinc from ZnSO₄ because copper is below zinc. Similarly, metals above hydrogen in the series react with dilute acids (such as HCl) to produce hydrogen gas, whereas metals below hydrogen — like copper, silver, and gold — do not.

Let's work through the reaction between iron metal and silver nitrate solution step by step. When iron nails are placed in AgNO₃(aq), a displacement reaction occurs and silver crystals form on the iron surface.

Electron transfer reactions can be harnessed in two main types of electrochemical cells. In a voltaic (galvanic) cell, a spontaneous redox reaction generates electrical energy — this is how batteries work. In an electrolytic cell, an external power source forces a non-spontaneous reaction to occur — this is how electroplating and the extraction of aluminium are carried out.

The qualitative reactivity series is useful, but IB Chemistry also expects you to work with quantitative data: standard electrode potentials (E°). Every half-equation in the IB Data Booklet is written as a reduction, and its E° value (in volts) tells you how strongly that species attracts electrons compared to the standard hydrogen electrode (SHE), which is defined as 0.00 V.

At Higher Level, you will also encounter the relationship ΔG° = −nFE°, which connects the cell potential to the Gibbs free energy change, where n is the number of moles of electrons transferred and F is the Faraday constant (96 485 C mol⁻¹). For now, the key idea is that a positive cell potential means the reaction is thermodynamically favoured under standard conditions.