Redox Reactions and Electric Current Production
A redox reaction is a reaction in which electrons are transferred from one substance (which is oxidized) to another (which is reduced), and when the two halves of the reaction are kept apart and connected by a wire, the electron transfer becomes an electric current. In brief: in Zn + Cu2+ → Zn2+ + Cu, zinc loses two electrons and copper(II) gains them; in a Daniell cell this gives a voltage of 1.10 V and a free energy of −212 kJ per mole of zinc.
Electrochemistry connects chemical reactions with electricity. Batteries convert chemical energy into electrical energy, and electrolysis uses electrical energy to drive chemical change. The common starting point is the redox reaction. This article reviews oxidation, reduction and oxidation numbers, explains how separating the half-reactions produces a current, introduces the units of charge and current and works through examples.
What are oxidation and reduction?
- Oxidation is the loss of electrons (an increase in oxidation number).
- Reduction is the gain of electrons (a decrease in oxidation number).
- The reducing agent gives electrons and is itself oxidized. The oxidizing agent takes electrons and is itself reduced.
A mnemonic is OIL RIG: Oxidation Is Loss, Reduction Is Gain (of electrons). The two processes always occur together, since electrons cannot be lost without being gained somewhere. The overall reaction is the sum of two half-reactions, for example:
- Oxidation half-reaction: Zn → Zn2+ + 2e−
- Reduction half-reaction: Cu2+ + 2e− → Cu
- Overall: Zn + Cu2+ → Zn2+ + Cu
How do you identify oxidation and reduction?
Assign oxidation numbers to the atoms in the reactants and products. The atoms whose oxidation number increases are oxidized, and those whose number decreases are reduced. The main rules are: elements have an oxidation number of 0; monatomic ions have their charge; oxygen is usually −2 and hydrogen +1 in compounds; the sum of the oxidation numbers equals the overall charge of the species.
How does a redox reaction produce a current?
If a zinc rod is put into a copper(II) sulfate solution, electrons pass directly from Zn to Cu2+ at the surface of the rod, copper deposits and the energy of the reaction is released as heat. If, however, the zinc and copper(II) are kept in separate compartments, connected by a wire, the electrons must pass through the wire to get from Zn to Cu2+, which makes a useful electric current. A salt bridge or a porous barrier lets ions move between the compartments to complete the circuit and to keep each solution electrically neutral. This arrangement is a galvanic (voltaic) cell, which converts chemical energy directly into electrical energy.
Which electrical units are used?
| Quantity | Unit | Definition or value |
|---|---|---|
| Charge (Q) | Coulomb (C) | Charge of 6.242 × 1018 electrons |
| Current (I) | Ampere (A) | 1 A = 1 C/s |
| Potential difference (E) | Volt (V) | 1 V = 1 J/C |
| Charge of one electron | C | 1.602 × 10−19 C |
| Faraday constant (F) | C/mol | Charge of 1 mol of electrons = 96,485 C/mol |
| Electrical work | Joule (J) | w = Q × E |
Worked examples
Example 1: identifying the oxidized and the reduced species
In Fe2O3 + 3CO → 2Fe + 3CO2, the oxidation number of iron falls from +3 to 0 (reduction), and the oxidation number of carbon rises from +2 in CO to +4 in CO2 (oxidation). Fe2O3 is the oxidizing agent, and CO is the reducing agent. This is the main reaction in a blast furnace.
Example 2: charge from current and time
A current of 0.50 A flows for 2.0 hours. Q = I × t = 0.50 A × 7200 s = 3600 C. In moles of electrons: n = Q / F = 3600 / 96,485 = 0.0373 mol.
Example 3: number of electrons transferred
In 2Mg + O2 → 2MgO, each Mg atom loses 2 electrons, and each O atom gains 2. For 2 mol of Mg, 4 mol of electrons are transferred, which corresponds to 4 × 96,485 = 3.86 × 105 C.
Redox reactions in everyday life
| Process | Species oxidized | Species reduced |
|---|---|---|
| Combustion of methane | C in CH4 (−4 → +4) | O in O2 (0 → −2) |
| Rusting of iron | Fe (0 → +2, then +3) | O2 |
| Cellular respiration | Glucose | O2 |
| Photosynthesis | O in H2O (−2 → 0) | C in CO2 (+4 → 0 in glucose, on average) |
| Bleaching with hypochlorite | Colored compound | Cl in ClO− (+1 → −1) |
| Alkaline battery | Zn | MnO2 |
How much energy can be obtained?
The maximum electrical work from a cell is related to the Gibbs free energy: ΔG = −nFE, where n is the number of moles of electrons transferred, F is the Faraday constant and E is the cell potential. For the Daniell cell, with n = 2 and E° = 1.10 V: ΔG° = −2 × 96,485 × 1.10 = −212,300 J = −212 kJ per mole of zinc. This is the same free energy as for the direct reaction in a beaker; only the way in which the energy is released is different. A negative ΔG (and a positive E) means a spontaneous reaction.
How did electrochemistry begin?
Luigi Galvani observed in the 1780s that a frog’s leg twitched when touched with two different metals, and he attributed it to “animal electricity”. Alessandro Volta showed that the electricity came from the contact of the metals, and in 1800 he built the voltaic pile, the first battery, from alternating discs of zinc and silver with moistened cloth between them. Within weeks William Nicholson and Anthony Carlisle used it to decompose water, and in 1807 Humphry Davy isolated potassium and sodium by electrolysis. Michael Faraday then formulated the quantitative laws in 1833–1834.
Which mistakes are common?
- Mixing up oxidation and reduction: oxidation is loss of electrons, even if oxygen is not involved.
- Mixing up the agent and the process: the species that is oxidized is the reducing agent.
- Forgetting that electrons are not free in solution: they flow through the wire, and ions carry the charge in the solution.
- Confusing charge and current: current is charge per second.
- Forgetting the unit conversions: hours must be changed to seconds.