Oxidation and reduction
Oxidation–reduction reactions redistribute electrons. Oxidation is electron loss or an increase in oxidation state; reduction is electron gain or a decrease in oxidation state. The two processes are inseparable because transferred electrons must be conserved.
Oxidizing and reducing agents
The reducing agent donates electrons and is itself oxidized. The oxidizing agent accepts electrons and is itself reduced.
In Zn + Cu²⁺ → Zn²⁺ + Cu, zinc is the reducing agent and Cu²⁺ is the oxidizing agent. Naming the agent by what happens to the other species is the source of many errors.
Oxidation states reveal hidden redox
Not every redox reaction displays free electrons. Combustion, corrosion and many organic reactions are recognized by changes in oxidation state.
In CH₄ + 2O₂ → CO₂ + 2H₂O, carbon changes from −4 to +4 and is oxidized; oxygen changes from 0 to −2 and is reduced.
Balancing in acidic solution
For MnO₄⁻ + Fe²⁺ → Mn²⁺ + Fe³⁺, split the reaction. Fe²⁺ → Fe³⁺ + e⁻. For permanganate, balance O with H₂O, H with H⁺ and charge with electrons:
MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O
Multiply the Fe half-reaction by 5, add, and cancel electrons.
Basic solution
A reliable method is to balance as if acidic first, then add OH⁻ to both sides to neutralize every H⁺. Combine H⁺ + OH⁻ into H₂O and cancel waters that appear on both sides.
This keeps the logic consistent instead of inventing a separate list of memorized steps.
Redox beyond obvious electron transfer
In covalent chemistry, electrons are shared rather than fully transferred, yet oxidation-state bookkeeping still tracks the change. Hydrogenation of an alkene reduces carbon; oxidation of an alcohol increases the oxidation state of carbon even though no free electron appears in the reaction equation.
Electrochemistry makes the electron transfer spatially explicit by separating oxidation and reduction into different electrodes.
A complete acidic half-reaction example
Balance MnO₄⁻ + Fe²⁺ → Mn²⁺ + Fe³⁺ in acid. The iron half-reaction is Fe²⁺ → Fe³⁺ + e⁻. For manganese, first balance oxygen with four H₂O on the product side, then hydrogen with 8H⁺ on the reactant side, and finally charge with 5e⁻:
MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O
5Fe²⁺ → 5Fe³⁺ + 5e⁻
MnO₄⁻ + 8H⁺ + 5Fe²⁺ → Mn²⁺ + 4H₂O + 5Fe³⁺
Final checks are mechanical: count every element and verify total charge on both sides. The electrons must cancel completely in the overall chemical equation.
Disproportionation and comproportionation
One species can sometimes be both oxidized and reduced. In disproportionation, atoms in one oxidation state split into products with higher and lower states. Hydrogen peroxide can disproportionate:
2H₂O₂ → 2H₂O + O₂
Oxygen starts at −1, becomes −2 in water and 0 in O₂. The reverse type, where two oxidation states combine to an intermediate state, is comproportionation.
Redox in batteries, corrosion and metabolism
A galvanic cell separates oxidation from reduction so electrons travel through an external circuit. Corrosion couples metal oxidation to a reduction reaction, often involving dissolved oxygen. Biological respiration transfers electrons through a chain of redox carriers rather than in one violent step.
The shared principle is electron-energy management. Different systems control where oxidation and reduction occur and how the released free energy is captured.
Oxidation state does not always mean literal electron transfer
In strongly ionic reactions, oxidation-state changes can correspond closely to electron transfer between species. In covalent molecules, oxidation state is often a formal accounting change produced by shifting ownership of shared electron pairs.
Oxidizing methane to CO₂ is described as carbon going from −4 to +4, but eight electrons do not sit on carbon and then physically jump as localized particles to oxygen. The oxidation-state method still correctly tracks the redox balance.
A basic-medium balancing example
Consider ClO⁻ → Cl⁻ in basic solution. Balance the half-reaction first as if acidic:
ClO⁻ + 2H⁺ + 2e⁻ → Cl⁻ + H₂O
Add 2OH⁻ to both sides, combine 2H⁺ + 2OH⁻ into 2H₂O, then cancel one water:
ClO⁻ + H₂O + 2e⁻ → Cl⁻ + 2OH⁻
The method works because atom and charge conservation are enforced at every stage.
Electron transfer and free energy
A balanced redox equation tells how many electrons are transferred, but not whether the reaction is spontaneous or fast. Thermodynamics determines the free-energy driving force; kinetics determines the rate.
In electrochemistry, cell potential is linked to reaction free energy by ΔG = −nFE. A positive galvanic-cell voltage corresponds to a negative ΔG for the written spontaneous reaction under those conditions.
Thermodynamics and kinetics are separate. A redox reaction can be favorable yet slow if bond breaking, solvent reorganization or another activation barrier makes electron transfer difficult. Catalysts can change the rate without changing the balanced electron count.
Conversely, a rapid redox reaction is not automatically strongly exergonic. Rate and driving force answer different questions.
Exercises
Agents
In 2Fe³⁺ + Sn²⁺ → 2Fe²⁺ + Sn⁴⁺, identify oxidizing and reducing agents.
Solution
Fe³⁺ is reduced, so it is the oxidizing agent. Sn²⁺ is oxidized, so it is the reducing agent.
Electron count
How many electrons are required to reduce Mn from +7 to +2?
Solution
A decrease of five oxidation units requires 5 electrons per Mn.
Conservation
What two quantities must a balanced redox equation conserve?
Solution
Atoms (mass) and net electric charge.