Oxidation–Reduction Reaction

An oxidation–reduction reaction, commonly called a redox reaction, is a chemical reaction in which the oxidation states of participating atoms change. Oxidation increases oxidation state; reduction decreases it. In reactions involving direct electron transfer, oxidation corresponds to electron loss and reduction to electron gain. The processes are coupled: electrons released by one component are accepted by another. Redox chemistry also includes reactions involving covalent substances, for which oxidation-state changes provide formal electron accounting rather than necessarily describing transfer of free electrons. (openstax.org)

Oxidation states and electron accounting

An oxidation state, or oxidation number, is a formal quantity assigned to an atom by treating bonds between different elements as ionic. Bonding electrons are assigned to the more electronegative partner, while bonds between identical atoms are divided equally. This convention, based on electronegativity, allows oxidation and reduction to be identified even when electrons remain shared in covalent bonds. Oxidation state is therefore not generally identical to an atom’s actual electrical charge. (goldbook.iupac.org)

Atoms in an uncombined element have oxidation state zero. A monatomic ion has an oxidation state equal to its charge. The sum of oxidation states is zero for a neutral compound and equals the overall charge for a polyatomic ion. Common introductory assignments include oxygen at −2 and hydrogen at +1, with exceptions such as oxygen at −1 in peroxides and hydrogen at −1 in metal hydrides. Comparing assignments before and after a reaction identifies the atoms oxidized and reduced. (openstax.org)

An oxidizing agent, or oxidant, causes another species to be oxidized and is itself reduced. A reducing agent, or reductant, causes reduction and is itself oxidized. These labels describe roles in a particular reaction rather than immutable properties of substances. (openstax.org)

Half-reactions and balancing

A redox equation can be separated into two half-reactions, one representing oxidation and the other reduction. For zinc reacting with aqueous copper(II) ions:

\mathrm{Zn(s)\rightarrow Zn^{2+}(aq)+2e^-}

\mathrm{Cu^{2+}(aq)+2e^-\rightarrow Cu(s)}

Adding them cancels the electrons:

\mathrm{Zn(s)+Cu^{2+}(aq)\rightarrow Zn^{2+}(aq)+Cu(s)}.

Zinc changes from 0 to +2 and is the reductant; copper changes from +2 to 0, so copper(II) ions are the oxidant. Both atom counts and electric charge must balance. The coefficients express the reaction’s stoichiometry, not necessarily its molecular mechanism. (openstax.org)

In aqueous acidic solution, the half-reaction method balances elements other than hydrogen and oxygen first. Oxygen is balanced with water, hydrogen with H⁺, and charge with electrons. The half-reactions are multiplied until their electron counts match, then added. In basic solution, an equation balanced this way can be converted by adding OH⁻ to both sides to neutralize H⁺ and canceling excess water. Solution pH can also affect which products form, not merely how an equation is written. (openstax.org)

Electrochemical reactions

Electrochemistry examines the relationship between redox reactions and electrical processes. When oxidation and reduction occur at separate electrodes, electrons can pass through an external circuit. Oxidation occurs at the anode, and reduction at the cathode, regardless of the electrodes’ electrical signs. Ionic transport through the electrolyte completes the circuit and maintains charge balance. (openstax.org)

A galvanic cell converts the driving force of a spontaneous redox reaction into electrical work. By contrast, electrolysis uses an external power source to drive a reaction that is otherwise nonspontaneous under the operating conditions. A rechargeable battery functions galvanically during discharge and electrolytically during charging; electrode roles reverse when the reaction reverses. (openstax.org)

Potentials, equilibrium, and rates

Reduction potentials quantify the thermodynamic tendency of specified redox couples to accept electrons relative to a reference electrode. Standard reduction potentials are conventionally measured against the standard hydrogen electrode, assigned zero volts. Using both electrode values as reduction potentials:

E^\circ_{\mathrm{cell}} =E^\circ_{\mathrm{cathode}}-E^\circ_{\mathrm{anode}}.

A more positive standard reduction potential indicates a stronger oxidant under the specified standard conditions. (openstax.org)

Cell potential is related to Gibbs free energy by

\Delta G=-nFE,

where $n$ is the number of moles of electrons transferred per mole of reaction as written, and $F$ is the Faraday constant. A positive cell potential corresponds to a negative reaction free-energy change. Under nonstandard conditions, the Nernst equation gives

E=E^\circ-\frac{RT}{nF}\ln Q,

where $Q$ is the reaction quotient, expressed using activities. At equilibrium, the net reaction has $E=0$ and $\Delta G=0$ . Thermodynamic favorability does not establish reaction speed: chemical kinetics and activation barriers determine how rapidly a favorable process proceeds. (openstax.org)

Biological and industrial roles

In cellular respiration, organic substrates are oxidized while electron acceptors are reduced. NAD⁺ accepts reducing equivalents to form NADH, which subsequently donates electrons to other reactions. During aerobic respiration, the electron transport chain ultimately transfers electrons to oxygen, producing water. Coupled proton transport establishes the gradient used in oxidative phosphorylation. In photosynthesis, light-driven reactions generate reducing power for carbon assimilation; oxygenic photosynthesis obtains electrons by oxidizing water. (openstax.org)

Industrial redox processes include electrolytic metal extraction, refining, and electroplating. Corrosion is an unwanted application of the same chemistry: metal oxidation is coupled to reduction reactions occurring elsewhere on a surface. Protective coatings and sacrificial anodes alter or interrupt these electrochemical pathways. (openstax.org)