Redox (/ˈrɛdɒks/ RED-oks, /ˈriːdɒks/ REE-doks, reduction–oxidation[2] or oxidation–reduction[3]: 150 ) is a type of chemical reaction in which the oxidation states of the reactants change.[4] Oxidation is the loss of electrons or an increase in the oxidation state, while reduction is the gain of electrons or a decrease in the oxidation state. The oxidation and reduction processes occur simultaneously in the chemical reaction.
Redox reactions fall into two classes. In electron transfer, usually a single electron flows from the atom, ion, or molecule being oxidized to the one being reduced, a case often described in terms of redox couples and electrode potentials. In atom transfer, an atom passes from one substrate to another; for example, in the rusting of iron the oxidation state of the iron atoms increases as the metal converts to an oxide, while oxygen is reduced as it accepts the released electrons. Although oxidation is commonly associated with forming oxides, other chemical species can serve the same function;[5] in hydrogenation, bonds such as C=C are reduced by the transfer of hydrogen atoms.
Redox reactions occur throughout nature and industry. Cellular respiration and photosynthesis, combustion, and the corrosion of metals all proceed through redox chemistry, as do the reactions that power batteries and other electrochemical cells. Industry uses redox reactions to extract metals from their ores by smelting, to electroplate objects, and to manufacture chemicals such as nitric acid; in soils, sediments, and water, redox gradients drive the biogeochemical cycling of elements.
Terminology
"Redox" is a portmanteau of "reduction" and "oxidation". The term was first used in a 1928 article by Leonor Michaelis and Louis B. Flexner.[6][7]
Oxidation is a process in which a substance loses electrons. Reduction is a process in which a substance gains electrons.
The processes of oxidation and reduction occur simultaneously and cannot occur independently.[5] In redox processes, the reductant transfers electrons to the oxidant. Thus, in the reaction, the reductant or reducer or reducing agent loses electrons and is oxidized, while the oxidant or oxidizer or oxidizing agent gains electrons and is reduced.
The pair of an oxidizing and reducing agent that is involved in a particular reaction is called a redox pair. A redox couple is a reducing species and its corresponding oxidizing form,[8] e.g., Fe2+
/ Fe3+
.The oxidation alone and the reduction alone are each called a half-reaction because two half-reactions always occur together to form a whole reaction.[5]
In electrochemical reactions the oxidation and reduction processes do occur simultaneously but are separated in space.
Oxidants
Oxidation originally implied a reaction with oxygen to form an oxide. Later, the term was expanded to encompass substances that accomplished chemical reactions similar to those of oxygen. Ultimately, the meaning was generalized to include all processes involving the loss of electrons or the increase in the oxidation state of a chemical species.[9]: A49 Substances that have the ability to oxidize other substances (cause them to lose electrons) are said to be oxidative or oxidizing, and are known as oxidizing agents, oxidants, or oxidizers. The oxidant removes electrons from another substance, and is thus itself reduced.[9]: A50 Because it "accepts" electrons, the oxidizing agent is also called an electron acceptor. Oxidants are usually chemical substances with elements in high oxidation states[3]: 159 (e.g., N
2O
4, MnO−
4, CrO
3, Cr
2O2−
7, OsO
4), or else highly electronegative elements (e.g. O2, F2, Cl2, Br2, I2) that can gain extra electrons by oxidizing another substance.[3]: 909
Oxidizers are oxidants, but the term is mainly reserved for sources of oxygen, particularly in the context of explosions. Nitric acid is a strong oxidizer.[10]
Reductants
Substances that have the ability to reduce other substances (cause them to gain electrons) are said to be reductive or reducing and are known as reducing agents, reductants, or reducers. The reductant transfers electrons to another substance and is thus itself oxidized.[3]: 159 Because it donates electrons, the reducing agent is also called an electron donor. Electron donors can also form charge transfer complexes with electron acceptors. The word reduction originally referred to the loss in weight upon heating a metallic ore such as a metal oxide to extract the metal. In other words, ore was "reduced" to metal.[11] Antoine Lavoisier demonstrated that this loss of weight was due to the loss of oxygen as a gas. Later, scientists realized that the metal atom gains electrons in this process. The meaning of reduction then became generalized to include all processes involving a gain of electrons.[11] Reducing equivalent refers to chemical species which transfer the equivalent of one electron in redox reactions. The term is common in biochemistry.[12] A reducing equivalent can be an electron or a hydrogen atom as a hydride ion.[13]
Reductants in chemistry are very diverse. Electropositive elemental metals, such as lithium, sodium, magnesium, iron, zinc, and aluminium, are good reducing agents. These metals donate electrons relatively readily.[14]
Hydride transfer reagents, such as NaBH4 and LiAlH4, reduce by atom transfer: they transfer the equivalent of hydride or H−. These reagents are widely used in the reduction of carbonyl compounds to alcohols.[15][16] A related method of reduction involves the use of hydrogen gas (H2) as sources of H atoms.[3]: 288
Electronation and de-electronation
The electrochemist John Bockris proposed the words electronation and de-electronation to describe reduction and oxidation processes, respectively, when they occur at electrodes.[17] These words are analogous to protonation and deprotonation.[18] IUPAC has recognized the terms electronation[19] and de-electronation.[20]
Rates, mechanisms, and energies
Redox reactions can occur slowly, as in the formation of rust, or rapidly, as in the case of burning fuel. Electron transfer reactions are generally fast, occurring within the time of mixing.[21]
The mechanisms of atom-transfer reactions are highly variable because many kinds of atoms can be transferred, and such reactions can involve several steps. Electron-transfer reactions, by contrast, proceed by two distinct pathways. In inner-sphere transfer, the two reactants share a bridging ligand through which the electron passes;[22] in outer-sphere transfer, the electron moves between reactants whose coordination shells remain intact.[23] Henry Taube received the 1983 Nobel Prize in Chemistry for distinguishing these pathways through experiments on metal complexes.[24]
The rate of an outer-sphere electron transfer is described by Marcus theory, developed by Rudolph A. Marcus.[25] The theory expresses the activation energy in terms of two quantities: the standard free-energy change of the reaction and the reorganization energy, the energy needed to distort the reactants and the surrounding solvent into the configuration of the products before the electron moves.[25] It predicts an "inverted region", in which the rate falls once the driving force exceeds the reorganization energy.[25] Marcus received the 1992 Nobel Prize in Chemistry for the theory.[26]
Analysis of bond energies and ionization energies in water allows calculation of the thermodynamic aspects of redox reactions.[27]
Standard electrode potentials (reduction potentials)
Each half-reaction has a standard electrode potential (Eo
cell), which is equal to the potential difference or voltage at equilibrium under standard conditions of an electrochemical cell in which the cathode reaction is the half-reaction considered, and the anode is a standard hydrogen electrode where hydrogen is oxidized:[28]
- 1⁄2 H2 → H+ + e−
The electrode potential of each half-reaction is also known as its reduction potential (Eo
red), or potential when the half-reaction takes place at a cathode. The reduction potential is a measure of the tendency of the oxidizing agent to be reduced. Its value is zero for H+ + e− → 1⁄2H2 by definition, positive for oxidizing agents stronger than H+ (e.g., +2.866 V for F2) and negative for oxidizing agents that are weaker than H+ (e.g., −0.763V for Zn2+).[9]: 873
For a redox reaction that takes place in a cell, the potential difference is:
- Eo
cell = Eo
cathode − Eo
anode
However, the potential of the reaction at the anode is sometimes expressed as an oxidation potential:
- Eo
ox = −Eo
red
The oxidation potential is a measure of the tendency of the reducing agent to be oxidized but does not represent the physical potential at an electrode. With this notation, the cell voltage equation is written with a plus sign
- Eo
cell = Eo
red(cathode) + Eo
ox(anode)
Examples of redox reactions
In the reaction between hydrogen and fluorine, hydrogen is being oxidized and fluorine is being reduced: