Redox


Redox is a type of chemical reaction in which the oxidation states of the reactants change. 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.
There are two classes of redox reactions:
  • Electron-transfer – Only one electron flows from the atom, ion, or molecule being oxidized to the atom, ion, or molecule that is reduced. This type of redox reaction is often discussed in terms of redox couples and electrode potentials.
  • Atom transfer – An atom transfers from one substrate to another. For example, in the rusting of iron, the oxidation state of iron atoms increases as the iron converts to an oxide, and simultaneously, the oxidation state of oxygen decreases as it accepts electrons released by the iron. Although oxidation reactions are commonly associated with forming oxides, other chemical species can serve the same function. In hydrogenation, bonds like C=C are reduced by transfer of hydrogen atoms.

    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.
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. In redox processes, the reductant transfers electrons to the oxidant. Thus, in the reaction, the reductant or reducing agent loses electrons and is oxidized, and the oxidant 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, e.g., /.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.
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. Substances that have the ability to oxidize other substances 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. Because it "accepts" electrons, the oxidizing agent is also called an electron acceptor. Oxidants are usually chemical substances with elements in high oxidation states, or else highly electronegative elements that can gain extra electrons by oxidizing another substance.
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.
File:GHS-pictogram-rondflam.svg|thumb|upright|The international pictogram for oxidizing chemicals

Reductants

Substances that have the ability to reduce other substances 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. 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. 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. Reducing equivalent refers to chemical species which transfer the equivalent of one electron in redox reactions. The term is common in biochemistry. A reducing equivalent can be an electron or a hydrogen atom as a hydride ion.
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.
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. A related method of reduction involves the use of hydrogen gas as sources of H atoms.

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. These words are analogous to protonation and deprotonation. IUPAC has recognized the terms electronation and de-electronation.

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.
The mechanisms of atom-transfer reactions are highly variable because many kinds of atoms can be transferred. Such reactions can also be quite complex, involving many steps. The mechanisms of electron-transfer reactions occur by two distinct pathways, inner sphere electron transfer and outer sphere electron transfer.
Analysis of bond energies and ionization energies in water allows calculation of the thermodynamic aspects of redox reactions.

Standard electrode potentials (reduction potentials)

Each half-reaction has a standard electrode potential, 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:
The electrode potential of each half-reaction is also known as its reduction potential, 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− → H2 by definition, positive for oxidizing agents stronger than H+ and negative for oxidizing agents that are weaker than H+.
For a redox reaction that takes place in a cell, the potential difference is:
However, the potential of the reaction at the anode is sometimes expressed as an oxidation potential:
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

Examples of redox reactions

In the reaction between hydrogen and fluorine, hydrogen is being oxidized and fluorine is being reduced:
This spontaneous reaction releases a large amount of energy because two H-F bonds are much stronger than one H-H bond and one F-F bond. This reaction can be analyzed as two half-reactions. The oxidation reaction converts hydrogen to protons:
The reduction reaction converts fluorine to the fluoride anion:
The half-reactions are combined so that the electrons cancel:
The protons and fluoride combine to form hydrogen fluoride in a non-redox reaction:
The overall reaction is:

Metal displacement

In this type of reaction, a metal atom in a compound or solution is replaced by an atom of another metal. For example, copper is deposited when zinc metal is placed in a copper sulfate solution:
In the above reaction, zinc metal displaces the copper ion from the copper sulfate solution, thus liberating free copper metal. The reaction is spontaneous and releases 213 kJ per 65 g of zinc.
The ionic equation for this reaction is:
As two half-reactions, it is seen that the zinc is oxidized:
And the copper is reduced:

Other examples

A disproportionation reaction is one in which a single substance is both oxidized and reduced. For example, thiosulfate ion with sulfur in oxidation state +2 can react in the presence of acid to form elemental sulfur and sulfur dioxide.
Thus one sulfur atom is reduced from +2 to 0, while the other is oxidized from +2 to +4.

Redox reactions in industry

is a technique used to control the corrosion of a metal surface by making it the cathode of an electrochemical cell. A simple method of protection connects protected metal to a more easily corroded "sacrificial anode" to act as the anode. The sacrificial metal, instead of the protected metal, then corrodes.
Oxidation is used in a wide variety of industries, such as in the production of cleaning products and oxidizing ammonia to produce nitric acid.
Redox reactions are the foundation of electrochemical cells, which can generate electrical energy or support electrosynthesis. Metal ores often contain metals in oxidized states, such as oxides or sulfides, from which the pure metals are extracted by smelting at high temperatures in the presence of a reducing agent. The process of electroplating uses redox reactions to coat objects with a thin layer of a material, as in chrome-plated automotive parts, silver plating cutlery, galvanization and gold-plated jewelry.