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Redox ( reduction–oxidation, , ) is a type of chemical reaction in which the of change.

  • Oxidation is the loss of electrons or an increase in the oxidation state of a chemical or atoms within it.
  • Reduction is the gain of electrons or a decrease in the oxidation state of a chemical or atoms within it.

There are two classes of redox reactions:

  • electron-transfer where only one (usually) electron flows from the reducing agent to the oxidant. This type of redox reaction is often discussed in term of redox couples and electrode potentials.
  • atom transfer, where an atom transfers from one substrate to another. For example, in the of iron, the oxidation state of iron atoms increases as it converts to an oxide and simultaneously the oxidation state of oxygen decreases as it accepts electrons released by iron. Although oxidation reactions are commonly associated with the formation of oxides, other chemical species can serve the same function. In , (and other) bonds are reduced by transfer of hydrogen atoms.


Terminology
"Redox" is a of the words "reduction" and "oxidation". The term "redox" was first used in 1928. 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 because two half-reactions always occur together to form a whole reaction.


Oxidants
Oxidation originally implied reaction with oxygen to form an oxide. Later, the term was expanded to encompass oxygen-like substances that accomplished parallel chemical reactions. Ultimately, the meaning was generalized to include all processes involving the loss of electrons. 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 , oxidants, or oxidizers. The oxidant (oxidizing agent) removes electrons from another substance, and is thus itself reduced. And, because it "accepts" electrons, the oxidizing agent is also called an electron acceptor. Oxidants are usually chemical substances with elements in high oxidation states (e.g., , , , , ), 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. is an oxidizer.

[[Oxygen]] is the quintessential oxidizer.
     


Reducers
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 , reductants, or reducers. The reductant (reducing agent) transfers electrons to another substance and is thus itself oxidized. And, because it donates electrons, the reducing agent is also called an . 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 such as a 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 which transfer the equivalent of one in reactions. The term is common on biochemistry.
(2022). 9788121924535, S. Chand.
A reducing equivalent can be an electron, a hydrogen atom, as a .
(2017). 9781464126116

Reductants in chemistry are very diverse. elemental , such as , , , , , and , are good reducing agents. These metals donate or give away electrons relatively readily. They transfer electrons.

Hydride transfer reagents, such as NaBH4 and LiAlH4, reduce by atom transfer: they transfer the equivalent of hydride or H-. These reagents widely used in [the reduction of compounds to .

(1996). 9780841233447, American Chemical Society.
(1990). 9780841217805, American Chemical Society. .
A related method of reduction involves the use of hydrogen gas (H2) as sources of H atoms.


Rates, mechanisms, and energies
Redox reactions can occur slowly, as in the formation of , or rapidly, as in the case of burning fuel. Electron transfer reactions are generally fast, occuring 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, i.e. involve 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 allow calculation of the thermodynamic aspects of redox reactions. For example, the fact that the combustion of methane yields more heat than the combustion of methanol.


Standard electrode potentials (reduction potentials)
Each half-reaction has a standard electrode potential ( E), which is equal to the potential difference or at equilibrium under of an electrochemical cell in which the reaction is the considered, and the is a standard hydrogen electrode where hydrogen is oxidized:
 H2 → H+ + e.

The electrode potential of each half-reaction is also known as its reduction potential E, 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+ (e.g., +2.866 V for F2) and negative for oxidizing agents that are weaker than H+ (e.g., −0.763 V for Zn2+).Electrode potential values from:

(2022). 9780130143297, Prentice Hall. .

For a redox reaction that takes place in a cell, the potential difference is:

E = EE

However, the potential of the reaction at the anode is sometimes expressed as an oxidation potential:

E = – E.
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
E = E + E


Examples of redox reactions
In the reaction between and , hydrogen is being oxidized and fluorine is being reduced:

+ → 2 HF

This reaction is spontaneous and releases 542 kJ per 2 g of hydrogen because the H-F bond is much stronger than the F-F bond. This reaction can be analyzed as two . The oxidation reaction converts hydrogen to protons:

→ 2  + 2 

The reduction reaction converts fluorine to the fluoride anion:

+ 2 e → 2 

The half reactions are summed such that the electrons cancel:

>
|align=right2 H+ + 2 e
+ 2 e2 F

H2 + F22 H+ + 2 F

The protons and fluoride combine to form hydrogen fluoride in a non-redox reaction:

2 H+ + 2 F → 2 HF
The overall reaction is:
+ → 2 HF


Metal displacement
In this type of reaction, a metal atom in a compound (or in a solution) is replaced by an atom of another metal. For example, is deposited when metal is placed in a copper(II) sulfate solution:

Zn(s)+ CuSO4(aq) → ZnSO4(aq) + Cu(s)

In the above reaction, zinc metal displaces the copper(II) ion from copper sulfate solution and thus liberates free copper metal. The reaction is spontaneous and releases 213 kJ per 65 g of zinc because relative to zinc, copper metal is lower in energy due to bonding via its partially filled d-orbitals.

The ionic equation for this reaction is:

Zn + Cu2+ → Zn2+ + Cu

As two , it is seen that the zinc is oxidized:

Zn → Zn2+ + 2 e

And the copper is reduced:

Cu2+ + 2 e → Cu


Other examples


Corrosion and rusting
  • The term refers to the electrochemical oxidation of metals in reaction with an oxidant such as oxygen. , the formation of , is a well-known example of electrochemical corrosion; it forms as a result of the oxidation of metal. Common rust often refers to iron(III) oxide, formed in the following chemical reaction:4Fe + 3O2 -> 2Fe2O3
  • The oxidation of iron(II) to iron(III) by hydrogen peroxide in the presence of an acid:Fe^{2+} -> {Fe^{3+}} + e-H2O2 + 2e- -> 2OH- Overall equation:{2Fe^{2+}} + {H2O2} + 2H+ -> {2Fe^{3+}} + 2H2O


Disproportionation
A disproportionation reaction is one in which a single substance is both oxidized and reduced. For example, ion with sulfur in oxidation state +2 can react in the presence of acid to form elemental sulfur (oxidation state 0) and (oxidation state +4).
S2O32- + 2 H+ → S + SO2 + H2O
Thus one sulfur atom is reduced from +2 to 0, while the other is oxidized from +2 to +4.
(2022). 9780130143297, Prentice Hall.


Redox reactions in industry
Cathodic protection 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. A common application of cathodic protection is in steel, in which a sacrificial coating of zinc on steel parts protects them from rust.

Oxidation is used in a wide variety of industries such as in the production of and oxidizing to produce .

Redox reactions are the foundation of electrochemical cells, which can generate electrical energy or support . Metal often contain metals in oxidized states such as oxides or sulfides, from which the pure metals are extracted by at high temperature in the presence of a reducing agent. The process of uses redox reactions to coat objects with a thin layer of a material, as in parts, silver plating , and .


Redox reactions in biology
Top: ( of )
Bottom: dehydroascorbic acid ( of )
Many important processes involve redox reactions. Before some of these processes can begin iron must be assimilated from the environment.
(2022). 9783110588903, De Gruyter.

Cellular respiration, for instance, is the oxidation of (C6H12O6) to and the reduction of to . The summary equation for cell respiration is:

C6H12O6 + 6 O2 → 6 CO2 + 6 H2O

The process of cell respiration also depends heavily on the reduction of NAD+ to NADH and the reverse reaction (the oxidation of NADH to NAD+). and cellular respiration are complementary, but photosynthesis is not the reverse of the redox reaction in cell respiration:

6 CO2 + 6 H2O + → C6H12O6 + 6 O2

Biological energy is frequently stored and released by means of redox reactions. Photosynthesis involves the reduction of into and the oxidation of water into molecular oxygen. The reverse reaction, respiration, oxidizes sugars to produce carbon dioxide and water. As intermediate steps, the reduced carbon compounds are used to reduce nicotinamide adenine dinucleotide (NAD+) to NADH, which then contributes to the creation of a , which drives the synthesis of adenosine triphosphate (ATP) and is maintained by the reduction of oxygen. In animal cells, perform similar functions. See the Membrane potential article.

reactions are redox reactions that occur as a part of and killing microorganisms, where an electron detaches from a molecule and then reattaches almost instantaneously. Free radicals are a part of redox molecules and can become harmful to the human body if they do not reattach to the redox molecule or an . Unsatisfied free radicals can spur the mutation of cells they encounter and are, thus, causes of cancer.

The term redox state is often used to describe the balance of , NAD+/NADH and NADP+/NADPH in a biological system such as a cell or organ. The redox state is reflected in the balance of several sets of metabolites (e.g., and , beta-hydroxybutyrate, and ), whose interconversion is dependent on these ratios. An abnormal redox state can develop in a variety of deleterious situations, such as hypoxia, shock, and . Redox mechanism also control some cellular processes. Redox proteins and their genes must be co-located for redox regulation according to the for the function of DNA in mitochondria and chloroplasts.


Redox cycling
Wide varieties of are reduced to form free radicals that contain one more electron than their parent compounds. In general, the electron donor is any of a wide variety of flavoenzymes and their . Once formed, these anion free radicals reduce molecular oxygen to and regenerate the unchanged parent compound. The net reaction is the oxidation of the flavoenzyme's coenzymes and the reduction of molecular oxygen to form superoxide. This catalytic behavior has been described as a or redox cycling.


Redox reactions in geology
Minerals are generally oxidized derivatives of metals. Iron is mined as its (Fe3O4). Titanium is mined as its dioxide, usually in the form of (TiO2). To obtain the corresponding metals, these oxides must be reduced, which is often achieved by heating these oxides with carbon or carbon monoxide as reducing agents. are the reactors where iron oxides and coke (a form of carbon) are combined to produce molten iron.The main chemical reaction producing the molten iron is:
Fe2O3 + 3CO → 2Fe + 3CO2


Redox reactions in soils
Electron transfer reactions are central to myriad processes and properties in soils, and electron "activity", quantified as Eh (platinum electrode potential (voltage) relative to the standard hydrogen electrode) or pe (analogous to pH as -log electron activity), is a master variable, along with pH, that controls and is governed by chemical reactions and biological processes. Early theoretical research with applications to flooded soils and paddy rice production was seminal for subsequent work on thermodynamic aspects of redox and plant root growth in soils.
(1992). 9780120007240
Later work built on this foundation, and expanded it for understanding redox reactions related to heavy metal oxidation state changes, pedogenesis and morphology, organic compound degradation and formation, free radical chemistry, wetland delineation, soil remediation, and various methodological approaches for characterizing the redox status of soils.
(2022). 9781439803059, CRC Press.


Mnemonics
The key terms involved in redox can be confusing.
(2022). 9781936137749, National Science Teachers Association. .
(2022). 9780028282107, Glencoe McGraw-Hill.
For example, a reagent that is oxidized loses electrons; however, that reagent is referred to as the reducing agent. Likewise, a reagent that is reduced gains electrons and is referred to as the oxidizing agent.
(2022). 9780840068460, Brooks/Cole, Cengage Learning. .
These are commonly used by students to help memorise the terminology:
(2022). 9780547054056, Houghton Mifflin. .

  • "" — oxidation is loss of electrons, reduction is gain of electrons
  • "LEO the lion says GER grr" — loss of electrons is oxidation, gain of electrons is reduction
  • "LEORA says GEROA" — the loss of electrons is called oxidation (reducing agent); the gain of electrons is called reduction (oxidizing agent).
  • "RED CAT" and "AN OX", or "AnOx RedCat" ("an ox-red cat") — reduction occurs at the cathode and the anode is for oxidation
  • "RED CAT gains what AN OX loses" – reduction at the cathode gains (electrons) what anode oxidation loses (electrons)
  • "PANIC" – Positive Anode and Negative is Cathode. This applies to electrolytic cells which release stored electricity, and can be recharged with electricity. PANIC does not apply to cells that can be recharged with redox materials. These , such as , produce electricity from internal redox reactions. Here, the positive electrode is the cathode and the negative is the anode.


See also


Further reading


External links

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