Sodium and fluorine bonding ionically to form sodium fluoride. Sodium loses its outer electron to give it a stable electron configuration, and this electron enters the fluorine atom exothermically. The oppositely charged ions are then attracted to each other. The sodium is oxidized, and the fluorine is reduced.

Redox (portmanteau of reduction and oxidation) reactions include all chemical reactions in which atoms have their oxidation state changed; in general, redox reactions involve the transfer of electrons between species.

The term "redox" comes from two concepts involved with electron transfer: reduction and oxidation.[18] It can be explained in simple terms:

  • Oxidation is the loss of electrons or an increase in oxidation state by a molecule, atom, or ion.
  • Reduction is the gain of electrons or a decrease in oxidation state by a molecule, atom, or ion.

Although oxidation reactions are commonly associated with the formation of oxides from oxygen molecules, these are only specific examples of a more general concept of reactions involving electron transfer.

Redox reactions, or oxidation-reduction reactions, have a number of similarities to acid–base reactions. Like acid–base reactions, redox reactions are a matched set, that is, there cannot be an oxidation reaction without a reduction reaction happening simultaneously. 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. When writing half-reactions, the gained or lost electrons are typically included explicitly in order that the half-reaction be balanced with respect to electric charge.

Though sufficient for many purposes, these descriptions are not precisely correct. Oxidation and reduction properly refer to a change in oxidation state — the actual transfer of electrons may never occur. Thus, oxidation is better defined as an increase in oxidation state, and reduction as a decrease in oxidation state. In practice, the transfer of electrons will always cause a change in oxidation state, but there are many reactions that are classed as "redox" even though no electron transfer occurs (such as those involving covalent bonds).

There are simple redox processes, such as the oxidation of carbon to yield carbon dioxide (CO
) or the reduction of carbon by hydrogen to yield methane (CH4), and more complex processes such as the oxidation of glucose (C6H12O6) in the human body through a series of complex electron transfer processes.


"Redox" is a portmanteau of "reduction" and "oxidation".

The word oxidation originally implied reaction with oxygen to form an oxide, since dioxygen (O2 (g)) was historically the first recognized oxidizing agent. 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 loss of electrons.

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 (1743-1794) showed 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 gain of electrons. Even though "reduction" seems counter-intuitive when speaking of the gain of electrons, it might help to think of reduction as the loss of oxygen, which was its historical meaning.

The electrochemist John Bockris has used the words electronation and deelectronation to describe reduction and oxidation processes respectively when they occur at electrodes.[19] These words are analogous to protonation and deprotonation, but they have not been widely adopted by chemists.

The term "hydrogenation" could be used instead of reduction, since hydrogen is the reducing agent in a large number of reactions, especially in organic chemistry and biochemistry. But, unlike oxidation, which has been generalized beyond its root element, hydrogenation has maintained its specific connection to reactions that add hydrogen to another substance (e.g., the hydrogenation of unsaturated fats in saturated fats, R-CH=CH-R + H2 → R-CH2-CH2-R).

Oxidizing and reducing agents

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 are involved in a particular reaction is called a redox pair. A redox couple is a reducing species and its corresponding oxidized form, e.g., Fe2+/Fe3+.


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. That is, 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, hence the name. Oxygen is the quintessential oxidizer.

Oxidants are usually chemical substances with elements in high oxidation states (e.g., H
, MnO
, CrO
, Cr
, OsO
), or else highly electronegative elements (O2, F2, Cl2, Br2) that can gain extra electrons by oxidizing another substance.


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 (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 donor. Electron donors can also form charge transfer complexes with electron acceptors.

Reductants in chemistry are very diverse.

  • Chemical Equation Balancer - An open source chemical equation balancer that handles redox reactions.
  • Video - Synthesis of Copper(II) Acetate 20 Feb. 2009
  • Redox reactions calculator
  • Redox reactions at Chemguide
  • Online redox reaction equation balancer, balances equations of any half-cell and full reactions

External links

  • Schüring, J., Schulz, H. D., Fischer, W. R., Böttcher, J., Duijnisveld, W. H. (editors)(1999). Redox: Fundamentals, Processes and Applications, Springer-Verlag, Heidelberg, 246 pp. ISBN 978-3-540-66528-1 (pdf 3,6 MB)
  • . January 2011. 


  1. ^  
  2. ^ a b c  
  3. ^ Gregory Girolami "Osmium weighs in" Nature Chemistry 2012, vol. 4, 954. doi:10.1038/nchem.1479
  4. ^ Cotton and Wilkinson, Advanced Inorganic Chemistry, p.1002
  5. ^ a b "Osmium tetroxide (as Os)". Documentation for Immediately Dangerous to Life or Health Concentrations (IDLHs). Centers for Disease Control. 
  6. ^ a b  
  7. ^ . p. 2−. 
  8. ^  
  9. ^ a b c  
  10. ^  
  11. ^  
  12. ^ a b c John O'M. Bockris and Amulya K. N. Reddy (1970). Modern Electrochemistry. Plenum Press. pp. 352–3. 
  13. ^ . January 2011. 
  14. ^  
  15. ^  
  16. ^ "Redox Reactions". 
  17. ^ Bozzola, J. J.; Russell, L. D. (1999). "Specimen Preparation for Transmission Electron Microscopy". Electron Microscopy : Principles and Techniques for Biologists. Sudbury, MA: Jones and Bartlett. pp. 21–31.  
  18. ^ Bozzola, J. J.; Russell, L. D. (1999). "Specimen Preparation for Transmission Electron Microscopy". Electron Microscopy : Principles and Techniques for Biologists. Sudbury, MA: Jones and Bartlett. pp. 21–31.  
  19. ^ Di Scipio, F.; Raimondo, S.; Tos, P.; Geuna, S. (2008). "A simple protocol for paraffin-embedded myelin sheath staining with osmium(VIII) oxide for light microscope observation". Microscopy Research and Technique 71 (7): 497–502.  
  20. ^ Di Scipio, F.; Raimondo, S.; Tos, P.; Geuna, S. (2008). "A simple protocol for paraffin-embedded myelin sheath staining with osmium(VIII) oxide for light microscope observation". Microscopy Research and Technique 71 (7): 497–502.  
  21. ^ Kiernan, J. A. """Re: "Disposal" of Osmium Tetroxide "Waste. Department of Anatomy & Cell Biology, The University of Western Ontario. 
  22. ^ Hawkins, J. M.; Meyer, A.; Lewis, T. A.; Loren, S.; Hollander, F. J. (1991). "Crystal Structure of Osmylated C60: Confirmation of the Soccer Ball Framework".  
  23. ^ Hawkins, J. M.; Meyer, A.; Lewis, T. A.; Loren, S.; Hollander, F. J. (1991). "Crystal Structure of Osmylated C60: Confirmation of the Soccer Ball Framework".  
  24. ^ "Chemical 'bomb plot' in UK foiled".  
  25. ^ R.H.Petrucci, W.S.Harwood and F.G.Herring (2002). General Chemistry (8th ed.). Prentice-Hall. p. 832. 
  26. ^ Bhattacharya, S. (2004-04-07). "Experts divided over poison bomb claim".  


See also

  • "RED CAT gains what AN OX loses" - Reduction occurs at the Cathode gains (electrons) what Anode Oxidation loses (electrons).
  • "RED CAT" and "AN OX", or "AnOx RedCat" ("an ox-red cat"), — Reduction occurs at the Cathode and the Anode is for Oxidation.
  • "LEORA says GEROA" — Loss of Electrons is Oxidation (Reducing Agent), Gain of Electrons is Reduction (Oxidizing Agent).[5]
  • "LEO the lion says GER" — Loss of Electrons is Oxidation, Gain of Electrons is Reduction.[3][4][5][6]
  • "OIL RIG"—Oxidation Is Loss of electrons, Reduction Is Gain of electrons.[3][4][5][6]

The key terms involved in redox are often confusing to students.[3][4] For example, an element that is oxidized loses electrons; however, that element is referred to as the reducing agent. Likewise, an element that is reduced gains electrons and is referred to as the oxidizing agent.[5] Acronyms or mnemonics are commonly used[6] to help remember what is happening:

Memory aids

2 + 3 + → 2 + 3 + 2 KOH

Adding these two half-cell reactions together gives the balanced equation:

6 e + 4 + 2 → 2 + 8 OH
6 OH + 3 2− → 3 2− + 3 + 6 e

Balancing the number of electrons in the two half-cell reactions gives:

Unbalanced reaction: + + → + + KOH
Reduction: 3 e + 2 + → + 4 OH
Oxidation: 2 OH + 2−2− + + 2 e

For example, in the reaction between potassium permanganate and sodium sulfite:

In basic media, OH ions and water are added to half reactions to balance the overall reaction.

Basic media

14 (aq) + 2 (aq) + 5 (s) → 7 (l) + 2 (aq) + 5 (aq) + 5 (aq)

Adding these two reactions eliminates the electrons terms and yields the balanced reaction:

8 (l) + 2 (aq) → 2 (aq) + 16 (aq) + 10 e
10 e + 30 + 5 (s) → 5 (aq) + 15 (l)

The reaction is balanced by scaling the two half-cell reactions to involve the same number of electrons (multiplying the oxidation reaction by the number of electrons in the reduction step and vice versa):

Unbalanced reaction: (aq) + (s) → (aq) + (aq)
Oxidation: 4 (l) + (aq) → (aq) + 8 (aq) + 5 e
Reduction: 2 e + 6 + (s) → (aq) + 3 (l)

For instance, when manganese(II) reacts with sodium bismuthate:

In acidic media, ions and water are added to half-reactions to balance the overall reaction.

Acidic media

Describing the overall electrochemical reaction for a redox process requires a balancing of the component half-reactions for oxidation and reduction. In general, for reactions in aqueous solution, this involves adding H+, OH, H2O, and electrons to compensate for the oxidation changes.

Balancing redox reactions

In geology, redox is important to both the formation of minerals and the mobilization of minerals, and is also important in some depositional environments. In general, the redox state of most rocks can be seen in the color of the rock. The rock forms in oxidizing conditions, giving it a red color. It is then "bleached" to a green—or sometimes white—form when a reducing fluid passes through the rock. The reduced fluid can also carry uranium-bearing minerals. Famous examples of redox conditions affecting geological processes include uranium deposits and Moqui marbles.

Mi Vida uranium mine, near Moab, Utah. The alternating red and white/green bands of sandstone correspond to oxidized and reduced conditions in groundwater redox chemistry.

Redox reactions in geology

Examples of redox cycling-inducing molecules are the herbicide paraquat and other viologens and quinones such as menadione.[2]

A wide variety of aromatic compounds are enzymatically 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 coenzymes. Once formed, these anion free radicals reduce molecular oxygen to superoxide, 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 futile cycle or redox cycling.

Redox cycling

The term redox state is often used to describe the balance of lactate and pyruvate, beta-hydroxybutyrate, and acetoacetate), whose interconversion is dependent on these ratios. An abnormal redox state can develop in a variety of deleterious situations, such as hypoxia, shock, and sepsis. Redox mechanism also control some cellular processes. Redox proteins and their genes must be co-located for redox regulation according to the CoRR hypothesis for the function of DNA in mitochondria and chloroplasts.

antioxidant. Unsatisfied free radicals can spur the mutation of cells they encounter and are, thus, causes of cancer.

Biological energy is frequently stored and released by means of redox reactions. Photosynthesis involves the reduction of carbon dioxide into sugars 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+), which then contributes to the creation of a proton gradient, which drives the synthesis of adenosine triphosphate (ATP) and is maintained by the reduction of oxygen. In animal cells, mitochondria perform similar functions. See the Membrane potential article.

6 CO2 + 6 H2O + light energy → C6H12O6 + 6 O2

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+). Photosynthesis and cellular respiration are complementary, but photosynthesis is not the reverse of the redox reaction in cell respiration:

C6H12O6 + 6 O2 → 6 CO2 + 6 H2O

Cellular respiration, for instance, is the oxidation of glucose (C6H12O6) to CO2 and the reduction of oxygen to water. The summary equation for cell respiration is:

Many important biological processes involve redox reactions.

Redox reactions in biology

The production of compact discs depends on a redox reaction, which coats the disc with a thin layer of metal film.

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, and gold-plated jewelry.

Redox reactions are the foundation of electrochemical cells.

Oxidation is used in a wide variety of industries such as in the production of cleaning products and oxidizing ammonia to produce nitric acid, which is used in most fertilizers.

The primary process of reducing ore at high temperature to produce metals is known as smelting.

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 galvanized steel, in which a sacrificial coating of zinc on steel parts protects them from rust.

Redox reactions in industry

2 Fe2+ + H2O2 + 2 H+ → 2 Fe3+ + 2 H2O
Overall equation:
Fe2+ → Fe3+ + e
H2O2 + 2 e → 2 OH
  • The oxidation of iron(II) to iron(III) by hydrogen peroxide in the presence of an acid:
4Fe + 3O2 → 2Fe2O3
  • The term corrosion refers to the electrochemical oxidation of metals in reaction with an oxidant such as oxygen. Rusting, the formation of iron oxides, is a well-known example of electrochemical corrosion; it forms as a result of the oxidation of iron metal. Common rust often refers to iron(III) oxide, formed in the following chemical reaction:
Iron rusting in pyrite cubes
Oxides, such as iron(III) oxide or rust, which consists of hydrated iron(III) oxides Fe2O3·nH2O and iron(III) oxide-hydroxide (FeO(OH), Fe(OH)3), form when oxygen combines with other elements

Corrosion and rusting

2 NO3 + 10 e + 12 H+ → N2 + 6 H2O

Other examples

Cu2+ + 2 e → Cu

And the copper is reduced:

Zn → Zn2+ + 2 e

As two half-reactions, it is seen that the zinc is oxidized:

Zn + Cu2+ → Zn2+ + Cu

The ionic equation for this reaction is:

In the above reaction, zinc metal displaces the copper(II) ion from copper sulfate solution and thus liberates free copper metal.

Zn(s)+ CuSO4(aq) → ZnSO4(aq) + Cu(s)

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, copper is deposited when zinc metal is placed in a copper(II) sulfate solution:

A redox reaction is the force behind an electrochemical cell like the Galvanic cell pictured. The battery is made out of a zinc electrode in a ZnSO4 solution connected with a wire and a porous disk to a copper electrode in a CuSO4 solution.

Metal displacement

+ → 2 HF

The overall reaction is:

2 H+ + 2 F → 2 HF

And the ions combine to form hydrogen fluoride:

2 H+ + 2 e
+ 2 e 2 F

+ 2 H+ + 2 F

When adding the reactions together the electrons are canceled:

Elements, even in molecular form, always have an oxidation state of zero. In the first half-reaction, hydrogen is oxidized from an oxidation state of zero to an oxidation state of +1. In the second half-reaction, fluorine is reduced from an oxidation state of zero to an oxidation state of −1.

Analyzing each half-reaction in isolation can often make the overall chemical process clearer. Because there is no net change in charge during a redox reaction, the number of electrons in excess in the oxidation reaction must equal the number consumed by the reduction reaction (as shown above).

+ 2 e → 2 F

and the reduction reaction:

→ 2 H+ + 2 e

the oxidation reaction:

We can write this overall reaction as two half-reactions:

+ F
→ 2 HF

A good example is the reaction between hydrogen and fluorine in which hydrogen is being oxidized and fluorine is being reduced:

Illustration of a redox reaction

Examples of redox reactions


  • Physical properties 1
  • Structure and electron configuration 2
  • Synthesis 3
  • Reactions 4
    • Oxidation of alkenes 4.1
    • Coordination chemistry 4.2
      • Oxofluorides 4.2.1
  • Uses 5
    • Organic synthesis 5.1
    • Biological staining 5.2
  • Etymology 6
  • Oxidizing and reducing agents 7
    • Oxidizers 7.1
    • Reducers 7.2
    • Polymer staining 7.3
    • Osmeth 7.4
    • Osmium ore refining 7.5
    • Buckminsterfullerene adduct 7.6
  • Safety considerations 8
  • Standard electrode potentials (reduction potentials) 9
  • References 10
  • External links 11

However, the potential of the reaction at the anode was sometimes expressed as an oxidation potential, E0ox = – E0. 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 E0cell = E0cathode + E0ox (anode)

For a redox reaction that takes place in a cell, the potential difference E0cell = E0cathode – E0anode

The electrode potential of each half-reaction is also known as its reduction potential E0red, 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 → ½ 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+).[1]

Each half-reaction has a standard electrode potential (E0cell), which is equal to the potential difference (or voltage) (E0cell) 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: ½ H2 → H+ + e-.

Standard electrode potentials (reduction potentials)

are used primarily in the reduction of carbon-carbon double or triple bonds.catalytic reductions. These catalyst nickel, or platinum, palladium) with a 2. Another method of reduction involves the use of hydrogen gas (Halcohols compounds to carbonyl primarily in the reduction of [24][22]