Redox

For other uses, seeRedox (disambiguation).

Redox(/ˈrɛdɒks/RED-oks,/ˈriːdɒks/REE-doks,reduction–oxidationoroxidation–reduction) is a type ofchemical reactionin which theoxidation statesof thereactantschange.Oxidation is the loss ofelectronsor 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.

Sodium"gives" one outer electron tofluorine,bonding them to formsodium fluoride.The sodium atom is oxidized, and fluorine is reduced.

When a few drops ofglycerol(mild reducing agent) are added to powderedpotassium permanganate(strong oxidizing agent), a violent redox reaction accompanied by self-ignition starts.

Example of areduction–oxidationreaction between sodium and chlorine, with theOIL RIGmnemonic

There are two classes of redox reactions:

• Electron-transfer– Only one (usually) 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 onesubstrateto another. For example, in therustingofiron,the oxidation state of iron atoms increases as the iron converts to anoxide,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.Inhydrogenation,bonds likeC=Care reduced bytransfer of hydrogen atoms.

Terminology

"Redox" is aportmanteauof the words "REDuction" and "OXidation." The term "redox" was first used in 1928.

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 orreducing agentloses electrons and is oxidized, and the oxidant oroxidizing agentgains 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.,Fe²⁺
/Fe³⁺
.The oxidation alone and the reduction alone are each called ahalf-reactionbecause two half-reactions always occur together to form a whole reaction.

Inelectrochemical reactionsthe oxidation and reduction processes do occur simultaneously but are separated in space.

Oxidants

Main article:Oxidizing agent

Oxidation originally implied a reaction with oxygen to form an oxide. Later, the term was expanded to encompasssubstancesthat 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 (cause them to lose electrons) are said to be oxidative or oxidizing, and are known asoxidizing 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 anelectron acceptor.Oxidants are usually chemical substances with elements in high oxidation states(e.g.,N
₂O
₄,MnO⁻
₄,CrO
₃,Cr
₂O²⁻
₇,OsO
₄), or else highlyelectronegativeelements (e.g.O₂,F₂,Cl₂,Br₂,I₂) 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 acidis a strong oxidizer.

Theinternationalpictogramfor oxidizing chemicals

Reductants

Main article:Reducing agent

Substances that have the ability to reduce other substances (cause them to gain electrons) are said to be reductive or reducing and are known asreducing 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 anelectron donor.Electron donors can also formcharge transfer complexeswith electron acceptors. The word reduction originally referred to the loss in weight upon heating a metallicoresuch as ametal oxideto extract the metal. In other words, ore was "reduced" to metal.Antoine Lavoisierdemonstrated 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 tochemical specieswhich transfer the equivalent of oneelectronin redox reactions. The term is common inbiochemistry.A reducing equivalent can be an electron or a hydrogen atom as ahydride ion.

Reductants in chemistry are very diverse.Electropositiveelementalmetals,such aslithium,sodium,magnesium,iron,zinc,andaluminium,are good reducing agents. These metals donate electrons relatively readily.

Hydride transfer reagents,such asNaBH₄andLiAlH₄,reduce by atom transfer: they transfer the equivalent of hydride or H⁻.These reagents are widely used in the reduction ofcarbonylcompounds toalcohols.A related method of reduction involves the use of hydrogen gas (H₂) as sources of H atoms.

Electronation and deelectronation

TheelectrochemistJohn Bockrisproposed the words electronation and de-electronation to describe reduction and oxidation processes, respectively, when they occur atelectrodes.These words are analogous toprotonationanddeprotonation.They have not been widely adopted by chemists worldwide,althoughIUPAChas recognized the terms electronationand de-electronation.

Rates, mechanisms, and energies

Redox reactions can occur slowly, as in the formation ofrust,or rapidly, as in the case of burningfuel.Electron transfer reactions are generally fast, occurring within the time of mi xing.

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 transferandouter sphere electron transfer.

Analysis of bond energies andionization energiesin water allows calculation of the thermodynamic aspects of redox reactions.

Standard electrode potentials (reduction potentials)

Each half-reaction has a standardelectrode potential(E^o
_cell), which is equal to the potential difference orvoltageat equilibrium understandard conditionsof anelectrochemical cellin which thecathodereaction is thehalf-reactionconsidered, and theanodeis astandard hydrogen electrodewhere hydrogen is oxidized:

1⁄2H₂→ H⁺+ e⁻

The electrode potential of each half-reaction is also known as its reduction potential (E^o
_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⁻→1⁄2H₂by 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⁺).

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

E^o
_cell=E^o
_cathode–E^o
_anode

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

E^o
_ox= –E^o
_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

E^o
_cell=E^o
_red(cathode)+E^o
_ox(anode)

Examples of redox reactions

Illustration of a redox reaction

In the reaction betweenhydrogenandfluorine,hydrogen is being oxidized and fluorine is being reduced:

H₂+ F₂→ 2 HF

This spontaneous reaction 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 twohalf-reactions.The oxidation reaction converts hydrogen toprotons:

H₂→ 2H⁺+ 2e⁻

The reduction reaction convertsfluorineto the fluoride anion:

F₂+ 2 e⁻→ 2F⁻

The half-reactions are combined so that the electrons cancel:

H 2	→	2 H++ 2 e−
F 2+ 2 e−	→	2 F−
H2+ F2	→	2 H++ 2 F−

The protons and fluoride combine to formhydrogen fluoridein a non-redox reaction:

2 H⁺+ 2 F⁻→ 2 HF

The overall reaction is:

H₂+ F₂→ 2 HF

Metal displacement

A redox reaction is the force behind anelectrochemical celllike theGalvanic cellpictured. The battery is made out of a zinc electrode in a ZnSO₄solution connected with a wire and a porous disk to a copper electrode in a CuSO₄solution.

In this type of reaction, ametalatom in a compound or solution is replaced by an atom of another metal. For example,copperis deposited whenzincmetal is placed in acopper(II) sulfatesolution:

Zn (s) + CuSO₄(aq) → ZnSO₄(aq) + Cu (s)

In the above reaction, zinc metal displaces the copper(II) 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:

Zn + Cu²⁺→ Zn²⁺+ Cu

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

Zn → Zn²⁺+ 2 e⁻

And the copper is reduced:

Cu²⁺+ 2 e⁻→ Cu

Other examples

• The reduction ofnitratetonitrogenin the presence of an acid (denitrification):

2 NO−3+ 10 e⁻+ 12 H⁺→ N₂+ 6 H₂O

• Thecombustionofhydrocarbons,such as in aninternal combustion engine,produceswater,carbon dioxide,some partially oxidized forms such ascarbon monoxide,and heatenergy.Complete oxidation of materials containingcarbonproduces carbon dioxide.
• The stepwise oxidation of a hydrocarbon by oxygen, inorganic chemistry,produces water and, successively: analcohol,analdehydeor aketone,acarboxylic acid,and then aperoxide.

Corrosion and rusting

Oxides, such asiron(III) oxideorrust,which consists of hydratediron(III) oxidesFe₂O₃·nH₂O andiron(III) oxide-hydroxide(FeO(OH), Fe(OH)₃), form when oxygen combines with other elements.

Iron rusting inpyritecubes

• The termcorrosionrefers to the electrochemical oxidation of metals in reaction with an oxidant such as oxygen.Rusting,the formation ofiron oxides,is a well-known example of electrochemical corrosion: it forms as a result of the oxidation ofironmetal. Common rust often refers toiron(III) oxide,formed in the following chemical reaction:

4 Fe + 3 O₂→ 2 Fe₂O₃

• The oxidation of iron(II) to iron(III) byhydrogen peroxidein the presence of anacid:

Fe²⁺→ Fe³⁺+ e⁻

H₂O₂+ 2 e⁻→ 2 OH⁻

Here the overall equation involves adding the reduction equation to twice the oxidation equation, so that the electrons cancel:

2 Fe²⁺+ H₂O₂+ 2 H⁺→ 2 Fe³⁺+ 2 H₂O

Disproportionation

Adisproportionationreaction is one in which a single substance is both oxidized and reduced. For example,thiosulfateion with sulfur in oxidation state +2 can react in the presence of acid to form elemental sulfur (oxidation state 0) andsulfur dioxide(oxidation state +4).

S₂O2−3+ 2 H⁺→ S + SO₂+ H₂O

Thus one sulfur atom is reduced from +2 to 0, while the other is oxidized from +2 to +4.

Redox reactions in industry

Cathodic protectionis a technique used to control the corrosion of a metal surface by making it thecathodeof anelectrochemical cell.A simple method of protection connects protected metal to a more easily corroded "sacrificial anode"to act as theanode.The sacrificial metal, instead of the protected metal, then corrodes. A common application of cathodic protection is ingalvanizedsteel, in which a sacrificial zinc coating on steel parts protects them from rust.

Oxidation is used in a wide variety of industries, such as in the production ofcleaning productsand oxidizingammoniato producenitric acid.

Redox reactions are the foundation of electrochemical cells, which can generate electrical energy or supportelectrosynthesis.Metaloresoften contain metals in oxidized states, such as oxides or sulfides, from which the pure metals are extracted bysmeltingat high temperatures in the presence of a reducing agent. The process ofelectroplatinguses redox reactions to coat objects with a thin layer of a material, as inchrome-platedautomotiveparts,silver platingcutlery,galvanizationandgold-platedjewelry.

Redox reactions in biology

Top:ascorbic acid(reduced formofVitamin C)
Bottom:dehydroascorbic acid(oxidized formofVitamin C)

Enzymatic browningis an example of a redox reaction that takes place in most fruits and vegetables.

Many essentialbiologicalprocesses involve redox reactions. Before some of these processes can begin, iron must beassimilatedfrom the environment.

Cellular respiration,for instance, is the oxidation ofglucose(C₆H₁₂O₆) toCO₂and the reduction ofoxygentowater.The summary equation for cellular respiration is:

C₆H₁₂O₆+ 6 O₂→ 6 CO₂+ 6 H₂O + Energy

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

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

Biological energyis frequently stored and released using redox reactions. Photosynthesis involves the reduction ofcarbon dioxideintosugarsand the oxidation ofwaterinto molecular oxygen. The reverse reaction, respiration, oxidizes sugars to produce carbon dioxide and water. As intermediate steps, the reduced carbon compounds are used to reducenicotinamide adenine dinucleotide(NAD⁺) to NADH, which then contributes to the creation of aproton gradient,which drives the synthesis ofadenosine triphosphate(ATP) and is maintained by the reduction of oxygen. In animal cells,mitochondriaperform similar functions.

See also:Membrane potential

Free radicalreactions are redox reactions that occur as part ofhomeostasisand killingmicroorganisms.In these reactions, an electron detaches from a molecule and then re-attaches almost instantly. Free radicals are part of redox molecules and can become harmful to the human body if they do not reattach to the redox molecule or anantioxidant.

The termredox stateis often used to describe the balance ofGSH/GSSG,NAD⁺/NADH andNADP⁺/NADPHin a biological system such as acellororgan.The redox state is reflected in the balance of several sets of metabolites (e.g.,lactateandpyruvate,beta-hydroxybutyrateandacetoacetate), whose interconversion is dependent on these ratios. Redox mechanisms also control some cellular processes. Redox proteins and their genes must be co-located for redox regulation according to theCoRR hypothesisfor the function ofDNAinmitochondriaandchloroplasts.

Redox cycling

Wide varieties ofaromatic compoundsareenzymaticallyreduced to formfree radicalsthat contain one more electron than their parent compounds. In general, the electron donor is any of a wide variety offlavoenzymesand theircoenzymes.Once formed, these anion free radicals reduce molecular oxygen tosuperoxideand 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 afutile cycleor redox cycling.

Redox reactions in geology

Blast furnaces ofTřinec Iron and Steel Works,Czech Republic

Minerals are generally oxidized derivatives of metals. Iron is mined as itsmagnetite(Fe₃O₄). Titanium is mined as its dioxide, usually in the form ofrutile(TiO₂). These oxides must be reduced to obtain the corresponding metals, often achieved by heating these oxides with carbon or carbon monoxide as reducing agents.Blast furnacesare 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+ 3 CO → 2 Fe + 3 CO₂

Redox reactions in soils

Electron transferreactions are central to myriad processes and properties in soils, andredox potential,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 andpaddy riceproduction was seminal for subsequent work on thermodynamic aspects of redox and plant root growth in soils.Later work built on this foundation, and expanded it for understanding redox reactions related to heavy metal oxidation state changes,pedogenesisand morphology, organic compound degradation and formation,free radicalchemistry,wetlanddelineation,soil remediation,and various methodological approaches for characterizing the redox status of soils.

Mnemonics

Main article:List of chemistry mnemonics

The key terms involved in redox can be confusing.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.Thesemnemonicsare commonly used by students to help memorise the terminology:

• "OIL RIG"—oxidationisloss of electrons,reductionisgain of electrons
• "LEO the lion says GER [grr]" —loss ofelectrons isoxidation,gain ofelectrons isreduction
• "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 toelectrolytic cellswhich release stored electricity, and can be recharged with electricity. PANIC does not apply to cells that can be recharged with redox materials. Thesegalvanic or voltaic cells,such asfuel cells,produce electricity from internal redox reactions. Here, the positive electrode is the cathode and the negative is the anode.

See also

• Anaerobic respiration
• Bessemer process
• Bioremediation
• Calvin cycle
• Chemical equation
• Chemical looping combustion
• Citric acid cycle
• Electrochemical series
• Electrochemistry
• Electrolysis
• Electron equivalent
• Electron transport chain
• Electrosynthesis
• Galvanic cell
• Hydrogenation
• Membrane potential
• Microbial fuel cell
• Murburn concept
• Nucleophilic abstraction
• Organic redox reaction
• Oxidative addition and reductive elimination
• Oxidative phosphorylation
• Partial oxidation
• Pro-oxidant
• Redox gradient
• Redox potential
• Redox therapy
• Reducing agent
• Reducing atmosphere
• Reduction potential
• Thermic reaction
• Transmetalation
• Sulfur cycle

Further reading

• Schüring, J.; Schulz, H. D.; Fischer, W. R.; Böttcher, J.; Duijnisveld, W. H., eds. (1999).Redox: Fundamentals, Processes and Applications.Heidelberg: Springer-Verlag. p. 246.hdl:10013/epic.31694.d001.ISBN978-3-540-66528-1.
• Tratnyek, Paul G.; Grundl, Timothy J.; Haderlein, Stefan B., eds. (2011).Aquatic Redox Chemistry.ACS Symposium Series. Vol. 1071.doi:10.1021/bk-2011-1071.ISBN978-0-8412-2652-4.