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–oxidation reaction between sodium and chlorine, with theOIL RIG mnemonic[1]
Redox (/ˈrɛdɒks/RED-oks,/ˈriːdɒks/REE-doks,reduction–oxidation[2] oroxidation–reduction[3]: 150 ) is a type ofchemical reaction in which theoxidation states of thereactants change.[4] Oxidation is the loss ofelectrons 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 (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 onesubstrate to another. For example, in therusting ofiron, 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.[5] Inhydrogenation, bonds likeC=C are reduced bytransfer of hydrogen atoms.
"Redox" is aportmanteau of the words "REDuction" and "OXidation." The term "redox" was first used in 1928.[6]
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 orreducing agent loses electrons and is oxidized, and the oxidant oroxidizing 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,[7] e.g.,Fe2+ /Fe3+ .The oxidation alone and the reduction alone are each called ahalf-reaction because two half-reactions always occur together to form a whole reaction.[5]
Inelectrochemical reactions the oxidation and reduction processes do occur simultaneously but are separated in space.
Oxidation originally implied a reaction with oxygen to form an oxide. Later, the term was expanded to encompasssubstances 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.[8]: 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 asoxidizing agents, oxidants, or oxidizers. The oxidant removes electrons from another substance, and is thus itself reduced.[8]: A50 Because it "accepts" electrons, the oxidizing agent is also called anelectron 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 highlyelectronegative 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.[9]
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.[3]: 159 Because it donates electrons, the reducing agent is also called anelectron donor. Electron donors can also formcharge transfer complexes with electron acceptors. The word reduction originally referred to the loss in weight upon heating a metallicore such as ametal oxide to extract the metal. In other words, ore was "reduced" to metal.[10]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.[10] Reducing equivalent refers tochemical species which transfer the equivalent of oneelectron in redox reactions. The term is common inbiochemistry.[11] A reducing equivalent can be an electron or a hydrogen atom as ahydride ion.[12]
Hydride transfer reagents, such asNaBH4 andLiAlH4, reduce by atom transfer: they transfer the equivalent of hydride or H−. These reagents are widely used in the reduction ofcarbonyl compounds toalcohols.[14][15] A related method of reduction involves the use of hydrogen gas (H2) as sources of H atoms.[3]: 288
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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 mixing.[16]
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[17] andouter sphere electron transfer.[18]
Analysis of bond energies andionization energies in water allows calculation of the thermodynamic aspects of redox reactions.[19]
Standard electrode potentials (reduction potentials)
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+).[8]: 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 anoxidation 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
In the reaction betweenhydrogen andfluorine, hydrogen is being oxidized and fluorine is being reduced:
H2 + F2 → 2 HF
This spontaneous reaction releases a large amount of energy (542 kJ per 2 g of hydrogen) because two H-F bonds are much stronger than one H-H bond and one F-F bond. This reaction can be analyzed as twohalf-reactions. The oxidation reaction converts hydrogen toprotons:
A redox reaction is the force behind anelectrochemical cell like theGalvanic 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.
In this type of reaction, ametal atom in a compound or solution is replaced by an atom of another metal. For example,copper is deposited whenzinc metal is placed in acopper(II) sulfate solution:
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:
The termcorrosion refers 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 ofiron metal. 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 peroxide in 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:
Adisproportionation 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 (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.[8]: 176
Cathodic protection is a technique used to control the corrosion of a metal surface by making it thecathode of 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.
Enzymatic browning is an example of a redox reaction that takes place in most fruits and vegetables.
Many essentialbiological processes involve redox reactions. Before some of these processes can begin, iron must beassimilated from the environment.[21]
Cellular respiration, for instance, is the oxidation ofglucose (C6H12O6) toCO2 and the reduction ofoxygen towater. 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+).Photosynthesis and cellular respiration are complementary, but photosynthesis is not the reverse of the redox reaction in cellular respiration:
Biological energy is frequently stored and released using redox reactions. Photosynthesis involves the reduction ofcarbon dioxide intosugars and the oxidation ofwater 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 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,mitochondria perform similar functions.
The termredox state is often used to describe the balance ofGSH/GSSG, NAD+/NADH andNADP+/NADPH in a biological system such as acell ororgan. The redox state is reflected in the balance of several sets of metabolites (e.g.,lactate andpyruvate,beta-hydroxybutyrate andacetoacetate), 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 hypothesis for the function ofDNA inmitochondria andchloroplasts.
Wide varieties ofaromatic compounds areenzymatically reduced to formfree radicals that contain one more electron than their parent compounds. In general, the electron donor is any of a wide variety offlavoenzymes and theircoenzymes. Once formed, these anion free radicals reduce molecular oxygen tosuperoxide 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 afutile cycle or redox cycling.
Minerals are generally oxidized derivatives of metals.Iron is mined as ores such asmagnetite (Fe3O4) andhematite (Fe2O3).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 furnaces 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:[22]
Electron transfer reactions 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 rice production was seminal for subsequent work on thermodynamic aspects of redox and plant root growth in soils.[23] 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.[24][25]
The key terms involved in redox can be confusing.[26][27] 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.[28] Thesemnemonics are commonly used by students to help memorise the terminology:[29]
"LEO the lion says GER [grr]" —loss ofelectrons isoxidation,gain ofelectrons isreduction[26][27][28][29]
"LEORA says GEROA" — the loss of electrons is called oxidation (reducing agent); the gain of electrons is called reduction (oxidizing agent).[28]
"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 cells which 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.
^abcHaustein, Catherine Hinga (2014)."Oxidation-reduction reaction". In K. Lee Lerner; Brenda Wilmoth Lerner (eds.).The Gale Encyclopedia of Science (5th ed.). Farmington Hills, MI: Gale Group.
^abcdPetrucci, Ralph H.; Harwood, William S.; Herring, F. Geoffrey (2017).General Chemistry: Principles and Modern applications (11th ed.). Toronto: Pearson.ISBN978-0-13-293128-1.
^"Nitric Acid Fact Sheet"(PDF).Department of Environmental Safety, Sustainability & Risk. University of Maryland. RetrievedFebruary 12, 2024.
^abWhitten, Kenneth W.; Gailey, Kenneth D.; Davis, Raymond E. (1992).General Chemistry (4th ed.). Saunders College Publishin. p. 147.ISBN0-03-072373-6.
^Jain JL (2004).Fundamentals of Biochemistry. S. Chand.ISBN81-219-2453-7.
^Lehninger AL, Nelson DL, Cox MM (January 1, 2017).Lehninger Principles of Biochemistry (Seventh ed.). New York, NY.ISBN9781464126116.OCLC986827885.{{cite book}}: CS1 maint: location missing publisher (link)
^Bartlett, Richmond J.; James, Bruce R. (1991). "Redox chemistry of soils".Advances in Agronomy.39:151–208.
^James, Bruce R.; Brose, Dominic A. (2012). "Oxidation-reduction phenomena". In Huang, Pan Ming; Li, Yuncong; Sumner, Malcolm E. (eds.).Handbook of soil sciences: properties and processes (second ed.). Boca Raton, Florida:CRC Press. pp. 14-1 -- 14-24.ISBN978-1-4398-0305-9.
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.
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