Redox

Redox (short for reduction–oxidation reaction) (pronunciation: /ˈrɛdɒks/ redoks or /ˈriːdɒks/ reedoks[1]) is a type of chemical reaction in which the oxidation states of atoms are changed. Redox reactions are characterized by the transfer of electrons between chemical species, most often with one species (the reducing agent) undergoing oxidation (losing electrons) while another species (the oxidizing agent) undergoes reduction (gains electrons).[2] The chemical species from which the electron is stripped is said to have been oxidized, while the chemical species to which the electron is added is said to have been reduced. In other words:

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

For example, during the combustion of wood, electrons are transferred from carbon atoms in the wood to oxygen atoms in the air. The oxygen atoms undergo reduction, gaining electrons, while the carbon atoms undergo oxidation, losing electrons. Thus oxygen is the oxidizing agent and carbon the reducing agent in this reaction.[3] Although oxidation reactions are commonly associated with the formation of oxides from oxygen molecules, oxygen is not necessarily included in such reactions, as other chemical species can serve the same function.[3]

Redox reactions can occur relatively slowly, as in the formation of rust, or much more rapidly, as in the case of burning fuel. There are simple redox processes, such as the oxidation of carbon to yield carbon dioxide (CO2) 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.

NaF
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.
Demonstration of the reaction between a strong oxidising and a reducing agent. When a few drops of glycerol (mild reducing agent) are added to powdered potassium permanganate (strong oxidising agent), a violent redox reaction accompanied by self-ignition starts.

Etymology

"Redox" is a portmanteau of the words "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 a 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. Since electrons are negatively charged, it is also helpful to think of this as reduction in electrical charge.

The electrochemist John Bockris has used the words electronation and deelectronation to describe reduction and oxidation processes respectively when they occur at electrodes.[4] These words are analogous to protonation and deprotonation, but they have not been widely adopted by chemists worldwide.

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 into saturated fats, R−CH=CH−R + H2 → R−CH2−CH2−R). The word "redox" was first used in 1928.[5]

Definitions

The processes of oxidation and reduction occur simultaneously and cannot happen independently of one another, similar to the acid–base reaction.[3] 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 general descriptions are not precisely correct. Although oxidation and reduction properly refer to a change in oxidation state — the actual transfer of electrons may never occur. The oxidation state of an atom is the fictitious charge that an atom would have if all bonds between atoms of different elements were 100% ionic. Thus, oxidation is best 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).

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 oxidizing form, e.g., Fe2+
/Fe3+

Oxidizers

GHS-pictogram-rondflam
The international pictogram for oxidizing chemicals.

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. Oxygen is the quintessential oxidizer.

Oxidants are usually chemical substances with elements in high oxidation states (e.g., H
2
O
2
, MnO
4
, CrO
3
, Cr
2
O2−
7
, OsO
4
), or else highly electronegative elements (O2, F2, Cl2, Br2) that can gain extra electrons by oxidizing another substance.

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 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. Electropositive elemental metals, such as lithium, sodium, magnesium, iron, zinc, and aluminium, are good reducing agents. These metals donate or give away electrons readily. Hydride transfer reagents, such as NaBH4 and LiAlH4, are widely used in organic chemistry,[6][7] primarily in the reduction of carbonyl compounds to alcohols. Another method of reduction involves the use of hydrogen gas (H2) with a palladium, platinum, or nickel catalyst. These catalytic reductions are used primarily in the reduction of carbon-carbon double or triple bonds.

Standard electrode potentials (reduction potentials)

Each half-reaction has a standard electrode potential (E0
cell
), 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:

12 H2 → H+ + e.

The electrode potential of each half-reaction is also known as its reduction potential E0
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 → ​12 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+).[8]

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

E0
cell
= E0
cathode
E0
anode

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

E0
ox
 = –E0
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

E0
cell
= E0
red(cathode)
+ E0
ox(anode)

Examples of redox reactions

Redox reaction
Illustration of a redox reaction

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

H
2
+ F
2
→ 2 HF

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

the oxidation reaction:

H
2
→ 2 H+ + 2 e

and the reduction reaction:

F
2
+ 2 e → 2 F

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).

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.

When adding the reactions together the electrons are canceled:

H
2
2 H+ + 2 e
F
2
+ 2 e
2 F

H2 + F2 2 H+ + 2 F

And the ions combine to form hydrogen fluoride:

2 H+ + 2 F → 2 HF

The overall reaction is:

H
2
+ F
2
→ 2 HF

Metal displacement

Galvanic cell with no cation flow
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.

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:

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 ionic equation for this reaction is:

Zn + Cu2+ → Zn2+ + Cu

As two half-reactions, 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

Rust screw
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
PyOx
Iron rusting in pyrite cubes
  • 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:
    4 Fe + 3 O2 → 2 Fe2O3
  • The oxidation of iron(II) to iron(III) by hydrogen peroxide in the presence of an acid:
    Fe2+ → Fe3+ + e
    H2O2 + 2 e → 2 OH
Overall equation:
2 Fe2+ + H2O2 + 2 H+ → 2 Fe3+ + 2 H2O

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 galvanized 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 cleaning products and oxidizing ammonia to produce nitric acid, which is used in most fertilizers.

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 temperature 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.

Redox reactions in biology

ascorbic acid
dehydroascorbic acid

Many important biological processes involve redox reactions.

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:

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

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

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+) to NADH, 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.

Free radical reactions are redox reactions that occur as a part of homeostasis 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 antioxidant. 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 GSH/GSSG, 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., 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.

Redox cycling

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 a futile cycle or redox cycling.

Redox reactions in geology

UraniumMineUtah
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.

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.

Balancing redox reactions

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.

Acidic media

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

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

Unbalanced reaction: Mn2+(aq) + NaBiO3(s) → Bi3+(aq) + MnO
4
(aq)
Oxidation: 4 H2O(l) + Mn2+(aq) → MnO
4
(aq) + 8 H+(aq) + 5 e
Reduction: 2 e + 6 H+ + BiO
3
(s) → Bi3+(aq) + 3 H2O(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):

8 H2O(l) + 2 Mn2+(aq) → 2 MnO
4
(aq) + 16 H+(aq) + 10 e
10 e + 30 H+ + 5 BiO
3
(s) → 5 Bi3+(aq) + 15 H2O(l)

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

14 H+(aq) + 2 Mn2+(aq) + 5 NaBiO3(s) → 7 H2O(l) + 2 MnO
4
(aq) + 5 Bi3+(aq) + 5 Na+
(aq)

Basic media

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

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

Unbalanced reaction: KMnO4 + Na2SO3 + H2O → MnO2 + Na2SO4 + KOH
Reduction: 3 e + 2 H2O + MnO
4
→ MnO2 + 4 OH
Oxidation: 2 OH + SO2−
3
SO2−
4
+ H2O + 2 e

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

6 e + 4 H2O + 2 MnO
4
→ 2 MnO2 + 8 OH
6 OH + 3 SO2−
3
→ 3 SO2−
4
+ 3 H2O + 6 e

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

2 KMnO4 + 3 Na2SO3 + H2O → 2 MnO2 + 3 Na2SO4 + 2 KOH

Mnemonics

The key terms involved in redox are often confusing.[9][10] 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.[11] These mnemonics are commonly used by students to help memorise the terminology:[12]

  • "OIL RIG" — oxidation is loss of electrons, reduction is gain of electrons.[9][10][11][12]
  • "LEO the lion says GER" — loss of electrons is oxidation, gain of electrons is reduction.[9][10][11][12]
  • "LEORA says GEROA" — loss of electrons is oxidation (reducing agent), gain of electrons is reduction (oxidizing agent).[11]
  • "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).

See also

References

Notes

  1. ^ "redox – definition of redox in English | Oxford Dictionaries". Oxford Dictionaries | English. Archived from the original on 2017-10-01. Retrieved 2017-05-15.
  2. ^ "Redox Reactions". wiley.com. Archived from the original on 2012-05-30. Retrieved 2012-05-09.
  3. ^ a b c Haustein, Catherine Hinga (2014). K. Lee Lerner and Brenda Wilmoth Lerner (eds.). Oxidation–reduction reaction. The Gale Encyclopedia of Science. 5th edition. Farmington Hills, MI: Gale Group.CS1 maint: Uses editors parameter (link)
  4. ^ Bockris, John O'M.; Reddy, Amulya K. N. (1970). Modern Electrochemistry. Plenum Press. pp. 352–3.
  5. ^ Harper, Douglas. "redox". Online Etymology Dictionary.
  6. ^ Hudlický, Miloš (1996). Reductions in Organic Chemistry. Washington, D.C.: American Chemical Society. p. 429. ISBN 978-0-8412-3344-7.
  7. ^ Hudlický, Miloš (1990). Oxidations in Organic Chemistry. Washington, D.C.: American Chemical Society. p. 456. ISBN 978-0-8412-1780-5.
  8. ^ Electrode potential values from: Petrucci, Ralph H.; Harwood, William S.; Herring, F. Geoffrey (2002). General chemistry: principles and modern applications (8th ed.). Upper Saddle River, N.J: Prentice Hall. p. 832. ISBN 978-0-13-014329-7. LCCN 2001032331. OCLC 46872308.
  9. ^ a b c Robertson, William (2010). More Chemistry Basics. National Science Teachers Association. p. 82. ISBN 978-1-936137-74-9.
  10. ^ a b c Phillips, John; Strozak, Victor; Wistrom, Cheryl (2000). Chemistry: Concepts and Applications. Glencoe McGraw-Hill. p. 558. ISBN 978-0-02-828210-7.
  11. ^ a b c d Rodgers, Glen (2012). Descriptive Inorganic, Coordination, and Solid-State Chemistry. Brooks/Cole, Cengage Learning. p. 330. ISBN 978-0-8400-6846-0.
  12. ^ a b c Zumdahl, Steven; Zumdahl, Susan (2009). Chemistry. Houghton Mifflin. p. 160. ISBN 978-0-547-05405-6.

Bibliography

  • 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. ISBN 978-3-540-66528-1.
  • Tratnyek, Paul G.; Grundl, Timothy J.; Haderlein, Stefan B., eds. (2011). Aquatic Redox Chemistry. ACS Symposium Series. 1071. doi:10.1021/bk-2011-1071. ISBN 978-0-8412-2652-4.

External links

Cofactor (biochemistry)

A cofactor is a non-protein chemical compound or metallic ion that is required for an enzyme's activity as a catalyst, a substance that increases the rate of a chemical reaction. Cofactors can be considered "helper molecules" that assist in biochemical transformations. The rates at which these happen are characterized by in an area of study called enzyme kinetics.

Cofactors can be divided into two types, either inorganic ions, or complex organic molecules called coenzymes. (Note that some scientists limit the use of the term "cofactor" to inorganic substances; both types are included here.) Coenzymes are mostly derived from vitamins and other organic essential nutrients in small amounts.

Coenzymes are further divided into two types. The first is called a "prosthetic group", which consists of a coenzyme that is tightly or even covalently, and permanently bound to a protein. The second type of coenzymes are called "cosubstrates", and are transiently bound to the protein. Cosubstrates may be released from a protein at some point, and then rebind later. Both prosthetic groups and cosubstrates have the same function, which is to facilitate the reaction of enzymes and protein. An inactive enzyme without the cofactor is called an apoenzyme, while the complete enzyme with cofactor is called a holoenzyme.Some enzymes or enzyme complexes require several cofactors. For example, the multienzyme complex pyruvate dehydrogenase at the junction of glycolysis and the citric acid cycle requires five organic cofactors and one metal ion: loosely bound thiamine pyrophosphate (TPP), covalently bound lipoamide and flavin adenine dinucleotide (FAD), cosubstrates nicotinamide adenine dinucleotide (NAD+) and coenzyme A (CoA), and a metal ion (Mg2+).Organic cofactors are often vitamins or made from vitamins. Many contain the nucleotide adenosine monophosphate (AMP) as part of their structures, such as ATP, coenzyme A, FAD, and NAD+. This common structure may reflect a common evolutionary origin as part of ribozymes in an ancient RNA world. It has been suggested that the AMP part of the molecule can be considered to be a kind of "handle" by which the enzyme can "grasp" the coenzyme to switch it between different catalytic centers.

Electrochemistry

Electrochemistry is the branch of physical chemistry that studies the relationship between electricity, as a measurable and quantitative phenomenon, and identifiable chemical change, with either electricity considered an outcome of a particular chemical change or vice versa. These reactions involve electric charges moving between electrodes and an electrolyte (or ionic species in a solution). Thus electrochemistry deals with the interaction between electrical energy and chemical change.

When a chemical reaction is caused by an externally supplied current, as in electrolysis, or if an electric current is produced by a spontaneous chemical reaction as in a battery, it is called an electrochemical reaction. Chemical reactions where electrons are transferred directly between molecules and/or atoms are called oxidation-reduction or (redox) reactions. In general, electrochemistry describes the overall reactions when individual redox reactions are separate but connected by an external electric circuit and an intervening electrolyte.

Electron transport chain

An electron transport chain (ETC) is a series of complexes that transfer electrons from electron donors to electron acceptors via redox (both reduction and oxidation occurring simultaneously) reactions, and couples this electron transfer with the transfer of protons (H+ ions) across a membrane. This creates an electrochemical proton gradient that drives the synthesis of adenosine triphosphate (ATP), a molecule that stores energy chemically in the form of highly strained bonds. The molecules of the chain include peptides, enzymes (which are proteins or protein complexes), and others. The final acceptor of electrons in the electron transport chain during aerobic respiration is molecular oxygen although a variety of acceptors other than oxygen such as sulfate exist in anaerobic respiration.

Electron transport chains are used for extracting energy via redox reactions from sunlight in photosynthesis or, such as in the case of the oxidation of sugars, cellular respiration. In eukaryotes, an important electron transport chain is found in the inner mitochondrial membrane where it serves as the site of oxidative phosphorylation through the action of ATP synthase. It is also found in the thylakoid membrane of the chloroplast in photosynthetic eukaryotes. In bacteria, the electron transport chain is located in their cell membrane.

In chloroplasts, light drives the conversion of water to oxygen and NADP+ to NADPH with transfer of H+ ions across chloroplast membranes. In mitochondria, it is the conversion of oxygen to water, NADH to NAD+ and succinate to fumarate that are required to generate the proton gradient.

Electron transport chains are major sites of premature electron leakage to oxygen, generating superoxide and potentially resulting in increased oxidative stress.

The electron transport chain consists of a spatially separated series of redox reactions in which electrons are transferred from a donor molecule to an acceptor molecule. The underlying force driving these reactions is the Gibbs free energy of the reactants and products. The Gibbs free energy is the energy available ("free") to do work. Any reaction that decreases the overall Gibbs free energy of a system is thermodynamically spontaneous.

The function of the electron transport chain is to produce a transmembrane proton electrochemical gradient as a result of the redox reactions. If protons flow back through the membrane, they enable mechanical work, such as rotating bacterial flagella. ATP synthase, an enzyme highly conserved among all domains of life, converts this mechanical work into chemical energy by producing ATP, which powers most cellular reactions.

A small amount of ATP is available from substrate-level phosphorylation, for example, in glycolysis. In most organisms the majority of ATP is generated in electron transport chains.

Flavin adenine dinucleotide

In biochemistry, flavin adenine dinucleotide (FAD) is a redox-active coenzyme associated with various proteins, which is involved with several important enzymatic reactions in metabolism. A flavoprotein is a protein that contains a flavin group, this may be in the form of FAD or flavin mononucleotide (FMN). There are many flavoproteins besides components of the succinate dehydrogenase complex, including α-ketoglutarate dehydrogenase and a component of the pyruvate dehydrogenase complex; some examples are shown in section 6.

FAD can exist in four different redox states, which are the flavin-N(5)-oxide, quinone, semiquinone, and hydroquinone. FAD is converted between these states by accepting or donating electrons. FAD, in its fully oxidized form, or quinone form, accepts two electrons and two protons to become FADH2 (hydroquinone form). The semiquinone (FADH·) can be formed by either reduction of FAD or oxidation of FADH2 by accepting or donating one electron and one proton, respectively. Some proteins, however, generate and maintain a superoxidized form of the flavin cofactor, the flavin-N(5)-oxide.

Flow battery

A flow battery, or redox flow battery (after reduction–oxidation), is a type of electrochemical cell where chemical energy is provided by two chemical components dissolved in liquids contained within the system and separated by a membrane. Ion exchange (accompanied by flow of electric current) occurs through the membrane while both liquids circulate in their own respective space. Cell voltage is chemically determined by the Nernst equation and ranges, in practical applications, from 1.0 to 2.2 volts.

A flow battery may be used like a fuel cell (where the spent fuel is extracted and new fuel is added to the system) or like a rechargeable battery (where an electric power source drives regeneration of the fuel). While it has technical advantages over conventional rechargeables, such as potentially separable liquid tanks and near unlimited longevity, current implementations are comparatively less powerful and require more sophisticated electronics.

The energy capacity is a function of the electrolyte volume (amount of liquid electrolyte), and the power is a function of the surface area of the electrodes.

Glutathione

Glutathione (GSH) is an antioxidant in plants, animals, fungi, and some bacteria and archaea. Glutathione is capable of preventing damage to important cellular components caused by reactive oxygen species such as free radicals, peroxides, lipid peroxides, and heavy metals. It is a tripeptide with a gamma peptide linkage between the carboxyl group of the glutamate side chain and the amine group of cysteine, and the carboxyl group of cysteine is attached by normal peptide linkage to a glycine.

Thiol groups are reducing agents, existing at a concentration around 5 mM in animal cells. Glutathione reduces disulfide bonds formed within cytoplasmic proteins to cysteines by serving as an electron donor. In the process, glutathione is converted to its oxidized form, glutathione disulfide (GSSG), also called L-(–)-glutathione.

Once oxidized, glutathione can be reduced back by glutathione reductase, using NADPH as an electron donor. The ratio of reduced glutathione to oxidized glutathione within cells is often used as a measure of cellular oxidative stress.

Hydroxylation

Hydroxylation is a chemical process that introduces a hydroxyl group (-OH) into an organic compound. In biochemistry, hydroxylation reactions are often facilitated by enzymes called hydroxylases. Hydroxylation is the first step in the oxidative degradation of organic compounds in air. It is extremely important in detoxification since hydroxylation converts lipophilic compounds into water-soluble (hydrophilic) products that are more readily removed by the kidneys or liver and excreted. Some drugs (for example, steroids) are activated or deactivated by hydroxylation.

Nicotinamide adenine dinucleotide

Nicotinamide adenine dinucleotide (NAD) is a cofactor found in all living cells. The compound is called a dinucleotide because it consists of two nucleotides joined through their phosphate groups. One nucleotide contains an adenine nucleobase and the other nicotinamide. Nicotinamide adenine dinucleotide exists in two forms: an oxidized and reduced form, abbreviated as NAD+ and NADH respectively.

In metabolism, nicotinamide adenine dinucleotide is involved in redox reactions, carrying electrons from one reaction to another. The cofactor is, therefore, found in two forms in cells: NAD+ is an oxidizing agent – it accepts electrons from other molecules and becomes reduced. This reaction forms NADH, which can then be used as a reducing agent to donate electrons. These electron transfer reactions are the main function of NAD. However, it is also used in other cellular processes, most notably a substrate of enzymes that add or remove chemical groups from proteins, in posttranslational modifications. Because of the importance of these functions, the enzymes involved in NAD metabolism are targets for drug discovery.

In organisms, NAD can be synthesized from simple building-blocks (de novo) from the amino acids tryptophan or aspartic acid. In an alternative fashion, more complex components of the coenzymes are taken up from food as niacin. Similar compounds are released by reactions that break down the structure of NAD. These preformed components then pass through a salvage pathway that recycles them back into the active form.

Some NAD is converted into the coenzyme nicotinamide adenine dinucleotide phosphate (NADP). The chemistry of NADP is similar to that of NAD, but it has different role, being predominantly a cofactor in anabolic metabolism.

NAD+ is written with a superscript plus (+) sign because of the formal charge on one of its nitrogen atoms; however, it is actually predominantly a singly charged anion (charge of minus 1) at physiological pH. NADH, on the other hand, is a doubly charged anion because of its two bridging phosphate groups.

Organic redox reaction

Organic reductions or organic oxidations or organic redox reactions are redox reactions that take place with organic compounds. In organic chemistry oxidations and reductions are different from ordinary redox reactions because many reactions carry the name but do not actually involve electron transfer in the electrochemical sense of the word. Instead the relevant criterion for organic oxidation is gain of oxygen and/or loss of hydrogen Simple functional groups can be arranged in order of increasing oxidation state. The oxidation numbers are only an approximation:

When methane is oxidized to carbon dioxide its oxidation number changes from −4 to +4. Classical reductions include alkene reduction to alkanes and classical oxidations include oxidation of alcohols to aldehydes. In oxidations electrons are removed and the electron density of a molecule is reduced. In reductions electron density increases when electrons are added to the molecule. This terminology is always centered on the organic compound. For example, it is usual to refer to the reduction of a ketone by lithium aluminium hydride, but not to the oxidation of lithium aluminium hydride by a ketone. Many oxidations involve removal of hydrogen atoms from the organic molecule, and the reverse, reduction adds hydrogens to an organic molecule.

Many reactions classified as reductions also appear in other classes. For instance conversion of the ketone to an alcohol by lithium aluminium hydride can be considered a reduction but the hydride is also a good nucleophile in nucleophilic substitution. Many redox reactions in organic chemistry have coupling reaction reaction mechanism involving free radical intermediates. True organic redox chemistry can be found in electrochemical organic synthesis or electrosynthesis. Examples of organic reactions that can take place in an electrochemical cell are the Kolbe electrolysis.In disproportionation reactions the reactant is both oxidised and reduced in the same chemical reaction forming two separate compounds.

Asymmetric catalytic reductions and asymmetric catalytic oxidations are important in asymmetric synthesis.

Oxidative stress

Oxidative stress reflects an imbalance between the systemic manifestation of reactive oxygen species and a biological system's ability to readily detoxify the reactive intermediates or to repair the resulting damage. Disturbances in the normal redox state of cells can cause toxic effects through the production of peroxides and free radicals that damage all components of the cell, including proteins, lipids, and DNA. Oxidative stress from oxidative metabolism causes base damage, as well as strand breaks in DNA. Base damage is mostly indirect and caused by reactive oxygen species (ROS) generated, e.g. O2− (superoxide radical), OH (hydroxyl radical) and H2O2 (hydrogen peroxide). Further, some reactive oxidative species act as cellular messengers in redox signaling. Thus, oxidative stress can cause disruptions in normal mechanisms of cellular signaling.

In humans, oxidative stress is thought to be involved in the development of ADHD, cancer, Parkinson's disease, Lafora disease, Alzheimer's disease, atherosclerosis, heart failure, myocardial infarction, fragile X syndrome, sickle-cell disease, lichen planus, vitiligo, autism, infection, chronic fatigue syndrome, and depression and seems to be characteristic of individuals with Asperger syndrome. However, reactive oxygen species can be beneficial, as they are used by the immune system as a way to attack and kill pathogens. Short-term oxidative stress may also be important in prevention of aging by induction of a process named mitohormesis.

Oxidizing agent

In chemistry, an oxidizing agent (oxidant, oxidizer) is a substance that has the ability to oxidize other substances — in other words to cause them to lose electrons. Common oxidizing agents are oxygen, hydrogen peroxide and the halogens.

In one sense, an oxidizing agent is a chemical species that undergoes a chemical reaction that removes one or more electrons from another atom. In that sense, it is one component in an oxidation–reduction (redox) reaction. In the second sense, an oxidizing agent is a chemical species that transfers electronegative atoms, usually oxygen, to a substrate. Combustion, many explosives, and organic redox reactions involve atom-transfer reactions.

Redox (operating system)

Redox is a Unix-like microkernel operating system written in the programming language Rust, a language with a strong focus on safety, stability, and high performance. Redox aims to be secure, usable, and free. Redox is inspired by prior kernels and operating systems, such as SeL4, MINIX, Plan 9, and BSD. It is similar to the GNU or BSD ecosystem, but in a memory-safe language and with modern technology. It is free and open-source software distributed under an MIT License.

Reducing agent

A reducing agent (also called a reductant or reducer) is an element (such as calcium) or compound that loses (or "donates") an electron to another chemical species in a redox chemical reaction. Since the reducing agent is losing electrons, it is said to have been oxidized.

If any chemical is an electron donor (reducing agent), another must be an electron recipient (oxidizing agent). A reducing agent is oxidized because it loses electrons in the redox reaction.

Thus, reducers (reducing agents) "reduce" (or, seen another way, are "oxidized" by) oxidizers (oxidizing agents), and oxidizers "oxidize" (that is, are "reduced" by) reducers.

In their pre-reaction states, reducers have more electrons (that is, they are by themselves reduced) and oxidizers have fewer electrons (that is, they are by themselves oxidized). A reducing agent typically is in one of its lower possible oxidation states and is known as the electron donor. Examples of reducing agents include the earth metals, formic acid, oxalic acid, and sulfite compounds.

For example, consider the overall reaction for aerobic cellular respiration:

C6H12O6(s) + 6O2(g) → 6CO2(g) + 6H2O(l)The oxygen (O2) is being reduced, so it is the oxidizing agent. The glucose (C6H12O6) is being oxidized, so it is the reducing agent.

In organic chemistry, reduction more specifically refers to the addition of hydrogen to a molecule, though the aforementioned definition still applies. For example, benzene is reduced to cyclohexane in the presence of a platinum catalyst:

C6H6 + 3 H2 → C6H12In organic chemistry, good reducing agents are reagents that deliver H2.

Historically, reduction referred to the removal of oxygen from a compound, hence the name 'reduction'. The modern sense of donating electrons is a generalisation of this idea, acknowledging that other components can play a similar chemical role to oxygen.

Reduction potential

Redox potential (also known as oxidation / reduction potential, ORP, pe, ε, or ) is a measure of the tendency of a chemical species to acquire electrons from or lose electrons to an electrode and thereby be reduced or oxidised, respectively. Redox potential is measured in volts (V), or millivolts (mV). Each species has its own intrinsic redox potential; for example, the more positive the reduction potential (reduction potential is more often used due to general formalism in electrochemistry), the greater the species' affinity for electrons and tendency to be reduced. ORP can reflect the antimicrobial potential of the water.

Titration

Titration (also called titrimetry and volumetric analysis) is a common laboratory method of quantitative chemical analysis to determine the concentration of an identified analyte (a substance to be analyzed). A reagent, called the titrant or titrator, is prepared as a standard solution of known concentration and volume. The titrant reacts with a solution of analyte (which may also be called the titrand) to determine the analyte's concentration. The volume of titrant that reacted with the analyte is called the titration volume.

Vanadium redox battery

The vanadium redox battery (VRB), also known as the vanadium flow battery (VFB) or vanadium redox flow battery (VRFB), is a type of rechargeable flow battery that employs vanadium ions in different oxidation states to store chemical potential energy. The vanadium redox battery exploits the ability of vanadium to exist in solution in four different oxidation states, and uses this property to make a battery that has just one electroactive element instead of two. For several reasons, including their relative bulkiness, most vanadium batteries are currently used for grid energy storage, i.e., attached to power plants or electrical grids.

The possibility of creating a vanadium flow battery was explored variously by Pissoort in the 1930s, NASA researchers in the 1970s, and Pellegri and Spaziante in the 1970s, but none of them were successful in demonstrating the technology. The first successful demonstration of the all-vanadium redox flow battery which employed vanadium in a solution of sulfuric acid in each half was by Maria Skyllas-Kazacos at the University of New South Wales in the 1980s. Her design used sulfuric acid electrolytes, and was patented by the University of New South Wales in Australia in 1986.The main advantages of the vanadium redox battery are that it can offer almost unlimited energy capacity simply by using larger electrolyte storage tanks; it can be left completely discharged for long periods with no ill effects; if the electrolytes are accidentally mixed, the battery suffers no permanent damage; a single state of charge between the two electrolytes avoids the capacity degradation due to a single cell in non-flow batteries; the electrolyte is aqueous and inherently safe and non-flammable; and the generation 3 formulation using a mixed acid solution developed by the Pacific Northwest National Laboratory operates over a wider temperature range allowing for passive cooling.

VRFBs can be used at depth of discharge (DOD) around 90% and more, i.e. deeper DODs than solid-state batteries (e.g. lithium-based and sodium-based batteries, which are usually specified with DOD=80%). In addition, VRFBs exhibit very long cycle lives: most producers specify cycle durability in excess of 15,000-20,000 charge/discharge cycles. These values are far beyond the cycle lives of solid-state batteries, which is usually in the order of 4,000-5,000 charge/discharge cycles. Consequently, the levelized cost of energy (LCOE, i.e. the system cost divided by the usable energy, the cycle life, and round-trip efficiency) of present VRFB systems is typically in the order of a few tens of $ cents or € cents, namely much lower than the LCOEs of equivalent solid-state batteries and close to the targets of $0.05 and €0.05, stated by the US Department of Energy and the European Commission Strategic Energy Technology (SET) Plan, respectively.The main disadvantages with vanadium redox technology are a relatively poor energy-to-volume ratio in comparison with standard storage batteries (although the Generation 3 formulation has doubled the energy density of the system), and the aqueous electrolyte makes the battery heavy and therefore only useful for stationary applications. Another disadvantage is the relatively high toxicity of oxides of vanadium (see vanadium § Safety).

Numerous companies and organizations involved in funding and developing vanadium redox batteries include Avalon Battery, Vionx (formerly Premium Power), UniEnergy Technologies and Ashlawn Energy in the United States; Renewable Energy Dynamics Technology in Ireland; Enerox GmbH (formerly Gildemeister energy storage) in Austria; Cellennium in Thailand; Rongke Power in China; Prudent Energy in China; Sumitomo in Japan; H2, Inc. in South Korea; redT in Britain., Australian Vanadium in Australia, and the now defunct Imergy (formerly Deeya). Lately, also several smaller size vanadium redox flow batteries were brought to market (for residential applications) mainly from StorEn Technologies (USA), Schmid Group, VoltStorage and Volterion (all three from Germany), VisBlue (Denmark) or Pinflow energy storage (Czechia).

Vitamin K epoxide reductase

Vitamin K epoxide reductase (VKOR) is an enzyme (EC 1.17.4.4) that reduces vitamin K after it has been oxidised in the carboxylation of glutamic acid residues in blood coagulation enzymes. VKOR is a member of a large family of predicted enzymes that are present in vertebrates, Drosophila, plants, bacteria and archaea. In some plant and bacterial homologues, the VKOR domain is fused with domains of the thioredoxin family of oxidoreductases.Four cysteine residues and one residue, which is either serine or threonine, are identified as likely active-site residues. Solved bacterial VKOR structures has enabled more insights into the catalytic mechanism. All VKORs are transmembrane proteins with at least three TM helices at the catalytic core. The quinone to be reduced is bound by a redox-active CXXC motif in the C-terminal helices, similar to the DsbB active site. Two other cysteines to the N-terminal are located in a loop outside of the transmembrane region; they relay electrons with a redox protein (or in the case of the bacterial homolog, its own fused domain).The human gene for VKOR is called VKORC1 (VKOR complex subunit 1). It is the target of anticoagulant warfarin. Its partner is a redox protein with an unknown identity. There is also a similar gene called VKORC1L1.

Working electrode

The working electrode is the electrode in an electrochemical system on which the reaction of interest is occurring. The working electrode is often used in conjunction with an auxiliary electrode, and a reference electrode in a three electrode system. Depending on whether the reaction on the electrode is a reduction or an oxidation, the working electrode is called cathodic or anodic, respectively. Common working electrodes can consist of materials ranging from inert metals such as gold, silver or platinum, to inert carbon such as glassy carbon, boron doped diamond or pyrolytic carbon, and mercury drop and film electrodes. Chemically modified electrodes are employed for the analysis of both organic and inorganic samples.

Zinc–cerium battery

Zinc–cerium batteries are a type of redox flow battery first developed by Plurion Inc. (UK) during the 2000s. In this rechargeable battery, both negative zinc and positive cerium electrolytes are circulated though an electrochemical flow reactor during the operation and stored in two separated reservoirs. Negative and positive electrolyte compartments in the electrochemical reactor are separated by a cation-exchange membrane, usually Nafion (DuPont). The Ce(III)/Ce(IV) and Zn(II)/Zn redox reactions take place at the positive and negative electrodes, respectively. Since zinc is electroplated during charge at the negative electrode this system is classified as a hybrid flow battery. Unlike in zinc–bromine and zinc–chlorine redox flow batteries, no condensation device is needed to dissolve halogen gases. The reagents used in the zinc-cerium system are considerably less expensive than those used in the vanadium flow battery.

Due to the high standard electrode potentials of both zinc and cerium redox reactions in aqueous media, the open-circuit cell voltage is as high as 2.43 V. Among the other proposed rechargeable aqueous flow battery systems, this system has the largest cell voltage and its power density per electrode area is second only to H2-Br2 flow battery. Methanesulfonic acid is used as supporting electrolyte, as it allows high concentrations of both zinc and cerium; the solubility of the corresponding methanesulfonates is 2.1 M for Zn, 2.4 M for Ce(III) and up to 1.0 M for Ce(IV). Methanesulfonic acid is particularly well suited for industrial electrochemical applications and is considered to be a green alternative to other support electrolytes.The Zn-Ce flow battery is still in early stages of development. The main technological challenge is the control of the inefficiency and self discharge (Zn corrosion via hydrogen evolution) at the negative electrode. In commercial terms, the need for expensive Pt-Ti electrodes increases the capital cost of the system in comparison to other RFBs.

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