Redox: From Wikipedia, the free encyclopedia Source: https://en.wikipedia.org/wiki/Redox From Wikipedia, the free encyclopedia Jump to navigation Jump to search For other uses, see  Redox (disambiguation) . 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. 0:47 Demonstration of the reaction between a strong oxidizing and a reducing agent. When a few drops of  glycerol  (mild reducing agent) are added to powdered  potassium permanganate  (strong oxidizing agent), a violent redox reaction accompanied by self-ignition starts. Example of a  reduction–oxidation  reaction between sodium and chlorine, with the  OIL RIG  mnemonic [1] Redox  ( reduction–oxidation ,  / ˈ r ɛ d ɒ k s /   RED-oks ,  / ˈ r iː d ɒ k s /   REE-doks [2] ) is a type of  chemical reaction  in which the  oxidation states  of  substrate  change. [3] Oxidation  is the  loss  of electrons or an  increase  in the oxidation state of a chemical or atoms within it. Reduction  is the  gain  of electrons or a  decrease  in the oxidation state of a chemical or atoms within it. There are two classes of redox reactions: electron-transfer  where only one (usually) electron flows from the reducing agent to the oxidant. This type of redox reaction is often discussed in term of redox couples and electrode potentials. atom transfer , where an atom transfers from one substrate to another. For example, in the  rusting  of iron, the oxidation state of iron atoms increases as it converts to an oxide and simultaneously the oxidation state of oxygen decreases as it accepts electrons released by iron. Although oxidation reactions are commonly associated with the formation of oxides, other chemical species can serve the same function. [4]  In  hydrogenation ,  C=C  (and other) bonds are reduced by transfer of hydrogen atoms. Contents 1 Terminology 1.1 Oxidants 1.2 Reducers 1.3 Electronation and deelectronation 2 Rates, mechanisms, and energies 3 Standard electrode potentials (reduction potentials) 4 Examples of redox reactions 4.1 Metal displacement 4.2 Other examples 4.3 Corrosion and rusting 4.4 Disproportionation 5 Redox reactions in industry 6 Redox reactions in biology 6.1 Redox cycling 7 Redox reactions in geology 8 Redox reactions in soils 9 Mnemonics 10 See also 11 References 12 Further reading 13 External links Terminology "Redox" is a  combination  of the words "reduction" and "oxidation". The term "redox" was first used in 1928. [5]  The processes of oxidation and reduction occur simultaneously and cannot occur independently. [4]  In redox processes, the reductant transfers electrons to the oxidant. Thus, in the reaction, the reductant or  reducing agent  loses electrons and is oxidized, and the oxidant or  oxidizing agent  gains electrons and is reduced. The pair of an oxidizing and reducing agent that is involved in a particular reaction is called a  redox pair . A  redox couple  is a reducing species and its corresponding oxidizing form, [6]  e.g.,  Fe 2+ /  Fe 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. Oxidants Oxidation  originally implied reaction with oxygen to form an oxide. Later, the term was expanded to encompass oxygen-like substances that accomplished parallel chemical reactions. Ultimately, the meaning was generalized to include all processes involving the loss of electrons. Substances that have the ability to  oxidize  other substances (cause them to lose electrons) are said to be  oxidative  or  oxidizing , and are known as  oxidizing agents , oxidants, or oxidizers. The oxidant (oxidizing agent) removes electrons from another substance, and is thus itself reduced. And, because it "accepts" electrons, the oxidizing agent is also called an  electron acceptor . Oxidants are usually chemical substances with elements in high oxidation states (e.g.,  H 2 O 2 ,  MnO − 4 ,  CrO 3 ,  Cr 2 O 2− 7 ,  OsO 4 ), or else highly  electronegative  elements ( O 2 ,  F 2 ,  Cl 2 ,  Br 2 ) that can gain extra electrons by oxidizing another substance. [ citation needed ] Oxidizers are oxidants but the term is mainly reserved for sources of oxygen, particularly in the context of explosions.  Nitric acid  is an oxidizer. The  international   pictogram  for oxidizing chemicals Main article:  Oxidizing agent Oxygen  is the quintessential oxidizer. Reducers 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 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. 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  demonstrated that this loss of weight was due to the loss of oxygen as a gas. Later, scientists realized that the metal atom gains electrons in this process. The meaning of  reduction  then became generalized to include all processes involving a gain of electrons.  Reducing equivalent  refers to  chemical species  which transfer the equivalent of one  electron  in  redox  reactions. The term is common in biochemistry. [7]  A reducing equivalent can be an electron, a hydrogen atom, as a  hydride ion . [8] 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 relatively readily. They transfer electrons. Hydride transfer reagents , such as  NaBH 4  and  LiAlH 4 , reduce by atom transfer: they transfer the equivalent of hydride or H - . These reagents widely used in [the reduction of  carbonyl  compounds to  alcohols . [9] [10]  A related method of reduction involves the use of hydrogen gas (H 2 ) as sources of H atoms. Electronation and deelectronation The electrochemist  John Bockris  proposed the words  electronation  and  deelectronation  to describe reduction and oxidation processes, respectively, when they occur at  electrodes . [11]  These words are analogous to  protonation  and  deprotonation . [12]  They have not been widely adopted by chemists worldwide, although  IUPAC  has recognized the term electronation. [13] Rates, mechanisms, and energies Redox reactions can occur slowly, as in the formation of  rust , or rapidly, as in the case of burning fuel. Electron transfer reactions are generally fast, occurring within the time of mixing. The mechanisms of atom-transfer reactions are highly variable because many kinds of atoms can be transferred. Such reactions can also be quite complex, i.e. involve many steps. The mechanisms of electron-transfer reactions occur by two distinct pathways,  inner sphere electron transfer  and  outer sphere electron transfer . Analysis of bond energies and ionization energies in water allow calculation of the thermodynamic aspects of redox reactions. Standard electrode potentials (reduction potentials) Each half-reaction has a  standard  electrode potential  ( E o 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: 1 ⁄ 2  H 2  → 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 ⁄ 2  H 2  by definition, positive for oxidizing agents stronger than H +  (e.g., +2.866 V for F 2 ) and negative for oxidizing agents that are weaker than H +  (e.g., −0.763 V for Zn 2+ ). [14] 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 an  oxidation 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 between  hydrogen  and  fluorine , hydrogen is being oxidized and fluorine is being reduced: H 2  +  F 2  → 2 HF This reaction is spontaneous and releases 542 kJ per 2 g of hydrogen because the H-F bond is much stronger than the F-F bond. This reaction can be analyzed as two  half-reactions . The oxidation reaction converts hydrogen to protons: H 2  → 2  H +  + 2  e − The reduction reaction converts fluorine to the fluoride anion: F 2  + 2 e −  → 2  F − The half reactions are combined so that the electrons cancel: H 2 → 2 H +  + 2 e − F 2  + 2 e − → 2 F − H 2  + F 2 → 2 H +  + 2 F − The protons and fluoride combine to form  hydrogen fluoride  in a non-redox reaction: 2 H +  + 2 F −  → 2 HF The overall reaction is: H 2  +  F 2  → 2 HF Metal displacement 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 ZnSO 4  solution connected with a wire and a porous disk to a copper electrode in a CuSO 4  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)+ CuSO 4 (aq) → ZnSO 4 (aq) + Cu(s) In the above reaction, zinc metal displaces the copper(II) ion from copper sulfate solution and thus liberates free copper metal. The reaction is spontaneous and releases 213 kJ per 65 g of zinc. The ionic equation for this reaction is: Zn + Cu 2+  → Zn 2+  + Cu As two  half-reactions , it is seen that the zinc is oxidized: Zn → Zn 2+  + 2 e − And the copper is reduced: Cu 2+  + 2 e −  → Cu Other examples The reduction of  nitrate  to  nitrogen  in the presence of an acid ( denitrification ): {\displaystyle {\ce {2NO3- + 10e- + 12H+ -> N2 + 6 H2O}}} The  combustion  of  hydrocarbons , such as in an  internal combustion engine , produces  water ,  carbon dioxide , some partially oxidized forms such as  carbon monoxide , and heat  energy . Complete oxidation of materials containing  carbon  produces carbon dioxide. The stepwise oxidation of a hydrocarbon by oxygen, in  organic chemistry , produces water and, successively: an  alcohol , an  aldehyde  or a  ketone , a  carboxylic acid , and then a  peroxide . Corrosion and rusting Oxides, such as  iron(III) oxide  or  rust , which consists of hydrated  iron(III) oxides  Fe 2 O 3 · n H 2 O and  iron(III) oxide-hydroxide  (FeO(OH), Fe(OH) 3 ), form when oxygen combines with other elements 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: {\displaystyle {\ce {4Fe + 3O2 -> 2Fe2O3}}} The oxidation of iron(II) to iron(III) by  hydrogen peroxide  in the presence of an acid: {\displaystyle {\ce {Fe^{2+}->{Fe^{3+}}+e-}}} {\displaystyle {\ce {H2O2 + 2e- -> 2OH-}}} Here the overall equation involves adding the reduction equation to twice the oxidation equation, so that the electrons cancel: {\displaystyle {\ce {{2Fe^{2+}}+{H2O2}+2H+->{2Fe^{3+}}+2H2O}}} Disproportionation A  disproportionation  reaction is one in which a single substance is both oxidized and reduced. For example,  thiosulfate  ion with sulfur in oxidation state +2 can react in the presence of acid to form elemental sulfur (oxidation state 0) and  sulfur dioxide  (oxidation state +4). S 2 O 3 2-  + 2 H +  → S + SO 2  + H 2 O Thus one sulfur atom is reduced from +2 to 0, while the other is oxidized from +2 to +4. [15] 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. [ citation needed ] Oxidation is used in a wide variety of industries such as in the production of  cleaning products  and oxidizing  ammonia  to produce  nitric acid . 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 . [ citation needed ] Redox reactions in biology Top:  ascorbic acid  ( reduced form  of  Vitamin C ) Bottom:  dehydroascorbic acid  ( oxidized form  of  Vitamin C ) Enzymatic browning  is an example of a redox reaction that takes place in most fruits and vegetables. Many important  biological  processes involve redox reactions. Before some of these processes can begin iron must be assimilated from the environment. [16] Cellular respiration , for instance, is the oxidation of  glucose  (C 6 H 12 O 6 ) to  CO 2  and the reduction of  oxygen  to  water . The summary equation for cell respiration is: C 6 H 12 O 6  + 6 O 2  → 6 CO 2  + 6 H 2 O 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 CO 2  + 6 H 2 O +  light energy  → C 6 H 12 O 6  + 6 O 2 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 Wide varieties 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 Blast furnaces of  Třinec Iron and Steel Works , Czech Republic Minerals are generally oxidized derivatives of metals. Iron is mined as its  magnetite  (Fe 3 O 4 ). Titanium is mined as its dioxide, usually in the form of  rutile  (TiO 2 ). To obtain the corresponding metals, these oxides must be reduced, which is often achieved by heating these oxides with carbon or carbon monoxide as reducing agents.  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: [17] Fe 2 O 3  + 3CO → 2Fe + 3CO 2 Redox reactions in soils [ edit ] Electron transfer reactions are central to myriad processes and properties in soils, and electron "activity", quantified as Eh (platinum electrode potential (voltage) relative to the standard hydrogen electrode) or pe (analogous to pH as -log electron activity), is a master variable, along with pH, that controls and is governed by chemical reactions and biological processes. Early theoretical research with applications to flooded soils and paddy rice production was seminal for subsequent work on thermodynamic aspects of redox and plant root growth in soils. [18]  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. [19] [20] Mnemonics [ edit ] Main article:  List of chemistry mnemonics The key terms involved in redox can be confusing. [21] [22]  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. [23]  These  mnemonics  are commonly used by students to help memorise the terminology: [24] " OIL RIG " —  o xidation  i s  l oss of electrons,  r eduction  i s  g ain of electrons [21] [22] [23] [24] "LEO the lion says GER [grr]" —  l oss of  e lectrons is  o xidation,  g ain of  e lectrons is  r eduction [21] [22] [23] [24] "LEORA says GEROA" — the loss of electrons is called oxidation (reducing agent); the gain of electrons is called reduction (oxidizing agent). [23] "RED CAT" and "AN OX", or "AnOx RedCat" ("an ox-red cat") — reduction occurs at the cathode and the anode is for oxidation "RED CAT gains what AN OX loses" – reduction at the cathode gains (electrons) what anode oxidation loses (electrons) "PANIC" – Positive Anode and Negative is Cathode. This applies to  electrolytic cells  which release stored electricity, and can be recharged with electricity. PANIC does not apply to cells that can be recharged with redox materials. These  galvanic or voltaic cells , such as  fuel cells , produce electricity from internal redox reactions. Here, the positive electrode is the cathode and the negative is the anode.