Redox

• 1. Introduction

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  Redox is two opposite events happening at the same instant, and neither can happen alone. One substance loses electrons. Another gains them. The two processes are locked together, oxidation and reduction, occurring simultaneously in a single chemical reaction. The word itself is a portmanteau, stitched from reduction and oxidation. It first appeared in a 1928 article by Leonor Michaelis and Louis B. Flexner. Hidden inside that compact name is a tangle that has confused students for generations. A reagent that loses electrons is oxidized, yet it is called the reducing agent. A reagent that gains electrons is reduced, yet it is called the oxidizing agent. How did a process named after oxygen come to describe reactions where no oxygen appears at all? Why does iron rusting on a fence and fuel burning in an engine belong to the same family as the electricity in a fuel cell? And what does it mean to give a reaction a voltage, measured against a single agreed-upon electrode? The answers reach from blast furnaces to flooded rice paddies, from the inside of a living cell to a sacrificial slab of metal bolted onto a pipeline.

• 2. Two Ways To Trade

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  Electron transfer is the cleaner of the two routes. Usually a single electron flows from the atom, ion, or molecule being oxidized to the one being reduced. Chemists discuss this kind of reaction in terms of redox couples and electrode potentials. A redox couple pairs a reducing species with its corresponding oxidizing form. The other route moves whole atoms. In atom transfer, an atom passes from one substrate to another, and the bookkeeping of electrons rides along with it. Rusting iron is the classic case. As iron converts to an oxide, the oxidation state of its atoms rises, while the oxidation state of oxygen falls as it accepts the electrons the iron releases. Oxidation does not require oxygen, despite the name. In hydrogenation, carbon-carbon double bonds are reduced through the transfer of hydrogen atoms instead. The two halves of any redox reaction each have their own name. A single oxidation, considered by itself, is a half-reaction, and so is a single reduction. Neither stands alone, because two half-reactions always occur together to form a whole. In electrochemical reactions the two processes still happen at the same time, but they are pulled apart in space, separated onto different electrodes.

• 3. Donors And Acceptors

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  The oxidant is the one that removes electrons from another substance, and in doing so it is itself reduced. Because it accepts electrons, the oxidizing agent is also called an electron acceptor. Oxidants tend to be substances with elements already in high oxidation states, or else highly electronegative elements such as O2, F2, Cl2, Br2, and I2, which can pull in extra electrons by oxidizing something else. The word oxidizer narrows the idea. Oxidizers are oxidants, but the term is mostly reserved for sources of oxygen, especially when explosions are involved, and nitric acid stands as a strong example. Reductants run the trade in the opposite direction. A reducing agent donates electrons to another substance and is itself oxidized, which makes it an electron donor. Electron donors can even form charge transfer complexes with electron acceptors. Reductants in chemistry are strikingly varied. Electropositive elemental metals such as lithium, sodium, magnesium, iron, zinc, and aluminium give up electrons readily and make good reducing agents. Some reductants work by atom transfer rather than bare electron flow. Hydride transfer reagents like NaBH4 and LiAlH4 hand over the equivalent of hydride, or H-, and are widely used to reduce carbonyl compounds to alcohols. A related approach uses hydrogen gas, H2, as the source of hydrogen atoms. The idea of a reducing equivalent, a species that transfers the equivalent of one electron, is common in biochemistry, where that equivalent may be an electron or a hydrogen atom carried as a hydride ion.

• 4. Reduced To Metal

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  Reduction once meant a loss in weight. Heating a metallic ore such as a metal oxide drove off mass, and the ore was said to be reduced to metal. Antoine Lavoisier showed that the missing weight was oxygen leaving as a gas. Only later did scientists recognize that the metal atom was gaining electrons in the bargain, and the meaning of reduction widened to cover all processes involving a gain of electrons. Oxidation traveled a parallel path. It originally implied a reaction with oxygen to form an oxide. The term then stretched to include substances that did chemically what oxygen did, and finally generalized to any loss of electrons or rise in oxidation state. The electrochemist John Bockris proposed fresh vocabulary for these events at electrodes, calling reduction electronation and oxidation de-electronation. The words mirror protonation and deprotonation, and IUPAC has recognized both.

• 5. Fast, Slow, And Measured

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  Rust forms slowly while burning fuel races, yet both are redox. Electron transfer reactions are generally fast, finishing within the time it takes to mix the reactants. Atom-transfer reactions are far more variable, because so many different kinds of atoms can be moved, and such reactions can grow complex across many steps. Electron-transfer reactions themselves run by two distinct pathways, inner sphere electron transfer and outer sphere electron transfer. Analysis of bond energies and ionization energies in water makes it possible to calculate the thermodynamic side of redox reactions. Each half-reaction carries a standard electrode potential, written E, equal to the voltage at equilibrium under standard conditions in a cell whose cathode runs the half-reaction in question. The anode in that setup is a standard hydrogen electrode, where hydrogen is oxidized as H2 turns into H+ plus an electron. This electrode potential is also the reduction potential, a measure of how strongly an oxidizing agent tends to be reduced. By definition its value is zero for H+ plus an electron forming H2. It is positive for oxidizing agents stronger than H+, such as +2.866 V for F2, and negative for weaker ones, such as -0.763 V for Zn2+. The cell potential is the difference between the two electrode potentials. The anode potential can instead be written as an oxidation potential, the negative of the electrode potential, which gauges how readily a reducing agent is oxidized but does not stand for any real physical potential at an electrode.

• 6. Spontaneous Energy

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  542 kJ pour out for every 2 g of hydrogen that reacts with fluorine. Hydrogen is oxidized to protons and fluorine is reduced to fluoride, and the reaction is spontaneous because two H-F bonds are far stronger than one H-H bond and one F-F bond together. Split into half-reactions, the electrons cancel, the protons and fluoride combine into hydrogen fluoride in a non-redox step. Metal displacement releases energy on the same principle. When zinc metal is placed in a copper(II) sulfate solution, the zinc displaces the copper(II) ion, and free copper metal deposits out. That spontaneous reaction gives off 213 kJ for every 65 g of zinc, the zinc oxidized and the copper reduced as two half-reactions. Some reactions turn a single substance against itself. In disproportionation, one substance is both oxidized and reduced at once. Thiosulfate ion, with sulfur in oxidation state +2, can react in the presence of acid to form elemental sulfur at oxidation state 0 and sulfur dioxide at oxidation state +4. One sulfur atom drops from +2 to 0 while another climbs from +2 to +4. Combustion shows the stepwise face of redox. Burning hydrocarbons in an internal combustion engine yields water, carbon dioxide, partially oxidized forms such as carbon monoxide, and heat, while the stepwise oxidation of a hydrocarbon by oxygen passes through an alcohol, an aldehyde or ketone, a carboxylic acid, and then a peroxide.

• 7. Industry And Earth

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  Cathodic protection turns a metal surface into the cathode of an electrochemical cell to hold off corrosion. The simple version wires the protected metal to a more easily corroded sacrificial anode, so the sacrificial metal corrodes instead. Corrosion itself is the electrochemical oxidation of metals by an oxidant such as oxygen, and rusting, the formation of iron oxides, is its best-known form, with common rust often meaning iron(III) oxide. Redox sits beneath much of heavy industry. Oxidation drives the making of cleaning products and the oxidizing of ammonia to produce nitric acid. Redox reactions are the foundation of electrochemical cells, which generate electrical energy or support electrosynthesis. Electroplating coats objects in a thin layer of material, seen in chrome-plated automotive parts, silver-plated cutlery, galvanization, and gold-plated jewelry. The same chemistry shapes geology. Minerals are generally oxidized derivatives of metals. Iron is mined as ores such as magnetite, Fe3O4, and hematite, Fe2O3, and titanium as its dioxide, usually rutile, TiO2. These oxides must be reduced to yield the metals, often by heating with carbon or carbon monoxide as reducing agents. Blast furnaces are the reactors where iron oxides and coke, a form of carbon, combine to produce molten iron. In soils, electron transfer reactions are central to countless processes, and redox potential, quantified as Eh or as pe, acts as a master variable alongside pH. Early theoretical work on flooded soils and paddy rice production opened a line of research that later reached into heavy metal oxidation states, pedogenesis, organic compound degradation, free radical chemistry, wetland delineation, and soil remediation.

• 8. The Engine Of Life

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  Iron must be assimilated from the environment before some essential biological processes can even begin. Aerobic cellular respiration is the oxidation of substrates, in this case glucose, paired with the reduction of oxygen to water. That process leans heavily on the reduction of NAD+ to NADH and the reverse oxidation of NADH back to NAD+. Photosynthesis and cellular respiration are complementary, yet photosynthesis is not simply the reverse of respiration's redox reaction. Biological energy is stored and released through redox. Photosynthesis reduces carbon dioxide into sugars and oxidizes water into molecular oxygen, while respiration runs the other way, oxidizing sugars to make carbon dioxide and water. Along the path, reduced carbon compounds help reduce nicotinamide adenine dinucleotide, NAD+, to NADH, which feeds a proton gradient that drives the synthesis of adenosine triphosphate, ATP, sustained by the reduction of oxygen. In animal cells, mitochondria carry out similar work. Chemists describe the balance of GSH/GSSG, NAD+/NADH, and NADP+/NADPH as the redox state, reflected in metabolite pairs such as lactate and pyruvate or beta-hydroxybutyrate and acetoacetate. The CoRR hypothesis holds that redox proteins and their genes must be co-located for redox regulation, the proposed function of DNA in mitochondria and chloroplasts. Some of this chemistry can spin uselessly. In redox cycling, aromatic compounds are enzymatically reduced into free radicals that hold one more electron than their parents, with flavoenzymes and their coenzymes serving as electron donors. These anion free radicals then reduce molecular oxygen to superoxide and regenerate the unchanged parent compound, a pattern described as a futile cycle. To keep the vocabulary straight, students lean on mnemonics. OIL RIG means oxidation is loss, reduction is gain. LEO the lion says GER carries the same lesson, while AnOx RedCat fixes oxidation at the anode and reduction at the cathode.