oxidation–reduction reaction

Introduction
also called  redox reaction 

      any chemical reaction in which the oxidation number of a participating chemical species changes. The term covers a large and diverse body of processes. Many oxidation– reduction reactions are as common and familiar as fire, the rusting and dissolution of metals, the browning of fruit, and respiration and photosynthesis—basic life functions.

General considerations

Major classifications
      Most oxidation–reduction (redox) processes involve the transfer of oxygen atoms, hydrogen atoms, or electrons, with all three processes sharing two important characteristics: (1) they are coupled—i.e., in any oxidation reaction a reciprocal reduction occurs, and (2) they involve a characteristic net chemical change—i.e., an atom or electron goes from one unit of matter to another. Both reciprocity and net change are illustrated below in examples of the three most common types of oxidation–reduction reactions.

Oxygen-atom (oxygen) transfer
      Carbon reacts with mercury(II) oxide (a compound in which mercury has a bonding capacity expressed as +2; see below Oxidation-state change (oxidation–reduction reaction)) to produce carbon dioxide and mercury metal. This reaction can be written in equation form:

      Carbon, receiving oxygen, is oxidized; mercury(II) oxide, losing oxygen, undergoes the complementary reduction; and the net change is the transfer of two oxygen atoms from mercury(II) oxide units to a carbon atom.

Hydrogen-atom (hydrogen) transfer
      Hydrogen atoms are transferred from hydrazine, a compound of nitrogen and hydrogen, to oxygen in the following reaction:

      Hydrazine, losing hydrogen, is oxidized to molecular nitrogen, while oxygen, gaining hydrogen, is reduced to water.

Electron transfer
      Zinc metal and copper(II) ion react in water solution, producing copper metal and an aqueous (denoted by aq) zinc ion according to the equation

      With the transfer of two of its electrons, the zinc metal is oxidized, becoming an aqueous zinc ion, while the copper(II) ion, gaining electrons, is reduced to copper metal. Net change is the transfer of two electrons, lost by zinc and acquired by copper.

      Because of their complementary nature, the oxidation and reduction processes together are referred to as redox reactions. The reactant that brings about the oxidation is called the oxidizing agent, and that reagent is itself reduced by the reducing agent. In the examples given above, mercury(II) oxide, oxygen, and the copper(II) ion are oxidizing agents, and carbon, hydrazine, and zinc are the reducing agents.

General theory
Stoichiometric basis
      Describing the redox processes as above conveys no information about the mechanism by which change takes place. A complete description of the net chemical change for a process is known as the stoichiometry of the reaction, which provides the characteristic combining proportions of elements and compounds. Reactions are classified as redox and nonredox on the basis of stoichiometry; oxygen-atom, hydrogen-atom, and electron transfer are stoichiometric categories.

Oxidation-state change
      Comprehensive definitions of oxidation and reduction have been made possible by modern molecular (molecule) structure theory. Every atom consists of a positive nucleus, surrounded by negative electrons, which determine the bonding characteristics of each element. In forming chemical bonds, atoms donate, acquire, or share electrons. This makes it possible to assign every atom an oxidation number, which specifies the number of its electrons that can be involved in forming bonds with other atoms. From the particular atoms in a molecule and their known bonding (chemical bonding) capacities, the bonding pattern within a molecule is determined, and each atom is regarded as being in a specific oxidation state, expressed by an oxidation number.

      Redox processes are defined as reactions accompanied by oxidation-state changes: an increase in an atom's oxidation number corresponds to an oxidation; a decrease, to a reduction. In this generalized theory, three examples of ways in which oxidation-state changes can occur are by oxygen-atom (gain, oxidation; loss, reduction), hydrogen-atom (loss, oxidation; gain, reduction), and electron (loss, oxidation; gain, reduction) transfer. The oxidation-state change definition is usually compatible with the above rules for applying the oxygen-atom-transfer and hydrogen-atom-transfer criteria and always compatible with the electron-transfer criterion when it is applicable. The oxidation state of any atom is indicated by a roman numeral following the name or symbol for the element. Thus, iron(III), or Fe(III), means iron in an oxidation state of +3. The uncombined Fe(III) ion is simply Fe3+.

Historical origins of the redox concept
      Of the chemical processes now regarded as redox reactions, combustion was the earliest focus of philosophical and scientific attention. The Greek scientific philosopher Empedocles listed fire as one of the four elements of matter. In more modern times the phlogiston theory enjoyed scientific popularity. This theory was first articulated in 1697 by G.E. Stahl of Germany. As noted earlier, it asserted that matter releases an elementary constituent, phlogiston, during combustion. Thus, the burning of charcoal was interpreted as the loss of phlogiston from carbon to the air. The theory was also applied to processes other than combustion; in the recovery of a metal from its oxide by heating with charcoal, for example, phlogiston was regarded as being transferred from carbon to the oxide.

      Phlogiston saturation was believed to be responsible for the limited ability of air in a closed container to support combustion. A notable consequence of the phlogiston theory was the notion that an oxide of a metal, such as mercury(II) oxide (HgO), was a chemically simpler substance than the metal itself: the metal could be obtained from the oxide only by the addition of phlogiston. The phlogiston theory, however, could provide no acceptable explanation of the gain in weight when an oxide is formed from a metal.

Combustion and oxide formation
      Late in the 18th century, the interrelated work of Joseph Priestley (Priestley, Joseph) and Antoine-Laurent Lavoisier (Lavoisier, Antoine-Laurent) led to the overthrow of the phlogiston theory. Lavoisier saw Priestley's discovery of oxygen in 1774 as the key to the weight gains known to accompany the burning of sulfur and phosphorus and the calcination of metals (oxide formation). In his Traité élémentaire de chimie, he clearly established that combustion consists of a chemical combination between oxygen from the atmosphere and combustible matter (see below Combustion and flame (oxidation–reduction reaction)). By the end of the century, his ideas were widely accepted and had been successfully applied to the more complex processes of respiration and photosynthesis. Reactions in which oxygen was consumed were classified as oxidations, while those in which oxygen was lost were termed reductions.

Electrochemical reactions (electrochemical reaction)
      During the 19th century, the evolving field of electrochemistry led to a broadened view of oxidation. It was possible, for instance, to produce the ferric, or iron(III), ion from the ferrous, or iron(II), ion at the anode (positive electrode, where electrons are absorbed from solution) of an electrochemical cell (a device in which chemical energy is converted to electrical energy), according to the equation:

      Molecular oxygen could effect a similar transformation, according to the equation:

      The similarity of the two processes led to a precursor of the electron-transfer explanation for redox reactions. After the discovery of the electron, the conviction that oxidation and reduction are accomplished through electron loss and gain became firmly entrenched. Thus, early in the 20th century chemists tended to attribute all redox reactions to the transfer of electrons. Later work on chemical bonding, however, demonstrated the incorrectness of that description. An electronegativity scale (listing of elements in descending order of their tendency to attract and hold bonding electrons) provided a firm basis for the oxidation-state assignments on which oxidation–reduction definitions have become based.

Examples of oxidation–reduction reactions
      Molecular oxygen is a conspicuously important oxidizing agent. It will directly oxidize all but a few of the metals and most of the nonmetals as well. Often these direct oxidations lead to normal oxides, such as those of lithium (Li), zinc (Zn), phosphorus (P), and sulfur (S).

      Organic foodstuffs are oxidized to carbon dioxide and water in respiration. The reaction stoichiometry can be illustrated for glucose, a simple sugar:

      Although the oxygen–glucose reaction is slow at ambient temperatures outside the living cell, it proceeds quickly under the influence of enzymatic catalysis within the body. Essentially all organic compounds react with oxygen under appropriate conditions, but the reaction rates at ordinary temperatures and pressures vary greatly.

      Many other oxidizing agents serve as oxygen-atom sources. Hydrogen peroxide (H2O2), acid chromate ion (HCrO4), and hypochlorous acid (HClO) are reagents often used in oxygen-atom-transfer reactions—for example, in the following reactions:

      In the simplest hydrogen-atom transfers, molecular hydrogen serves as the hydrogen-atom source. The hydrogenations (hydrogenation) of ethylene and of molecular nitrogen are illustrative in the following equations:

      Reactions of molecular hydrogen are characteristically slow at ordinary temperatures. The hydrogenation of molecular nitrogen and of olefins such as ethylene (an olefin is an unsaturated hydrocarbon compound; it has at least two adjacent carbon atoms joined by a double bond to which other atoms or groups of atoms can be joined directly) is a process of extraordinary commercial importance and requires catalysts to occur at useful rates.

      Hydrogen-atom transfer from an organic molecule to a suitable acceptor is a common mode of organic oxidation. The oxidation of formic acid by permanganate and that of ethanol by acid chromate share stoichiometry that features hydrogen-atom loss by the organic species, as shown in the following equations:

      The oxidizing agents permanganate and acid chromate, typical of many hydrogen-atom acceptors, undergo complicated changes rather than simple hydrogen-atom addition.

      Electron-transfer stoichiometry is usually associated with metal ions in aqueous solution, as shown in the following equations:

      Many positively charged metal ions have been shown to be bonded to water molecules, so that their electron-transfer reaction occurs between rather complex molecular groups. The iron ion formulas above, for example, are more properly written as [Fe(H2O)6]2+ and [Fe(H2O)6]3+ to reflect the presence of six water molecules bonded to the metal ion. Simple electron transfer between free ions is known only in the gas phase, as in this argon–sodium reaction:

      Several other types of redox reactions do not fall in the oxygen-atom, hydrogen-atom, or electron-transfer categories. Among these are reactions of fluorine, chlorine, bromine, and iodine. These four elements, known as the halogens (halogen element), form diatomic molecules, which are versatile oxidizing agents. The following examples are typical:

      Such reactions often qualify as redox processes only in the broad sense that oxidation-state changes occur. The oxidation-state characterization extends oxidation–reduction chemistry to include examples from the reactions of all the chemical elements.

Significance of redox reactions
      Oxidation–reduction reactions have vast importance not only in chemistry but in geology and biology as well. The surface of the Earth is a redox boundary between the planet's reduced metallic core and an oxidizing atmosphere. The Earth's crust is largely composed of metal oxides, and the oceans are filled with water, an oxide of hydrogen. The tendency of nearly all surface materials to be oxidized by the atmosphere is reversed by the life process of photosynthesis. Because they are constantly renewed by the photosynthetic reduction of carbon dioxide, life's complex compounds can continue to exist on the Earth's surface.

      For similar reasons, much of chemical technology hinges on the reduction of materials to oxidation states lower than those that occur in nature. Such basic chemical products as ammonia, hydrogen, and nearly all the metals are produced by reductive industrial processes. When not used as structural materials, these products are reoxidized in their commercial applications. The weathering of materials, including wood, metals, and plastics, is oxidative, since, as the products of technological or photosynthetic reductions, they are in oxidation states lower than those stable in the atmosphere.

      Solar radiation is converted to useful energy by a redox cycle that operates continually on a global scale. Photosynthesis converts radiant energy into chemical potential energy by reducing carbon compounds to low oxidation states, and this chemical energy is recovered either through enzymatic oxidations at ambient temperatures or during combustion at elevated temperatures.

Theoretical aspects

Oxidation states
      The idea of assigning an oxidation state to each of the atoms in a molecule evolved from the electron-pair concept of the chemical bond. Atoms within a molecule are held together by the force of attraction that the nuclei of two or more of them exert on electrons (valence electron) in the space between them. In many cases this sharing of electrons can be regarded as involving electron-pair bonds between adjacent nuclei. Electron-pair bonding is often diagrammed so as to show all the bonding and nonbonding valence electrons; e.g., the structures of atomic hydrogen, atomic chlorine, and hydrogen chloride shown below (each dot represents one valence electron):

       Pauling electronegativities of selected elementsThe hydrogen chloride diagram reflects the presence, in the internuclear region, of two electrons that are under the mutual attractive influence of both the hydrogen and chlorine nuclei. Oxidation states for the hydrogen and chlorine in HCl are assigned according to the net charges that remain on H and Cl when the shared electrons are assigned to the atom that has the greater attraction for them. Through physical measurements on isolated atoms and simple molecules, these relative attractive powers have been determined. The Table (Pauling electronegativities of selected elements) lists the electronegativity values for some important elements.

 In the hydrogen chloride molecule the chlorine is more electronegative than hydrogen and is, therefore, assigned both shared electrons. Chlorine has seven valence electrons in its neutral state. Having acquired an eighth electron in its reaction with hydrogen, chlorine is considered to have an oxidation state of −1. Hydrogen, on the other hand, is assigned +1, having lost the single valence electron that it has in its neutral state. Charges arrived at in this way are the basis for oxidation-state assignments, conventionally represented by roman numerals, such as in H(I) and Cl(−I) for the constituents of HCl. Because determination of oxidation states is simply a method of conceptually distributing shared electrons to individual atoms, the same number of electrons must be accounted for, before and after such assignment. The Table—> includes examples of molecules that have multiple bonds. The oxidation states of the atoms involved are added up algebraically in the table, and their sum must always equal the net charge on the molecule. There is no physical reality to oxidation states; they simply represent the results of calculations based on a formal rule.

      Oxidation states can be assigned for most common molecules with the help of a few guidelines. First, electrons shared by two atoms of the same element are divided equally; accordingly, elements are always in oxidation state of 0, regardless of their allotropic (allotropy) form (allotropic refers to the phenomenon of an element's having two or more forms; e.g., carbon can exist as diamond or graphite and in both cases is in the 0 oxidation state). Second, only fluorine is more electronegative than oxygen. Therefore, except in compounds containing oxygen–oxygen or oxygen–fluorine bonds, oxygen can be reliably assigned the oxidation state −2. Similarly, hydrogen is less electronegative than fluorine, oxygen, nitrogen, chlorine, sulfur, and carbon (F, O, N, Cl, S, and C), so it is in the +1 oxidation state in its combinations with those elements. For many common compounds containing only hydrogen, oxygen, and a third element, the third element's oxidation state can be calculated, assuming oxidation numbers of +1 for hydrogen and −2 for oxygen. When bonds are present between two elements that differ little in electronegativity, however, oxidation-state assignments become doubtful, and the distinction between redox and nonredox processes is not evident.

      There is a general reluctance, particularly regarding organic systems, to assume oxidation-state changes when the reaction results can be accounted for by the transfer or addition of water (H2O), ammonia (NH3), the hydroxide ion (OH), or the ions of hydrogen (H+), chlorine (Cl), bromine (Br), or iodine (I), or combinations of these species; e.g., the ammonium ion (NH4+), hydrogen chloride (HCl). The reason is that, in these molecules and ions, the elements are present in their most typical oxidation states: hydrogen(I), chlorine(−I), oxygen(−II), bromine (−I), iodine(−I), and nitrogen(−III).

Oxygen-atom transfer reactions
      The oxidation-state concept clarifies the relationship between oxygen-atom, hydrogen-atom, and electron transfer. The oxygen- and hydrogen-transfer criteria apply only when oxygen and hydrogen occur in their typical oxidation states. An example of an appropriate reaction involving oxygen-atom transfer is the reduction of ferrous oxide by carbon monoxide:

      In terms of oxidation-state changes, this oxygen-atom transfer is equivalent to the two-electron reduction of iron and complementary two-electron oxidation of carbon:

      Oxygen, which occurs in the oxidation state −2 in both reactants and products in the first equation, is not shown in the second. In transferring, the oxygen atom leaves two electrons behind, causing the reduction of iron, and acquires two electrons from the carbon atom, oxidizing the carbon.

      In a similar way, the hydrogenation of ethylene corresponds to a two-electron reduction of the two-carbon skeleton:

      In this example also, the second equation includes only the atoms that change oxidation states: the four hydrogen atoms initially present in ethylene are in the +1 oxidation state in both reactants and products and are therefore omitted. Each of the two neutral hydrogen atoms can be regarded as giving up an electron to, and thereby reducing, one of the carbon atoms. This example also demonstrates clearly that the oxidation that complements the reduction of ethylene is that of the two hydrogen atoms in H2—namely, from the 0 to the +1 oxidation state. General application of the oxidation-state concept leads to a formal viewpoint toward all redox reactions as electron-transfer reactions.

Half reactions
      One of the basic reasons that the concept of oxidation–reduction reactions helps to correlate chemical knowledge is that a particular oxidation or reduction can often be carried out by a wide variety of oxidizing or reducing agents. Reduction of the iron(III) ion to the iron(II) ion by four different reducing agents provides an example:

      Production of the same change in the aqueous iron(III) ion by different reductants emphasizes the fact that the reduction is a characteristic reaction of the iron system itself, and, therefore, the process may be written without specifying the identity of the reducing agent in the following way:

      Hypothetical equations of this type are known as half reactions. The symbol e, which stands for an electron, serves as a reminder that an unspecified reducing agent is required to bring about the change. Half reactions can be written, equally, for the reducing agents in the four reactions with ferric ion:

      Although hypothetical, half reactions are properly balanced chemical processes. Since V2+(aq) increases its oxidation number by one, from +2 to +3, in the first half reaction, an electron is shown as a product of the change. Similarly, two electrons are produced when the oxidation number of zinc increases from 0 to +2 in the second half reaction. When half reactions for hypothetical isolated oxidations and reductions are combined, the electrons must cancel if the equation for a possible overall chemical reaction is to result.

      The use of half reactions is a natural outgrowth of the application of the electron-transfer concept to redox reactions. Since the oxidation-state principle allows any redox reaction to be analyzed in terms of electron transfer, it follows that all redox reactions can be broken down into a complementary pair of hypothetical half reactions. Electrochemical cells (electrolytic cell) (in which chemical energy can be converted to electrical energy, and vice versa) provide some physical reality to the half-reaction idea. Oxidation and reduction half reactions can be carried out in separate compartments of electrochemical cells, with the electrons flowing through a connecting wire and the circuit completed by some arrangement for ion migration between the two compartments (but the migration need not involve any of the materials of the oxidation–reduction reactions themselves).

Redox potentials for common half reactions
 The analysis of the electrical potential, or voltage, developed by pairing various half reactions in electrochemical cells has led to the determination of redox potentials for a substantial number of common half reactions. While a detailed description of redox potentials requires the methods of thermodynamics (the branch of physics concerned with the role played by heat in the transformation of matter or energy), a great deal of useful information can be obtained from redox potentials with minimal recourse to formal theory. Basically, a table of half-cell potentials is a summary of the relative tendencies of different oxidations and reductions to occur. The Table—> lists selected half reactions and their corresponding reduction potentials (which are symbolized by E°).

 The physical significance of the values is directly linked to several agreements about their use. First, the greater the value of E° (the reduction potential), the greater the tendency of a half reaction to proceed from left to right (as written). The half reactions in the Table—> are listed from top to bottom in order of decreasing E°: the higher a reaction's position on the list, the greater the tendency of the reactants to accept electrons. In other words, reagents high on the list, such as fluorine gas (F2) and permanganate ion (MnO4), are strong oxidizing agents. Second, the reduction of hydrogen ions (hydrogen ion) (H+) to hydrogen gas (H2) is arbitrarily assigned the value 0 volts. Half cells with positive reduction potentials involve reactants that are more readily reduced than H+; conversely, those with negative potentials involve reactants that are more difficult to reduce than hydrogen ions.

      With the aid of reduction potentials, it is possible to predict whether a particular oxidation–reduction reaction can occur. The predictions require breaking down the overall reaction into two half reactions of known reduction potentials. For example, if a strip of zinc metal is dipped into a solution containing copper(II) ion, the possibility exists for a redox process, which can be regarded as the sum of the half reactions aqueous zinc ion (Zn2+[aq]) to zinc metal (ZnCu2+[aq]) to copper metal (CuPauling electronegativities of selected elementsCombining these two half reactions requires writing the zinc ion to zinc metal half reaction the reverse of the way it appears in the Table (Pauling electronegativities of selected elements). When the direction of a half reaction is reversed, so that it can be added to another half reaction, the sign of its redox potential is also reversed (in this case, from negative to positive), and the two reduction potential values are then added.

      The resulting E° value for the net reaction, +1.10 volts, measures the tendency of the net reaction to occur. If Eo for a particular net reaction is positive, the process may be expected to occur spontaneously when the reactants are mixed at specified concentrations (one mole per litre; see below Oxidation–reduction equilibria (oxidation–reduction reaction)). Therefore, it is predicted that copper metal should be deposited on a strip of zinc metal when the latter is immersed in a solution of a copper(II) salt. This reaction is, in fact, readily observed in the laboratory. A more specific physical interpretation of the +1.10 volt value is that it represents the voltage that would be produced by an ideal electrochemical cell based on the copper(II) ion to copper metal and zinc(II) ion to zinc metal half reactions with all the reagents at specified concentrations.

      When the same two half cells are combined, with both their directions (and therefore the signs of their redox potentials) reversed, it is predicted that the reverse reaction, the depositing of zinc metal from a zinc(II) ion solution onto a copper strip, will not occur spontaneously. As in the case of E° values for half reactions, those for net redox reactions also change sign when the direction of the reaction is reversed.

       Pauling electronegativities of selected elements Pauling electronegativities of selected elementsThe results of the copper–zinc system can be applied more generally to the half reactions in the Table (Pauling electronegativities of selected elements). For example, copper(II) ion in water (Cu2+[aq]) is an oxidant strong enough to force a half reaction lower on the table to proceed spontaneously in the opposite direction of that written. Therefore, not only is copper(II) ion expected to oxidize zinc metal (ZnII) ion (Zn2+[aq]); it is also predicted to oxidize hydrogen gas (H2[g]) to hydrogen ion (H+) and sodium metal (NaNa+). Similarly, fluorine gas (F2[g]), the strongest oxidant listed in the Table (Pauling electronegativities of selected elements), is predicted to oxidize spontaneously the products of all the other half reactions in the table. In contrast, the strongest reducing agent is solid sodium metal (NaOxidation–reduction equilibria
      In practice many chemical (chemical equilibrium) reactions can be carried out in either direction, depending on the conditions. The spontaneous direction predicted for a particular redox reaction by half-cell potentials is appropriate to a standard set of reaction conditions. Specifically, the temperature is assumed to be 25° C with reagents at specified concentrations. Gases are present at one atmosphere pressure and solutes at one mole per litre (one molecular weight in grams dissolved in one litre of solution) concentration (1M). Solids are assumed to be in contact with the reaction solution in their normal stable forms, and water is always taken to be present as the solvent. Many practical problems can be solved directly with standard reduction potentials.

      The usefulness of reduction potentials is greatly extended, however, by a thermodynamic relationship known as the Nernst equation, which makes it possible to calculate changes in half-cell potentials that will be produced by deviations from standard concentration conditions. In the reaction between zinc metal and copper(II) ion, standard conditions for zinc and copper metal require simply that both solids be present in contact with the solution; the E° values are not affected by either the total or proportionate amounts of the two metals. The calculation that the overall reaction is spontaneous by +1.10 volts is based on standard one mole per litre (1M) concentrations for aqueous zinc(II) ion (Zn2+[aq]) and aqueous copper(II) ion (Cu2+[aq]). Using the Nernst equation it is found that E° for the overall reaction will be +1.10 volts as long as both ions are present in equal concentrations, regardless of the concentration level.

      On the other hand, if the ratio of the zinc(II) to copper(II) ion concentrations is increased, the reduction potential (E°) falls until, at a very high preponderance of zinc ion, E° becomes 0 volt. At this point, there is no net tendency for the reaction to proceed spontaneously in either direction. If the zinc(II) to copper(II) ion ratio is increased further, the direction of spontaneity reverses, and zinc ion spontaneously oxidizes copper metal. In practice, such high zinc(II) to copper(II) ion concentration ratios are unattainable, which means that the reaction can only be carried out spontaneously with copper(II) ion oxidizing zinc metal. Many reactions with E° values smaller than +1.10 volts under standard conditions can be carried out in either direction by adjusting the ratio of product and reactant concentrations. The point at which E° = 0 volt represents a state of chemical equilibrium. When chemical reactions are at equilibrium, the concentrations of the reagents do not change with time, since net reaction is not spontaneous in either direction. Measurements of half-cell potentials combined with Nernst-equation calculations are a powerful technique for determining the concentration conditions that correspond to chemical equilibrium.

Reaction rates (reaction rate)
      There are practical limitations on predictions of the direction of spontaneity for a chemical reaction, the most important arising from the problem of reaction rates. An analogy can be made with the simple physical system of a block on a sloping plane. Because of the favourable energy change, the block tends spontaneously to slide down, rather than up, the slope, and, at mechanical equilibrium, it will be at the bottom of the slope, since that is the position of lowest gravitational energy. How rapidly the block slides down is a more complex question, since it depends on the amount and kind of friction present. The direction of spontaneity for a chemical reaction is analogous to the downhill direction for a sliding block, and chemical equilibrium is analogous to the position at the bottom of the slope; the rate at which equilibrium is approached depends on the efficiency of the available reaction processes. Between zinc metal and aqueous copper(II) ion, the reaction proceeds without observable delay, but various other spontaneous redox processes proceed at imperceptibly slow rates under ordinary conditions.

Biological processes
      A particularly significant illustration of the role of mechanisms in determining the rates of redox reactions concerns respiration, the central energy-producing process of life. Foodstuffs that are oxidized by molecular oxygen during respiration are quite unreactive with oxygen before ingestion. Such high-energy foods as grains and sugar can resist the atmosphere indefinitely but are rapidly converted to carbon dioxide and water through combination with oxygen during respiratory metabolism. The situation is exemplified by the behaviour of glucose at ambient temperatures.

      The significance of the different rate behaviour of high-energy foods inside and outside the cell has been dramatized by Albert Szent-Györgyi (Szent-Györgyi, Albert), a Hungarian biochemist resident in the United States, a pioneering researcher into the chemical mechanism of respiration:

You remember the exciting story of the grave of the Egyptian emperor. At its opening the breakfast of the emperor was found unburned though it had been exposed to the action of oxygen during several thousand years at a temperature that was not very different from 37° C [98.6° F]. Had the king risen and consumed his breakfast, as he had anticipated doing, the food would have been oxidized in no time, that is to say the cells of the emperor would have made reactions take place that would not run spontaneously (from Albert V. Szent-Györgyi, On Oxidation, Fermentation, Vitamins, Health and Disease; the Williams and Wilkins Company, 1939).

      Living systems are able to use respiratory oxidation as an energy source only because the same reactions are slow outside the cell. In return for providing an efficient mechanism for the oxidation of foods, the cell gains control over the disposition of the liberated chemical energy.

      Examples such as the chemistry of respiration make clear the importance of determining the rates and mechanisms of redox reactions. Often questions are difficult to answer even in regard to relatively simple reactions. It has been pointed out that many redox processes can be categorized as oxygen-atom-, hydrogen-atom-, or electron-transfer processes. These categories describe the net changes that are involved but provide no insight into the mechanisms of the reactions.

Mechanisms of redox reactions
      Some of the problems associated with formulating descriptions of the mechanisms are illustrated by the reaction between two metal ions that undergo complementary, one-unit changes in oxidation state:

      There are many different metal ions, designated with the letters M and N, which participate in redox reactions with this basic stoichiometry. To define the possible mechanisms with any precision, it is necessary to identify the groups that are directly bonded to the metals, as well as to specify which particular metals take part in the reactions. One relatively simple example is the chromium(II)–iron(III) reaction. Both chromium and iron aquo ions are surrounded by six water molecules in both the +2 and +3 oxidation states.

      The reaction appears from its stoichiometry to entail simple electron transfer from Cr(II) to Fe(III). There are, however, other possibilities. Their variety is most evident if it is assumed that the oxidizing agent, hexaaquoiron(III) (Fe[H2O]63+), dissociates a hydrogen ion before reacting with hexaaquochromium(II) (Cr[H2O]62+). The hexaaquoiron(III) ion is then converted to a hydroxopentaaquoiron(III) ion, known to be an important reactant in the Fe(III)–Cr(II) reaction. Its formation requires no change in the oxidation state of iron and should be regarded as a simple acid–base reaction. With hydroxopentaaquoiron(III) ion, (H2O)5FeOH2+, as the oxidizing agent, simple electron transfer does provide one plausible means for the reaction. All the steps in the mechanism must add up to give the overall reaction. Intermediate species such as hydroxopentaaquoiron(III) must cancel out in the addition process by occurring as products in one step and reactants in another:

      An alternative to simple electron transfer is the transfer of a hydrogen atom from a water molecule bound to chromium(II) to the hydroxide ion (OH) bound to iron(III). A third possibility is the transfer of a neutral OH group (called a hydroxyl radical) from iron(III) to chromium(II). Oxygen and hydrogen occur as oxygen(−II) and hydrogen(I) in both reactants and products. These oxidation states correspond to those present in the hydroxide ion OH. Therefore, when an OH group transfers, it must leave an electron behind, reducing the iron, and then reaccept an electron from chromium, oxidizing it. Since both the aqueous iron and chromium ions have a strong tendency to remain surrounded by six oxygen molecules, the most likely method of OH transfer would be the sharing of the transferred oxygen between the two metals at an intermediate stage of the reaction. But iron(III) is not a strong enough oxidizing agent to produce a significant concentration of neutral OH groups from OH ions. A more acceptable description of the OH transfer mechanism, therefore, is to regard OH as a bridging group through which electron transfer is accomplished as the two metal ion groups share a bonded oxygen atom.

      Experimentally, it is difficult to distinguish between the alternative mechanisms posed above. Classification schemes for reaction mechanisms are of little value unless they offer alternatives that can be experimentally verified. For that reason the most successful system for classifying the mechanisms of redox reactions between metal ions has been the inner sphere–outer sphere dichotomy. Outer-sphere reactions are those that take place without breaking any bonds between a metal and a group such as water or hydroxide ion bound to it. Both the simple electron- and hydrogen-atom-transfer mechanisms are outer-sphere processes. If, on the other hand, the oxidation-state changes take place within an intermediate product in which the two metals are directly bonded to a common bridging group (such as shown in the equation above), the mechanism is inner-sphere. Formation of the intermediate requires loss of a bound water molecule by one of the metals. The inner sphere–outer sphere distinction can be tested experimentally in some redox systems.

Maynard V. Olson

      Combustion is a chemical reaction between substances, usually including oxygen and usually accompanied by the generation of heat and light in the form of flame. The rate or speed at which the reactants combine is high, in part because of the nature of the chemical reaction itself and in part because more energy is generated than can escape into the surrounding medium, with the result that the temperature of the reactants is raised to accelerate the reaction even more. A familiar example is a lighted match. When a match is struck, friction heats (heat) the head to a temperature at which the chemicals react and generate more heat than can escape into the air, and they burn with a flame. If a wind blows away the heat or the chemicals are moist and friction does not raise the temperature sufficiently, the match goes out; properly ignited, the heat from the flame raises the temperature of a nearby layer of the matchstick and of oxygen in the air adjacent to it, and the wood and oxygen react in a combustion reaction. When equilibrium between the total heat energies of the reactants and the total heat energies of the products (including the actual heat and light emitted) is reached, combustion stops. Flames have a definable composition and a complex structure; they are said to be multiform and are capable of existing at quite low temperatures, as well as at extremely high temperatures. The emission of light in the flame results from the presence of excited particles and, usually, of charged atoms and molecules and of electrons.

General principles
      Combustion encompasses a great variety of phenomena with wide application in industry, the sciences, professions, and the home, and the application is based on knowledge of physics, chemistry, and mechanics; their interrelationship becomes particularly evident in treating flame propagation.

      In general terms, combustion is one of the most important of chemical reactions and may be considered a culminating step in the oxidation of certain kinds of substances. Though oxidation was once considered to be simply the combination of oxygen with any compound or element, the meaning of the word has been expanded to include any reaction (oxidation–reduction reaction) in which atoms lose electrons, thereby becoming oxidized. As has been pointed out, in any oxidation process the oxidizer takes electrons from the oxidizable substance, thereby itself becoming reduced (gaining electrons). Any substance at all can be an oxidizing agent. But these definitions, clear enough when applied to atomic structure to explain chemical reactions, are not as clearly applicable to combustion, which remains, generally speaking, a type of chemical reaction involving oxygen as the oxidizing agent but complicated by the fact that the process includes other kinds of reactions as well, and by the fact that it proceeds at an unusually fast pace. Furthermore, most flames have a section in their structure in which, instead of oxidations, reduction reactions occur. Nevertheless, the main event in combustion is the combining of combustible material with oxygen.

History
      Combustion, fire, and flame have been observed and speculated about from earliest times. Every civilization had its own explanation for them. The Greeks (ancient Greek civilization) interpreted combustion in terms of philosophical doctrines, one of which was that a certain “inflammable principle” was contained in all combustible bodies and this principle escaped when the body was burned to react with air. A generalization of the concept was provided by the phlogiston theory, formulated in the 17th century (see above). Treated at first as a purely metaphysical quality, phlogiston was conceived later as a material substance having weight, and, sometimes, negative weight. The inadequacy of the phlogiston theory became apparent only in the late 18th century when it proved unable to explain a host of new facts about combustion that were being observed for the first time as the result of increasing accuracy in laboratory experiments.

      The English natural philosopher Sir Francis Bacon (Bacon, Francis, Viscount Saint Alban (or Albans), Baron of Verulam) observed in 1620 that a candle flame has a structure at about the same time that Robert Fludd (Fludd, Robert), an English mystic, described an experiment on combustion in a closed container in which he determined that an amount of air was used up thereby. A German, Otto von Guericke (Guericke, Otto von), using an air pump he had invented in 1650, demonstrated that a candle would not burn in a container from which the air had been pumped. Robert Hooke (Hooke, Robert), an English scientist, in 1665 suggested that air had an active component that, upon heating, combined with combustible substances, giving rise to flame. Another idea ascribed the high temperature of flame to the fast motion of active air particles, and it was learned that sulfur mixed with nitre can burn in the absence of air (nitre is a compound of oxygen which releases oxygen to the sulfur).

      The first approximation of the true nature of combustion was posited by Lavoisier (Lavoisier, Antoine-Laurent), who discovered in 1772 that the products of burned sulfur or phosphorus, in effect their ashes, outweighed the initial substances, and postulated that the increased weight was due to their having combined with air. Interestingly, it was already known that metals transformed by heat to metallic ash weighed less than the metallic ash, but the theory was that in certain cases phlogiston in metals had a negative weight, and upon escaping during combustion, left the ash of the metal heavier than it had been with the phlogiston in it. Later, Lavoisier concluded that the “fixed” air that had combined with the sulfur was identical to a gas obtained by Priestley (Priestley, Joseph) on heating the metallic ash of mercury—that is, the “ashes” obtained when mercury was burned could be made to release the gas with which the metal had combined. This gas was also identical to that described by the Swedish chemist, Carl Wilhelm Scheele, as an active fraction of air that sustained combustion. Lavoisier called the gas “oxygen.”

      Lavoisier's theory that combustion was a reaction between the burning substance and the gas oxygen, present only to a limited extent in the atmosphere, was based on scientific principles, the most important of which was the law of the conservation (mass, conservation of) of matter (after Einstein's relativity theory, of matter and energy): the total amount of matter in the universe is constant. Even ancient philosophers had guessed this law and it was substantiated in the 17th century. Lavoisier also clarified the concept of element into a modern generalization, that it was a substance that could not be broken down, and this, too, supported his theory. Soon after, studies of gases by John Dalton (Dalton, John), and the first table of atomic weights that Dalton compiled, as well as many new gases discovered by other scientists, were important in supporting not only Lavoisier's theory of combustion but his whole new system of chemistry based on accurate measurement. The discoveries of nitrogen and hydrogen in the latter half of the 18th century, added to the earlier discoveries of carbon dioxide and carbon monoxide, and the discovery that the composition of air is remarkably constant though it is a mixture, all supported Lavoisier's theory. The proper explanation of combustion, perhaps the oldest recognized chemical reaction, is usually said to have been a keystone in the development of modern science.

      From 1815 to 1819, Sir Humphry Davy (Davy, Sir Humphry, Baronet) experimented on combustion, including measurements of flame temperatures, investigations of the effect on flames of rarefied gases, and dilution with various gases; he also discovered catalytic combustion—the oxidation of combustibles on a catalytic surface accompanied by the release of heat but without flame.

      Despite these discoveries, the materialistic theory of combustion lacked a clear concept of energy and, therefore, of the critical role that energy considerations play in an accurate explanation of combustion. It was Sir Benjamin Thompson's (Thompson, Sir Benjamin, Count Von Rumford) experiments with heat in 1798 that revealed evidence for the concept of heat as a movement of particles. Development of a kinetic (kinetic energy) theory of gases, based on the premise that heat results from the motion of molecules and atoms, of thermodynamics, and of thermochemistry, all in the 19th century, finally elucidated the energy aspects of combustion.

      Investigation of burning velocities, experiments on the order of events in the combustion of gas mixtures, and study of the breaking down of gas molecules by heat (thermal dissociation), in the last half of the 19th century, played a vital part in the refinement of theories concerning combustion mechanism. Studies of light emitted by flames led to its analysis in the spectroscope, a device that separates a mixture of light waves into the component waves, and to spectral analysis generally, including theories of atomic and molecular spectra, which, in turn, contributed to an understanding of the nature of flames. The Bunsen burner was also of importance in the study of flame structure. Progress in industry was a powerful stimulus in the search for clarification of flame phenomena. Explosion hazards in coal mines had drawn attention to flame propagation as far back as 1815, when Davy invented his safety lamp. In 1881 detonation was discovered and led at the beginning of the 20th century to a detonation theory, based on the assumption that a gas behaved as a fluid under certain conditions. Chemical kinetics after the 1930s became an indispensable part of flame propagation theory.

Physical and chemical aspects of combustion
The chemical reactions
      Combustion, with rare exceptions, is a complex chemical process involving many steps that depend on the properties of the combustible substance. It is initiated by external factors such as heat, light, and sparks. The reaction sets in as the mixture of combustibles attains the ignition temperature, and several aspects of this step can be defined.

 First, a relationship exists between the ignition temperature and the pressure of the mixture under specific conditions. Figure 1—> shows the relationship for a mixture of hydrogen and oxygen. Only one temperature corresponds to a given pressure, whereas one or three pressures, called the explosion limits, may correspond to one temperature. The mechanism of the reaction determines the explosion limits: the reaction can proceed only when the steps in the sequence of reactions occur faster than the terminal steps. Thus, for combustion to be initiated with light, or with a spark, the light intensity or the spark energy must exceed certain minimal values.

      The complexity of the combustion reaction mechanism and the rapidly varying temperatures and concentrations in the mixture make it difficult and often impossible to derive an equation that would be useful for predicting combustion phenomena over wide temperature and concentration ranges. Instead, use is made of empirical expressions derived for specific reaction conditions.

      Most reactions terminate when what is called thermal equilibrium has been attained—i.e., when the energy of the reactants equals the energy of the products.

Special combustion reactions
      Reactions of oxygen with hydrogen, with carbon monoxide, and with hydrocarbons are most important from the standpoint of theory and, at the same time, are of great practical value.

      Hydrogen combustion proceeds by complicated branched-chain reactions involving the interaction of hydrogen and oxygen atoms with oxygen and hydrogen molecules, respectively, to produce hydroxyl radicals (radical). The final reaction product is water, formed by the combination of hydroxyl with hydrogen molecules. Reactions that terminate the chain, such as recombination of atoms and fragments of molecules, and the adsorption of active particles on solid surfaces, also play an important part in the mechanism of hydrogen combustion. The knowledge of rate constants for these processes derived empirically makes possible a quantitative description of all combustion characteristics, such as explosion limits, delay of ignition, and burning velocity.

      Combustion of carbon monoxide is mainly restricted to mixtures containing hydrogen or hydrogen compounds. In this case the reaction mechanism differs from that of hydrogen combustion chiefly in that it involves a step of fast interaction between hydroxyl and carbon monoxide.

      Pure mixtures of carbon monoxide and oxygen (or air) can be ignited only with sparks of high energy, or under high pressures and temperatures. The chemical mechanism of their combustion is not yet clear, probably because oxidation of carbon monoxide, a reaction that is part of the combustion of practically all natural fuels, usually occurs in the presence of hydrogen or hydrogen compounds: the breakdown of wood, coal, petroleum, etc., during burning produces carbon monoxide, hydrogen, and compounds of carbon and hydrogen.

      The mechanisms of combustion of hydrocarbons and of other organic compounds are known in general outline only. Many elementary steps of hydrocarbon combustion involving oxygen and hydrogen atoms, hydroxyl, and organic radicals are similar to those for hydrogen; however, the overall mechanism of hydrocarbon combustion is complicated by the diversity of molecules and radicals involved. Moreover, with oxidation, thermal decomposition, a breakdown of complex organic compounds without oxidation, also takes place.

      Two types of hydrocarbon combustion have been defined: (1) slow combustion at temperatures below 500° C, including cool flames observed at certain pressures, and (2) combustion at higher temperatures, accompanied by hot flames.

      Ignition in two stages is characteristic of higher hydrocarbons, in which a cool-flame stage yielding readily oxidizable products precedes that of the hot flame.

Physical processes
      In addition to chemical reactions, physical processes that transfer mass and energy by diffusion or convection occur in gaseous combustion. In the absence of external forces, the rate of component diffusion depends upon the concentration of the constituents, pressure, and temperature changes, and on diffusion coefficients (a measure of the speed of diffusion). The latter are either measured or calculated in terms of the kinetic theory of gases. The process of diffusion is of great importance in combustion reactions, in flames—that is, in gaseous mixtures, and in solids or liquids. Diffusion heat transfer (by molecular means) follows a law (Fourier law) stating that the heat flux (a measure of the quantity of heat) is proportional to the temperature gradient (the difference in temperature between two limiting temperatures). The coefficient of proportionality, called the thermal conductivity coefficient, is also measured or calculated in terms of the kinetic theory of gases, like the diffusion coefficient.

      Convective transport of mass and energy may be accounted for by buoyant forces, external forces, and turbulent and eddy motions. Convection is in the main responsible for the mixing of gases (e.g., in furnaces).

       sublimation (the direct evaporation of a solid without intermediate liquid stage), evaporation of liquids, and the mechanical destruction of burning samples, all methods of transforming solids and liquids into gases, greatly contribute to their ease of combustion. In general, the combustion of condensed systems—i.e., liquids and solids—is more complex than that of gases, which accounts for the greater development of the gas combustion theory.

      Energy transport in combustion may also occur by the emission of light, mostly in the infrared.

Combustion phenomena and classification
      All flames can be classified either as premixed flames or as flames that burn without premixing.

Premixed flames
      Flame combustion is most prominent with fuels that have been premixed with an oxidant, either oxygen or a compound that provides oxygen, for the reaction. The temperature of flames with this mixture is often several thousand degrees. The chemical reaction in such flames occurs within a narrow zone several microns thick. This combustion zone is usually called the flame front.

      Dilution of the burning mixture with an inert gas, such as helium or nitrogen, lowers the temperature and, consequently, the reaction rate. Great amounts of inert gas extinguish the flame, and the same result is achieved when substances that remove any of the active species are added to the flame. Conditions must be such that the flame is fixed at the burner nozzle or in the combustion chamber; this positioning is required in many practical uses of combustion. Various devices, such as pilot flames and recirculation methods, are designed for this purpose.

      The principal quantitative characteristic of a flame is its normal, or fundamental, burning velocity, which depends on the chemical and thermodynamic properties of the mixture, and on pressure and temperature, under given conditions of heat loss. The burning velocity value ranges from several centimetres even to tens of metres, per second. The dependence of the burning velocity on molecular structure, which is responsible for fuel reactivity, is known for a great many fuel–air mixtures.

      A widely applied thermal theory, one of the first flame propagation theories, implies that combustion proceeds primarily at temperatures close to the maximum the flame can achieve. A set of differential equations developed for thermal conductivity and diffusion is reduced to one equation that yields the burning velocity value. Further development of the theory has been made and computers now make most simplifications unnecessary.

      According to thermal theories, flame propagation is accounted for by heat energy transport from the combustion zone to the unburned mixture, to raise the temperature of the mixture. Diffusion theory assumes that thermodynamic equilibrium sets in behind the flame front at a maximum temperature, and that radicals produced in this zone diffuse into the unburned mixture and ignite it. Both heat transport and diffusion of active particles must be considered essential for ignition.

Diffusion flames
      Diffusion flames, smoothly flowing (laminar) or turbulent, belong to the class of flames whose ingredients are not mixed prior to entering the burning zone. Molecular or turbulent diffusion is responsible for the mixing of the gases in such flames. The distribution of the combustible material and of oxygen over various flame cross sections is regular in laminar flames but becomes much more complicated in turbulent gas flows. The transition from laminar to turbulent flames occurs at a certain point of the flow and depends upon the actual conditions.

      Industrial flames, including those activated in furnaces, belong to the turbulent diffusion type. The theory of diffusion flames, however, is less advanced than that of premixed flames and, since gas intermixing is mainly responsible for the structure and the features of diffusion flames, the theory operates more in terms of physics and of thermodynamics than chemical reactions. Flames such as those of candles, of liquid droplets, and of many propellants and condensed fuels are diffusion flames.

Oxidizing and reducing flame
 When a premixed flame burns in open air with an excess of fuel, there appears in addition to the flame zone a zone of diffusion flame; this is accounted for by the diffusion of atmospheric oxygen, as, for example, in the Bunsen flame produced by a burner to which the air intake can be regulated, thereby altering the flow from an intensely hot one—in which most of the fuel gases are oxidized to carbon dioxide and water—to a relatively low temperature flux, in which most of the fuel gas is only partially oxidized. Such flames consist of an inner and outer cone—two zones in which different chemical reactions take place, the reducing and oxidizing zones, respectively. The oxidizing nature of the outer cone is due to excess oxygen (see Figure 2—>).

Explosions
      The transition from combustion to explosion is caused by an acceleration of the reaction, induced either by a rise in temperature or by increasing lengths of the reaction chain. The first is called thermal explosion, and the second, chain explosion.

Thermal explosions
      Thermal explosion theory is based on the idea that progressive heating raises the rate at which heat is released by the reaction until it exceeds the rate of heat loss from the area. At a given composition of the mixture and a given pressure, explosion will occur at a specific ignition temperature that can be determined from the calculations of heat loss and heat gain.

      The thermal explosion theory accounts for temperature rise and fuel consumption during the induction period. At sufficiently high rates of consumption the explosion will not occur.

Chain-branch reactions
      It follows from the theory of branched-chain reactions that there is a limit to ignition, or to explosion, without a rise of temperature. In this case, what is called a chain explosion will occur when the probabilities of chain branching and of termination are equal. Usually, however, explosions are of a chain-thermal nature (i.e., both heat accumulation and chain auto-acceleration contribute to explosion).

      The progressive acceleration of reaction accounted for by growth of the flame front area and by transition from laminar to turbulent flow gives rise to a shock wave. The increase in temperature due to compression in the shock wave results in self-ignition of the mixture and detonation sets in. The shock wave-combustion zone complex forms the detonation wave. Detonation differs from normal combustion in its ignition mechanism and in the supersonic velocity of 2–5 kilometres per second for gases and 8–9 kilometres per second for solid and liquid explosives.

      Detonation is impossible when the energy loss from the reaction zone exceeds a certain limit. The detonation limits observed for high explosives are also eventually dependent on factors responsible for the chemical reaction rate: the charge and diameter of the grain, etc.

Special aspects
      Emission of light is a characteristic feature of combustion. Infrared, visible, and ultraviolet bands of molecules, and atomic lines are usually observed in flame spectra. Moreover, continuous spectra from incandescent particles, or from recombination of atoms, radicals, and ions are frequently observed. The sources of flame radiation are the thermal energy of gas ( thermoluminescence) and the chemical energy released in exothermic elementary reactions ( chemiluminescence). In a Bunsen burner fed with a sufficient amount of air, up to 20 percent of the reaction heat is released as infrared energy and less than one percent as visible and ultraviolet radiation, the infrared being mostly thermoluminescence. Radiation from the inner cone of the Bunsen flame in the visible and ultraviolet regions represents chemiluminescence.

      Many flames contain electrons and positive and negative ions, a fact evident from the electrical conductivity of flames, and also from the deviation of the flame cone in an electric field (the charges are attracted or repelled, distorting the flame), a phenomenon usually interpreted as a mechanical effect called electric wind. The resulting change of the flame shape can affect the burning velocity. ionization, like the emission of light, can be the result of equilibrium processes, when it is called thermal ionization, or it can be related to chemical processes and called chemical ionization. Thermal ionization may be expected in very hot flames containing alkali metals or alkaline-earth metals (for example, sodium and calcium) as impurities because of their low-ionization potentials. The high concentrations of ions and electrons in the flames of organic species are undoubtedly due to chemical, rather than to thermal ionization. This is seen in the fact that electroconductivity in the inner cone of the Bunsen flame is several times higher than that of the outer cone. The reactions of ions and electrons may yield atoms and radicals and in this way affect the burning velocity.

      Formation of soot is a feature of all hydrocarbon flames. It makes the flame luminous and nontransparent. The mechanism of soot formation is accounted for by simultaneous polymerization, a process whereby molecules or molecular fragments are combined into extremely large groupings, and dehydrogenation, a process that eliminates hydrogen from molecules.

Applications
      The uses of combustion and flame phenomena can be categorized under five general heads.

In heating devices
      Heating devices for vapour production (steam, etc.), in metallurgy, and in industry generally, utilize the combustion of gases, wood, coal, and liquid fuels. Control of the combustion process to obtain optimal efficiency is ensured by proper ratio and distribution of the fuel and the oxidant in the furnace, stove, kiln, etc., choice of conditions for heat transport from the combustion products to the heated bodies, and by appropriate aerodynamics of gas flows in the furnace. Radiation contributes to a certain extent to heat exchange. Combustion in furnaces being a very complicated process, only general ideas can be given by the combustion theory, so that the optimal conditions and the furnace design are usually decided empirically.

In explosives (explosive)
      The combustion and detonation of explosives are widely used in all sorts of work with mechanical action or explosion as the eventual goals. Practical applications of explosives are based on the theory of their combustion and detonation. The combustion of condensed explosives occurs mostly in the gas phase due to their evaporation, sublimation, or decomposition and can be treated in terms of the theory of gas combustion, which provides for the burning velocity, its dependence on temperature, pressure, and on the parameters determining the combustion regime and the nature of explosives, etc. Control of combustion and detonation in their practical applications is made possible by use of the theory, together with the results of experimental investigations on combustion and detonation.

In internal combustion (internal-combustion engine) engines
      These comprise various engines, gas turbines, turbojets, and ramjets. The Otto engine operates with a mixture compressed in a cylinder by a piston. Shortly before the piston reaches the top the mixture is ignited with a spark, and the flame propagates at a normal velocity into the unburned mixture, increasing the pressure and moving the piston. There is a maximum of compression for any mixture composition and any engine design. Detonation occurs beyond this maximum because of the appearance of centres where self-ignition takes place before the flame front. Loss of power is one result of detonation; compounds hindering self-ignition are used to prevent it.

      The diesel engine operates with a fuel spray injected into the engine cylinder as liquid droplets that mix with air by turbulent diffusion and that evaporate. At normal operations of the engine the temperature of compressed air is sufficiently high for self-ignition of the fuel.

      In gas turbines (turbine), compressed air enters the combustion chamber where it mixes with the fuel. The expanding combustion products impart their energy to the turbine blades.

      Two kinds of jet engines (jet engine) are used in aircraft: the turbojet and the ramjet. The turbine of a turbojet engine is used to operate the compressor. Thrust comes from the repulsion of products flowing out of a nozzle. In a ramjet engine, air is compressed and slowed down in the diffuser without any compressor or turbine device.

In rocket propulsion
      The products of combustion of gaseous, liquid, or solid propellants in rockets are ejected from the combustion chamber through the (de Laval) nozzle at a high velocity. Knowledge of the kinetics of chemical processes in the nozzle is essential to determine the thrust required. The thrust decreases with the increasing mean molecular weight of the combustion products. Mixtures of low molecular weight and high heat of combustion, therefore, are used for rockets.

In chemical reactions
      Flames are used in various ways to produce chemical reactions. The bead test in analytical chemistry is one example. The reducing power of a flame that has insufficient oxygen is utilized in limited ways. The soot produced by some flames is commercially useful. The manufacture of coke and charcoal is dependent on the judicious control of combustion and flame.

Victor Nikolaevich Kondratiev

Additional Reading
Eduard Farber, Oxygen and Oxidation Theories and Techniques in the 19th Century and the First Part of the 20th (1967), presents a short history of oxidation concepts. F. Albert Cotton and Geoffrey Wilkinson, Advanced Inorganic Chemistry, 5th ed. (1988), a comprehensive reference work, contains examples of inorganic redox reactions. Kenneth L. Rinehart, Jr., Oxidation and Reduction of Organic Compounds (1973), provides information on organic reactions. Wendell M. Latimer, The Oxidation States of the Elements and Their Potentials in Aqueous Solutions, 2nd ed. (1952), surveys the redox behaviour of the elements with an emphasis on half-reaction potentials. W. Mansfield Clark, Oxidation-Reduction Potentials of Organic Systems (1960, reissued 1972), emphasizes biologically important reactions. Linus Pauling, The Nature of the Chemical Bond and the Structure of Molecules and Crystals, 3rd ed. (1960, reissued 1989), contains a detailed treatment of electronegativities. Ross Stewart, Oxidation Mechanisms (1964), is a concise monograph on organic oxidation-reduction mechanisms; while inorganic mechanisms are treated by Graham Lappin, Redox Mechanisms in Inorganic Chemistry (1994). Oxidation-reduction reactions brought about by absorption of light are discussed in Lennart Eberson, Electron Transfer Reactions in Organic Chemistry (1987); Marye Anne Fox and Michel Chanon (eds.), Photoinduced Electron Transfer, 4 vol. (1988); and in two parts of the Topics in Current Chemistry series: Electron Transfer (irregular); and Photoinduced Electron Transfer (irregular). Eugene Rabinowitch and Govindjee, Photosynthesis (1969), includes a good overview of the global redox cycle of respiration and photosynthesis. Other studies of photosynthesis are Govindjee (ed.), Photosynthesis, 2 vol. (1982); and Christine H. Foyer, Photosynthesis (1984). G. Scott (ed.), Atmospheric Oxidation and Antioxidants, 3 vol. (1993), presents atmospheric examples.Maynard V. Olson Ed.Joshua C. Gregory, Combustion from Heracleitos to Lavoisier (1934), recounts the history of combustion science to the 18th century; while William A. Bone and Donald T.A. Townend, Flame and Combustion in Gases (1927), covers the period from the late 17th century to the early 20th century. Wilhelm Jost, Explosion and Combustion Processes in Gases (1946; originally published in German, 1939), reviews the main theoretical and experimental research on various problems of combustion and explosions. R.M. Fristrom and A.A. Westenberg, Flame Structure (1965), contains a general review and critical treatment of experimental and theoretical data on flame structure. J.A. Barnard and John N. Bradley, Flame and Combustion, 2nd ed. (1985), is also useful. More advanced and detailed monographs include Bernard Lewis and Guenther Von Elbe, Combustion, Flames, and Explosions of Gases, 3rd ed. (1987); N. Semenoff, Chemical Kinetics and Chain Reactions, trans. from Russian (1935), and Some Problems of Chemical Kinetics and Reactivity, 2 vol. (1958–59; originally published in Russian, 1954); and A.G. Gaydon and H.G. Wolfhard, Flames: Their Structure, Radiation, and Temperature, 4th ed. rev. (1979).Victor Nikolaevich Kondratiev Ed.

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Universalium. 2010.

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