Redox is a chemical reaction in which the oxidation states of atoms are changed. Any such reaction involves both a reduction process and a complementary oxidation process, two key concepts involved with electron transfer processes. Redox reactions include all chemical reactions; the chemical species from which the electron is stripped is said to have been oxidized, while the chemical species to which the electron is added is said to have been reduced. It can be explained in simple terms: Oxidation is the loss of electrons or an increase in oxidation state by a molecule, atom, or ion. Reduction is a decrease in oxidation state by a molecule, atom, or ion; as an example, during the combustion of wood, oxygen from the air is reduced, gaining electrons from carbon, oxidized. Although oxidation reactions are associated with the formation of oxides from oxygen molecules, oxygen is not included in such reactions, as other chemical species can serve the same function; the reaction can occur slowly, as with the formation of rust, or more in the case of fire.
There are simple redox processes, such as the oxidation of carbon to yield carbon dioxide or the reduction of carbon by hydrogen to yield methane, more complex processes such as the oxidation of glucose in the human body. "Redox" is a portmanteau of the words "reduction" and "oxidation". The word oxidation implied reaction with oxygen to form an oxide, since dioxygen was the first recognized oxidizing agent; the term was expanded to encompass oxygen-like substances that accomplished parallel chemical reactions. The meaning was generalized to include all processes involving loss of electrons; the word reduction 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 showed. Scientists realized that the metal atom gains electrons in this process; the meaning of reduction became generalized to include all processes involving a gain of electrons. Though "reduction" seems counter-intuitive when speaking of the gain of electrons, it might help to think of reduction as the loss of oxygen, its historical meaning.
Since electrons are negatively charged, it is helpful to think of this as reduction in electrical charge. The electrochemist John Bockris has used the words electronation and deelectronation to describe reduction and oxidation processes when they occur at electrodes; these words are analogous to protonation and deprotonation, but they have not been adopted by chemists worldwide. The term "hydrogenation" could be used instead of reduction, since hydrogen is the reducing agent in a large number of reactions in organic chemistry and biochemistry. But, unlike oxidation, generalized beyond its root element, hydrogenation has maintained its specific connection to reactions that add hydrogen to another substance; the word "redox" was first used in 1928. The processes of oxidation and reduction occur and cannot happen independently of one another, similar to the acid–base reaction; the oxidation alone and the reduction alone are each called a half-reaction, because two half-reactions always occur together to form a whole reaction.
When writing half-reactions, the gained or lost electrons are included explicitly in order that the half-reaction be balanced with respect to electric charge. Though sufficient for many purposes, these general descriptions are not correct. Although oxidation and reduction properly refer to a change in oxidation state — the actual transfer of electrons may never occur; the oxidation state of an atom is the fictitious charge that an atom would have if all bonds between atoms of different elements were 100% ionic. Thus, oxidation is best defined as an increase in oxidation state, reduction as a decrease in oxidation state. In practice, the transfer of electrons will always cause a change in oxidation state, but there are many reactions that are classed as "redox" though no electron transfer occurs. In redox processes, the reductant transfers electrons to the oxidant. Thus, in the reaction, the reductant or reducing agent loses electrons and is oxidized, the oxidant or oxidizing agent gains electrons and is reduced.
The pair of an oxidizing and reducing agent that are involved in a particular reaction is called a redox pair. A redox couple is a reducing species and its corresponding oxidizing form, e.g. Fe2+/Fe3+ Substances that have the ability to oxidize other substances are said to be oxidative or oxidizing and are known as oxidizing agents, oxidants, or oxidizers; that is, the oxidant removes electrons from another substance, is thus itself reduced. And, because it "accepts" electrons, the oxidizing agent is called an electron acceptor. Oxygen is the quintessential oxidizer. Oxidants are chemical substances with elements in high oxidation states, or else electronegative elements that can gain extra electrons by oxidizing another substance. Substances that have the ability to reduce other substances are said to be reductive or reducing and are known as
An epoxide is a cyclic ether with a three-atom ring. This ring approximates an equilateral triangle, which makes it strained, hence reactive, more so than other ethers, they are produced on a large scale for many applications. In general, low molecular weight epoxides are colourless and nonpolar, volatile. A compound containing the epoxide functional group can be called an epoxy, epoxide and ethoxyline. Simple epoxides are referred to as oxides. Thus, the epoxide of ethylene is ethylene oxide. Many compounds have trivial names, ethylene oxide is called "oxirane"; some names emphasize the presence of the epoxide functional group, as in the compound 1,2-epoxyheptane, which can be called 1,2-heptene oxide. A polymer formed from epoxide precursors is called an epoxy, but such materials do not contain epoxide groups; the dominant epoxides industrially are ethylene oxide and propylene oxide, which are produced on the scales of 15 and 3 million tonnes/year. The epoxidation of ethylene involves its reaction of oxygen according to the following stoichiometry: 7 H2C=CH2 + 6 O2 → 6 C2H4O + 2 CO2 + 2 H2OThe direct reaction of oxygen with alkenes is useful only for this epoxide.
Modified heterogeneous silver catalysts are employed. Other alkenes fail to react usefully propylene, though TS-1 supported Au catalysts can perform propylene epoxidation selectively. Aside from ethylene oxide, most epoxides are generated by treating alkenes with peroxide-containing reagents, which donate a single oxygen atom. Safety considerations weigh on these reactions because organic peroxides are prone to spontaneous decomposition or combustion. Metal complexes are useful catalysts for epoxidations involving hydrogen peroxide and alkyl hydroperoxides. Peroxycarboxylic acids, which are more electrophilic, convert alkenes to epoxides without the intervention of metal catalysts. In specialized applications, other peroxide-containing reagents are employed, such as dimethyldioxirane. Depending on the mechanism of the reaction and the geometry of the alkene starting material, cis and/or trans epoxide diastereomers may be formed. In addition, if there are other stereocenters present in the starting material, they can influence the stereochemistry of the epoxidation.
Metal-catalyzed epoxidations were first explored using tert-butyl hydroperoxide. Association of TBHP with the metal generates the active metal peroxy complex containing the MOOR group, which transfers an O center to the alkene. Organic peroxides are used for the production of propylene oxide from propylene. Catalysts are required as well. Both t-butyl hydroperoxide and ethylbenzene hydroperoxide can be used as oxygen sources. More for laboratory operations, the Prilezhaev reaction is employed; this approach involves the oxidation of the alkene with a peroxyacid such as m-CPBA. Illustrative is the epoxidation of styrene with perbenzoic acid to styrene oxide: The reaction proceeds via what is known as the "Butterfly Mechanism"; the peroxide is viewed as an electrophile, the alkene a nucleophile. The reaction is considered to be concerted; the butterfly mechanism allows ideal positioning of the O-O sigma star orbital for C-C Pi electrons to attack. Because two bonds are broken and formed to the epoxide oxygen, this is formally an example of a coarctate transition state.
Hydroperoxides are employed in catalytic enantioselective epoxidations, such as the Sharpless epoxidation and the Jacobsen epoxidation. Together with the Shi epoxidation, these reactions are useful for the enantioselective synthesis of chiral epoxides. Oxaziridine reagents may be used to generate epoxides from alkenes. Arene oxides are intermediates in the oxidation of arenes by cytochrome P450. For prochiral arenes, the epoxides are obtained in high enantioselectivity. Chiral epoxides can be derived enantioselectively from prochiral alkenes. Many metal complexes give active catalysts, but the most important involve titanium and molybdenum; the Sharpless epoxidation reaction is one of the premier enantioselective chemical reactions. It is used to prepare 2,3-epoxyalcohols from secondary allylic alcohols; this method involves dehydrohalogenation. It is a variant of the Williamson ether synthesis. In this case, an alkoxide ion intramolecularly displaces chloride; the precursor compounds are called halohydrins.
Starting with propylene chlorohydrin, most of the world's supply of propylene oxide arises via this route. An intramolecular epoxide formation reaction is one of the key steps in the Darzens reaction. In the Johnson–Corey–Chaykovsky reaction epoxides are generated from carbonyl groups and sulfonium ylides. In this reaction, a sulfonium is the leaving group instead of chloride. Electron-deficient olefins, such as enones and acryl derivatives can be epoxidized using nucleophilic oxygen compounds such as peroxides; the reaction is a two-step mechanism. First the oxygen performs a nucleophilic conjugate addition to give a stabilized carbanion; this carbanion attacks the same oxygen atom, displacing a leaving group from it, to close the epoxide ring. Epoxides are uncommon in nature, they arise via oxygenation of alkenes by the action of cytochrome P450. Ring-opening reactions dominate the reactivity of epoxides. Alcohols, amines and many other reagents add to epoxides; this reaction is the basi
Enthalpy, a property of a thermodynamic system, is equal to the system's internal energy plus the product of its pressure and volume. In a system enclosed so as to prevent matter transfer, for processes at constant pressure, the heat absorbed or released equals the change in enthalpy; the unit of measurement for enthalpy in the International System of Units is the joule. Other historical conventional units still in use include the calorie. Enthalpy comprises a system's internal energy, the energy required to create the system, plus the amount of work required to make room for it by displacing its environment and establishing its volume and pressure. Enthalpy is defined as a state function that depends only on the prevailing equilibrium state identified by the system's internal energy and volume, it is an extensive quantity. Enthalpy is the preferred expression of system energy changes in many chemical and physical measurements at constant pressure, because it simplifies the description of energy transfer.
In a system enclosed so as to prevent matter transfer, at constant pressure, the enthalpy change equals the energy transferred from the environment through heat transfer or work other than expansion work. The total enthalpy, H, of a system cannot be measured directly; the same situation exists in classical mechanics: only a change or difference in energy carries physical meaning. Enthalpy itself is a thermodynamic potential, so in order to measure the enthalpy of a system, we must refer to a defined reference point; the ΔH is a positive change in endothermic reactions, negative in heat-releasing exothermic processes. For processes under constant pressure, ΔH is equal to the change in the internal energy of the system, plus the pressure-volume work p ΔV done by the system on its surroundings; this means that the change in enthalpy under such conditions is the heat absorbed or released by the system through a chemical reaction or by external heat transfer. Enthalpies for chemical substances at constant pressure refer to standard state: most 1 bar pressure.
Standard state does not speaking, specify a temperature, but expressions for enthalpy reference the standard heat of formation at 25 °C. Enthalpy of ideal gases and incompressible solids and liquids does not depend on pressure, unlike entropy and Gibbs energy. Real materials at common temperatures and pressures closely approximate this behavior, which simplifies enthalpy calculation and use in practical designs and analyses; the word enthalpy was coined late, in the early 20th century, in analogy with the 19th-century terms energy and entropy. Where energy uses the root of the Greek word ἔργον "work" to express the idea of "work-content" and where entropy uses the Greek word τροπή "transformation" to express the idea of "transformation-content", so by analogy, enthalpy uses the root of the Greek word θάλπος "warmth, heat" to express the idea of "heat-content"; the term does in fact stand in for the older term "heat content", a term, now deprecated as misleading, as dH refers to the amount of heat absorbed in a process at constant pressure only, but not in the general case.
Josiah Willard Gibbs used the term "a heat function for constant pressure" for clarity. Introduction of the concept of "heat content" H is associated with Benoît Paul Émile Clapeyron and Rudolf Clausius; the term enthalpy first appeared in print in 1909. It is attributed to Heike Kamerlingh Onnes, who most introduced it orally the year before, at the first meeting of the Institute of Refrigeration in Paris, it gained currency only in the 1920s, notably with the Mollier Steam Tables and Diagrams, published in 1927. Until the 1920s, the symbol H was used, somewhat inconsistently, for "heat" in general; the definition of H as limited to enthalpy or "heat content at constant pressure" was formally proposed by Alfred W. Porter in 1922; the enthalpy of a thermodynamic system is defined as H = U + p V, where H is enthalpy U is the internal energy of the system p is pressure V is the volume of the systemEnthalpy is an extensive property. This means, it is convenient to introduce the specific enthalpy h = H m, where m is the mass of the system, or the molar enthalpy H m = H n, where n is the number of moles.
For inhomogeneous systems the enthalpy is the sum of the enthalpies of the composing subsystems: H = ∑ k H k, where H is the total enthalpy of all the subsystems k refers to the various subsystems H k refers to the enthalpy of each subsystem ∑ k
Catalysis is the process of increasing the rate of a chemical reaction by adding a substance known as a catalyst, not consumed in the catalyzed reaction and can continue to act repeatedly. Because of this, only small amounts of catalyst are required to alter the reaction rate in principle. In general, chemical reactions occur faster in the presence of a catalyst because the catalyst provides an alternative reaction pathway with a lower activation energy than the non-catalyzed mechanism. In catalyzed mechanisms, the catalyst reacts to form a temporary intermediate, which regenerates the original catalyst in a cyclic process. A substance which provides a mechanism with a higher activation energy does not decrease the rate because the reaction can still occur by the non-catalyzed route. An added substance which does reduce the reaction rate is not considered a catalyst but a reaction inhibitor. Catalysts may be classified as either heterogeneous. A homogeneous catalyst is one whose molecules are dispersed in the same phase as the reactant's molecules.
A heterogeneous catalyst is one whose molecules are not in the same phase as the reactant's, which are gases or liquids that are adsorbed onto the surface of the solid catalyst. Enzymes and other biocatalysts are considered as a third category. In the presence of a catalyst, less free energy is required to reach the transition state, but the total free energy from reactants to products does not change. A catalyst may participate in multiple chemical transformations; the effect of a catalyst may vary due to the presence of other substances known as inhibitors or poisons or promoters. Catalyzed reactions have a lower activation energy than the corresponding uncatalyzed reaction, resulting in a higher reaction rate at the same temperature and for the same reactant concentrations. However, the detailed mechanics of catalysis is complex. Catalysts may bind to the reagents to polarize bonds, e.g. acid catalysts for reactions of carbonyl compounds, or form specific intermediates that are not produced such as osmate esters in osmium tetroxide-catalyzed dihydroxylation of alkenes, or cause dissociation of reagents to reactive forms, such as chemisorbed hydrogen in catalytic hydrogenation.
Kinetically, catalytic reactions are typical chemical reactions. The catalyst participates in this slowest step, rates are limited by amount of catalyst and its "activity". In heterogeneous catalysis, the diffusion of reagents to the surface and diffusion of products from the surface can be rate determining. A nanomaterial-based catalyst is an example of a heterogeneous catalyst. Analogous events associated with substrate binding and product dissociation apply to homogeneous catalysts. Although catalysts are not consumed by the reaction itself, they may be inhibited, deactivated, or destroyed by secondary processes. In heterogeneous catalysis, typical secondary processes include coking where the catalyst becomes covered by polymeric side products. Additionally, heterogeneous catalysts can dissolve into the solution in a solid–liquid system or sublimate in a solid–gas system; the production of most industrially important chemicals involves catalysis. Most biochemically significant processes are catalysed.
Research into catalysis is a major field in applied science and involves many areas of chemistry, notably organometallic chemistry and materials science. Catalysis is relevant to many aspects of environmental science, e.g. the catalytic converter in automobiles and the dynamics of the ozone hole. Catalytic reactions are preferred in environmentally friendly green chemistry due to the reduced amount of waste generated, as opposed to stoichiometric reactions in which all reactants are consumed and more side products are formed. Many transition metals and transition metal complexes are used in catalysis as well. Catalysts called. A catalyst works by providing an alternative reaction pathway to the reaction product; the rate of the reaction is increased as this alternative route has a lower activation energy than the reaction route not mediated by the catalyst. The disproportionation of hydrogen peroxide creates oxygen, as shown below. 2 H2O2 → 2 H2O + O2This reaction is preferable in the sense that the reaction products are more stable than the starting material, though the uncatalysed reaction is slow.
In fact, the decomposition of hydrogen peroxide is so slow that hydrogen peroxide solutions are commercially available. This reaction is affected by catalysts such as manganese dioxide, or the enzyme peroxidase in organisms. Upon the addition of a small amount of manganese dioxide, the hydrogen peroxide reacts rapidly; this effect is seen by the effervescence of oxygen. The manganese dioxide is not consumed in the reaction, thus may be recovered unchanged, re-used indefinitely. Accordingly, manganese dioxide catalyses this reaction. Catalytic activity is denoted by the symbol z and measured in mol/s, a unit, called katal and defined the SI unit for catalytic activity since 1999. Catalytic activity is not a kind of reaction rate, but a property of the catalyst under certain conditions, in relation to a specific chemical reaction. Catalytic activity of one katal of a catalyst means one mole of that catalyst will catalyse 1 mole of the reactant to product in one second. A catalyst may and will have different catalytic activity for di
Three-center two-electron bond
A 3-center 2-electron bond is an electron-deficient chemical bond where three atoms share two electrons. The combination of three atomic orbitals form three molecular orbitals: one bonding, one non-bonding, one anti-bonding; the two electrons go into the bonding orbital, resulting in a net bonding effect and constituting a chemical bond among all three atoms. In many common bonds of this type, the bonding orbital is shifted towards two of the three atoms instead of being spread among all three. An example of a 3c–2e bond is the trihydrogen cation H+3; this type of bond is called banana bond. An extended version of the 3c–2e bond model features in cluster compounds described by the polyhedral skeletal electron pair theory, such as boranes and carboranes; these molecules derive their stability from having a filled set of bonding molecular orbitals as outlined by Wade's rules. The monomer BH3 is unstable. A B−H−B 3-center-2-electron bond is formed when a boron atom shares electrons with a B−H bond on another boron atom.
The two electrons in a B−H−B bonding molecular orbital are spread out across three internuclear spaces. In diborane, there are two such 3c-2e bonds: two H atoms bridge the two B atoms, leaving two additional H atoms in ordinary B−H bonds on each B; as a result, the molecule achieves stability since each B participates in a total of four bonds and all bonding molecular orbitals are filled, although two of the four bonds are 3-centre B−H−B bonds. The reported bond order for each B−H interaction in a bridge is 0.5, so that the bridging B−H−B bonds are weaker and longer than the terminal B−H bonds, as shown by the bond lengths in the structural diagram. This bonding pattern is seen in trimethylaluminium, which forms a dimer Al26 with the carbon atoms of two of the methyl groups in bridging positions; this type of bond occurs in carbon compounds, where it is sometimes referred to as hyperconjugation. The first stable subvalent Be complex observed contains a three-center two-electron π-bond that consists of donor-acceptor interactions over the C-Be-C core of a Be-carbene adduct.
Carbocation rearrangement reactions occur through three-center bond transition states. Because the three center bond structures have about the same energy as carbocations, there is virtually no activation energy for these rearrangements so they occur with extraordinarily high rates. Carbonium ions such as ethanium C2H+7 have three-center two-electron bonds; the best known and studied structure of this sort is the 2-Norbornyl cation. Three-center four-electron bond 2-Norbornyl cation Dihydrogen complex
Enzymes are macromolecular biological catalysts. Enzymes accelerate chemical reactions; the molecules upon which enzymes may act are called substrates and the enzyme converts the substrates into different molecules known as products. All metabolic processes in the cell need enzyme catalysis in order to occur at rates fast enough to sustain life. Metabolic pathways depend upon enzymes to catalyze individual steps; the study of enzymes is called enzymology and a new field of pseudoenzyme analysis has grown up, recognising that during evolution, some enzymes have lost the ability to carry out biological catalysis, reflected in their amino acid sequences and unusual'pseudocatalytic' properties. Enzymes are known to catalyze more than 5,000 biochemical reaction types. Most enzymes are proteins; the latter are called ribozymes. Enzymes' specificity comes from their unique three-dimensional structures. Like all catalysts, enzymes increase the reaction rate by lowering its activation energy; some enzymes can make their conversion of substrate to product occur many millions of times faster.
An extreme example is orotidine 5'-phosphate decarboxylase, which allows a reaction that would otherwise take millions of years to occur in milliseconds. Chemically, enzymes are like any catalyst and are not consumed in chemical reactions, nor do they alter the equilibrium of a reaction. Enzymes differ from most other catalysts by being much more specific. Enzyme activity can be affected by other molecules: inhibitors are molecules that decrease enzyme activity, activators are molecules that increase activity. Many therapeutic drugs and poisons are enzyme inhibitors. An enzyme's activity decreases markedly outside its optimal temperature and pH, many enzymes are denatured when exposed to excessive heat, losing their structure and catalytic properties; some enzymes are used commercially, in the synthesis of antibiotics. Some household products use enzymes to speed up chemical reactions: enzymes in biological washing powders break down protein, starch or fat stains on clothes, enzymes in meat tenderizer break down proteins into smaller molecules, making the meat easier to chew.
By the late 17th and early 18th centuries, the digestion of meat by stomach secretions and the conversion of starch to sugars by plant extracts and saliva were known but the mechanisms by which these occurred had not been identified. French chemist Anselme Payen was the first to discover an enzyme, diastase, in 1833. A few decades when studying the fermentation of sugar to alcohol by yeast, Louis Pasteur concluded that this fermentation was caused by a vital force contained within the yeast cells called "ferments", which were thought to function only within living organisms, he wrote that "alcoholic fermentation is an act correlated with the life and organization of the yeast cells, not with the death or putrefaction of the cells."In 1877, German physiologist Wilhelm Kühne first used the term enzyme, which comes from Greek ἔνζυμον, "leavened" or "in yeast", to describe this process. The word enzyme was used to refer to nonliving substances such as pepsin, the word ferment was used to refer to chemical activity produced by living organisms.
Eduard Buchner submitted his first paper on the study of yeast extracts in 1897. In a series of experiments at the University of Berlin, he found that sugar was fermented by yeast extracts when there were no living yeast cells in the mixture, he named the enzyme that brought about the fermentation of sucrose "zymase". In 1907, he received the Nobel Prize in Chemistry for "his discovery of cell-free fermentation". Following Buchner's example, enzymes are named according to the reaction they carry out: the suffix -ase is combined with the name of the substrate or to the type of reaction; the biochemical identity of enzymes was still unknown in the early 1900s. Many scientists observed that enzymatic activity was associated with proteins, but others argued that proteins were carriers for the true enzymes and that proteins per se were incapable of catalysis. In 1926, James B. Sumner crystallized it; the conclusion that pure proteins can be enzymes was definitively demonstrated by John Howard Northrop and Wendell Meredith Stanley, who worked on the digestive enzymes pepsin and chymotrypsin.
These three scientists were awarded the 1946 Nobel Prize in Chemistry. The discovery that enzymes could be crystallized allowed their structures to be solved by x-ray crystallography; this was first done for lysozyme, an enzyme found in tears and egg whites that digests the coating of some bacteria. This high-resolution structure of lysozyme marked the beginning of the field of structural biology and the effort to understand how enzymes work at an atomic level of detail. An enzyme's name is derived from its substrate or the chemical reaction it catalyzes, with the word ending in -ase. Examples are alcohol dehydrogenase and DNA polymerase. Different enzymes that catalyze the same chemical reaction are called isozymes; the International Union of Biochemistry and Molecular Biology have developed a nomenclature for enzymes, the EC numbers. The first number broadly classifies the enzyme based on its mechanism; the top-level classification is: EC 1, Oxidoreductases: catalyze oxidation/reducti
A hypervalent molecule is a molecule that contains one or more main group elements bearing more than eight electrons in their valence shells. Phosphorus pentachloride, sulfur hexafluoride, chlorine trifluoride, the chlorite ion, the triiodide ion are examples of hypervalent molecules. Hypervalent molecules were first formally defined by Jeremy I. Musher in 1969 as molecules having central atoms of group 15–18 in any valence other than the lowest. Several specific classes of hypervalent molecules exist: Hypervalent iodine compounds are useful reagents in organic chemistry Tetra-, penta- and hexavalent phosphorus and sulfur compounds Noble gas compounds Halogen polyfluorides N-X-L nomenclature, introduced collaboratively by the research groups of Martin and Kochi in 1980, is used to classify hypervalent compounds of main group elements, where: N represents the number of valence electrons X is the chemical symbol of the central atom L the number of ligands to the central atomExamples of N-X-L nomenclature include: XeF2, 10-Xe-2 PCl5, 10-P-5 SF6, 12-S-6 IF7, 14-I-7 The debate over the nature and classification of hypervalent molecules goes back to Gilbert N. Lewis and Irving Langmuir and the debate over the nature of the chemical bond in the 1920s.
Lewis maintained the importance of the two-center two-electron bond in describing hypervalence, thus using expanded octets to account for such molecules. Using the language of orbital hybridization, the bonds of molecules like PF5 and SF6 were said to be constructed from sp3dn orbitals on the central atom. Langmuir, on the other hand, upheld the dominance of the octet rule and preferred the use of ionic bonds to account for hypervalence without violating the rule. In the late 1920s and 1930s, Sugden argued for the existence of a two-center one-electron bond and thus rationalized bonding in hypervalent molecules without the need for expanded octets or ionic bond character. In the 1940s and 1950s, Rundle and Pimentel popularized the idea of the three-center four-electron bond, the same concept which Sugden attempted to advance decades earlier; the attempt to prepare hypervalent organic molecules began with Hermann Staudinger and Georg Wittig in the first half of the twentieth century, who sought to challenge the extant valence theory and prepare nitrogen and phosphorus-centered hypervalent molecules.
The theoretical basis for hypervalency was not delineated until J. I. Musher's work in 1969. In 1990, Magnusson published a seminal work definitively excluding the significance of d-orbital hybridization in the bonding of hypervalent compounds of second-row elements; this had long been a point of contention and confusion in describing these molecules using molecular orbital theory. Part of the confusion here originates from the fact that one must include d-functions in the basis sets used to describe these compounds, the contribution of the d-function to the molecular wavefunction is large; these facts were interpreted to mean that d-orbitals must be involved in bonding. However, Magnusson concludes in his work that d-orbital involvement is not implicated in hypervalency. A 2013 study showed that although the Pimentel ionic model best accounts for the bonding of hypervalent species, the energetic contribution of an expanded octet structure is not null. In this modern valence bond theory study of the bonding of xenon difluoride, it was found that ionic structures account for about 81% of the overall wavefunction, of which 70% arises from ionic structures employing only the p orbital on xenon while 11% arises from ionic structures employing an s d z 2 hybrid on xenon.
The contribution of a formally hypervalent structure employing an orbital of sp3d hydridization on xenon accounts for 11% of the wavefunction, with a diradical contribution making up the remaining 8%. The 11% sp3d contribution results in a net stabilization of the molecule by 7.2 kcal mol-1, a minor but significant fraction of the total energy of the total bond energy. Other studies have found minor but non-negligible energetic contributions from expanded octet structures in SF6 and XeF6. Despite the lack of chemical realism, the IUPAC recommends the drawing of expanded octet structures for functional groups like sulfones and phosphoranes, in order to avoid the drawing of a large number of formal charges or partial single bonds. Both the term and concept of hypervalency still fall under criticism. In 1984, in response to this general controversy, Paul von Ragué Schleyer proposed the replacement of'hypervalency' with use of the term hypercoordination because this term does not imply any mode of chemical bonding and the question could thus be avoided altogether.
The concept itself has been criticized by Ronald Gillespie who, based on an analysis of electron localization functions, wrote in 2002 that "as there is no fundamental difference between the bonds in hypervalent and non-hypervalent molecules there is no reason to continue to use the term hypervalent."Fo