The electron-pair bond

In 1916, G. N. Lewis, professor of chemistry at the University of California in Berkeley, postulated that a bond between two atoms A and B can arise by their sharing a pair of electrons. Each atom usually contributes one electron. This electron pair bond is called a covalent bond. On this basis, he pictured the 

An examination of the Lewis diagrams shows CH4 and NH3 to be similar in that there are two electrons (represented by dots) adjacent to each hydrogen, whereas carbon and nitrogen are each associated with eight electrons. An important difference is that one electron pair on nitrogen is not shared by a hydrogen. This permits ammonia to react in such a way as to share its Iree electron pair with some other atom. The resulting bond is also an electron pair 

The reaction of ammonia with acids to form ammonium salts, equation (1), produces a coordinate covalent hond. The four N bonds in are, 

These reactions are Lewis acid-base reactions. The Lewis theory of acids and bases defines an acid as a substance capable of accepting a pair of electrons and a base as a substance that donates a pair of electrons. The terms acceptor and donor are sometimes used for acid and base, respectively. A Lewis acid-base reaction results in the formation of a coordinate bond, equation (9). 

The Lewis acid-base concept classifies metal ions as acids. Furthermore, compounds such as BF3, AICI3, SO3, and SiF4, which can accept electron 

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The years from 1923 to 1938 were relatively unproductive for G. N. Lewis insofar as his own research was concerned. The applications of the electron-pair bond came largely in the areas of organic and quantum chemistry in neither of these fields did Lewis feel at home. In the early 1930s. he published a series of relatively minor papers dealing with the properties of deuterium. Then in 1939 he began to publish in the field of photochemistry. Of approximately 20 papers in this area, several were of fundamental importance, comparable in quality to the best work of his early years. Retired officially in 1945, Lewis died a year later while carrying out an experiment on fluorescence. 

Properties of the Electron-Pair Bond.—From the foregoing discussion we infer the following properties of the election-pair bond. 

The electron-pair bond is formed through the interaction of an unpaired electron on each of two atoms. 

It is not proposed to develop a complete proof of the above rules at this place, for even the formal justification of the electron-pair bond in the simplest cases (diatomic molecule, say) requires a formidable array of symbols and equations. The following sketch outlines the construction of an inclusive proof. 

The whole question is clarified when considered in relation to the foregoing quantum mechanical treatment of the electron-pair bond. For the iron-group elements the following rules follow directly from that treatment and from the rules of line spectroscopy. 

Compounds of several different types might be formed by introduction of nitric oxide into complex ions. If the metal atom provides one of the electrons of the electron-pair bond, NO should assume the 02-like 82 structure. If both bond electrons come originally from NO (which then... 

It is shown that a stable shared-electron bond involving one eigenfunction for each of two atoms can be formed under certain circumstances with either one, two, or three electrons. An electron-pair bond can be formed by two arbitrary atoms. A one-electron bond and a three-electron bond, however, can be formed only when a certain criterion involving the nature of the atoms concerned is satisfied. Of these bonds the electron-pair bond is the most stable, with a dissociation energy of 2-4 v. e. The one-electron bond and the three-electron bond have a dissociation energy... 

The electron-pair bond as postulated by Lewis consists of two electrons held jointly by two atoms. By assuming that atoms tend to surround themselves with an outer shell of either shared or unshared electron pairs, usually four in number, but sometimes more or less, Lewis... 

In Sections 42 and 43 we shall describe the accurate and reliable wave-mechanical treatments which have been given the hydrogen molecule-ion and hydrogen molecule. These treatments are necessarily rather complicated. In order to throw further light on the interactions involved in the formation of these molecules, we shall preface the accurate treatments by a discussion of various less exact treatments. The helium molecule-ion, He , will be treated in Section 44, followed in Section 45 by a general discussion of the properties of the one-electron bond, the electron-pair bond, and the three-electron bond. 

The curves for HC1, HBr and HI do not cross, in the main because the ionic radii of Cl-, Br and I are much larger than that of F-. Accordingly the normal states of these molecules are essentially of the electron-pair bond type, and the formulas H Cl , H Br , and H I maybe used as giving a reasonably accurate picture of the state of the molecules. This conclusion had been reached before on the basis of other arguments, especially the tendency of fluorine alone of the halogens to form hydrogen bonds. 

Instead of formulating the wave function for a crystal as a sum of functions describing various ways of distributing the electron-pair bonds among the interatomic positions, as was done in the first section of this paper, let us formulate it in terms of two-electron functions describing a single resonating valence bond. A bond between two adjacent atoms ai and cq- may be described by a function < i3-(l, 2) in which 1 and 2 represent two electrons and the function i may have the simple Heitler-London form... 

Here r is the radius vector from the origin to a point R in the crystal, t is the electron-pair-bond function in the region near R, Pfc is the momentum vector corresponding to the three quantum numbers k (the density of states being calculated in the usual way), h is Planck s constant, and G is the normalizing factor. 

This term involves correlation and polarization on the CO portion of the system with the electron pair bonds to the Hs still in place. 

The study of the reactions of the simple free radicals begun by Bodenstein and Lind in 1906 on the kinetics of gas phase reactions showed that the reactions of H2 with CI2 and Bt2 were complex processes/ and a radical chain mechanism for these reactions (equations 14-18) was proposed in 1919 by Christiansen, Herzfeld, and Polanyi/ The theoretical basis for understanding these reactions in terms of free radicals was presented by G.N. Lewis in 1916, with the theory of the electron pair bond, and free radicals, or odd molecules / Further studies on chain reactions including the extension to explosions in gaseous systems were made by Hinshelwood and by Semenovwho shared the Nobel Prize in 1956. 

The theoretical basis for the understanding of free radicals was first provided by G.N. Lewis in 1916. His clear recognition of the electron pair bond and the possibility of odd electron systems was heavily influenced by the work of pioneers such as Gomberg, Schlenk, and Wieland, who had showed... 

These were bold and simple statements. To put them in a modern context, the discovery of triphenylmethyl combined the novelty of something like bucky balls with the controversial nature of something like polywater or cold fusion. Thus Gomberg was soon to find that the triphenylmethyl problem was attractive and complex enough to occupy him and many others for a long time. A first period lasted until about 1911 when the phenomena observed had been clarified to the satisfaction of a majority of the research community. Theoretically, little understanding was possible before the advent of the electron pair bond and, in particular, theory based on quantum mechanical concepts. This meant that the theory available... 

In the following sections of this chapter there are given, after an introductory survey of the types of chemical bonds, discussions of the concept of resonance and of the nature of the one-electron bond and the electron-pair bond. 

The second method of discussing the electronic structure of molecules, usually called the valence-bond method, involves the use of a wave function of such a nature that the two electrons of the electron-pair bond between, two atoms tend to remain on the two different atoms. The prototype of this method is the Heitler-London treatment of the hydrogen a olecule, which we shall now discuss. 

Hence we see that a very simple treatment of the system of two hydrogen atoms leads to an explanation of the formation of a stable molecule, the energy of the electron-pair bond being in the main the resonance energy corresponding to the interchange of the two electrons between the two atomic orbitals. 

The one-electron bond in the hydrogen molecule-ion is about half as strong as the electron-pair bond in the hydrogen molecule (Do . 60.95 kcal/mole for H/, 102.62 kcal/mcle for Hr—Secs. 1-4, 1-5) and, since the same number of atomic orbitals is needed for a one-electron ixmd as for an electron-pair bond, it is to be expected that in general molecules containing one-electron bonds will be less stable than those in which all the stable bond orbitals are used in electron-pair-bond formation. Moreover, there is a significant condition that must be satisfied in order for a stable one-electron bond to be formed between two atoms namely, that the two atoms be identical or closely similar (Sec. 1-4). For these reasons one-electron bonds are rare—much rarer, indeed, than three-electron bonds, to which similar restrictions apply. 

It may be pointed out that the one-electron bond, the electron-pair bond, and the three-electron bond use one stable bond orbital of each of two atoms, and one, two, and three electrons, respectively. 

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