20th centuryIn the early 20th century, the American chemist Gilbert N. Lewis began to use dots in lecture, while teaching undergraduates at Harvard, to represent the electrons around atoms. His students favored these drawings, which stimulated him in this direction. From these lectures, Lewis noted that elements with a certain number of electrons seemed to have a special stability.
Moreover, noting that a cube has eight corners Lewis envisioned an atom as having eight sides available for electrons, like the corner of a cube. Subsequently, in 1902 he devised a conception in which cubic atoms can bond on their sides to form cubic-structured molecules.
In other words, electron-pair bonds are formed when two atoms share an edge, as in structure C below. This results in the sharing of two electrons. Similarly, charged ionic-bonds are formed by the transfer of an electron from one cube to another, without sharing an edge A. An intermediate state B where only one corner is shared was also postulated by Lewis.
In 1913, while working as the chair of the department of chemistry at the University of California, Berkeley, Lewis read a preliminary outline of paper by an English graduate student, Alfred Lauck Parson, who was visiting Berkeley for a year. In this paper, Parson suggested that the electron is not merely an electric charge but is also a small magnet (or "magneton" as he called it) and furthermore that a chemical bond results from two electrons being shared between two atoms. This, according to Lewis, meant that bonding occurred when two electrons formed a shared edge between two complete cubes.
On these views, in his famous 1916 article The Atom and the Molecule, Lewis introduced the “Lewis structure” to represent atoms and molecules, where dots represent electrons and lines represent covalent bonds. In this article, he developed the concept of the electron-pair bond, in which two atoms may share one to six electrons, thus forming the single electron bond, a single bond, a double bond, or a triple bond.
|“||An electron may form a part of the shell of two different atoms and cannot be said to belong to either one exclusively.||”|
In 1927, the physicits Fritz London and Walter Heitler applied the new quantum mechanics to the deal with the saturable, nondynamic forces of attraction and repulsion, i.e., exchange forces, of the hydrogen molecule. Their valence bond treatment of this problem, in their joint paper, was a landmark in that it brought chemistry under quantum mechanics. Their work was an influence on Pauling, who had just received his doctorate and visited Heitler and London in Zürich on a Guggenheim Fellowship .
Subsequently, in 1931, building on the work of Heitler and London and on theories found in Lewis' famous article, Pauling published his ground-breaking article "The Nature of the Chemical Bond" (see: manuscript) in which he used quantum mechanics to calculate properties and structures of molecules, such as angles between bonds and rotation about bonds. On these concepts, Pauling developed hybridization theory to account for bonds in molecules such as CH4, in which four sp³ hybridised orbitals are overlapped by hydrogen's 1s orbital, yielding four sigma (σ) bonds. The four bonds are of the same length and strength, which yields a molecular structure as shown below:
In 1926, French physicist Jean Perrin received the Nobel Prize in physics for proving, conclusively, the existence of molecules. He did this by calculating Avogadro's number using three different methods, all involving liquid phase systems. First, he used a gamboge soap-like emulsion, second by doing experimental work on Brownian motion, and third by confirming Einstein’s theory of particle rotation in the liquid phase.