The questions raised by this fund of knowledge remained unanswered until the internal structure of atoms began to be unraveled at the end of the 19th century. The classic view, proposed by John Dalton, that atoms are irreducible, unchangeable entities, virtually eliminated the prospect of understanding their properties, for it implied the absence of internal structure. The mutability of atoms, and hence the first glimmerings of an understanding of their constitution and their properties, came with the discovery of the electron as a universal constituent of matter. The electron was the first subatomic particle to be discovered and in due course proved to be the most important one for the explanation of the chemical bond. This importance stems in large part from the ease with which electrons can be removed from one atom and transferred to another. As will be seen below, this transferability of electrons is the key to bond formation, and all theories of the chemical bond focus on the redistribution of an atom’s electrons when it links to another atom.

More will be said about the essential features of the arrangement of electrons in atoms in the following section. The key to understanding the structure of the periodic table and hence the pattern of bonding between atoms was the realization that electrons are arranged in shells that surround a central positively charged nucleus. Each shell can contain a characteristic maximum number of electrons. The outermost shell contains the electrons that are involved in bond formation, for they are the least tightly bound to the nucleus and thus can be removed most readily. This shell is called the valence shell. The most important feature of the valence shell is that for the noble gases it is complete (in the sense explained below) with its full complement of electrons (i.e., eight, excepting the case of helium). Thus, the formation of chemical bonds appears to be related to the incompleteness of the valence shell.

Contributions of Lewis

The role of the valence shell in bond formation was expounded by the American chemist Gilbert N. Lewis about 1916. Important independent studies were made by German physicist Walther Kossel, and later contributions followed from American chemist Irving Langmuir. First, Lewis proposed that ionic bonds are formed by the complete transfer of electrons from the valence shell of one atom into the valence shell of another atom and that the transfer proceeds until the valence shells of both have reached the electronic composition characteristic of the nearest noble gas atom in the periodic table. Thus, sodium has one electron in its valence shell, and its loss results in a singly charged cation, Na⁺, with a neonlike arrangement of electrons. Chlorine, on the other hand, has a valence shell that needs one more electron to achieve the closed shell characteristic of its noble gas neighbour, argon, and so readily forms the singly charged anion Cl⁻. Thus, it is easy to comprehend the formation of sodium chloride as a collection of Na⁺ ions and Cl⁻ ions.

Lewis proposed that a covalent bond consists of two electrons that are shared between atoms rather than being fully donated by one atom to another. He had no means of knowing why a pair of electrons should be so important (that understanding would come only later with the introduction of quantum mechanics), but his insight rationalized a great body of chemical facts. As in the formation of ionic bonds, Lewis emphasized the importance of the nearest-noble-gas valence shell and proposed that, as in the formation of ionic bonds, electron sharing continues until each atom possesses a noble gas configuration.

In summary, Lewis’s ideas are expressed by his celebrated octet rule, which states that electron transfer or electron sharing proceeds until an atom has acquired an octet of electrons (i.e., the eight electrons characteristic of the valence shell of a noble gas atom). When complete transfer occurs, the bonding is ionic. When electrons are merely shared, the bonding is covalent, and each shared electron pair constitutes one chemical bond.

Such is the basis of the theory of chemical bonding that is still widely held. There is much to explain and more to understand, however, and there are many important exceptions to Lewis’s ideas, which cannot as a consequence provide a complete explanation of bonding. The following sections step back from this historical account and put Lewis’s important ideas in a broader context that will show more of their power. At the same time the more-advanced treatment of bonding will transcend Lewis’s ideas and account for features of bonding that his views could not embrace.