Chemical Bond
The term chemical bond refers to any force of attraction between two atoms or ions. Today, chemists recognize the existence of at least four types of chemical bonds: ionic (electrovalent), covalent, metallic, and hydrogen.
The concept of chemical bonds goes back—at least in its simplest and most general form--to the ancient Greeks. Philosophers who thought of matter as being composed of individual particles often considered the possibility that those particles might join, or bond with, each other in some way. As early as 100 b.c. Asklepiades of Prusa introduced the concept of "clusters" of atoms, somewhat similar to those that make up a molecule.
In 1789 the English chemist William Higgins (1763-1825) speculated about the way in which the fundamental particles of matter might combine with each other. When John Dalton proposed his atomic theory around 1803, he specifically hypothesized that the atoms of elements would combine with each other to form "compound atoms." No modern theory of bonding was possible, however, until the concept of a molecule was introduced by Amedeo Avogadro and then clarified by Stanislao Cannizarro and until the formulas of compounds could be agreed upon. Those steps were finally taken in the mid-1800s. Shortly thereafter, in 1858, Friedrich Kekulé attempted to describe the way in which atoms might combine with each other in forming organic compounds.
Kekulé hypothesized that a carbon atom was able to join with, or bond to, four other atoms. He also suggested that carbon atoms could bond with each other endlessly. The drawings he made to depict the bonding of atoms to each other nicely conveyed his ideas, but they were too clumsy to be used by chemists. They were quickly replaced by a system of chemical symbols joined by short dashed lines, introduced at about the same time by A. S. Couper (1831-1892).
The bonding in organic compounds was further clarified in 1874 by Jacobus van't Hoff. Van't Hoff emphasized the importance of thinking of molecules in three-dimensional terms. He suggested that the bonds in a carbon atom are directed to the four corners of a tetrahedron, with the carbon atom at its center. By placing the four other atoms at the corners of the tetrahedron, van 't Hoff was able to explain the existence of two or more forms of a compound that had identical molecular formulas, but different optical properties (optical isomers). The French chemist Joseph Le Bel (1847-1930) proposed a similar theory at about the same time.
An important first step in explaining precisely how a bond forms was offered in 1904 by the German chemist Richard Abegg (1869-1910). Abegg came to the conclusion that the electronic structure of the atoms of inert gases--a complete outer shell of eight electrons--constituted a stable configuration. He suggested that atoms with more or less than eight electrons in their outer shell would gain or lose electrons in order to attain that stable configuration. This suggestion essentially describes the process now known as ionic bonding.
Abegg's theory of bonding was extended by Gilbert Newton Lewis, Walther Kossel (1888-1956), and Irving Langmuir in the late 1910s. Those investigators pointed out a second way in which atoms might bond with each other. Rather than losing or gaining electrons, they said, two atoms might share electrons with each other. In each case, the contribution of one electron from each atom would make up a shared pair that would provide a stable configuration for both atoms. Langmuir coined the term covalent bond for this kind of arrangement.
Another possibility was also evident. Both electrons in the shared pair might be contributed by only one of the atoms. This kind of bond explained the existence of coordination complexes first studied in detail by Alfred Werner and was called, therefore, a coordinate covalent bond.
The theory of chemical bonding attained its highest development in the work of Linus Pauling in the late 1930s. Before Pauling's work, most discussion of chemical bonds had assumed that the electrons in an atom are stationary. Pauling knew that they were actually in motion, and he developed a new explanation of bonding based on that notion. He first used the wave mechanical theory developed by Louis Victor de Broglie to write equations for the motion of electrons in an atom. Then he showed how the wave patterns of two electrons from adjacent atoms might overlap to produce a new pattern with less energy than that of the two original electrons. The lower energy of the combined pattern provided an explanation for the bond formed between the two atoms. Pauling also showed that the combined pattern is usually not one that resulted in a purely ionic or purely covalent bond, but a hybrid of the two.
In 1994 scientists for the first time successfully studied diradicals, the molecular species hypothesized to be archetypal of chemical bond transformations in many classes of reactions. This research used femtosecond laser techniques with mass spectrometry to freezing the diradicals in time so they could be observed. These studies contributed to scientific knowledge about chemical bonds because their results established diradicals as intermediate agents in the formation of certain bonds.
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