James Richard Fromm
When John Dalton developed the atomic-molecular theory of matter in the last century, he visualized the atoms as hard, massive particles much like billiard balls. The billiard ball analogy was very useful in the development of the kinetic-molecular theory of gases. In order to account for the attachment of atoms to each other when compounds were formed, Dalton and others visualized atoms as having hooks or fasteners attached so that they could react with other atoms. It became clear almost immediately that the nature of these points of attachment varied considerably from one element to another and were characteristic of the atoms of the elements themselves. Hydrogen seemed to have only one point of attachment, oxygen usually two, and carbon usually four. Most other elements had a quite variable capacity for attachment; their number of attachment points depended upon what other elements they were combining with to form a particular molecule or ion. A point of attachment of an atom came to be called a valence.
The atomic-molecular theory led chemists to make ever more accurate determinations of the atomic mass of atoms (atomic weights) and the molar mass of molecules. From molar masses and atomic masses the stoichiometry of many compounds was determined. Stoichiometry alone, however, does not determine the actual three-dimensional arrangement of the atoms of a molecule in space which is the molecular structure. Moreover, it became clear that the bonding between different atoms in the same molecule was not identical. For example, if the salt NaClO4 is dissolved in water and the equally soluble salt KBr is added, a precipitate of KClO4 will form. Removal of the water will then leave the salt NaBr. The atoms of K and of Na can be interchanged in this way, but the oxygen atoms will remain attached to the chlorine. Perbromate salts are known to exist, but they cannot be formed in this way. The bonding between Na or K and perchlorate must be different in kind from the bonding within perchlorate between Cl and O.
Chemists now recognize three qualitatively different types of chemical bonding which are differentiated on the basis of the type of interaction of the electrons with other atoms. In ionic bonding, electrons are transferred from one originally neutral atom to another. This produces a negative ion or anion and a positive ion or cation. Anions and cations, having opposite charge, are attracted to each other electrostatically; ionic bonds as such do not exist. In the salt NaClO4, the sodium ions and perchlorate ions are associated with each other in a the lattice structure of an ionic crystal. This type of "bonding" is discussed in another section.
The most common type of chemical bonding is covalent bonding. Covalent bonding arises when two, or less often a different number of, electrons are shared between two or more atoms. For example, in the molecule of water (H2O), each hydrogen atom originally has one electron; in the molecule, it shares one pair of electrons with the oxygen atom. The second electron of each pair originally belonged to the oxygen. The remaining four electrons in the outer orbitals of the oxygen, as well as the two electrons in the inner 1s subshell of the oxygen, are not directly involved in the covalent bonding. In most normal covalent bonding, each atom contributes one electron per bond of the two electrons required per bond. In the salt NaClO4, the chlorine atoms and the oxygen atoms are attached to each other by covalent bonds which do not dissociate in aqueous solutions. Covalent bonding is the subject of this section and those following it.
When covalent bonding occurs in which both of the electrons of a bond are contributed by the same atom, the covalent bond is called a coordinate covalent bond. Coordinate covalent bonding is usually found in complex ions or molecules in which the central atom is a transition metal ion. The coordinating groups or molecules, called ligands, contribute both of the electrons involved in the coordinate covalent bond to the central metal atom. Coordinate covalent bonding is the subject of other sections.
Many atoms which cannot react by transfer of electrons form stable compounds. Hydrogen atoms, reacting with each other to form the quite stable diatomic gas, cannot have electron affinities greater than their own ionization potentials. The stability of the diatomic hydrogen molecule is the simplest example of the sharing of electrons between two or more atoms that is called covalent bonding. Covalent bonding is so much more common in chemistry than is ionic bonding that the term chemical bonding is sometimes used in place of the term covalent bonding.