A strong force of attraction holding atoms together in a molecule or crystal. Typically chemical bonds have energies of about 1000 kJ mol⁻¹ and are distinguished from the much weaker forces between molecules (see van der Waals’ force). There are various types. Ionic (or electrovalent) bonds can be formed by transfer of electrons. For instance, the calcium atom has an electron configuration of [Ar]4s², i.e. it has two electrons in its outer shell. The chlorine atom is [Ne]3s²3p⁵, with seven outer electrons. If the calcium atom transfers two electrons, one to each chlorine atom, it becomes a Ca²⁺ ion with the stable configuration of an inert gas [Ar]. At the same time each chlorine, having gained one electron, becomes a Cl⁻ ion, also with an inert-gas configuration [Ar]. The bonding in calcium chloride is the electrostatic attraction between the ions. Covalent bonds are formed by sharing of valence electrons rather than by transfer. For instance, hydrogen atoms have one outer electron (1s¹). In the hydrogen molecule, H₂, each atom contributes 1 electron to the bond. Consequently, each hydrogen atom has control of 2 electrons – one of its own and the second from the other atom – giving it the electron configuration of an inert gas [He]. In the water molecule, H₂O, the oxygen atom, with six outer electrons, gains control of an extra two electrons supplied by the two hydrogen atoms. This gives it the configuration [Ne]. Similarly, each hydrogen atom gains control of an extra electron from the oxygen, and has the [He] electron configuration.

A particular type of covalent bond is one in which one of the atoms supplies both the electrons. These are known as coordinate (semipolar or dative) bonds, and written A→B, where the direction of the arrow denotes the direction in which electrons are donated.

Covalent or coordinate bonds in which one pair of electrons is shared are electron-pair bonds and are known as single bonds. Atoms can also share two pairs of electrons to form double bonds, three pairs in triple bonds, four pairs in quadruple bonds, five pairs in quintuple bonds, and six pairs in sextuple bonds. Quadruple and higher bonds occur in certain compounds with metal–metal bonds. Compounds with sextuple bonds are extremely rare, and it is thought that molecules with higher than sextuple bonds could occur only with atoms with an atomic number of about 100. See orbital.

In a compound such as sodium chloride, Na⁺Cl⁻, there is probably complete transfer of electrons in forming the ionic bond (the bond is said to be heteropolar). Alternatively, in the hydrogen molecule H–H, the pair of electrons is equally shared between the two atoms (the bond is homopolar). Between these two extremes, there is a whole range of intermediate bonds, which have both ionic and covalent contributions. Thus, in hydrogen chloride, H–Cl, the bonding is predominantly covalent with one pair of electrons shared between the two atoms. However, the chlorine atom is more electronegative than the hydrogen and has more control over the electron pair; i.e. the molecule is polarized with a positive charge on the hydrogen and a negative charge on the chlorine, forming a dipole. See also banana bond; hydrogen bond; metallic bond; multicentre bond; multiple bond.