text: If the structure is not known, the average bond valence, S can be calculated from the atomic valence, V, if the coordination number, N, of the atom is known using Eq. 3.
  :                           (Eq. 3)
  If the coordination number is not known, a typical coordination number for the atom can be used instead.  Some atoms, such as sulfur(VI), are only found with one coordination number with oxygen, in this case 4, but others, such as sodium, are found with a range of coordination numbers, though most lie close to the average, which for sodium is 6.2.  In the absence of any better information, the average coordination number observed with oxygen is a convenient approximation, and when this number is used in Eq. 3, the resulting average bond valence is known as the bonding strength of the atom.
  Since the bonding strength of an atom is the valence expected for a bond formed by that atom, it follows that the most stable bonds will be formed between atoms with the same bonding strengths.  In practice some tolerance is allowed, but bonds are rarely formed if the ratio of the bonding strengths of the two atoms exceeds two, a condition expressed by the inequality shown in Eq. 4.  This is known and the valence matching rule.
  :                  (Eq. 4)
  Atoms with non-bonding valence electrons, i.e., with lone pairs, have more flexibility in their bonding strength than those without lone pairs depending on whether the lone pairs are stereoactive or not.  If the lone pairs are not stereoactive, they are spread uniformly around the valence shell, if they are stereoactive they are concentrated in one portion of the coordination sphere preventing that portion from forming bonds.  This results in the atom having a smaller coordination number, hence a higher bonding strength, when the lone pair is stereoactive.  Ions with lone pairs have a greater ability to adapt their bonding strength to match that of the counter-ion.  The lone pairs become stereoactive when the bonding strength of the counter-ion exceeds twice the bonding strength of the ion when its lone pairs are inactive.
  Compounds that do not satisfy Eq. 4 are difficult, if not impossible, to prepare, and chemical reactions tend to favour the compounds that provide the best valence match.  For example, the aqueous solubility of a compound  depends on whether its ions are better matched to water than they are to each other.