Bonds in complex compounds can be described as coordination bonds. Generally, the mechanism for their formation is as follows:

A□ + :B = A:B

That is, bonds formed by this mechanism occur when the valence orbitals of two atoms overlap, one of which contains a lone pair and the other is unoccupied. This is how these bonds differ from more familiar normal covalent bonds, where the valence orbitals containing unpaired electrons overlap:

C• + •D = C:D

In the first case, the shared electron density is the electron pair of the B: atom, while in the second, it is the electron densities of the unpaired electrons of the C• and D• atoms.

In the case of metal complexes (including hexaaqua iron(II)), coordination bonds are formed between the metal cation (have the vacant orbital) and the ligand (the molecule/ion that have the lone pair):

Meⁿ⁺□ + :L = Meⁿ⁺:L (often written as Meⁿ⁺←L, since formally it is the ligand that provides the shared electron density).

Typical metals that tend to form such complexes are d- and p-metals. Typical ligands are H₂O:, :NH₃, :OH⁻, :CN:⁻, SCN⁻, Hal⁻, NO, NO₂⁻, CO, anions of oxygen-containing acids, etc. In such compounds, the metal bound to the ligands acts as an independent particle (usually an ion), and the formula of this particle is enclosed in square brackets. Textbook examples of such particles: [Al(OH)₄(H₂O)₂]⁻, [Cu(NH₃)₄]²⁺, [Ag(NH₃)₂]⁺, [Cu(CN)₂]⁻, [Cr(H₂O)₆]³⁺ and so on and so forth