Coordination chemistry

(chemistry) The chemistry of metal ions in their interactions with other molecules or ions.

A field which, in its broadest usage, is acid-base chemistry as defined by G. N. Lewis. However, the term coordination chemistry is generally used to describe the chemistry of metals and metal ions in their interactions with other molecules or ions. For example, reactions (1)–(3)
1

2

3
show acid-base-type reactions; the products formed are coordination ions or compounds, and this area of chemistry is known as coordination chemistry.

Thus, it follows that coordination compounds are compounds that contain a central atom or ion and a group of ions or molecules surrounding it. Such a compound tends to retain its identity, even in solution, although partial dissociation may occur. The charge on the coordinated species may be positive, zero, or negative, depending on the charges carried by the central atom and the coordinated groups. These groups are called ligands, and the total number of attachments to the central atom is called the coordination number. Other names commonly used for these compounds include complex compounds, complex ions, Werner complexes, coordinated complexes, chelate compounds, or simply complexes. See also Acid and base; Chelation.

Experimental observations as early as the middle of the 18th century reported the isolation of coordination compounds, but valence theory could not adequately account for such materials. The correct interpretation of these compounds was given by Alfred Werner in 1893. He introduced the concept of residual or secondary valence, and suggested that elements have this type of valence in addition to their normal or primary valence. Thus, platinum(IV) has a normal valence of 4 but a secondary valence or coordination number of 6. This then led to the formulaton of PtCl₄ · 6NH₃ as [Pt(NH₃)₆]⁴⁺, 4Cl⁻ and of PtCl₄ · 5NH₃ as [Pt(NH₃)₅Cl]³⁺, 3Cl⁻. The compound with five ammonias has only three ionic chlorides, the fourth is inside the coordination sphere, and therefore is not readily precipitated upon the addition of silver ion. Although the exact nature of the coordinate bond between metal and ligand remains the subject of considerable discussion, it is agreed that the formulations of Werner are essentially correct.

Three theories have been used to explain the nature of the coordinate bond. These are the valence bond theory, the electrostatic theory, including crystal field corrections, and the molecular orbital theory. Currently, the theory used almost exclusively is the molecular orbital theory. The valence bond theory for metal complexes considers that the pair of electrons on the ligand enter the hybridized atomic orbitals of the metal and that the bond is either essentially covalent or essentially ionic. Several of the properties of these substances can be explained on the basis of this theory. The electrostatic theory, plus the crystal field theory for the transition metals, assumes that the metal-ligand bond is caused by electrostatic interactions between point charges and dipoles and that there is no sharing of electrons. In addition to explaining the structure and magnetic properties, the crystal field theory affords an adequate interpretation of the visible spectra of metal complexes. The molecular orbital theory assumes that the electrons move in molecular orbitals which extend over all the nuclei of the metalligand system. In this manner, it serves to make use of both the valence bond theory and crystal field theory. The molecular orbital theory is therefore the best approximation to the nature of the coordinate bond because it is sufficiently flexible to permit both covalent and ionic bonding as well as the splitting of d orbitals into various energy levels. See also Molecular orbital theory.

The stability of metal complexes depends both on the metal ion and the ligand. In general the stability of metal complexes increases if the central ion increases in charge, decreases in size, and increases in electron affinity. Several characteristics of the ligand are known to influence the stability of complexes: (1) basicity of the ligand, (2) the number of metal-chelate rings per ligand, (3) the size of the chelate ring, (4) steric effects, (5) resonance effects, and (6) the ligand atom. Since coordination compounds are formed as a result of acid-base reactions where the metal ion is the acid and the ligand is the base, it follows that generally the more basic ligand will tend to form the more stable complex. The size of the chelate ring is likewise an important factor. For saturated ligands such as ethylenediamine, five-membered rings are the most stable for chelates containing one or more double bonds.

Steric factors often have a very large effect on the stability of metal complexes. This is most frequently observed with ligands having a large group attached to the ligand atom or near it. Thus complexes, of the type shown in the illustration, with alkyl groups R in the position designated are much less stable than the parent complex where R = H. This results from the steric strain introduced by the size of the alkyl group on or adjacent to the ligand atom. In contrast to this, alkyl substitution at any other position results in the formation of more stable complexes because the ligand becomes more basic, and the bulky group is now removed from a position near the coordination site.

Structural formulas of metal complexes which are affected by steric factors.

Finally, the ligand atom itself plays a significant role in controlling the stability of metal complexes. For most of the metal ions, the smallest ligand atom with the largest electron density will form the most stable complex.

Often the most stable complex is also the least reactive or most inert. Several factors, such as the electronic configuration of the central metal ion, its coordination number, and the extent of chelation, all have a marked effect on the rate of reaction of a given compound. See also Chemical bonding; Magnetochemistry; Solid-state chemistry; Stereochemistry.