Coordination Theory

COORDINATION THEORY concerns the structure of so-called complex compounds (see) or more broadly, molecular compounds. Its basic principles were first formulated by the German chemist Werner (A. Werner: 1893), and at present, in the doctrine of chemical affinity, it plays no less important a role than the theory of valence. According to Werner's theory, when atoms combine with each other, they do not completely exhaust all their affinity forces; they only saturate their "main valence." There remain forces of residual affinity, called by Werner "auxiliary valence," through which the simplest molecules can further combine with each other and give more complex compounds, for example: PtCl₄·2HCl; CoCl₃·6NH₃; AuCl₃·HCl; NHg·HCl and so on. In these complex molecules, Werner distinguishes the "central atom," or "central core," around which the other atoms or groups of atoms are grouped in "zones," or "spheres," both in the form of complete neutral molecules and in the form of charged residues-ions. The number of atoms, molecules, or radicals located in close proximity to the central atom, in the first zone, was called by Werner the "coordination number," or "number of coordination places." It has acquired in the theory of molecular compounds the same important significance that the concept of valence has in the theory of atomic compounds. The most typical coordination numbers for elements are 4 and 6, followed by 2 and 8, and among odd numbers, 3; less common are numbers 1, 5, 7, 12. Just as the valence of an element cannot be greater than a certain number, so the coordination number of an element cannot exceed some limiting value. Compounds with "maximum coordination number" are called "coordinationally saturated." From this point of view, for example, NH₃ is only valence-saturated but not coordinationally saturated, while the ammonium ion NH₄ is both valence and coordinationally saturated, since the coordination number of nitrogen = 4. In carbon, both the valence and coordination number equal 4, and consequently its compounds can be considered valence and coordinationally saturated, which apparently determines the characteristic features of carbon compounds. All coordination places are completely equivalent, and each of them can be replaced by individual atoms as well as by whole molecules or molecular residues. As experience shows, individual atoms can occupy only one coordination place, while some molecules (or radicals) can occupy several coordination places. For example, residues of carbonic and sulfuric acids (CO₃" and SO₄") sometimes occupy one, sometimes two places; the residue of phosphoric acid (PO₄") occupies 3 coordination places; the molecule ethylenediamine occupies 2 coordination places, and so on. Neutral whole molecules are usually located in the first coordination sphere, while atoms or molecular residues can be in both the first and second spheres. Their transition from the first sphere to the second is associated with their simultaneous transition into an ionized state and vice versa (see Complex compounds). All these principles can be expressed in the form of coordination formulas, enclosing the first zone in rectangular brackets and indicating the coordination bond with dots, for example: Cl Cl -| Cl Cl NH₃ -| H₃N -| NH₃ Pt K₂ Co Cl -| Cl Cl H₃N -| NH₃ NH₄ Chloroplatinate Hexaamminecobalt chloride Chloroauric acid Ammonium chloride Chloroplatinic acid The maximum coordination numbers of various elements in most cases turn out to be the same (mostly 4, 6) and therefore do not depend on the chemical nature of the coordinating atoms. This remarkable fact led Werner to the idea that the coordination number must have spatial significance. It shows how many atoms (or molecules) can be arranged in space around another atom and determines the symmetry of this arrangement. For atoms with coordination number 4, by analogy with carbon compounds, one can assume either a tetrahedral arrangement of the four coordinating atoms or a planar configuration in the form of a flat quadrangle. For the numerous series of compounds with coordination number 6, one can assume that the atoms or molecules coordinating around the central atom are located at the corners of an octahedron. Such a spatial interpretation of the coordination number leads to the possibility of spatial isomerism, or stereoisomerism, in those complexes in which two of the 4 or 6 coordinating groups are different. These isomers are called cis- and trans-isomers (see figure). For the octahedron a-trans-position; b-cis-position. In some cases, such an arrangement of groups (for example, ethylenediamine, residues of oxalic acid) is possible, which leads to the appearance of isomers possessing mirror symmetry with respect to each other and, as a consequence, characterized by the possibility of optical activity. All these cases of spatial isomerism have actually been found in a number of complex compounds of Co, Cr, Pt, Ir. This is one of the most brilliant confirmations of coordination theory. The spatial symmetrical arrangement of coordinating groups has now also been proven by X-ray studies of the crystal structure of complex compounds. The presence of the same complex groupings is also detected by X-ray analysis in the crystals of the simplest compounds. Thus, in the crystal lattice of NaCl, one can distinguish the group (NaCl₆) or (ClNa₆). This makes it possible to significantly expand the application of coordination theory and, on the one hand, to consider crystals as a kind of complex complex molecules, and on the other hand, to apply the laws of symmetry of crystal lattices to the composition of complex compounds. From the point of view of coordination theory, one can also consider those compounds which until now fit completely within the framework of valence theory, for example: SO₃ + H₂O → [SO₄] H₂, SO₃ + HCl → [SO₃Cl] H, and so on, and thus unite all complex chemical compounds into a single whole. Coordination theory does not prejudge the question of the nature of auxiliary valence forces. From the point of view of the modern electronic theory, these forces are due to the electric field that exists around every polar molecule, and even more so around an ion. The stronger this field, the stronger the attraction of molecules to each other or to individual ions should be, and the more stable the resulting complex. Atoms with a small volume and a large number of charges should have the strongest field, and therefore they should be most capable of complex formation. And indeed, the most typical complex-forming elements are those that in the periodic system are in the middle of large periods (Cr, Mn, Fe, Co, Ni, Cu, Rh, Pt, Ir)—elements with small atomic volume and high charge. In addition to this theory of the electrostatic field, developed mainly by Kossel, there are also attempts to explain the structure of molecular compounds by electronic bonds similar to the bonds of homopolar molecules, as we have, for example, in the molecules H₂, Cl₂, O₂. Lit.—see literature for the article Complex compounds.

L. L sin', N. Shilov.