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coordination compounds, chemical compounds that do not fit the unpaired-electron concept of the formation of chemical bonds.
More complicated complexes are usually formed from the reaction of simple chemical compounds. Thus, the reaction of cyano salts of iron and potassium produces the complex potassium ferrocyanide: Fe(CN)2 + 4KCN = K4[Fe(CN)6]. Complexes are widely distributed. The overall number of complexes already synthesized apparently exceeds that of simple compounds. Complexes exist in solutions and melts and in crystalline and gaseous states. The transition of a substance from one physical state to another may lead to a change in the composition and structure of a complex, to the decomposition of some complex groupings, and to the formation of new groupings.
The center of a complex is made up of a central atom called the complex-former (in the example given, iron) and the coordinated molecules or ions, that is, the molecules bound to it, called ligands (in this case, the acidic cyano residue).
The ligands make up the inner sphere of the complex. There are complexes that consist only of a central atom and ligands, such as the metal carbonyls Ti(CO)7, Cr(CO)6, and Fe(CO)5. Ions not bound directly to the central atom are part of the outer sphere of the complex. The outer-sphere ions may be either cations, such as K+ in K4[Fe(CN)6], or anions, such as SO42-, in [Cu(NH3)4SO4. The formulas of complexes are written with the outer-sphere ions outside brackets.
A complex grouping bearing an excess positive electric charge, such as [Cu(NH3)4]2+, or an excess negative charge, such as [Fe(CN)6]4-, is called a complex ion. In solution, complexes with outer-sphere ions are practically completely dissociated. For example,
K2[CoCl4] = 2K+ [CoCl4]2- [Cu(NH3)4]SO4 = [Cu(NH3)4]2+ + SO42-
Complex ions may also be dissociated in solution. For example,
[CoCl4]2- ⇄ Co2+ + 4C1-
The stability of complexes in solution is determined by the dissociation constant K of its complex ion:
(In writing the dissociation constant the equilibrium ion concentrations are put inside the brackets.) The dissociation constant characterizes the thermodynamic stability of the complex and depends on the energy of the bond between the central atom and the ligand.
Another concept is the kinetic stability, or inertness, of a complex grouping—the inability of a complex ion to exchange inner-sphere ions or molecules rapidly for other addends. For example, [Fe(H2O)6]3+ and [Cr(H2O)6]3+ have practically the same Me—H2O bond energies (116 and 122 kcal per mole); however, the former complex exchanges ligands rapidly, and the latter, which is inert, exchanges ligands slowly.
The number of ions or molecules directly bound to the central atom is called the atom’s coordination number. For example, the coordination numbers of the complexes K4[Fe(CN)6], Ti(CO)7, and [Cu(NH3)4]SO4 are 6, 7, and 4, respectively. The coordination numbers of various complex-formers differ. The value of the coordination number varies with the size and chemical nature of the central atoms and ligands. At present, coordination numbers from 1 to 12 are known, although numbers 4 and 6 are most common.
The components of complexes are extremely varied. Central-atom complex-formers are most often transition metals (Ti, V, Cr, Mn, Fe, Co, Ni, Cu, Zn, Zr, Nb, Mo, Ru, Rh, Pd, Ag, Cd, Hf, Ta, W, Re, Os, Ir, Pt, Au, and Hg), rare earth elements, actinides, and some nonmetals, such as B, P, and Si. The ligands may be the anions of acids, such as F-, Cl-, Br-, I-, CN-, NO2-, SO42-, and PO43-, and the most varied types of neutral organic and inorganic molecules and radicals containing oxygen, nitrogen, phosphorus, sulfur, selenium, and carbon atoms.
The most typical inorganic complexes have the anions of acids in the inner sphere (acido complexes). Water is the most common ligand. Aquo complexes form from dissolving simple salts in water (for example, CoCl2 + 6H2O = [Co(H2O)6]2++ 2C1-). Crystalline aquo complexes are called crystal hydrates.
A number of solvato complexes are formed by dissolving salts in various organic and inorganic liquids. Crystalline solvato complexes are called crystal solvates; these include the products of ammonia addition, called ammines (for example, [Ni (NH3)6]Cl2), the products of alcohol addition, or alcoholates, and the products of ether addition, or etherates.
Complicated molecules add to the central atom through oxygen atoms (water, alcohols, ethers and esters), nitrogen atoms (ammonia and its organic derivatives, the amines), phosphorus atoms (PCI3 and phosphine derivatives), and carbon atoms. Often the ligand adds to the central atom through several of its atoms; such ligands are called polydentate. These are ligands among the complicated organic derivatives that coordinate through two, three, four, five, six, or even eight atoms (corresponding to bi-, tri-, tetra-, penta-, hexa-, and octa-dentate ligands). Polydentate organic ligands may form cyclic complexes of the nonelectrolyte type (chelate compounds). For example,
The best ligands, from the viewpoint of the stability of the complexes they form, are complexons, or aminopolycarboxylic acids. The most common such compound is ethylenediaminetetraacetic acid (HOOCCH2)2NCH2CH2N(CH2COOH)2 (complexon II, or EDTA).
Inorganic acido ligands are usually monodentate (less often, bidentate). For example, in the compound (NH4)2[Ce-(NO3)6], each NO3 group is attached to the cerium atom through two oxygen atoms and is thus bidentate. The coordination number of cerium in this compound is 12.
There is no distinct boundary between complexes and ordinary (simple) compounds. The same compounds, depending on the research problems posed, may be considered either simple compounds or complexes. For example, in any crystalline inorganic compound, atoms usually considered complex-formers possess a definite coordination number and, consequently, a neighboring sphere no different in principle from the analogous grouping in an ordinary complex.
The theory of the structure of complexes originated with A. Werner (1893), who proposed the concepts of primary and secondary valence, coordination, the coordination number, and the geometry of complex molecules. The importance of these concepts lasted for a considerable period in the subsequent history of chemistry. L. A. Chugaev, I. I. Cherniaev, and other Soviet scientists have made considerable contributions to the chemistry of complexes—in particular, to the relationship between the structure of the complex and the reactivity of the coordinated groups.
However, classicial coordination theory was unable to explain the formation of certain new classes of complexes, to predict their structure, or to establish the interrelationship between the structure and the physical properties of such complexes. A satisfactory solution to these questions became possible only by using modern quantum-mechanical concepts of the nature of the chemical bond.
Complexes are used widely in separating and purifying platinum metals, gold, silver, nickel, cobalt, and copper, sorting rare earth elements and alkali metals, and a number of other industrial processes. Complexes are widely used in the qualitative and quantitative analysis of the most diverse elements. Complexes found in living organisms consist of compounds of metal ions (Fe, Cu, Mg, Mn, Mo, Zn, Co) with proteins (known as metalloproteins), vitamins, coenzymes, transport substances, or other compounds that carry out specific metabolic functions. Natural complexes play an especially important role in respiration, photosynthesis, biological oxidation, and enzyme catalysis.
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Bersuker, I. B., and A. V. Ablov. Khimicheskaia sviaz’ v kompleksnykh soedineniiakh. Kishinev, 1962.
Grinberg, A. A. Vvedenie v khimiiu kompleksnykh soedinenii,2nd ed. Leningrad-Moscow, 1951.
Day, C, and J. Selbin. Teoreticheskaia neorganicheskaia khimiia. Moscow, 1971. (Translated from English.)
Golovnia, V. A., and I. A. Fedorov. Osnovnye poniatiia khimii kompleksnykh soedinenii. Moscow, 1961.
Iatsimirskii, K. B. Termokhimiia kompleksnykh soedinenii. Moscow, 1951.
Cotton, F., and G. Wilkinson. Sovremennaia neorganicheskaia khimiia, parts 1–3. Moscow, 1969. (Translated from English.)
B. F. DZHURINSKII