coordination compound

coordination compound
complex (def. 10). Also called coordination complex.

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Introduction
 any of a class of substances with chemical structures in which a central metal atom is surrounded by nonmetal atoms or groups of atoms, called ligands (ligand), joined to it by chemical bonds. Coordination compounds include such substances as 12 (vitamin B12), hemoglobin, and chlorophyll, dyes (dye) and pigments (pigment), and catalysts (catalyst) used in preparing organic substances.

 A major application of coordination compounds is their use as catalysts (catalyst), which serve to alter the rate of chemical reactions. Certain complex metal catalysts, for example, play a key role in the production of polyethylene and polypropylene. In addition, a very stable class of organometallic coordination compounds has provided impetus to the development of organometallic chemistry. Organometallic coordination compounds are sometimes characterized by “sandwich” structures, in which two molecules of an unsaturated cyclic hydrocarbon, which lacks one or more hydrogen atoms, bond on either side of a metal atom. This results in a highly stable aromatic system.

      The following article covers the history, applications, and characteristics (including structure and bonding, principle types of complexes, and reactions and syntheses) of coordination compounds. For more information about specific properties or types of coordination compounds, see the articles isomerism; coordination number; chemical reaction; and organometallic compound.

Coordination compounds in nature
 Naturally occurring coordination compounds are vital to living organisms. Metal complexes play a variety of important roles in biological systems. Many enzymes (enzyme), the naturally occurring catalysts that regulate biological processes, are metal complexes (metalloenzymes); for example, carboxypeptidase, a hydrolytic enzyme important in digestion, contains a zinc ion coordinated to several amino acid residues of the protein. Another enzyme, catalase, which is an efficient catalyst for the decomposition of hydrogen peroxide, contains iron- porphyrin complexes. In both cases, the coordinated metal ions are probably the sites of catalytic activity. hemoglobin also contains iron-porphyrin complexes, its role as an oxygen carrier being related to the ability of the iron atoms to coordinate oxygen molecules reversibly. Other biologically important coordination compounds include chlorophyll (a magnesium-porphyrin complex) and 12 (vitamin B12), a complex of cobalt with a macrocyclic ligand known as corrin.

Coordination compounds in industry
      The applications of coordination compounds in chemistry and technology (technology, history of) are many and varied. The brilliant and intense colours of many coordination compounds, such as Prussian blue, render them of great value as dyes (dye) and pigments (pigment). Phthalocyanine complexes (e.g., copper phthalocyanine), containing large-ring ligands closely related to the porphyrins (porphyrin), constitute an important class of dyes for fabrics.

      Several important hydrometallurgical processes utilize metal complexes. nickel, cobalt, and copper can be extracted from their ores (ore) as ammine complexes using aqueous ammonia. Differences in the stabilities and solubilities of the ammine complexes can be utilized in selective precipitation procedures that bring about separation of the metals. The purification of nickel can be effected by reaction with carbon monoxide to form the volatile tetracarbonylnickel complex, which can be distilled and thermally decomposed to deposit the pure metal. Aqueous cyanide solutions usually are employed to separate gold from its ores in the form of the extremely stable dicyanoaurate(−1) complex. Cyanide complexes also find application in electroplating.

      There are a number of ways in which coordination compounds are used in the analysis of various substances. These include (1) the selective precipitation of metal ions as complexes—for example, nickel(2+) ion as the dimethylglyoxime complex (shown below),

      (2) the formation of coloured complexes, such as the tetrachlorocobaltate(2−) ion, which can be determined spectrophotometrically—that is, by means of their light absorption properties, and (3) the preparation of complexes, such as metal acetylacetonates, which can be separated from aqueous solution by extraction with organic solvents.

      In certain circumstances, the presence of metal ions (ion) is undesirable, as, for example, in water, in which calcium (Ca2+) and magnesium (Mg2+) ions cause hardness. In such cases the undesirable effects of the metal ions frequently can be eliminated by “sequestering” the ions as harmless complexes through the addition of an appropriate complexing reagent. Ethylenediaminetetraacetic acid (EDTA) forms very stable complexes, and it is widely used for this purpose. Its applications include water softening (by tying up Ca2+ and Mg2+) and the preservation of organic substances, such as vegetable oils and rubber, in which case it combines with traces of transition metal ions that would catalyze oxidation of the organic substances.

      A technological and scientific development of major significance was the discovery in 1954 that certain complex metal catalysts (catalyst)—namely, a combination of titanium trichloride, or TiCl3, and triethylaluminum, or Al(C2H5)3—bring about the polymerizations (polymerization) of organic compounds (chemical compound) with carbon-carbon double bonds under mild conditions to form polymers (polymer) of high molecular weight and highly ordered (stereoregular) structures. Certain of these polymers are of great commercial importance because they are used to make many kinds of fibres, films, and plastics (plastic). Other technologically important processes based on metal complex catalysts include the catalysis by metal carbonyls (metal carbonyl), such as hydridotetracarbonylcobalt, of the so-called hydroformylation of olefins (olefin)—i.e., of their reactions with hydrogen and carbon monoxide to form aldehydes (aldehyde)—and the catalysis by tetrachloropalladate(2−) ions of the oxidation of ethylene in aqueous solution to acetaldehyde (see chemical reaction and catalysis).

History of coordination compounds
      Perhaps the earliest known coordination compound is the bright red alizarin dye first used in India and known to the ancient Persians and Egyptians. It is a calcium aluminum chelate complex of hydroxyanthraquinone. The first scientifically recorded observation of a completely inorganic coordination compound is German chemist, physician, and alchemist Andreas Libavius (Libavius, Andreas)'s description in 1597 of the blue colour (due to [Cu(NH3)4]2+) formed when lime water containing sal ammoniac (ammonium chloride) (NH4Cl) comes into contact with brass.

      Another example of a coordination compound is the substance Prussian blue, with formula KFe[Fe(CN)6], which has been used as an artist's pigment since the beginning of the 18th century. Another early example of the preparation of a coordination compound is the use in 1760 of a sparingly soluble compound, potassium hexachloroplatinate(2−), K2[PtCl6], to refine the element platinum.

      The sustained and systematic development of modern coordination chemistry, however, usually is considered to have begun with the discovery by the French chemist B.M. Tassaert in 1798 that ammoniacal solutions of cobalt chloride, CoCl3, develop a brownish mahogany colour. He failed to follow up on his discovery, however. It remained for others to isolate orange crystals with the composition CoCl3 ∙ 6NH3, the correct formulation of which is recognized to be [Co(NH3)6]Cl3; this shows that the six ammonia molecules are associated with the cobalt(3+) ion and the positive charge is balanced by three chloride anions (anion). The particularly significant feature of this observation was the recognition that two independently stable compounds (i.e., cobalt chloride and ammonia) could combine to form a new chemical compound with properties quite different from those of the constituent compounds.

      In the 19th century, as more complexes were discovered, a number of theories were proposed to account for their formation and properties. The most successful and widely accepted of these theories was the so-called chain theory (1869) of the Swedish chemist Christian Wilhelm Blomstrand, as modified and developed by the Danish chemist Sophus Mads Jørgensen. Jørgensen's extensive preparations of numerous complexes provided the experimental foundation not only for the Blomstrand-Jørgensen chain theory but for Alsatian-born Swiss chemist Alfred Werner's (Werner, Alfred) coordination theory (1893) as well.

      Blomstrand proposed that ammonia molecules could link together as −NH3− chains, similar to −CH2− chains in hydrocarbons (hydrocarbon). The number of NH3 molecules associated with the metal (i.e., the length of the chain) depends on the metal and its oxidation state. Werner later explained this number more adequately with his concept of coordination number. Jørgensen proposed that atoms (atom) or groups that dissociated into ions (ion) in solution were bonded through the NH3 chain, whereas those that did not were bonded directly to the metal ion.

      Werner called these two types of bonding ionogenic and nonionogenic, respectively. He proposed that the first occurred outside the coordination sphere and the second inside it. In his first experimental work in support of his coordination theory, Werner, together with the Italian Arturo Miolati, determined the electrical conductivities of solutions of several series of coordination compounds and claimed that the number of ions formed agreed with the constitutions (manners of bonding of the ligands (ligand)) predicted by his theory rather than those predicted by Jørgensen.

      Werner also established the configuration (the spatial arrangement of ligands around the metal ion) of complexes by comparing the number and type of isomers (isomerism) (see below Isomerism (coordination compound)) that he actually prepared for various series of compounds with the number and type theoretically predicted for various configurations. In this way he was able not only to refute the rival Blomstrand-Jørgensen chain theory but also to demonstrate unequivocally that hexacoordinate cobalt(+3) possesses an octahedral configuration. Shortly after he and his American student Victor L. King resolved (split) [CoCl(NH3)(en)2]Cl2 into its optical isomers (see below Enantiomers and Diastereomers (coordination compound)) in 1911, Werner received the 1913 Nobel Prize for Chemistry. The zenith of his quarter-century experimental achievements was attained with his resolution of the completely inorganic tetranuclear compound, [tris(tetraammine-μ-dihydroxocobalt(+3))cobalt(+3)](6+) bromide,

      first prepared by Jørgensen, which effectively silenced even Werner's most vociferous opponents. Today he is universally recognized as the founder not only of coordination chemistry but of structural inorganic chemistry as well.

Characteristics of coordination compounds
      Coordination compounds have been studied extensively because of what they reveal about molecular structure and chemical bonding, as well as because of the unusual chemical nature and useful properties of certain coordination compounds. The general class of coordination compounds—or complexes, as they are sometimes called—is extensive and diverse. The substances in the class may be composed of electrically neutral molecules or of positively or negatively charged species (ions (ion)).

      Among the many coordination compounds having neutral molecules is uranium(+6) fluoride, or uranium hexafluoride (UF6). The structural formula of the compound represents the actual arrangement of atoms (atom) in the molecules (molecule):

      In this formula the solid lines, which represent bonds between atoms, show that four of the fluorine (F) atoms are bonded to the single atom of uranium (U) and lie in a plane with it, the plane being indicated by dotted lines (which do not represent bonds), whereas the remaining two fluorine atoms (also bonded to the uranium atom) lie above and below the plane, respectively.

      An example of an ionic coordination complex is the hydrated ion of nickel, (Ni), hexaaquanickel(2+) ion, [Ni(H2O)6]2+, the structure of which is shown below. In this structure, the symbols and lines are used as above, and the brackets and the “two plus” (2+) sign show that the double positive charge is assigned to the unit as a whole.

      The central metal atom in a coordination compound itself may be neutral or charged (ionic). The coordinated groups—or ligands (ligand)—may be neutral molecules such as water (in the above example), ammonia (NH3), or carbon monoxide (CO); negatively charged ions (anions (anion)) such as the fluoride (in the first example above) or cyanide ion (CN); or, occasionally, positively charged ions (cations (cation)) such as the hydrazinium (N2H5+) or nitrosonium (NO+) ion.

      Complex ions—that is, the ionic members of the family of coordination substances—may exist as free ions in solution, or they may be incorporated into crystalline materials (salts) with other ions of opposite charge. In such salts, the complex ion may be either the cationic (positively charged) or the anionic (negatively charged) component (or, on occasion, both). The hydrated nickel ion (above) is an example of a cationic complex. An anionic complex is the hexacyanide of the ferric iron (Fe+3) ion, the hexacyanoferrate(3−) ion, [Fe(CN)6]3−, or

      Crystalline salts containing complex ions include potassium hexacyanoferrate(3−) (potassium ferricyanide), K3[Fe(CN)6], and the hexahydrate of nickel chloride, hexaaquanickel(2+) chloride, [Ni(H2O)6]Cl2. In each case the charge on the complex ion is balanced by ions of opposite charge. In the case of potassium ferricyanide, three positively charged potassium ions, K+, balance the negative charge on the complex, and in the nickel complex the positive charges are balanced by two negative chloride ions, Cl. The oxidation state of the central metal is determined from the charges on the ligands and the overall charge on the complex. For example, in hexaaquanickel(2+), water is electrically neutral and the charge on the complex ion is +2; thus, the oxidation state of Ni is +2. In hexacyanoferrate(3−), all six cyano ligands have a charge of –1; thus, the overall charge of –3 dictates that the oxidation state of Fe is +3.

      The distinction between coordination compounds and other substances is, in fact, somewhat arbitrary. The designation coordination compound, however, is generally restricted to substances whose molecules or ions are discrete entities and in which the central atom is metal. Accordingly, molecules such as sulfur(+6) fluoride (sulfur hexafluoride; SF6) and carbon(+4) fluoride (carbon tetrafluoride; CF4) are not normally considered coordination compounds, because sulfur (S) and carbon (C) are nonmetallic elements (chemical element). Yet there is no great difference between these compounds and, say, uranium hexafluoride. Furthermore, such simple ionic salts as sodium chloride (NaCl) or nickel(+2) fluoride (nickel difluoride; NiF2) are not considered coordination compounds, because they consist of continuous ionic lattices rather than discrete molecules. Nevertheless, the arrangement (and bonding) of the anions surrounding the metal ions in these salts is similar to that in coordination compounds. Coordination compounds generally display a variety of distinctive physical and chemical properties, such as colour, magnetic susceptibility, solubility and volatility, an ability to undergo oxidation-reduction reactions (oxidation–reduction reaction), and catalytic activity.

      A coordination compound is characterized by the nature of the central metal atom or ion, the oxidation state of the latter (that is, the gain or loss of electrons (electron) in passing from the neutral atom to the charged ion, sometimes referred to as the oxidation number), and the number, kind, and arrangement of the ligands. Because virtually all metallic elements form coordination compounds—sometimes in several oxidation states and usually with many different kinds of ligands—a large number of coordination compounds are known.

      Coordination number is the term proposed by Werner to denote the total number of bonds from the ligands to the metal atom. Coordination numbers generally range between 2 and 12, with 4 (tetracoordinate) and 6 (hexacoordinate) being the most common. Werner referred to the central atom and the ligands surrounding it as the coordination sphere. Coordination number should be distinguished from oxidation number (defined in the previous paragraph). The oxidation number, designated by an Arabic number with an appropriate sign (or, sometimes, by a Roman numeral in parentheses), is an index derived from a simple and formal set of rules and is not a direct indicator of electron distribution or of the charge on the central metal ion or compound as a whole. For the hexaamminecobalt(3+) ion, [Co(NH3)6]3+, and the neutral molecule triamminetrinitrocobalt(3+), [Co(NO2)3(NH3)3], the coordination number of cobalt is 6 while its oxidation number is +3.

Ligands and chelates (chelate)
      Each molecule or ion of a coordination compound includes a number of ligands (ligand), and, in any given substance, the ligands may be all alike, or they may be different. The term ligand was proposed by the German chemist Alfred Stock in 1916. Attachment of the ligands to the metal atom may be through only one atom, or it may be through several atoms. When only one atom is involved, the ligand is said to be monodentate; when two are involved, it is didentate, and so on. In general, ligands utilizing more than one bond are said to be polydentate. Because a polydentate ligand is joined to the metal atom in more than one place, the resulting complex is said to be cyclic—i.e., to contain a ring of atoms. Coordination compounds containing polydentate ligands are called chelates (chelate) (from Greek chele, “claw”), and their formation is termed chelation. Chelates are particularly stable and useful. An example of a typical chelate is bis(1,2-ethanediamine)copper(2+), the complex formed between the cupric ion (Cu2+) and the organic compound ethylenediamine (NH2CH2CH2NH2, often abbreviated as en in formulas). The formula of the complex is

      [Cu(NH2CH2CH2NH2)2]2+

      and the structural formula is

Mononuclear, monodentate
      The simplest types of coordination compounds are those containing a single metal atom or ion (mononuclear compounds) surrounded by monodentate ligands. Most of the coordination compounds already cited belong to this class. Among the ligands forming such complexes are a wide variety of neutral molecules (such as ammonia, water, carbon monoxide, and nitrogen), as well as monoatomic and polyatomic anions (such as the hydride, fluoride, chloride, oxide, hydroxide, nitrite, thiocyanate, carbonate, sulfate, and phosphate ions). Coordination of such ligands to the metal virtually always occurs through an atom possessing an unshared pair of electrons, which it donates to the metal to form a coordinate bond with the latter. Among the atoms that are known to coordinate to metals are those of virtually all the nonmetallic elements (nonmetal) (such as hydrogen, carbon, oxygen, nitrogen, and sulfur), with the exception of the noble gases (noble gas) (helium [He], neon [Ne], argon [Ar], krypton [Kr], and xenon [Xe]).

Polydentate
      The chelate complex of a copper ion and ethylenediamine mentioned above is an example of a compound formed between a metal ion and a didentate ligand. Two further examples of chelate complexes are shown below.

      These are a nickel complex with a tetradentate large-ring ligand, known as a porphyrin, and a calcium complex with a hexadentate ligand, ethylenediaminetetraacetate (EDTA). Because metal-ligand attachment in such chelate complexes is through several bonds, such complexes tend to be very stable.

      The commonest and most stable complexes of the lanthanoid metals (the series of 14 f-block elements following lanthanum [atomic number 57]) are those with chelating oxygen ligands, such as EDTA-type anions or hydroxo acids (e.g., tartaric (tartaric acid) or citric (citric acid) acids). The formation of such water-soluble complexes is employed in the separation of lanthanoids by ion-exchange chromatography. Lanthanoid β-diketonates are well known because some fluorinated β-diketonates yield volatile complexes amenable to gas-chromatographic separations. Neutral complexes can complex further to yield anionic species such as octacoordinated tetrakis(thenoyltrifluoroacetato)neodymate(1–), [Nd(CF3COCHCOCF3)4].

      Certain ligands may be either monodentate or polydentate, depending on the particular compound in which they occur. The carbonate ion, (CO3)2−, for example, is coordinated to the cobalt (Co3+) ions in two cobalt complexes, pentaamminecarbonatocobalt(+), [Co(CO3)(NH3)5]+, and tetraamminecarbonatocobalt(+), [Co(CO3)(NH3)4]+, through one and two oxygen atoms, respectively.

      Polynuclear complexes are coordination compounds containing two or more metal atoms, or ions, in a single coordination sphere. The two atoms may be held together through direct metal-metal bonds, through bridging ligands, or both. Examples of each are shown above (see above Polydentate (coordination compound)), along with a unique metal-cluster complex having six metal atoms in its nucleus (see organometallic compound).

      Generally, the systematic naming of coordination compounds is carried out by rules recommended by the International Union of Pure and Applied Chemistry (IUPAC). Among the more important of these are the following:
● Neutral and cationic complexes are named by first identifying the ligands, followed by the metal; its oxidation number may be given in Roman numerals enclosed within parentheses. Alternatively, the overall charge on the complex may be given in Arabic numbers in parentheses. This convention is generally followed here. In formulas, anionic ligands (ending in -o; in general, if the anion name ends in -ide, -ite, or -ate, the final e is replaced by -o, giving -ido, -ito, and -ato) are cited in alphabetical order ahead of neutral ones also in alphabetical order (multiplicative prefixes are ignored). When the complex contains more than one ligand of a given kind, the number of such ligands is designated by one of the prefixes di-, tri-, tetra-, penta-, and so on or, in the case of complex ligands, by bis-, tris-, tetrakis-, pentakis-, and so on. In names (as opposed to formulas) the ligands are given in alphabetical order without regard to charge. The oxidation number of the metal is defined in the customary way as the residual charge on the metal if all the ligands were removed together with the electron pairs involved in coordination to the metal. The following examples are illustrative (aqua is the name of the water ligand):

● Anionic complexes are similarly named, except that the name is terminated by the suffix -ate; for example:

● In the case of salts, the cation is named first and then the anion; for example:

● Polynuclear complexes are named as follows, bridging ligands being identified by a prefix consisting of the Greek letter mu (μ-):

      In addition to their systematic designations, many coordination compounds are also known by names reflecting their discoverers or colours. Examples are

Structure and bonding of coordination compounds
      Werner originally postulated that coordination compounds can be formed because the central atoms carry the capacity to form secondary, or coordinate, bonds, in addition to the normal, or valence, bonds (valence). A more complete description of coordinate bonding, in terms of electron pairs (covalent bond), became possible in the 1920s, following the introduction of the concept that all covalent bonds (covalent bond) consist of electron pairs shared between atoms, an idea advanced chiefly by the American physical chemist Gilbert N. Lewis (Lewis, Gilbert N.). In Lewis's formulation, when both electrons are contributed by one of the atoms, as in the boron-nitrogen bond formed when the substance boron trifluoride (BF3) combines with ammonia, the bond is called a coordinate bond:

      In Lewis's formulas, the valence (or bonding) electrons are indicated by dots, with each pair of dots between two atomic symbols representing a bond between the corresponding atoms.

      Following Lewis's ideas, the suggestion was made that the bonds between metals and ligands were of this same type, with the ligands acting as electron donors and the metal ions as electron acceptors. This suggestion provided the first electronic interpretation of bonding in coordination compounds. The coordination reaction between silver ions and ammonia illustrates the resemblance of coordination compounds to the situation in the boron-nitrogen compound. According to this view, the metal ion can be regarded as a so-called Lewis acid and the ligands as Lewis bases (base):

      A coordinate bond may also be denoted by an arrow pointing from the donor to the acceptor.

      Many coordination compounds have distinct geometric structures. Two common forms are the square planar, in which four ligands are arranged at the corners of a hypothetical square around the central metal atom, and the octahedral, in which six ligands are arranged, four in a plane and one each above and below the plane. Altering the position of the ligands relative to one another can produce different compounds with the same chemical formula. Thus, a cobalt ion linked to two chloride ions and four molecules of ammonia can occur in both green and violet forms according to how the six ligands are placed. Replacing a ligand also can affect the colour. A cobalt ion linked to six ammonia molecules is yellow. Replacing one of the ammonia molecules with a water molecule turns it rose red. Replacing all six ammonia molecules with water molecules turns it purple.

      Among the essential properties of coordination compounds are the number and arrangement of the ligands attached to the central metal atom or ion—that is, the coordination number and the coordination geometry, respectively. The coordination number of a particular complex is determined by the relative sizes of the metal atom and the ligands, by spatial (steric) constraints governing the shapes (conformations) of polydentate ligands, and by electronic factors, most notably the electronic configuration of the metal ion. Although coordination numbers from 1 to 16 are known, those below 3 and above 8 are rare. Possible structures and examples of species for the various coordination numbers are as follows: three, trigonal planar ([Au {P(C6H5)3}3]+; four, tetrahedral ([CoCl4]2−) or square planar ([PtCl4]2−); five, trigonal bipyramid ([CuCl5]>}]3−) or square pyramid (VO(acetylacetonate)2); six, octahedral ([Co(NO2)6]3−) or trigonal prismatic ([Re {S2C2(C6H5)2}3]); seven, pentagonal bipyramid (Na5[Mo(CN)7].10H2O), capped trigonal prism (cation in [Ca(H2O)7]2[Cd6Cl16(H2O)2].H2O), or capped octahedron (cation in [Mo(CNC6H5)7][PF6]2); eight, square antiprism or dodecahedron ([Zr(acetylacetonate)4]; and nine, capped square antiprism (La(NH3)9]3+) or tricapped trigonal prism ([ReH9]2−).

       Coordination numbers and geometries of metal cyanide complexesThe influence of the electronic configuration of the metal ion is illustrated by the examples in the table. The numbers labeled “total number of valence electrons” in this table comprise the d electrons of the metal ion together with the pair of electrons donated by each of the ligands. Coordination numbers and geometries of metal cyanide complexes

      Coordination numbers are also affected by the 18-electron rule (sometimes called the noble gas rule), which states that coordination compounds in which the total number of valence electrons (valence electron) approaches but does not exceed 18 (the number of electrons in the valence shells of the noble gases (noble gas)) are most stable. The stabilities of 18-electron valence shells are also reflected in the coordination numbers of the stable mononuclear carbonyls of different metals that have oxidation number 0—e.g., tetracarbonylnickel, pentacarbonyliron, and hexacarbonylchromium (each of which has a valence shell of 18).

      The 18-electron rule applies particularly to covalent complexes, such as the cyanides, carbonyls, and phosphines (phosphine). For more ionic (also called outer-orbital) complexes, such as fluoro or aqua complexes, electronic factors are less important in determining coordination numbers, and configurations corresponding to more than 18 valence electrons are not uncommon. Several nickel(+2) complexes, for example—including the hexafluoro, hexaaqua, and hexaammine complexes—each have 20 valence electrons.

      Any one metal ion tends to have the same coordination number in different complexes—e.g., generally six for chromium(+3)—but this is not invariably so. Differences in coordination number may result from differences in the sizes of the ligands; for example, the iron(+3) ion is able to accommodate six fluoride ions in the hexafluoro complex [FeF6]>]3− but only four of the larger chloride ions in the tetrachloro complex [FeCl4]. In some cases, a metal ion and a ligand form two or more complexes with different coordination numbers—e.g., tetracyanonickelate [Ni(CN)4]>]2− and pentacyanonickelate [Ni(CN)5]>]3−, both of which contain Ni in the +2 oxidation state.

      Coordination compounds often exist as isomers (isomerism)—i.e., as compounds with the same chemical composition but different structural formulas. Many different kinds of isomerism occur among coordination compounds. The following are some of the more common types.

Cis-trans isomerism
      Cis-trans (geometric) isomers of coordination compounds differ from one another only in the manner in which the ligands are distributed spatially; for example, in the isomeric pair of diamminedichloroplatinum compounds

      the two ammonia molecules and the two chlorine atoms are situated next to one another in one isomer, called the cis (Latin for “on this side”) isomer, and across from one another in the other, the trans (Latin for “on the other side”) isomer. A similar relationship exists between the cis and trans forms of the tetraamminedichlorocobalt(1+) ion:

Enantiomers (enantiomorph) and diastereomers
      So-called optical isomers (or enantiomers) have the ability to rotate plane-polarized light in opposite directions. Enantiomers exist when the molecules of the substances are mirror images but are not superimposable upon one another. In coordination compounds, enantiomers can arise either from the presence of an asymmetric ligand, such as one isomer of the amino acid, alanine (aminopropionic acid),

      or from an asymmetric arrangement of the ligands. Familiar examples of the latter variety are octahedral complexes carrying three didentate ligands, such as ethylenediamine, NH2CH2CH2NH2. The two enantiomers corresponding to such a complex are depicted by the structures below.

      The ethylenediamine ligands above are indicated by a curved line between the symbols for the nitrogen atoms.

      Diastereomers (diastereoisomer), on the other hand, are not superimposable and also are not mirror images. Using AB as an example of a chelating ligand, in which the symbol AB implies that the two ends of the chelate are different, there are six possible isomers of a complex cis-[M(AB)2X2]. For example, AB might correspond to alanine [CH3CH(NH2)C(O)O], where both N and O are attached to the metal. Alternatively, AB could represent a ligand such as propylenenediamine, [NH2CH2C(CH3)HNH2], where the two ends of the molecule are distinguished by the fact that one of the Hs on a C is substituted with a ) group (methyl group).

Ionization isomerism
      Certain isomeric pairs occur that differ only in that two ionic groups exchange positions within (and without) the primary coordination sphere. These are called ionization isomers and are exemplified by the two compounds, pentaamminebromocobalt sulfate, [CoBr(NH3)5]SO4, and pentaamminesulfatocobalt bromide, [Co(SO4)(NH3)5]Br. In the former the bromide ion is coordinated to the cobalt(3+) ion, and the sulfate ion is outside the coordination sphere; in the latter the sulfate ion occurs within the coordination sphere, and the bromide ion is outside it.

Linkage isomerism
      Isomerism also results when a given ligand is joined to the central atom through different atoms of the ligand. Such isomerism is called linkage isomerism. A pair of linkage isomers are the ions [Co(NO2)(NH3)5]2+and [Co(ONO)(NH3)5]2+, in which the anionic ligand is joined to the cobalt atom through nitrogen or oxygen, as shown by designating it with the formulas NO2(nitro) and ONO(nitrito), respectively. Another example of this variety of isomerism is given by the pair of ions [Co(CN)5(NCS)]3− and [Co(CN)5(SCN)]3−, in which an isothiocyanate (NCS) and a thiocyanate group (SCN) are bonded to the cobalt(3+) ion through a nitrogen or sulfur atom, respectively.

Coordination isomerism
      Ionic coordination compounds that contain complex cations (cation) and anions (anion) can exist as isomers if the ligands associated with the two metal atoms are exchanged, as in the pair of compounds, hexaamminecobalt(3+) hexacyanochromate(3–), [Co(NH3)6][Cr(CN)6], and hexaamminechromium(3+) hexacyanocobaltate(3–), [Cr(NH3)6][Co(CN)6]. Such compounds are called coordination isomers, as are the isomeric pairs obtained by redistributing the ligands between the two metal atoms, as in the doubly coordinated pair, tetraammineplatinum(2+) hexachloroplatinate(2–), [Pt(NH3)4][PtCl6], and tetraamminedichloroplatinum(2+) tetrachloroplatinate(2–), [PtCl2(NH3)4][PtCl4].

Ligand isomerism
      Isomeric coordination compounds are known in which the overall isomerism results from isomerism solely within the ligand groups. An example of such isomerism is shown by the ions, bis(1,3-diaminopropane)platinum(2+) and bis(1,2-diaminopropane)platinum(2+),

Bonding theories
Valence bond theory
      Several theories currently are used to interpret bonding in coordination compounds. In the valence bond (VB) theory, proposed in large part by the American scientists Linus Pauling (Pauling, Linus) and John C. Slater, bonding is accounted for in terms of hybridized orbitals of the metal ion, which is assumed to possess a particular number of vacant orbitals (orbital) available for coordinate bonding that equals its coordination number. (See the article chemical bonding for a discussion of the theories of chemical bonding.) Each ligand donates an electron pair to form a coordinate-covalent bond, which is formed by the overlap of an unoccupied orbital of the metal ion and a filled orbital of a ligand. The configuration of the complex depends on the type and number of orbitals involved in the hybridization—e.g., sp (linear), sp3 (tetrahedral), dsp2 (square planar), and d2sp3 (octahedral), in which the superscripts denote the number of orbitals of a particular type. In many cases, the number of unpaired electrons, as determined by magnetic susceptibility measurements, agrees with the theoretical prediction. The theory was modified in 1952 by the Canadian-born American Nobel chemistry laureate Henry Taube (Taube, Henry), who distinguished between inner orbital complexes (d2sp3) and outer orbital complexes (sp3d2) to account for discrepancies between octahedral complexes. The main defect of the simple VB theory lies in its failure to include the antibonding molecular orbitals produced during complex formation. Thus, it fails to offer an explanation for the striking colours of many complexes, which arise from their selective absorption of light of only certain wavelengths (wavelength). From the early 1930s through the early 1950s, VB theory was used to interpret almost all coordination phenomena, for it gave simple answers to the questions of geometry and magnetic susceptibility with which chemists of that time were concerned.

Crystal field theory
      Considerable success in understanding certain coordination compounds also has been achieved by treating them as examples of simple ionic or electrostatic bonding. The German theoretical physicist Walther Kossel's ionic model of 1916 was revitalized and developed by the American physicists Hans Bethe (Bethe, Hans) and John H. Van Vleck (Van Vleck, John H.) into the crystal field theory (CFT) of coordination, used by physicists as early as the 1930s but not generally accepted by chemists until the 1950s. This view attributes the bonding in coordination compounds to electrostatic forces between the positively charged metal ions and negatively charged ligands—or, in the case of neutral ligands (e.g., water and ammonia), to charge separations (dipoles (electric dipole)) that appear within the molecules. Although this approach meets with considerable success for complexes of metal ions with small electronegative (electronegativity) ligands, such as fluoride or chloride ions or water molecules, it breaks down for ligands of low polarity (charge separation), such as carbon monoxide. It also requires modification to explain why the spectral (light-absorption) and magnetic properties of coordinated metal ions generally differ from those of the free ions and why, for a given metal ion, these properties depend on the nature of the ligands.

Ligand field (ligand field theory) and molecular orbital theories
 Since 1950 it has been apparent that a more complete theory, which incorporates contributions from both ionic and covalent bonding, is necessary to give an adequate account of the properties of coordination compounds. Such a theory is the so-called ligand field theory (LFT), which has its origin in the more general, but more complicated, theory of chemical bonding called the molecular orbital (MO) theory. (Molecular orbitals describe the spatial distributions of electrons in molecules (molecule), just as atomic orbitals describe the distributions in atoms (atom).) This theory accounts with remarkable success for most properties of coordination compounds.

      The magnetic properties of a coordination compound can provide indirect evidence of the orbital energy levels used in bonding. Hund rules, which describe the order in which electrons fill atomic shells (see crystal: Magnetism (crystal)), require that the maximum number of unpaired electrons in energy levels have equal or almost equal energies. Compounds that contain no unpaired electrons are slightly repelled by a magnetic field and are said to be diamagnetic. Because unpaired electrons behave like tiny magnets (magnet), compounds that contain unpaired electrons are attracted by a magnetic field and are said to be paramagnetic. The measure of a compound's magnetism is called its magnetic moment. The complex ion hexafluoroferrate(3–) (FeF63−) has a magnetic moment to be expected from a substance with five unpaired electrons, as does the free iron(3+) ion (Fe3+), whereas the magnetic moment of the closely related hexacyanoferrate(3–) ([Fe(CN)6]3−), which also contains Fe3+, corresponds to only one unpaired electron.

      LFT is able to account for this difference in magnetic properties. For octahedral complexes the electrons of the ligands fill all six bonding molecular orbitals, whereas any electrons from the metal cation occupy the nonbonding (t2g) and antibonding (eg) orbitals. The orbital splitting between the two sets of orbitals (t2g and eg) is designated as the orbital ligand field parameter, δo(where o stands for octahedral). Ligands whose orbitals interact strongly with the metal cation's orbitals are called strong-field ligands. For such ligands the orbital splitting is between the t2g and eg orbitals, and consequently the δovalue is large. Ligands whose orbitals interact only weakly with the metal cation's orbitals are called weak-field ligands. For such ligands the orbital splitting is between the t2g and eg orbitals, and consequently the δovalue is small. For transition metal ions with electron configurations d0 through d3 and d8 through d10, only one configuration is possible, so the net spin of the electrons in the complex is the same for both strong-field and weak-field ligands. In contrast, for transition metal ions with electron configurations d4 through d7 (Fe3+ is d5), both high-spin and low-spin states are possible depending on the ligand involved. Strong-field ligands, such as the cyanide ion, result in low-spin complexes, whereas weak-field ligands, such as the fluoride ion, result in high-spin complexes. Therefore, in the [Fe(CN) 6] 3− ion, all five electrons occupy the t2g orbitals, resulting in a magnetic moment indicating one unpaired electron; in the [FeF6] 3− ion, three electrons occupy the t2g orbitals and two electrons occupy the eg orbitals, resulting in a magnetic moment indicating five unpaired electrons.

 An important conclusion from LFT is that two types of bonds, called sigma (σ) bonds (sigma bond) and pi (π) bonds (pi bond), occur in coordination compounds just as they do in ordinary covalent (organic) compounds. The more usual of the two are σ bonds, which are symmetrical about the axis of the bond; π bonds, which are less common, are unsymmetrical with regard to the bond axis. In coordination compounds, π bonding may result from donation of electrons from ligands, such as fluorine or oxygen atoms, to empty d orbitals of the metal atoms. An example of this type of bonding occurs in the chromate ion, (CrO4)2−, in which the oxygen atoms donate electrons to the central chromium ion (Cr6+). Alternatively, electrons from d orbitals of the metal atom may be donated to empty orbitals of the ligand. This is the case in the compound tetracarbonylnickel, Ni(CO)4, in which empty π orbitals in the carbon monoxide molecules accept d-orbital electrons from the nickel atom.

      Ligands may be classified according to their donor and acceptor abilities. Some ligands that possess no orbitals with symmetry appropriate for π bonding, such as ammonia, are σ donors only. On the other hand, ligands with occupied p orbitals are potential π donors and may donate these electrons along with the σ-bonding electrons. For ligands with vacant π* or d orbitals, there is a possibility of π back bonding, and the ligands may be π acceptors. Ligands can be arranged in a so-called spectrochemical series in order from strong π acceptors (correlated with low spin, strong field, and large δ values) to strong π donors (correlated with high spin, weak field, and small δ values) as follows: CO, CN > 1,10-phenanthroline > NO2 > en > NH3 > NCS > H2O > F > RCOO (where R is an alkyl group) > OH > Cl > Br > I. Additional ligands could be added here, but such an expanded list would not be very useful, because the order of the ligands is affected by the nature and charge on the metal ion, the presence of other ligands, and other factors.

      The energy of the light absorbed as electrons are raised to higher levels is the difference in energy between the d orbital levels of transitional metal complexes. As a result, electronic spectra can provide direct evidence of orbital energy levels and information about bonding and electronic configurations in complexes. In some cases, these spectra can also provide information about the magnitude of the effect of ligands on the d orbitals of the metal (δo). The energy levels of d-electron configurations, as opposed to the energies of individual electrons, are complicated, since electrons in atomic orbitals can interact with each other. Tetrahedral complexes give more intense absorption spectra than do octahedral complexes. For f-orbital systems (lanthanoids, 4fn, and actinoids, 5fn) the LFT treatment is similar to that for d-orbital systems. However, the number of parameters is greater, and, even in complexes with cubic symmetry, two parameters are needed to describe the splittings of the f orbitals. Furthermore, f-orbital wave functions (wave function) are not well known, and interpretation of the properties of f-electron systems is much more difficult than it is for d systems. In an effort to overcome such difficulties with f-orbital systems, an approach called the angular overlap model (AOM) was developed, but it proved of relatively little value for these systems.

Principal types of complexes
      The tendency for complexes to form between a metal ion and a particular combination of ligands (ligand) and the properties of the resulting complexes depend on a variety of properties of both the metal ion and the ligands. Among the pertinent properties of the metal ion are its size, charge, and electron configuration. Relevant properties of the ligand include its size and charge, the number and kinds of atoms available for coordination, the sizes of the resulting chelate rings formed (if any), and a variety of other geometric (steric) and electronic factors.

       Chromium complexes of various oxidation statesMany elements, notably certain metals, exhibit a range of oxidation states (oxidation number)—that is, they are able to gain or lose varying numbers of electrons. The relative stabilities of these oxidation states are markedly affected by coordination of different ligands. The highest oxidation states correspond to empty or nearly empty d subshells (as the patterns of d orbitals are called). These states are generally stabilized most effectively by small negative ligands, such as fluorine and oxygen atoms, which possess unshared electron pairs. Such stabilization reflects, in part, the contribution of π bonding caused by electron donation from the ligands to empty d orbitals of the metal ions in the complexes. Conversely, neutral ligands, such as carbon monoxide and unsaturated hydrocarbons (hydrocarbon), which are relatively poor electron donors but which can accept π electrons from filled d orbitals of the metal, tend to stabilize the lowest oxidation states of metals. Intermediate oxidation states are most effectively stabilized by ligands such as water, ammonia, and cyanide ion, which are moderately good σ−electron donors but relatively poor π−electron donors or acceptors (see above Structure and bonding (coordination compound)). Chromium complexes of various oxidation states

Aqua complexes
      Few ligands equal water with respect to the number and variety of metal ions with which they form complexes. Nearly all metallic elements form aqua complexes, frequently in more than one oxidation state. Such aqua complexes include hydrated ions in aqueous solution as well as hydrated salts such as hexaaquachromium(3+) chloride, [Cr(H2O)6]Cl3. For metal ions with partially filled d subshells (i.e., transition metals), the coordination numbers (coordination number) and geometries of the hydrated ions in solution can be inferred from their light-absorption spectra, which are generally consistent with octahedral coordination by six water molecules. Higher coordination numbers probably occur for the hydrated rare-earth ions such as lanthanum(3+).

      When other ligands are added to an aqueous solution of a metal ion, replacement of water molecules in the coordination sphere may occur, with the resultant formation of other complexes. Such replacement is generally a stepwise process, as illustrated by the following series of reactions that results from the progressive addition of ammonia to an aqueous solution of a nickel(2+) salt:

      [Ni(H2O)6]2++ NH3⇌ [Ni(NH3)(H2O)5]2++ H2O

      With increasing additions of ammonia, the equilibria are shifted toward the higher ammine complexes (those with more ammonia and less water) until ultimately the hexaamminenickel(2+) ion predominates:

      [Ni(NH3)5(H2O)]2++ NH3⇌ [Ni(NH3)6]2++ H2O

      The tendency of metal ions in aqueous solution to form complexes with ammonia as well as with organic amines (amine) (derivatives of ammonia, with chains of carbon atoms attached to the nitrogen atom) is widespread. The stabilities of such complexes exhibit a considerable range of dependence on the nature of the metal ion as well as on that of the amine. The marked enhancement of stability that results from chelation is reflected in the equilibrium constants of the reactions—values that indicate the relative proportions of the starting materials and the products at equilibrium. Complexes of hexaaquanickel(2+) ions can be formed with a series of polyamines—i.e.,

      [Ni(H2O)6]2++ nL ⇌ [NiLn(H2O)6 −n]2++ nH2O,

       Equilibrium constants for the formation of various nickel-amine complexesin which L is the ligand and n the number of water molecules displaced from the complex. In this series the equilibrium constants, KL, increase dramatically as the possibilities for chelation increase (that is, as the number of nitrogen atoms available for bonding to the metal atom increases). Equilibrium constants for the formation of various nickel-amine complexes

      It should be noted that, in the particular examples cited above, the coordination number of the metal ion is invariant throughout the substitution process, but this is not always the case. Thus, the ultimate products of the addition of the cyanide ion to an aqueous solution of hexaaquanickel(2+) ion are tetracyanonickelate(2−) and pentacyanonickelate(3−), both containing nickel in the +2 oxidation state. Similarly, addition of the chloride ion to a solution of hexaaquairon(3+) yields tetrachloroferrate(3−). Both complexes contain iron in the same oxidation state of +3.

Halo complexes
      Probably the most widespread class of complexes involving anionic ligands is that of the complexes of the halide ions—i.e., the fluoride, chloride, bromide, and iodide ions. In addition to forming simple halide salts, such as sodium chloride and nickel difluoride (in which the metal ions are surrounded by halide ions, these in a sense being regarded as coordinated to them), many metals form complex halide salts—such as potassium tetrachloroplatinate(2−), K2[PtCl4]—that contain discrete complex ions. Most metal ions also form halide complexes in aqueous solution. The stabilities of such complexes span an enormous range—from the alkali-metal ions (lithium, sodium, potassium, and so on), whose formation of halide complexes in aqueous solution can barely be detected, to extremely stable halide complexes, such as the tetraiodomercurate(2−), tetrachlorothallate(1−), and tetrachloropalladate(2−) ions, the extent of whose dissociation is extremely small.

      The stabilities of halide complexes reflect a pattern by which metal ions can be divided into two general classes, designated as A and B or as hard and soft, respectively. (Generally, the electrons in the atoms of the hard elements are considered to form a compact and not easily deformable group, whereas those in the atoms of the soft elements form a looser group—that is, one more easily deformed.) For the former class—which includes Be, Mg, Sc, Cr, Fe, Ni, Cu, In, and Sn—the order of increasing stability of the halide complexes in aqueous solution is iodides < bromides < chlorides < fluorides. Conversely, for the class B (or soft) ions—such as Pt, Ag, Cd, Hg, Tl, and Pb—the order of increasing stability of the halide complexes is fluorides < chlorides < bromides < iodides. In contrast to class-A metals, those of class B also tend to form more stable complexes with sulfur-containing ligands than with oxygen-containing ligands and more stable complexes with phosphorus ligands than with nitrogen ligands.

Carbonyl complexes
      Following the discovery of the first metal carbonyl complex, tetracarbonylnickel, Ni(CO)4, in 1890, many compounds containing carbon monoxide coordinated to transition metals have been prepared and characterized. For reasons already discussed, such compounds generally contain metal atoms or ions in low oxidation states. The following are some of the more common types of metal carbonyl compounds: (1) simple mononuclear carbonyls of metals in the zero oxidation state, such as tetracarbonylnickel, pentacarbonyliron, and hexacarbonylchromium—highly toxic volatile compounds, the most stable of which have filled valence shells of 18 electrons (electron), (2) salts of anionic and cationic carbonyls, such as tetracarbonylcobaltate(−1) and hexacarbonylmanganese(+1), (3) dinuclear and polynuclear carbonyls, such as bis(tetracarbonylcobalt), the structural formula of which was shown earlier (see above Polynuclear (coordination compound)), and (4) mixed complexes containing other ligands in addition to CO: pentacarbonylchloromanganese, tetracarbonylhydridocobalt, and tricarbonylnitrosylcobalt (see organometallic compound).

      Although molecular nitrogen, N2, is isoelectronic with carbon monoxide (that is, it has the same number and arrangement of electrons), its tendency to form complexes with metals is much smaller. The first complex containing molecular nitrogen as a ligand—i.e., pentaamminenitrogenruthenium(2+), [Ru(NH3)5(N2)]2+—was prepared in 1965, and many others have been discovered subsequently. Such complexes have attracted considerable interest because of their possible roles in the chemical and biological fixation of nitrogen (nitrogen fixation).

Nitrosyl complexes
      Nitrosyl complexes can be formed by the reaction of nitric oxide (NO) with many transition metal compounds or by reactions involving species containing nitrogen and oxygen. Some of these complexes have been known for many years—e.g., pentaaquanitrosyliron(2+) ion, [Fe(H2O)5NO]2+, which formed in the classical brown-ring test for the qualitative detection of nitrate ion; Roussin's red (K2[Fe2S2(NO)4]) and black (K[Fe4S3(NO)7]) salts; and sodium pentacyanonitrosylferrate(3−) dihydrate (sodium nitroprusside), Na2[Fe(CN)5NO]∙2H2O. Such complexes, which can be cationic, neutral, or anionic and which are usually deeply coloured (red, brown, purple, or black), have been extensively studied because they pose unique problems of structure and bonding and because they have potential uses as homogeneous catalysts (catalyst) for a variety of reactions. More recently, the research field has been expanded to include organometallic species (see organometallic compound).

      Because the nitrosonium ion (NO+) is isoelectronic with carbon monoxide and because its mode of coordination to transition metals is potentially similar to that of carbon monoxide, metal nitrosyls have been recognized as similar to carbonyls and are sometimes formulated as NO+ complexes. Carbonyl ligands can be replaced by nitric oxide in substitution reactions (substitution reaction). Such similarities may be deceptive, however, for the additional electron in neutral nitric oxide requires a more complicated treatment of M-NO bond formation. The NO ligand exhibits several geometries of coordination—linear (e.g., [IrH(NO){P(C6H5)3}3]+, [Mn(CO)2(NO){P(C6H5)3}3], and Na2[Ru(OH)(NO2)4(NO)].2H2O); bent (e.g., [CoNO(NH3)5]2+and [IrCl2(NO){P(C6H5)3}2]); or both (e.g., [RuCl(NO)2{P(C6H5)3}2]+). Like CO, NO also can act as a bridging ligand between two (e.g., [{Cr(η5−C5H5)(NO)}22−NH2)(μ2−NO)]) or three (e.g., [Mn35−C5H5)32−NO)33−NO)]) metal atoms. (The η5 indicates that five carbon atoms of the C5H5 group are bonded to the chromium atom.)

Cyano (nitrile) and isocyano complexes
      Cyano complexes, such as Prussian blue, mentioned above, are among the oldest coordination compounds. In addition to being a pseudohalide, the CN ion is isoelectronic with CO, RCN, RNC, N2, and NO+ (R is an alkyl group), and metal carbonyls (metal carbonyl) and cyanide complexes are structurally similar. Also, like CO, CN enters into π as well as σ bonding with transition metal atoms or ions. Cyano complexes are among the most stable transition metal complexes; the extreme toxicity of CN (like that of CO) is due to its irreversible formation of a strong complex with hemoglobin, which prevents oxygen from binding reversibly to hemoglobin, thereby prohibiting the transport and release of oxygen in the body. Similarly, the ability of CN to form very stable complexes with silver (Ag(CN)2) and gold (Au(CN)2) is the basis for its use in the extraction and purification of these metals. As a monodentate ligand, CN coordinates (bonds) through carbon as the donor atom, but, as a didentate ligand, it usually coordinates at both ends (C and N) and acts as a bridging ligand (−CN−) to form infinite linear (chain) polymers (polymer) as in Prussian blue, AgCN, AuCN, Zn(CN)2, and Cd(CN)2.

      The cyanide ion forms complexes with transition metals and with zinc, cadmium, and mercury, usually by substitution in aqueous solution with no change in oxidation state. The most important complexes are anionic with the formula [Mn+(CN)x](xn)−, where Mn+represents a transition metal ion. Examples are [Ni(CN)4]2−, [Pt(CN)4]2−, [Fe(CN)6]4− or 3−, [Co(CN)6]3−, [Pt(CN)6]2−, and [Mo(CN)8]5−, 4−, or 3−. The free anhydrous parent acids (acid) of many of these anions—for example, H4[Fe(CN)6] and H3[Rh(CN)6]—have been isolated.

      Cyanide complexes exhibit a variety of coordination numbers and configurations. Metal ions with a d10structure form linear complexes of coordination number 2—as, for example, [M(CN)2] (where M = Hg, Ag, or Au)—while the isoelectronic complexes [Cu(CN)4]3−, [Ag(CN)4]3−, [Zn(CN)4]2−, [Cd(CN)4]2−, and [Hg(CN)4]2− are tetrahedral. All the hexacoordinate complexes are octahedral, while octacoordinate complexes are cubic, dodecahedral, or square antiprismatic. (The dodecahedron and square antiprism are two structures that can be obtained by distorting the simple cube.) For d2, d4, d6, d8, and d10 transition metal ions, the octa-, hepta-, hexa-, penta-, and tetracoordinate complexes, respectively, are species with maximum coordination number.

      Mixed complexes of type [M(CN)5X]n (where X = H2O, NH3, CO, NO, H, or a halogen (halogen element)) also exist. The cyanide ion has the ability to stabilize metal ions in low oxidation states (probably by accepting electron density into its π*orbitals)—e.g., [Ni(CN)4]4−, which contains nickel in the formal oxidation state of zero. (See the article chemical bonding for a discussion of π*orbitals.) Cyanide complexes have figured prominently in numerous kinetic (kinetics) studies. For example, fast electron-transfer reactions between [Fe(CN)6]3− and [Fe(CN)6]4− and between [Mo(CN)8]3− and [Mo(CN)8]4− established the outer-sphere mechanism for redox reactions (see oxidation-reduction reaction (oxidation–reduction reaction)); replacement of water in [Co(CN)5H2O]2− established the dissociative mechanism for substitution at a Co(3+) ion.

      Transition metals also form complexes with organic cyanides (RCN or ArCN, called nitriles (nitrile)) and organic isocyanides (isocyanide) (RNC or ArNC, called isonitriles)—where R and Ar are alkyl and aryl groups, respectively—by reaction of a metal halide, carbonyl, or other complex with the nitrile or isonitrile, respectively. Nitriles and isonitriles appear to be stronger donors of σ electrons than carbon monoxide, but they are capable of extensive back acceptance of π electrons from metals in lower oxidation states—as in Cr(CNR)6 or Cr(CNAr)6 and Ni(CNR)4 or Ni(CNAr)4, which are analogous to the corresponding carbonyls Cr(CO)6 and Ni(CO)4, respectively. Although a bridging isonitrile group has been reported in (π−C5H5)2Fe2(CO)3(CNC6H5), this type of bonding is unusual.

Organometallic (organometallic compound) complexes
      Organometallic complexes are complexes formed between organic groups and metal atoms. They can be divided into two general classes: (1) complexes containing metal-carbon σ bonds and (2) π-bonded metal complexes of unsaturated hydrocarbons (hydrocarbon)—that is, compounds with multiple bonds between carbon atoms (see organometallic compound).

Isopoly and heteropoly anions
      The amphoteric (amphoterism) metals of groups VB ( vanadium, niobium, and tantalum) and VIB ( chromium, molybdenum, and tungsten) in the +5 and +6 oxidation states, respectively, form weak acids that readily condense (polymerize) to form anions containing several molecules of the acid anhydride. If these condensed acids contain only one type of acid anhydride, they are called isopoly acids, and their salts are called isopoly salts. The acid anhydrides also can condense with other acids (e.g., phosphoric or silicic acids) to form heteropoly acids, which can form heteropoly salts. The condensation reactions, which occur reversibly in dilute aqueous solution, involve formation of oxo bridges by elimination of water from two molecules of the weak acid. The best-known and simplest example is the condensation of yellow chromate ion (CrO42−) to form the orange isopoly dichromate ion (Cr2O72−), an equilibrium reaction the extent of which depends on the pH. In acidic solution the isopoly anion Cr2O72−, predominates while in basic solution the simple ion CrO42− predominates.

      Heteropoly acids and their salts may be formed by coordination of the central atom with four to six oxo anions, which may be mononuclear (containing one metal ion each), as in H7[P(MoO4)6], or trinuclear (containing three metal ions each), as in H3[P(W3O10)4]. Incomplete replacement of oxygen atoms in PO43− ions by MoO3 groups can result in dimers (two-molecule polymers), as, for example, {OP[O(MoO3)3]3}26−. About 70 elements (chemical element) can act as central (hetero) atoms in heteropoly anions. Because each element may form more than one heteropoly anion and some of these anions can contain several different heteroatoms, thousands of heteropoly acids exist. Heteroatoms may be primary (these atoms are essential to the polyanion structure and thus not susceptible to chemical exchange) or secondary (these atoms can be removed by chemical reaction from the polyanion structure without destroying it). Heteropoly anions can be regarded as coordination compounds with polyanion ligands; e.g., [(H3N)5Cr(OH2)]3+ can be considered the parent of [(SiW11O39)Cr(OH2)]5−.

      A variety of synthetic procedures are available for the preparation of isopoly acids and salts, which are usually less stable than heteropoly compounds. Heteropolymolybdates and heteropolytungstates are always prepared in solution, usually after acidifying and heating the theoretical amounts of reactants. In general, free heteropoly acids and salts, of which the heteropolymolybdates and heteropolytungstates are the best known, have very high molecular weights (some above 4,000) as compared with other inorganic electrolytes (electrolyte), are very soluble in water and organic solvents, are almost always highly hydrated with several hydrates existing, and are highly coloured. Some are strong oxidizing agents that can be reduced to stable, intensely deep blue species (heteropoly blues), which in turn can act as reducing agents, restoring the original colour on oxidation. The stoichiometry, oxidation-reduction potentials, and other characteristics of these reactions have been investigated by various methods. The free acids, which are polyprotic (contain several replaceable hydrogen ions (hydrogen ion)), are fairly strong and nearly always stronger than the corresponding acids from which they are derived.

      All heteropolymolybdate and heteropolytungstate anions are decomposed in strongly basic solution to form simple molybdate or tungstate ions and either an oxy anion or a hydrous metal oxide of the central metal atom, e.g.:

      [P2Mo18O62]6− + 34OH ——-> 18MoO42− + 2HPO42− + 16H2O [NiW6O24H6]4− + 8OH ——-> 6WO42− + Ni(OH) 2 + 6H2O

      Throughout specific ranges of pH and other conditions, most solutions of heteropolymolybdates and heteropolytungstates appear to contain predominantly one distinct species of anion, many of which are remarkably stable and nonlabile.

      The first heteropoly compound, (NH4)3[PMo12O40], was obtained by the Swedish chemist Jöns Jacob Berzelius (Berzelius, Jöns Jacob) in 1826 as a yellow, crystalline precipitate, the formation of which is still used for the classical qualitative detection and quantitative estimation of phosphorus (after conversion to phosphate). By the beginning of the 20th century, hundreds of isopoly and heteropoly compounds were reported, many of which were based on incorrect analyses or failure to detect mixed crystals. Formulas were reported in terms of the old Berzelius dualistic theory as a combination of oxides (oxide), such as 3Na2O∙Cr2O3∙12MoO3∙20H2O for Na3CrMo6O24H6∙7H2O, and often merely expressed analytical results rather than structure. In addition to their use in analytical chemistry, heteropoly compounds have found use as catalysts, molecular sieves, corrosion inhibitors, photographic fixing agents, and precipitants for basic dyes (dye).

      Few structural studies of such compounds were carried out, but this lack did not prevent the elaboration of various unsuccessful theories to account for their structures. In 1907 Werner applied his coordination theory to the structure of 12-tungstosilicic acid, H4 [SiW12O40], and its salts by assuming that the central group is an SiO44− ion surrounded octahedrally by six RW2O6+ groups (R = a unipositive ion), four linked by primary (ionic) and two linked by secondary (coordinate covalent) valences. Difficulties were encountered by this system as well as by the later (1910–21), more elaborate Miolati-Rosenheim theory. Modern conclusive knowledge of the structures of heteropoly compounds did not begin until 1934, with J.F. Keggin's determination of the structure of H3 [PO4W12O36]∙5H2O by the most direct means, X-ray diffraction.

      The structures of isopoly and heteropoly compounds consist of polyhedrons sharing corners and edges with one another. In heteropolymolybdates or heteropolytungstates, each molybdenum or tungsten atom is located at the centre of an octahedron, each vertex of which is occupied by an oxygen atom. These octahedrons can share corners or edges or both with other MoO6 or WO6 octahedrons. In [Mo8O26]4− eight MoO6 octahedrons share edges. In [PMo12O40]3− the central phosphorus atom is located at the centre of a PO4 tetrahedron, which is surrounded by 12 MoO6 octahedrons, which share corners so that the correct number of oxygen atoms is utilized.

Important types of reactions of coordination compounds

      Coordination to a positive metal ion usually enhances the acidity (i.e., the tendency to lose protons (proton)) of hydrogen-containing ligands, such as water and ammonia. Thus, many metal ions in aqueous solution commonly exhibit acidic behaviour. Such behaviour is exemplified by hydrolysis reactions of the type shown in the following equilibrium:

      [M(H2O)x]n+⇌ [M(OH)(H2O)x− 1](n− 1)++ H+,QC

      in which M represents the metal ion, n its charge, and x the number of coordinated water molecules.

      The acidities of such aqua ions depend on the charge, size, and electronic configuration of the metal ion. This dependence is reflected in the values of acid dissociation constants, which range from about 10−14 (a value only slightly larger than for pure water, for which the dissociation constant = 10−15.7) for the hydrated lithium ion, to about 10−2 (a value equivalent to that of a fairly strong acid) for the hydrated uranium(4+) ion. Acid-base equilibria are rapidly established in solution, generally within a fraction of a second (see chemical reaction).

      In some cases, hydrolysis of a metal ion may be accompanied by polymerization to form dinuclear or polynuclear hydroxo- or oxygen-bridged complexes.

      Even very weakly acidic ligands, such as ammonia, can acquire appreciable acidity through coordination to a metal ion. Thus, the hexaammineplatinum(4+) ion dissociates according to the following equilibrium:

      [Pt(NH3)6]4+⇌ [Pt(NH2)(NH3)5]3++ H+.

      In addition to intrinsic strength, acids and bases have other properties that determine the extent of reactions. According to the hard and soft acids and bases (HSAB) theory, the metal cation and anion are considered to be acids and bases, respectively. Hard acids and bases are small and nonpolarizable, whereas soft acids and bases are larger and more polarizable. Interactions between two hard or soft acids or bases are stronger than ones between one hard and one soft acid or base. The theory can be used to explain solubilities, formation of metallic ores, and some reactions of metal cations with ligands.

Substitution (substitution reaction)
      One of the most general reactions exhibited by coordination compounds is that of substitution (substitution reaction), or replacement, of one ligand by another. This process is depicted in a generalized manner by the equation MLx− 1Y + Z → MLx− 1Z + Y for a metal complex of coordination number x. The ligands L, Y, and Z may be chemically similar or different. (Charges have been omitted here for simplicity.)

      A class of substitution reactions that affords the widest possible comparison of different metal ions is the replacement of water in the coordination spheres of metal-aqua complexes in aqueous solution. The substitution may be by another water molecule (which can be labeled with the isotope oxygen-18 to permit the reaction to be followed) or by a different ligand, such as the chloride ion. Reactions of both types occur as shown below (oxygen-18 is indicated by the symbol

).

      Many such reactions are extremely fast, and it has been only since 1950, following the development of appropriate experimental methods (including stopped flow, nuclear magnetic resonance, and relaxation spectrometry), that the kinetics and mechanisms of this class of reactions have been extensively investigated. Rates of substitution of metal-aqua ions have been found to span a wide range, the characteristic times required for substitution ranging from less than 10−9 second for monopositive ions, such as hydrated potassium ions, to several days for certain more highly charged ions, such as hexaaquachromium(3+) and hexaaquarhodium(3+). The rate of substitution parallels the ease of loss of a water molecule from the coordination sphere of the aqua complex and thus increases with increasing size and with decreasing charge of the metal ion. For transition metal ions, electronic factors also have an important influence on rates of substitution.

      There are two limiting mechanisms (or pathways) through which substitution may occur—namely, dissociative and associative mechanisms. In the dissociative mechanism, a ligand is lost from the complex to give an intermediate compound of lower coordination number. This type of reaction path is typical of octahedral complexes, many aqua complexes, and metal carbonyls such as tetracarbonylnickel. An example of a dissociative reaction pathway for an octahedral complex of cobalt is as follows:

      The associative mechanism for substitution reactions, on the other hand, involves association of an extra ligand with the complex to give an intermediate of higher coordination number; one of the original ligands is then lost to restore the initial coordination number. Substitution reactions of square planar complexes, such as those of the nickel(2+), palladium(2+), and platinum(2+) ions, usually proceed through associative pathways involving intermediates with coordination number five. An example of a reaction following such a pathway is

      A characteristic feature of this class of reactions is the sensitivity of the rate of substitution of a given ligand to the nature of the ligand in the trans position. The trans ligand activates a ligand for replacement as follows, in decreasing order:

      TBCO, CN, C2H43, H2, I, SCN−, Cl3, H2O.TL

      The trans effect may be used for synthetic purposes; thus, the reaction of the tetrachloroplatinate(2−) ion with ammonia yields cis-diamminedichloroplatinum, whereas the reaction of the tetraammineplatinum(2+) ion with the chloride ion gives the trans isomer, trans-diamminedichloroplatinum. The reactions are shown below.

      In both reactions, the trans effect causes the introduction of the ligand trans to chloride rather than trans to ammonia.

Lability and inertness
      In considering the mechanisms of substitution (exchange) reactions, Canadian-born American chemist Henry Taube (Taube, Henry) distinguished between complexes that are labile (reacting completely in about one minute in 0.1 M solution at room temperature [25 °C, or 77 °F]) and those that are inert (under the same conditions, reacting either too slowly to measure or slowly enough to be followed by conventional techniques). These terms refer to kinetics (reaction rates) and should not be confused with the thermodynamic terms unstable and stable, which refer to equilibrium. For example, as mentioned above, most cyanide complexes are extremely stable (they possess very small dissociation constants); yet, if their rate of exchange with carbon-14-labeled cyanide, as represented in the following equation,

      [M(CN)x]y+ x14CN⇌ [M(14CN)x]y+ xCN,

      is measured, [Ni(CN)4]2− and [Hg(CN)4]2− are found to be labile, whereas [Mn(CN)6]3−, [Fe(CN)6]4−, [Fe(CN)6]3−, and [Cr(CN)6]3− are inert. On the other hand, [Co(NH3)6]3+, a kinetically inert complex, is thermodynamically stable in acidic solution. Inertness may result from the lack of a suitable low-energy pathway for the reaction. In short, stable complexes possess large positive free energies (free energy) of reaction (ΔG), whereas inert complexes merely possess large positive free energies of activation (ΔG*).

      While the existence of geometric or optical isomers (see above Isomerism (coordination compound)) in the solid state or in solution at nonequilibrium concentrations is evidence supporting the inertness of the complex, this does not constitute absolute proof. Conversely, the possibility of intramolecular rearrangement means that failure to isolate geometric isomers or to resolve the racemic mixture into optical isomers is not absolute proof of lability.

      Taube has interpreted lability of complexes according to their electronic configuration in terms of VB theory. Labile complexes are either of the outer orbital type (outer d orbitals involved in bonding—e.g., sp3d2 as opposed to d2sp3 [inner orbital] for octahedral complexes) or of the inner orbital type with at least one vacant d orbital (available for accommodation of a seventh group during the [associative] substitution reaction).

      Coordination compounds that exist in two or more isomeric forms (see above Isomerism (coordination compound)) may undergo reactions that convert one isomer to another. Examples are the linkage isomerization and cis-trans isomerization reactions depicted below.

      The first of these has been shown to proceed intramolecularly (i.e., without dissociation of the nitrite ligand), whereas the second probably occurs through dissociation of one of the water-molecule ligands.

Oxidation-reduction (oxidation–reduction reaction)
      Transition metals (transition element) commonly exhibit two or more stable oxidation states, and their complexes accordingly are able to undergo oxidation-reduction reactions (oxidation–reduction reaction). The simplest such reactions involve electron transfer between two complexes, with little if any accompanying rearrangement or chemical change. An example is shown below:

      In other cases, oxidation-reduction is accompanied by significant chemical rearrangement. An example is

      Two limiting mechanisms of electron transfer, commonly designated outer-sphere and inner-sphere mechanisms, have been recognized. Outer-sphere electron transfer occurs without dissociation or disruption of the coordination sphere of either complex—i.e., through both intact coordination spheres. The first reaction above is of this type. On the other hand, inner-sphere electron transfer—e.g., the second reaction above—proceeds by formation of a dinuclear complex in which the two metal ions are joined by a common bridging ligand (in this case the chloride ion) through which the electron is transferred. Such electron transfer also may occur through polyatomic bridging ligands to which the two metal ions are attached at different sites separated by several atoms, as in the reduction of pentaammine(isonicotinamide)cobalt(3+) by hexaaquachromium(2+) ion through a bridged intermediate:

      Strikingly large differences in rates of electron transfer are observed even between closely related reactions. Thus, the rate of reduction of the pentaamminebromocobalt(3+) ion by the hexaaquachromium(2+) ion is about 107 times higher than that of the acetatopentaamminecobalt(2+) ion by the same chromium ion.

Synthesis of coordination compounds
      The great variety of coordination compounds is matched by the diversity of methods through which such compounds can be synthesized. Complex halides, for example, may be prepared by direct combination of two halide salts (either in the molten state or in a suitable solvent). Palladium chloride and potassium chloride, for example, react to give the complex potassium tetrachloropalladate(2−), as shown in the following equation:

      Another widely used route to coordination compounds is through the direct combination of a metal ion and appropriate ligands in solution. Thus, the addition of a sufficiently high concentration of ammonia to an aqueous solution of a nickel(2+) salt leads, through a series of reactions (see above Aqua complexes (coordination compound)), to the formation of the hexaamminenickel(2+) ion, which can be precipitated, for example, as the sulfate salt, [Ni(NH3)6]SO4.

      Complexes of metal ions in high oxidation states are sometimes more readily formed by adding the ligands to a solution of the metal ion in a lower oxidation state in the presence of an oxidizing agent. Thus, addition of ammonia to an aqueous solution of a cobalt(2+) salt in the presence of air or oxygen leads to the formation of cobalt(3+)-ammine complexes such as hexaamminecobalt(3+), [Co(NH3)6]3+, and pentaammineaquacobalt(3+), [Co(NH3)5(H2O)]3+, ions.

      Complexes of metals in low oxidation states, such as the carbonyls of metals in their zero oxidation states, can sometimes be prepared by direct combination of the metal with the ligand, as, for example, in the reaction of nickel metal with carbon monoxide.

      More commonly, a salt of the metal is reduced in the presence of the ligand. An example of this type of synthesis is the reduction of cobalt carbonate with hydrogen in the presence of carbon monoxide to give bis(tetracarbonylcobalt).

      Similar procedures are applicable to the synthesis of metal sandwich compounds containing cyclopentadienyl and benzene ligands. Dibenzenechromium, for example, can be prepared from chromic chloride, benzene, and aluminum, as shown in the following equation.

      Hydrido complexes of transition metals can be prepared by reactions of suitable precursors either with molecular hydrogen or with suitable reducing agents such as hydrazine or sodium borohydride; for example,

      Transition metal complexes containing metal-carbon bonds can be prepared by a variety of routes, some of the more important of which are illustrated by the following examples (for further treatment of carbonyl synthesis, see organometallic compound).

George B. Kauffman Jack Halpern

Additional Reading

History
George B. Kauffman (trans. and ed.), Classics in Coordination Chemistry, 3 vol. (1968–78), contains annotated translations of the most important contributions from 1798 to 1935; George B. Kauffman (ed.), Werner Centennial (1967), is a collection of papers surveying historical and research aspects; and George B. Kauffman, Coordination Chemistry: A Century of Progress (1994), is a collection of historical, review, and research papers.

Chemistry
The chemistry of coordination compounds is covered in Jon A. McCleverty and Thomas Meyer (eds.), Comprehensive Coordination Chemistry II: From Biology to Nanotechnology, 10 vol. (2003), a compendium of synthesis, reactions, properties, and applications; Arthur E. Martell and Melvin Calvin, Chemistry of the Metal Chelate Compounds (1971), a text that describes the chelate effect; and Michael Thor Pope, Heteropoly and Isopoly Oxometalates, 2nd ed. (1987).George B. Kauffman

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Universalium. 2010.

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