COMPLEX COMPOUND IN BIOLOGICAL SYSTEMS

COMPLEX COMPOUND IN BIOLOGICAL SYSTEMS.

 

Coordination compound, any of a class of substances with chemical structures in which a central metal atom is surrounded bynonmetal atoms or groups of atoms, called ligands, joined to it by chemical bonds. Coordination compounds include such substances as vitamin B₁₂, hemoglobin, and chlorophyll, dyesand pigments, and catalysts used in preparing organic substances.

A major application of coordination compounds is their use as catalysts, which serve to alter the rate of chemical reactions. Certaincomplex metal catalysts, for example, play a key role in the production ofpolyethylene 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,  the articlesisomerism; coordinatioumber; chemical reaction; and organometallic compound.

Coordination compounds iature

Naturally occurring coordination compounds are vital to living organisms. Metal complexes play a variety of important roles in biological systems. Many enzymes, 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 acidresidues of the protein. Another enzyme, catalase, which is an efficient catalyst for the decomposition ofhydrogen 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 anoxygen 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 vitamin B₁₂, a complex of cobalt with a macrocyclic ligand known as corrin.

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).

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

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—may be neutral molecules such as water (in the above example),ammonia (NH₃), or carbon monoxide (CO); negatively charged ions (anions) such as the fluoride (in the first example above) or cyanide ion (CN⁻); or, occasionally, positively charged ions (cations) 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 nonmetallicelements. 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 undergooxidation-reduction reactions, 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 in passing from the neutral atom to the charged ion, sometimes referred to as the oxidatioumber), 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.

Coordinatioumber is the term proposed by Werner to denote the total number of bonds from the ligands to the metal atom. Coordinatioumbers 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. Coordinatioumber should be distinguished from oxidatioumber (defined in the previous paragraph). The oxidatioumber, designated by an Arabic number with an appropriate sign (or, sometimes, by a Romaumeral 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 coordinatioumber of cobalt is 6 while its oxidation number is +3.

Ligands and chelates

Each molecule or ion of a coordination compound includes a number of ligands, 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 (from Greek chele, “claw”), and their formation is termed chelation. Chelates are particularly stable and useful. An example of a typicalchelate is bis(1,2-ethanediamine)copper(2+), the complex formed between the cupric ion (Cu²⁺) and the organic compoundethylenediamine(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 (such as hydrogen, carbon, oxygen, nitrogen, and sulfur), with the exception of the noble gases (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 calciumcomplex 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 or citric 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 (I), [Nd(CF₃COCHCOCF₃)₄]⁻.

Certain ligands may be either monodentate or polydentate, depending on the particular compound in which they occur. The carbonateion, (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

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 ( above Polydentate), along with a unique metal-cluster complex having six metal atoms in its nucleus (organometallic compound).

Nomenclature

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:

1.  Neutral and cationic complexes are named by first identifying the ligands, followed by the metal; itsoxidatioumber may be given in Romaumerals 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 anioame 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. Iames (as opposed to formulas) the ligands are given in alphabetical order without regard to charge. The oxidatioumber 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):

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

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

4.  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.

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. A more complete description of coordinate bonding, in terms of electron pairs, became possible in the 1920s, following the introduction of the concept that all covalent bonds consist of electron pairs shared between atoms, an idea advanced chiefly by the American physical chemist Gilbert N. Lewis. 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:

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

Geometry

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 thecoordination 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 coordinatioumbers from 1 to 16 are known, those below 3 and above 8 are rare. Possible structures and examples of species for the various coordinatioumbers 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⁻).

The 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.

Coordinatioumbers 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 approaches but does not exceed 18 (the number of electrons in the valence shells of the noble gases) are most stable. The stabilities of 18-electron valence shells are also reflected in the coordinatioumbers of the stable mononuclear carbonyls of different metals that have oxidatioumber 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, andphosphines. For more ionic (also called outer-orbital) complexes, such as fluoro or aqua complexes, electronic factors are less important in determining coordinatioumbers, 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 coordinatioumber in different complexes – e.g., generally six for chromium(+3) – but this is not invariably so. Differences in coordinatioumber 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 coordinatioumbers—e.g., tetracyanonickelate [Ni(CN)₄]²⁻ and pentacyanonickelate [Ni(CN)₅]³⁻, both of which contain Ni in the +2 oxidation state.

Isomerism

Coordination compounds often exist as isomers—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:

Principal types of complexes: The tendency for complexes to form between a metal ion and a particular combination of ligands 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.

Many elements, notably certain metals, exhibit a range of oxidation states—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, 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.

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 dsubshells (i.e., transition metals), the coordinatioumbers 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 coordinatioumbers 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 (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 = [NiL_n(H₂O)_{6 −n}]²⁺+ nH₂O, in which L is the ligand and n the number of water molecules displaced from the complex. In this series the equilibrium constants, K_L, increase dramatically as the possibilities for chelation increase (that is, as the number of nitrogen atoms available for bonding to the metal atom increases).

It should be noted that, in the particular examples cited above, the coordinatioumber 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.