Coordination complex

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Cisplatin, PtCl2(NH3)2
A platinum atom with four ligands

In chemistry, a coordination complex or metal complex consists of a central atom or ion, which
is usually metallic and is called the coordination centre, and a surrounding array of bound
molecules or ions, that are in turn known as ligands or complexing agents.[1][2] Many metal-
containing compounds, especially those of transition metals, are coordination complexes.[3]

Contents
 [hide] 

 1 Nomenclature and terminology

 2 History
 3 Structures
o 3.1 Geometry
o 3.2 Isomerism
 3.2.1 Stereoisomerism
 3.2.1.1 Cis–trans isomerism and facial–meridional isomerism
 3.2.1.2 Optical isomerism
 3.2.2 Structural isomerism
 4 Electronic properties
o 4.1 Color of transition metal complexes
o 4.2 Colors of Lanthanide complexes
o 4.3 Magnetism
o 4.4 Reactivity
 5 Classification
o 5.1 Older classifications of isomerism
 6 Naming complexes
 7 Stability Constant
 8 Application of coordination compounds
  9 See also
 10 External links
 11 References

Nomenclature and terminology[edit]

Coordination complexes are so pervasive that the structure and reactions are described in many
ways, sometimes confusingly. The atom within a ligand that is bonded to the central atom or ion
is called the donor atom. In a typical complex, a metal ion is bound to several donor atoms,
which can be the same or different. Polydentate (multiple bonded) ligands consist of several
donor atoms, several of which are bound to the central atom or ion. These complexes are called
chelate complexes, the formation of such complexes is called chelation, complexation, and
coordination.

The central atom or ion, together with all ligands comprise the coordination sphere.[4][5] The
central atoms or ion and the donor atoms comprise the first coordination sphere.

Coordination refers to the "coordinate covalent bonds" (dipolar bonds) between the ligands and
the central atom. Originally, a complex implied a reversible association of molecules, atoms, or
ions through such weak chemical bonds. As applied to coordination chemistry, this meaning has
evolved. Some metal complexes are formed virtually irreversibly and many are bound together
by bonds that are quite strong.[6][7]

The number of ligands attached to a metal ion is called the coordination number. The most
common coordination numbers are 6, 4, 2. A hydrated ion is one kind of a complex ion (or,
simply, complex), a species formed between a central metal ion and one or more surrounding
ligands, molecules or ions that contain at least one lone pair of electrons,

For example, Cobalt(II) hexahydrated ion or Hexaaquacobolt(II) ion Co(H20) 62+, is a hydrated-
complex ion that consists of 6 ligands attached to a metal ion Co. The oxidation state and the
coordination number reflects the number of bonds formed between the metal ion and the ligands
in the complex ion.

History[edit]
Main article: Coordination complex

Coordination complexes have been a commonly known chemical substance since the early years
of chemistry. Its early uses included dyes such as Prussian blue. The first time that their function
first came to be understood however was in the late 1800s beginning in 1869 with the work of
Christian Wilhelm Blomstrand. Blomstrand developed what has come to be known as the
complex ion chain theory. The theory claimed that the reason coordination complexes form is
because in solution, ions would be bound via ammonia chains. He compared this effect to the
way that various carbohydrate chains form.
Alfred Werner

Following this theory, Danish scientist Sophus Mads Jorgensen made improvements to it. In his
version of the theory, Jorgensen claimed that when a molecule dissociates in a solution there
were two possible outcomes: the ions would bind via the ammonia chains Blomstrand had
described or the ions would bind directly to the metal.

It was not until 1893 that the most widely accepted version of the theory today was published by
Alfred Werner. Werner’s work included two important changes to the Blomstrand theory. The
first was that Werner described the two different ion possibilities in terms of location in the
coordination sphere. He claimed that if the ions were to form a chain this would occur outside of
the coordination sphere while the ions that bound directly to the metal would do so within the
coordination sphere.[8] In one of Werner’s most important discoveries however he disproved the
majority of the chain theory. Werner was able to discover the spatial arrangements of the ligands
that were involved in the formation of the complex hexacoordinate cobalt. His theory allows one
to understand the difference between coordinated and ionic in a compound, for example chloride
in the cobaltammine chlorides and to explain many of the previously inexplicable isomers.

Structure of hexol

In 1914, Werner resolved the first coordination complex, called hexol, into optical isomers,
overthrowing the theory that only carbon compounds could possess chirality.

Structures[edit]
The ions or molecules surrounding the central atom are called ligands. Ligands are generally
bound to the central atom by a coordinate covalent bond (donating electrons from a lone electron
pair into an empty metal orbital), and are said to be coordinated to the atom. There are also
organic ligands such as alkenes whose pi bonds can coordinate to empty metal orbitals. An
example is ethene in the complex known as Zeise's salt, K+[PtCl3(C2H4)]−.
Geometry[edit]

In coordination chemistry, a structure is first described by its coordination number, the number
of ligands attached to the metal (more specifically, the number of donor atoms). Usually one can
count the ligands attached, but sometimes even the counting can become ambiguous.
Coordination numbers are normally between two and nine, but large numbers of ligands are not
uncommon for the lanthanides and actinides. The number of bonds depends on the size, charge,
and electron configuration of the metal ion and the ligands. Metal ions may have more than one
coordination number.

Typically the chemistry of transition metal complexes is dominated by interactions between s

and p molecular orbitals of the ligands and the d orbitals of the metal ions. The s, p, and d
orbitals of the metal can accommodate 18 electrons (see 18-Electron rule). The maximum
coordination number for a certain metal is thus related to the electronic configuration of the
metal ion (to be more specific, the number of empty orbitals) and to the ratio of the size of the
ligands and the metal ion. Large metals and small ligands lead to high coordination numbers, e.g.
[Mo(CN)8]4−. Small metals with large ligands lead to low coordination numbers, e.g.
Pt[P(CMe3)]2. Due to their large size, lanthanides, actinides, and early transition metals tend to
have high coordination numbers.

Different ligand structural arrangements result from the coordination number. Most structures
follow the points-on-a-sphere pattern (or, as if the central atom were in the middle of a
polyhedron where the corners of that shape are the locations of the ligands), where orbital
overlap (between ligand and metal orbitals) and ligand-ligand repulsions tend to lead to certain
regular geometries. The most observed geometries are listed below, but there are many cases that
deviate from a regular geometry, e.g. due to the use of ligands of different types (which results in
irregular bond lengths; the coordination atoms do not follow a points-on-a-sphere pattern), due to
the size of ligands, or due to electronic effects (see, e.g., Jahn–Teller distortion):

 Linear for two-coordination

 Trigonal planar for three-coordination
 Tetrahedral or square planar for four-coordination
 Trigonal bipyramidal or square pyramidal for five-coordination
 Octahedral (orthogonal) or trigonal prismatic for six-coordination
 Pentagonal bipyramidal for seven-coordination
 Square antiprismatic for eight-coordination
 Tri-capped trigonal prismatic (Triaugmented triangular prism) for nine-coordination.

Some exceptions and provisions should be noted:

 The idealized descriptions of 5-, 7-, 8-, and 9- coordination are often not geometrically
distinct from alternative structures with slightly different L–M–L (ligand–metal–ligand)
angles. The classic example of this is the difference between square pyramidal and
trigonal bipyramidal structures.
 Due to special electronic effects such as (second-order) Jahn–Teller stabilization, certain
geometries are stabilized relative to the other possibilities, e.g. for some compounds the
 trigonal prismatic geometry is stabilized relative to octahedral structures for six-
coordination.

Isomerism[edit]

The arrangement of the ligands is fixed for a given complex, but in some cases it is mutable by a
reaction that forms another stable isomer.

There exist many kinds of isomerism in coordination complexes, just as in many other
compounds.

Stereoisomerism[edit]

Stereoisomerism occurs with the same bonds in different orientations relative to one another.
Stereoisomerism can be further classified into:

Cis–trans isomerism and facial–meridional isomerism[edit]

Cis–trans isomerism occurs in octahedral and square planar complexes (but not tetrahedral).
When two ligands are adjacent they are said to be cis, when opposite each other, trans. When
three identical ligands occupy one face of an octahedron, the isomer is said to be facial, or fac. In
a fac isomer, any two identical ligands are adjacent or cis to each other. If these three ligands and
the metal ion are in one plane, the isomer is said to be meridional, or mer. A mer isomer can be
considered as a combination of a trans and a cis, since it contains both trans and cis pairs of
identical ligands.

cis-[CoCl2(NH3)4]+

trans-[CoCl2(NH3)4]+
 

fac-[CoCl3(NH3)3]

mer-[CoCl3(NH3)3]

Optical isomerism[edit]

Optical isomerism occurs when a molecule is not superimposable with its mirror image. It is so
called because the two isomers are each optically active, that is, they rotate the plane of polarized
light in opposite directions. The symbol Λ (lambda) is used as a prefix to describe the left-
handed propeller twist formed by three bidentate ligands, as shown. Likewise, the symbol Δ
(delta) is used as a prefix for the right-handed propeller twist.[9]


Λ-[Fe(ox) ]3−
3

Δ-[Fe(ox)3]3−
 
Λ-cis-[CoCl (en) ]+
2 2

Δ-cis-[CoCl2(en)2]+

Structural isomerism[edit]

Structural isomerism occurs when the bonds are themselves different. There are four types of
structural isomerism: ionisation isomerism, solvate or hydrate isomerism, linkage isomerism and
coordination isomerism.

1. Ionisation isomerism – the isomers give different ions in solution although they have the
same composition. This type of isomerism occurs when the counter ion of the complex is
also a potential ligand. For example pentaamminebromidocobalt(III)sulfate
[CoBr(NH3)5]SO4 is red violet and in solution gives a precipitate with barium chloride,
confirming the presence of sulfate ion, while pentaamminesulfatecobalt(III)bromide
[CoSO4(NH3)5]Br is red and tests negative for sulfate ion in solution, but instead gives a
precipitate of AgBr with silver nitrate.
2. Solvate or hydrate isomerism – the isomers have the same composition but differ with
respect to the number of solvent ligand molecules as well as the counter ion in the crystal
lattice. For example [Cr(H2O)6]Cl3 is violet colored, [CrCl(H2O)5]Cl2·H2O is blue-green,
and [CrCl2(H2O)4]Cl·2H2O is dark green
3. Linkage isomerism occurs with ambidentate ligands that can bind in more than one
place. For example, NO2 is an ambidentate ligand: It can bind to a metal at either the
N atom or an O atom.
4. Coordination isomerism – this occurs when both positive and negative ions of a salt are
complex ions and the two isomers differ in the distribution of ligands between the cation
and the anion. For example [Co(NH3)6][Cr(CN)6] and [Cr(NH3)6][Co(CN)6]

Electronic properties[edit]
Many of the properties of transition metal complexes are dictated by their electronic structures.
The electronic structure can be described by a relatively ionic model that ascribes formal charges
to the metals and ligands. This approach is the essence of crystal field theory (CFT). Crystal field
theory, introduced by Hans Bethe in 1929, gives a quantum mechanically based attempt at
understanding complexes. But crystal field theory treats all interactions in a complex as ionic and
assumes that the ligands can be approximated by negative point charges.

More sophisticated models embrace covalency, and this approach is described by ligand field
theory (LFT) and Molecular orbital theory (MO). Ligand field theory, introduced in 1935 and
built from molecular orbital theory, can handle a broader range of complexes and can explain
complexes in which the interactions are covalent. The chemical applications of group theory can
aid in the understanding of crystal or ligand field theory, by allowing simple, symmetry based
solutions to the formal equations.

Chemists tend to employ the simplest model required to predict the properties of interest; for this
reason, CFT has been a favorite for the discussions when possible. MO and LF theories are more
complicated, but provide a more realistic perspective.

The electronic configuration of the complexes gives them some important properties:

Color of transition metal complexes[edit]

Synthesis of copper(II)-tetraphenylporphyrin, a metal complex, from tetraphenylporphyrin and

copper(II) acetate monohydrate.

Transition metal complexes often have spectacular colors caused by electronic transitions by the
absorption of light. For this reason they are often applied as pigments. Most transitions that are
related to colored metal complexes are either d–d transitions or charge transfer bands. In a d–
d transition, an electron in a d orbital on the metal is excited by a photon to another d orbital of
higher energy. A charge transfer band entails promotion of an electron from a metal-based
orbital into an empty ligand-based orbital (Metal-to-Ligand Charge Transfer or MLCT). The
converse also occurs: excitation of an electron in a ligand-based orbital into an empty metal-
based orbital (Ligand to Metal Charge Transfer or LMCT). These phenomena can be observed
with the aid of electronic spectroscopy; also known as UV-Vis.[10] For simple compounds with
high symmetry, the d–d transitions can be assigned using Tanabe–Sugano diagrams. These
assignments are gaining increased support with computational chemistry.

Colours of Various Example Coordination Complexes

  Fe2+ Fe3+ Co2+ Cu2+ Al3+ Cr3+
[Fe(H2O)6]3+
[Fe(H2O)6]2+ [Co(H2O)6]2+ [Cu(H2O)6]2+ [Al(H2O)6]3+ [Cr(H2O)6]3+
Hydrated Yellow/brow
Pale green Pink Blue Colourless Green
Ion n
Solution Solution Solution Solution Solution
Solution
[Fe(H2O)4(O [Fe(H2O)3(O [Co(H2O)4(O [Cu(H2O)4(OH) [Al(H2O)3(O [Cr(H2O)3(O
OH ,−
H)2] H)3] H)2] 2] H)3] H)3]
dilute Dark green Brown Blue/green Blue White Green
Precipitate Precipitate Precipitate Precipitate Precipitate Precipitate
[Fe(H 2O) 4(O [Fe(H 2O) 3(O [Co(H 2 O) 4 (O [Cu(H 2O) 4(OH)
OH−, [Al(OH)4]− [Cr(OH)6]3−
H)2] H)3] H)2] 2]
concentrat Colourless Green
Dark green Brown Blue/green Blue
ed Solution Solution
Precipitate Precipitate Precipitate Precipitate
[Fe(H2O)4(O [Fe(H2O)3(O [Co(H2O)4(O [Cu(H2O)4(OH) [Al(H2O)3(O [Cr(H2O)3(O
NH3, H)2] H)3] H)2] 2] H)3] H)3]
dilute Dark green Brown Blue/green Blue White Green
Precipitate Precipitate Precipitate Precipitate Precipitate Precipitate
[Fe(H2O)4(O [Fe(H2O)3(O [Co(NH3)6] [Cu(NH3)4(H2O [Al(H2O)3(O
2+
NH3, [Cr(NH3)6]3+
H) ] H)3] Straw )2]2+ H)3]
concentrat 2 Green
Dark green Brown coloured Deep blue White
ed Solution
Precipitate Precipitate Solution Solution Precipitate
[Fe(H2O)3(O
FeCO3 H)3] CoCO3 CuCO3
CO3 2−
Dark green Brown Pink Blue/green
Precipitate Precipitate + Precipitate Precipitate
bubbles

Colors of Lanthanide complexes[edit]

Superficially lanthanide complexes are similar to those of the transition metals in that some are
coloured. However for the common Ln3+ ions (Ln = lanthanide) the colors are all pale, and hardly
influenced by the nature of the ligand. The colors are due to 4f electron transitions. As the 4f
orbitals in lanthanides are “buried” in the xenon core and shielded from the ligand by the 5s and
5p orbitals they are therefore not influenced by the ligands to any great extent leading to a much
smaller crystal field splitting than in the transition metals. The absorption spectra of an Ln 3+ ion
approximates to that of the free ion where the electronic states are described by spin-orbit
coupling (also called L-S coupling or Russell-Saunders coupling). This contrasts to the transition
metals where the ground state is split by the crystal field. Absorptions for Ln 3+ are weak as
electric dipole transitions are parity forbidden (Laporte Rule forbidden) but can gain intensity
due to the effect of a low-symmetry ligand field or mixing with higher electronic states (e.g. d
orbitals). Also absorption bands are extremely sharp which contrasts with those observed for
transition metals which generally have broad bands.[11][12] This can lead to extremely unusual
effects, such as significant color changes under different forms of lighting.

Magnetism[edit]

Main article: magnetochemistry

Metal complexes that have unpaired electrons are magnetic. Considering only monometallic
complexes, unpaired electrons arise because the complex has an odd number of electrons or
because electron pairing is destabilized. Thus, monomeric Ti(III) species have one "d-electron"
and must be (para)magnetic, regardless of the geometry or the nature of the ligands. Ti(II), with
two d-electrons, forms some complexes that have two unpaired electrons and others with none.
This effect is illustrated by the compounds TiX 2[(CH3)2PCH2CH2P(CH3)2]2: when X = Cl, the
complex is paramagnetic (high-spin configuration), whereas when X = CH3, it is diamagnetic
(low-spin configuration). It is important to realize that ligands provide an important means of
adjusting the ground state properties.

In bi- and polymetallic complexes, in which the individual centers have an odd number of
electrons or that are high-spin, the situation is more complicated. If there is interaction (either
direct or through ligand) between the two (or more) metal centers, the electrons may couple
(antiferromagnetic coupling, resulting in a diamagnetic compound), or they may enhance each
other (ferromagnetic coupling). When there is no interaction, the two (or more) individual metal
centers behave as if in two separate molecules.

Reactivity[edit]

Complexes show a variety of possible reactivities:

 Electron transfers

A common reaction between coordination complexes involving ligands are inner and
outer sphere electron transfers. They are two different mechanisms of electron transfer
redox reactions, largely defined by the late Henry Taube. In an inner sphere reaction, a
ligand with two lone electron pairs acts as a bridging ligand, a ligand to which both
coordination centres can bond. Through this, electrons are transferred from one centre to
another.

 (Degenerate) ligand exchange

One important indicator of reactivity is the rate of degenerate exchange of ligands. For
example, the rate of interchange of coordinate water in [M(H 2O)6]n+ complexes varies over
20 orders of magnitude. Complexes where the ligands are released and rebound rapidly
are classified as labile. Such labile complexes can be quite stable thermodynamically.
Typical labile metal complexes either have low-charge (Na +), electrons in d-orbitals that
are antibonding with respect to the ligands (Zn2+), or lack covalency (Ln3+, where Ln is
 any lanthanide). The lability of a metal complex also depends on the high-spin vs. low-
spin configurations when such is possible. Thus, high-spin Fe(II) and Co(III) form labile
complexes, whereas low-spin analogues are inert. Cr(III) can exist only in the low-spin
state (quartet), which is inert because of its high formal oxidation state, absence of
electrons in orbitals that are M–L antibonding, plus some "ligand field stabilization"
associated with the d3 configuration.

 Associative processes

Complexes that have unfilled or half-filled orbitals often show the capability to react with
substrates. Most substrates have a singlet ground-state; that is, they have lone electron
pairs (e.g., water, amines, ethers), so these substrates need an empty orbital to be able to
react with a metal centre. Some substrates (e.g., molecular oxygen) have a triplet ground
state, which results that metals with half-filled orbitals have a tendency to react with such
substrates (it must be said that the dioxygen molecule also has lone pairs, so it is also
capable to react as a 'normal' Lewis base).

If the ligands around the metal are carefully chosen, the metal can aid in (stoichiometric or
catalytic) transformations of molecules or be used as a sensor.

Classification[edit]
Metal complexes, also known as coordination compounds, include all metal compounds, aside
from metal vapors, plasmas, and alloys. The study of "coordination chemistry" is the study of
"inorganic chemistry" of all alkali and alkaline earth metals, transition metals, lanthanides,
actinides, and metalloids. Thus, coordination chemistry is the chemistry of the majority of the
periodic table. Metals and metal ions exist, in the condensed phases at least, only surrounded by
ligands.

The areas of coordination chemistry can be classified according to the nature of the ligands, in
broad terms:

 Classical (or "Werner Complexes"): Ligands in classical coordination chemistry bind to

metals, almost exclusively, via their "lone pairs" of electrons residing on the main group
atoms of the ligand. Typical ligands are H2O, NH3, Cl−, CN−, en. Some of the simplest
members of such complexes are described in metal aquo complexes, metal ammine
complexes,

Examples: [Co(EDTA)]−, [Co(NH3)6]Cl3, [Fe(C2O4)3]K3

 Organometallic Chemistry: Ligands are organic (alkenes, alkynes, alkyls) as well as

"organic-like" ligands such as phosphines, hydride, and CO.

Example: (C5H5)Fe(CO)2CH3
  Bioinorganic Chemistry: Ligands are those provided by nature, especially including the
side chains of amino acids, and many cofactors such as porphyrins.

Example: hemoglobin contains heme, a porphyrin complex of iron

Example: chlorophyll contains a porphyrin complex of magnesium
Many natural ligands are "classical" especially including water.

 Cluster Chemistry: Ligands are all of the above also include other metals as ligands.

Example Ru3(CO)12

 In some cases there are combinations of different fields:

Example: [Fe4S4(Scysteinyl)4]2−, in which a cluster is embedded in a biologically active species.

Mineralogy, materials science, and solid state chemistry – as they apply to metal ions – are
subsets of coordination chemistry in the sense that the metals are surrounded by ligands. In many
cases these ligands are oxides or sulfides, but the metals are coordinated nonetheless, and the
principles and guidelines discussed below apply. In hydrates, at least some of the ligands are
water molecules. It is true that the focus of mineralogy, materials science, and solid state
chemistry differs from the usual focus of coordination or inorganic chemistry. The former are
concerned primarily with polymeric structures, properties arising from a collective effects of
many highly interconnected metals. In contrast, coordination chemistry focuses on reactivity and
properties of complexes containing individual metal atoms or small ensembles of metal atoms.

Older classifications of isomerism[edit]

Traditional classifications of the kinds of isomer have become archaic with the advent of modern
structural chemistry. In the older literature, one encounters:

 Ionisation isomerism describes the possible isomers arising from the exchange between
the outer sphere and inner sphere. This classification relies on an archaic classification of
the inner and outer sphere. In this classification, the "outer sphere ligands," when ions in
solution, may be switched with "inner sphere ligands" to produce an isomer.
 Solvation isomerism occurs when an inner sphere ligand is replaced by a solvent
molecule. This classification is obsolete because it considers solvents as being distinct
from other ligands. Some of the problems are discussed under water of crystallization.

Naming complexes[edit]
The basic procedure for naming a complex:

1. When naming a complex ion, the ligands are named before the metal ion.
2. Write the names of the ligands in alphabetical order. (Numerical prefixes do not affect
the order.)
 o Multiple occurring monodentate ligands receive a prefix according to the number
of occurrences: di-, tri-, tetra-, penta-, or hexa. Polydentate ligands (e.g.,
ethylenediamine, oxalate) receive bis-, tris-, tetrakis-, etc.
o Anions end in ido. This replaces the final 'e' when the anion ends with '-ate', e.g.
sulfate becomes sulfato. It replaces 'ide': cyanide becomes cyanido.
o Neutral ligands are given their usual name, with some exceptions: NH 3 becomes
ammine; H2O becomes aqua or aquo; CO becomes carbonyl; NO becomes
nitrosyl.
3. Write the name of the central atom/ion. If the complex is an anion, the central atom's
name will end in -ate, and its Latin name will be used if available (except for mercury).

Ligands Coding name Old name

H20 aqua aquo
NH3 ammine ammino
OH- hydroxo hydroxy
Cl- chlorido chloro
F- fluorido fluoro
CN- cyanido cyano

1. If the central atom's oxidation state needs to be specified (when it is one of several
possible, or zero), write it as a Roman numeral (or 0) in parentheses.
2. Name cation then anion as separate words (if applicable, as in last example)

Examples:

metal changed to
cobalt cobaltate
aluminium aluminate
chromium chromate
vanadium vanadate
copper cuprate
iron ferrate
[NiCl4]2− → tetrachloridonickelate(II) ion
[CuCl5NH3]3− → amminepentachloridocuprate(II) ion
[Cd(CN)2(en)2] → dicyanidobis(ethylenediamine)cadmium(II)
[CoCl(NH3)5]SO4 → pentaamminechloridocobalt(III) sulfate

The coordination number of ligands attached to more than one metal (bridging ligands) is
indicated by a subscript to the Greek symbol μ placed before the ligand name. Thus the dimer of
aluminium trichloride is described by Al2Cl4(μ2-Cl)2.

Stability Constant[edit]
Main article: Stability constants of complexes
The affinity of metal ions for ligands is described by stability constant. This constant, also
referred to as the formation constant, is given the notation of K f and can be calculated through
the following method for simple cases:

(X)Metal(aq)+(Y)Lewis Base(aq) = (Z)Complex Ion(aq)

Kf = [Complex Ion]Z / [Metal]X[Lewis Base]Y

Formation constant vary widely. Large values indicate that the metal has high affinity for the
ligand, provided the system is at equilibrium.[13]

Sometimes the stability constant will be in a different form known as the constant of destability.
This constant is expressed as the inverse of the constant of formation and is denoted as Kd =
1/Kf .[14] This constant represents the reverse reaction for the decomposition of a complex ion into
its individual metal and ligand components. When comparing the values for Kd, the larger the
value is the more unstable the complex ion is.

As a result of these complex ions forming in solutions they also can play a key role in solubility
of other compounds. When a complex ion is formed it can alter the concentrations of its
components in the solution. For example:

Ag(aq)++2NH4OH(aq) = Ag(NH3)2++H2O

AgClS+H2Ol = Ag(aq)++Cl(aq)-

In these reactions which both occurred in the same reaction vessel, the solubility of the silver
chloride would be increased as a result of the formation of the complex ion. The complex ion
formation is favorable takes away a significant portion of the silver ions in solution, as a result
the equilibrium for the formation of silver ions from silver chloride will shift to the right to make
up for the deficit.

This new solubility can be calculated given the values of Kf and Ksp for the original reactions.
The solubility is found essentially by combining the two separate equilibria into one combined
equilibrium reaction and this combined reaction is the one that determines the new solubility. So
Kc , the new solubility constant, is denoted by Kc = Ksp * Kf.

Application of coordination compounds[edit]

Metals only exist in solution as coordination complexes, it follows then that this class of
compounds are useful. Coordination compounds are found both in the natural world and
artificially in industry. Some common complex ions include such substances as vitamin B12 ,
hemoglobin , chlorophyll, and some dyes and pigments. One major use of coordination
compounds is in homogeneous catalysis for the production of organic substances.
Coordination compounds have uses in both nature and in industry. Coordination compounds are
vital to many living organisms. For example many enzymes are metal complexes, like
carboxypeptidase, a hydrolytic enzyme important in digestion. This enzyme consists of a zinc
ion surrounded by many amino acid residues. Another complex ion enzyme is catalase, which
decomposes the cell waste hydrogen peroxide. This enzyme contains iron-porphyrin complexes,
similar to that in hemoglobin. Chlorophyll contains a magnesium-porphyrin complexes, and
vitamin B12 is a complex with cobalt and corrin.

Coordination compounds are also widely used in industry. The intense colors of many
compounds render them of great use as dyes and pigments. Specifically Phthalocyanine
complexes are an important class of dyes for fabrics. Nickel, cobalt, and copper can be extracted
using hydrometallurgical processes involving complex ions. They are extracted from their ores
as ammine complexes with aqueous ammonia. Metals can also be separated using the selective
precipitation and solubility of complex ions, as explained in later sections. Cyanide complexes
are often used in electroplating.

Coordination compounds can also be used to identify unknown substances in a solution. This
analysis can be done by utilizing the selective precipitation of the complex ions, the formation of
color complexes which can be measured spectrophotometrically, or the preparation of
complexes, such as metal acetylacetonates, which can be separated with organic solvents.

A combination of titanium trichloride and triethylaluminum brings about the polymerization of

organic compounds with carbon-carbon double bonds to form polymers of high molecular
weight and ordered structures. Many of these polymers are of great commercial importance
because they are used in common fibers, films, and plastics.

Other common uses of coordination compounds in industry include the following:

1. They are used in photography, i.e., AgBr forms a soluble complex with sodium
thiosulfate in photography.
2. K[Ag(CN)2] is used forelectroplating of silver, and K[Au(CN)2] is used for gold plating.
3. Some ligands oxidise Co2+ to Co3+ ion.
4. Ethylenediaminetetraacetic acid (EDTA) is used for estimation of Ca2+ and Mg2+ inhard
water.
5. Silver and gold are extracted by treating zinc with their cyanide complexes

See also[edit]

Wikimedia Commons has media related to Coordination compounds.

 IUPAC nomenclature of inorganic chemistry

 Coordination geometry
  Inclusion compounds
 Organometallic chemistry deals with a special class of coordination compounds where
organic fragments are bonded to a metal at least through one C atom.
 Metallaprism
 Coordination polymers, in which coordination complexes are the repeating units.

External links[edit]

The Wikibook A-level Chemistry/OCR (Salters) has a page on the topic of: Complexes

 Naming Coordination Compounds

 Transition Metal Complex Colors

References[edit]
1. Jump up ^ IUPAC, Compendium of Chemical Terminology, 2nd ed. (the "Gold Book") (1997).
Online corrected version:  (2006–) "complex".
2. Jump up ^ IUPAC, Compendium of Chemical Terminology, 2nd ed. (the "Gold Book") (1997).
Online corrected version:  (2006–) "coordination entity".
3. Jump up ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd
ed.). Butterworth-Heinemann. ISBN 0080379419.
4. Jump up ^ chemistry-dictionary.com - Definition of coordination sphere
5. Jump up ^ What Is A Coordination Compound?
6. Jump up ^ Cotton, Frank Albert; Geoffrey Wilkinson; Carlos A. Murillo (1999). Advanced
Inorganic Chemistry. p. 1355. ISBN 978-0-471-19957-1.
7. Jump up ^ Miessler, Gary L.; Donald Arthur Tarr (1999). Inorganic Chemistry. p. 642.
ISBN 978-0-13-841891-5.
8. Jump up ^ "Coordination Compound".
9. Jump up ^ Miessler, Gary L.; Donald Arthur Tarr (1999). "9". Inorganic Chemistry. pp. 315,
316. ISBN 978-0-13-841891-5.
10. Jump up ^ Harris, D., Bertolucci, M., Symmetry and Spectroscopy. 1989 New York, Dover
Publications
11. Jump up ^ Cotton, F. Albert; Wilkinson, Geoffrey; Murillo, Carlos A.; Bochmann, Manfred
(1999), Advanced Inorganic Chemistry (6th ed.), New York: Wiley-Interscience, ISBN 0-471-
19957-5
12. Jump up ^ Cotton, Simon (2006). Lanthanide and Actinide Chemistry. John Wiley & Sons Ltd.
13. Jump up ^ "Complex Ion Equilibria".
14. Jump up ^ Stretton, Tom. "Solubility and Complex-ion Equilibria" (PDF).

 De Vito, D.; Weber, J. ; Merbach, A. E. “Calculated Volume and Energy Profiles for
Water Exchange on t2g 6 Rhodium(III) and Iridium(III) Hexaaquaions: Conclusive
Evidence for an Ia Mechanism” Inorganic Chemistry, 2004, Volume 43, pages 858–863.
doi:10.1021/ic035096n
  Zumdahl, Steven S. Chemical Principles, Fifth Edition. New York: Houghton Mifflin,
2005. 943–946, 957. OCLC 77760970
 Harris, D., Bertolucci, M., Symmetry and Spectroscopy. 1989 New York, Dover
Publications

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