The materials to prepare students for practical lessons of inorganic chemistry

The materials to prepare students for practical lessons of inorganic chemistry

LESSON № 12.

Theme: COMPLEX COMPOUND

Plan

 

Reactions of complex-formation. Co-ordinating theory of A. Werner and modern pictures of complex compound structure. Concept about complex-formatters ion (central ion). Nature, co-ordinating number, orbital hybridization of complex-formatters ion. Concept about ligands. Co-ordinating capacity (dentity) of ligands. Internal and external spheres of complexes. Geometry of complex ion. Nature of chemical bonds in complex compounds. Classification of complex compounds by charge of internal sphere and by nature of ligands. Polynuclease complexes.

Iron-, cobalt-, magnesium- and zinc-contain biocomplex compounds. Concept about a metaloligand homoeostasis. Complexons and their using in medicine as toxicides at poisoning by heavy metals (chelatotherapy) and as antioxidants at storage of medicinal preparations.

Introduction

Coordination chemistry emerged from the work of Alfred Werner, a Swiss chemist who examined different compounds composed of cobalt(III) chloride and ammonia. Upon the addition of hydrochloric acid, Werner observed that ammonia could not be completely removed. He then proposed that the ammonia must be bound more tightly to the central cobalt ion. However, when aqueous silver nitrate was added, one of the products formed was solid silver chloride. The amount of silver chloride formed was related to the number of ammonia molecules bound to the cobalt(III) chloride. For example, when silver nitrate was added to CoCl₃·6NH_3, all three chlorides were converted to silver chloride. However, when silver nitrate was added to CoCl₃·5NH₃, only 2 of the 3 chlorides formed silver chloride. When CoCl₃·4NH₃ was treated with silver nitrate, one of the three chlorides precipitated as silver chloride.

The resulting observations suggested the formation of complex or coordination compounds. In the inner coordination sphere, which is also referred to in some texts as the first sphere, ligands are directly bound to the central metal. In the outer coordination sphere, sometimes referred to as the second sphere, other ions are attached to the complex ion. Werner was awarded the Nobel Prize in 1913 for his coordination theory. The following table is a summary of Werner’s observations:

As the table above shows, the complex ion [Co(NH₃)₆]³⁺ is countered by the three chloride ions. The multi-level binding of coordination complexes play an important role in determining the dissociation of these complexes in aqueous solution. For example, [Co(NH₃)₅Cl]²⁺(Cl^–)₂ dissociates into 3 ions while [Co(NH₃)₄Cl₂ ]⁺(Cl^–) dissociates into 2 ions.ot further dissociate. By applying a current through the aqueous solutions of the resulting complex compounds, Werner measured the electrical conductivity and thus the dissociation properties of the complex compounds. The results confirmed his hypothesis of the formation of complex compounds. It is important to note that the above compounds have a coordinatioumber of 6, which is a common coordination number for many inorganic complexes. Coordinatioumbers for complex compounds typically range from 1 to 16.

Properties of Coordination Complexes

Some methods of verifying the presence of complex ions include studying its chemical behavior. This can be achieved by observing the compounds’ color, solubility, absorption spectrum, magnetic properties, etc. The properties of complex compounds are separate from the properties of the individual atoms. By forming coordination compounds, the properties of both the metal and the ligand are altered.

Metal-ligand bonds are typically thought of Lewis acid-base interactions. The metal atom acts as an electron pair acceptor (Lewis acid), while the ligands act as electron pair donors (Lewis base). The nature of the bond between metal and ligand is stronger thanintermolecular forces because they form directional bonds between the metal ion and the ligand, but are weaker than covalent bonds and ionic bonds.

Common Ligands

Monodentate ligands donate one pair of electrons to the central metal atoms. An example of these ligands are the haldide ions (F^–, Cl^–, Br^–, I^–). Polydentate ligands, also called chelates or chelating agents, donate more than one pair of electrons to the metal atom forming a stronger bond and a more stable complex. A common chelating agent is ethylenediamine (en), which, as the name suggests, contains two ammines or :NH₂ sites which can bind to two sites on the central metal. An example of a tridentate ligand is bis-diethylenetriammine. An example of such a coordination complex is bis-diethylenetriamine cobalt ^III.

 

Complex ions can form many compounds by binding with other complex ions in multiple ratios. This leads to many combinations of coordination compounds. The structures of certain coordination compounds can also have isomers, which can change their interactions with other chemical agents. The binding between metal and ligands is studied in metals, tetrahedral, and octahedral structures. There are many pharmaceutical and biological applications of coordination complexes and their isomers.

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’s description in 1597 of the blue colour (due to [Cu(NH₃)₄]²⁺) formed when lime water containing sal ammoniac (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. 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 foundatioot only for the Blomstrand-Jørgensen chain theory but for Alsatian-born Swiss chemist Alfred Werner’s coordination theory (1893) as well.

Blomstrand proposed that ammonia molecules could link together as −NH₃− chains, similar to −CH₂− chains in hydrocarbons. 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 coordinatioumber. Jørgensen proposed that atoms or groups that dissociated into ions 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) 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 (see below Isomerism) 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) 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.

 

 

Coordination chemistry is the study of compounds formed between metal ions and other neutral or negatively charged molecules such as [Co(NH₂CH₂CH₂NH₂)₂ClNH₃]²⁺ Cl₂^2-. In this formulation, Co(NH₂CH₂CH₂NH₂)₂ClNH₃]²⁺ is known as a metal complex, which is a charged species consisting of metal ion bonded to one or more groups of molecules. The bonded molecules are called ligand. The little picture shown here depicts a structure of a 6-coordinated complex.

A common metal complex is Ag(NH₃)₂⁺, formed when Ag⁺ ions are mixed with neutral ammonia molecules.

Ag⁺ + 2 NH₃ -> Ag(NH₃)₂⁺

A complex Ag(S₂O₃)₂^3- is formed between silver ions and negative thiosulfate ions:

Ag⁺ + 2 S₂O₃^2- -> Ag(S₂O₃)₂^3-

Metal complexes are also called coordination compounds. Their structures are important data and properties. Compounds having the same chemical formula but different structures are called isomers. Isomers with different geometic arrangements of ligands are called geometric isomers whereas isomers whose structures are mirror images of each other are called optical isomers. When a beam of polarized light passes optical isomers or their solutions, the plane of polarization rotates in different directions. The beam rotates to the left for one isomer, and right for its mirror image.

How did the study of coordination compounds started?

The structures of the complexes were proposed based on a coordination sphere of 6. The 6 ligands can be amonia molecules or chloride ions. Two different structures were proposed for the last two compounds, thetrans compound has two chloride ions on opposit vertices of an octahedral, whereas the the two chloride ions are adjacent to each other in the cis compound. The cis and trans compounds are known as geometric isomers.

Other cobalt complexes studied by Werner are also interesting. It has been predicted that the complex Co(NH₂CH₂CH₂NH₂)₂ClNH₃]²⁺ should exist in two forms, which are mirror images of each other. Werner isolated solids of the two forms, and structural studies confirmed his interpretations. The ligand NH₂CH₂CH₂NH₂ is ethylenediamine (en) often represented by en.

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 and John H. Van Vleck 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) that appear within the molecules. Although this approach meets with considerable success for complexes of metal ions with small electronegative 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.

At almost exactly the same time that chemists were developing the valence-bond model for coordination complexes, physicists such as Hans Bethe, John Van Vleck, and Leslie Orgel were developing an alternative known as crystal field theory. This theory tried to describe the effect of the electrical field of neighboring ions on the energies of the valence orbitals of an ion in a crystal. Crystal field theory was developed by considering two compounds: manganese(II) oxide, MnO, and copper(I) chloride, CuCl.

 

Octahedral Crystal Fields

Each Mn²⁺ ion in manganese(II) oxide is surrounded by six O^2- ions arranged toward the corners of an octahedron, as shown in the figure below. MnO is therefore a model for an octahedral complex in which a transition-metal ion is coordinated to six ligands.

What happens to the energies of the 4s and 4p orbitals on an Mn²⁺ion when this ion is buried in an MnO crystal? Repulsion between electrons that might be added to these orbitals and the electrons on the six O^2- ions that surround the metal ion in MnO increase the energies of these orbitals. The three 4p orbitals are still degenerate, however. These orbitals still have the same energy because each 4porbital points toward two O^2- ions at the corners of the octahedron.

Repulsion between electrons on the O^2- ions and electrons in the 3dorbitals on the metal ion in MnO also increases the energy of these orbitals. But the five 3d orbitals on the Mn²⁺ ion are no longer degenerate. Let’s assume that the six O^2- ions that surround each Mn²⁺ ion define an XYZ coordinate system. Two of the 3d orbitals (3d_x²_-y² and 3d_z²) on the Mn²⁺ ion point directly toward the six O^2-ions, as shown in the figure below. The other three orbitals (3d_xy, 3d_xz, and 3d_yz) lie between the O^2- ions.

The energy of the five 3d orbitals increases when the six O^2- ions are brought close to the Mn²⁺ ion. However, the energy of two of these orbitals (3d_x²_-y² and 3d_z²) increases much more than the energy of the other three (3d_xy, 3d_xz, and 3d_yz), as shown in the figure below. The crystal field of the six O^2- ions in MnO therefore splits the degeneracy of the five 3d orbitals. Three of these orbitals are now lower in energy than the other two.

By convention, the d_xy, d_xz, and d_yz orbitals in an octahedral complex are called the t_2g orbitals. The d_x²_-y² and d_z² orbitals, on the other hand, are called the e_g orbitals.

The easiest way to remember this convention is to note that there are three orbitals in the t_2g set.

The difference between the energies of the t_2g and e_g orbitals in an octahedral complex is represented by the symbol  _o. This splitting of the energy of the d orbitals is not trivial;  _o for the Ti(H₂O)₆³⁺ ion, for example, is 242 kJ/mol.

The magnitude of the splitting of the t_2g and e_g orbitals changes from one octahedral complex to another. It depends on the identity of the metal ion, the charge on this ion, and the nature of the ligands coordinated to the metal ion.

Tetrahedral Crystal Fields

Each Cu⁺ ion in copper(I) chloride is surrounded by four Cl^– ions arranged toward the corners of a tetrahedron, as shown in the figure below. CuCl is therefore a model for a tetrahedral complex in which a transition-metal ion is coordinated to four ligands.

Once again, the negative ions in the crystal split the energy of the datomic orbitals on the transition-metal ion. The tetrahedral crystal field splits these orbitals into the same t_2g and e_g sets of orbitals as does the octahedral crystal field.

But the two orbitals in the e_g set are now lower in energy than the three orbitals in the t_2g set, as shown in the figure below.

 

To understand the splitting of d orbitals in a tetrahedral crystal field, imagine four ligands lying at alternating corners of a cube to form a tetrahedral geometry, as shown in the figure below. The d_x²_-y² andd_z² orbitals on the metal ion at the center of the cube lie between the ligands, and the d_xy, d_xz, and d_yz orbitals point toward the ligands. As a result, the splitting observed in a tetrahedral crystal field is the opposite of the splitting in an octahedral complex.

Because a tetrahedral complex has fewer ligands, the magnitude of the splitting is smaller. The difference between the energies of the t_2gand e_g orbitals in a tetrahedral complex (∆_t) is slightly less than half as large as the splitting in analogous octahedral complexes (∆_o).

∆t = 4/9 ∆o

 

Square-Planar Complexes

The crystal field theory can be extended to square-planar complexes, such as Pt(NH₃)₂Cl₂. The splitting of the d orbitals in these compounds is shown in the figure below.

The Spectrochemical Series

The splitting of d orbitals in the crystal field model not only depends on the geometry of the complex, it also depends on the nature of the metal ion, the charge on this ion, and the ligands that surround the metal. When the geometry and the ligands are held constant, this splitting decreases in the following order.

Metal ions at one end of this continuum are called strong-field ions, because the splitting due to the crystal field is unusually strong. Ions at the other end are known as weak-field ions.

When the geometry and the metal are held constant, the splitting of the d orbitals decreases in the following order.

CO

CN–

>

NO2–

>

NH3

>

-NCS–

>

H2O

>

OH–

F–

-SCN–

Cl–

>

Br–

strong-field ligands

weak-field ligands

Ligands that give rise to large differences between the energies of the t_2g and e_g orbitals are called strong-field ligands. Those at the opposite extreme are known as weak-field ligands.

Because they result from studies of the absorption spectra of transition-metal complexes, these generalizations are known as thespectrochemical series. The range of values of   for a given geometry is remarkably large. The value of  _o is 100 kJ/mol in the Ni(H₂O)₆²⁺ ion, for example, and 520 kJ/mol in the Rh(CN)₆^3- ion.

High-Spin Versus Low-Spin Octahedral Complexes

Once we know the relative energies of the d orbitals in a transition-metal complex, we have to worry about how these orbitals are filled. Degenerate orbitals are filled according to Hund’s rules.

·      One electron is added to each of the degenerate orbitals in a subshell before a second electron is added to any orbital in the subshell.

·      Electrons are added to a subshell with the same value of the spin quantum number until each orbital in the subshell has at least one electron.

Octahedral transition-metal ions with d¹, d², or d³ configurations can therefore be described by the following diagrams.

 

When we try to add a fourth electron, we are faced with a problem. This electron could be used to pair one of the electrons in the lower energy (t_2g) set of orbitals or it could be placed in one of the higher energy (e_g) orbitals. One of these configurations is called high-spinbecause it contains four unpaired electrons with the same spin. The other is called low-spin because it contains only two unpaired electrons. The same problem occurs with octahedral d⁵, d⁶, and d⁷complexes.

For octahedral d⁸, d⁹, and d¹⁰ complexes, there is only one way to write satisfactory configurations.

As a result, we have to worry about high-spin versus low-spin octahedral complexes only when there are four, five, six, or seven electrons in the d orbitals.

The choice between high-spin and low-spin configurations for octahedral d⁴, d⁵, d⁶, or d⁷ complexes is easy. All we have to do is compare the energy it takes to pair electrons with the energy it takes to excite an electron to the higher energy (e_g) orbitals. If it takes less energy to pair the electrons, the complex is low-spin. If it takes less energy to excite the electron, the complex is high-spin.

The amount of energy required to pair electrons in the t_2g orbitals of an octahedral complex is more or less constant. The amount of energy needed to excite an electron into the higher energy (e_g) orbitals, however, depends on the value of  _o for the complex. As a result, we expect to find low-spin complexes among metal ions and ligands that lie toward the high-field end of the spectrochemical series. High-spin complexes are expected among metal ions and ligands that lie toward the low-field end of these series.

Compounds in which all of the electrons are paired are diamagnetic they are repelled by both poles of a magnet. Compounds that contain one or more unpaired electrons are paramagnetic they are attracted to the poles of a magnet. The force of attraction between paramagnetic complexes and a magnetic field is proportional to the number of unpaired electrons in the complex. We can therefore determine whether a complex is high-spin or low-spin by measuring the strength of the interaction between the complex and a magnetic field.

 

LIGAND FIELD 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, just as atomic orbitals describe the distributions in atoms.) 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, require that the maximum number of unpaired electrons in energy levels have equal or almost equal energies. Compounds that contaio unpaired electrons are slightly repelled by a magnetic field and are said to be diamagnetic. Because unpaired electrons behave like tiny magnets, 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 (t_2g) and antibonding (e_g) orbitals. The orbital splitting between the two sets of orbitals (t_2g and e_g) is designated as the orbital ligand field parameter, δ_o(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 t_2g and e_g orbitals, and consequently the δ_ovalue 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 t_2g and e_g orbitals, and consequently the δ_ovalue 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 t_2g orbitals, resulting in a magnetic moment indicating one unpaired electron; in the [FeF₆] ³⁻ ion, three electrons occupy the t_2g orbitals and two electrons occupy the e_g 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 and pi (π) bonds, 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 compoundtetracarbonylnickel, 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^* ord 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 (δ_o). 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 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.

The valence-bond model and the crystal field theory explain some aspects of the chemistry of the transition metals, but neither model is good at predicting all of the properties of transition-metal complexes. A third model, based on molecular orbital theory, was therefore developed that is known as ligand-field theory. Ligand-field theory is more powerful than either the valence-bond or crystal-field theories. Unfortunately it is also more abstract.

The ligand-field model for an octahedral transition-metal complex such as the Co(NH₃)₆³⁺ ion assumes that the 3d, 4s, and 4p orbitals on the metal overlap with one orbital on each of the six ligands to form a total of 15 molecular orbitals, as shown in the figue below.

Six of these orbitals are bonding molecular orbitals, whose energies are much lower than those of the original atomic orbitals. Another six are antibonding molecular orbitals, whose energies are higher than those of the original atomic orbitals. Three are best described asnonbonding molecular orbitals, because they have essentially the same energy as the 3d atomic orbitals on the metal.

Ligand-field theory enables the 3d, 4s, and 4p orbitals on the metal to overlap with orbitals on the ligand to form the octahedral covalent bond skeleton that holds this complex together. At the same time, this model generates a set of five orbitals in the center of the diagram that are split into t_2g and e_g subshells, as predicted by the crystal-field theory. As a result, we don’t have to worry about “inner-shell” versus “outer-shell” metal complexes. In effect, we can use the 3d orbitals in two different ways. We can use them to form the covalent bond skeleton and then use them again to form the orbitals that hold the electrons that were originally in the 3dorbitals of the transition metal.

 

Example 1

Sketch the structures of isomers Co(en)₃³⁺ complex ion to show that they are mirror images of each other.

Solution

The images are shown on page 242 Inorganic Chemistry by Swaddle. If the triangular face of the end-amino group lie on the paper, you can draw lines to represent the en bidentate ligand. These lines will show that the two images are similar to the left-hand and right-hand screws.

From the description above, sketch the structures.

Discussion

Answer these questions:

· How many isomers does the complex Co(NH₂CH₂CH₂NH₂)₂ClNH₃]²⁺ have? Draw the structures of the isomers.

· How many isomers does Co(en)₂Cl₂⁺ have? Sketch the structures of the isomers.

· How many isomers does Co(NH₃)₄Cl₂⁺ have? Sketch the structures of the isomers.

How are coordination compounds named?

Structures of coordination compounds can be very complicated, and their names long because the ligands may already have long names. Knowing the rules of nomenclature not only enable you to understand what the complex is, but also let you give appropriate names to them.

Often, several groups of the ligand are involved in a complex. The number of ligand molecules per complex is indicated by a Greek prefix: mono-, di- (or bis), tri-, tetra-, penta-, hexa, hepta-, octa-, nona-, (ennea-), deca- etc for 1, 2, 3, … 10 etc. If the names of ligands already have one of these prefixes, the names are placed in parentheses. The prefices for the number of ligands become bis-, tris-, tetrakis, pentakis- etc.

For neutral ligands, their names are not changed, except the following few:

H₂O, aqua

NH₃, ammine (not two m’s, amine is for organic compounds)

CO, carbonyl

NO, nitrosyl

Normal names that will not change

C₅H₅N, pyradine

NH₂CH₂CH₂NH₂, ethylenediamine

C₅H₄N-C₅H₄N, dipyridyl

P(C₆H₅)₃, triphenylphosphine

NH₂CH₂CH₂NHCH₂CH₂NH₂, diethylenetriamine

The last “e” iames of negative ions are changed to “o” iames of complexes. Sometimes “ide” is changed to “o”. Note the following:

Cl^–, chloride -> chloro

OH^–, hydroxide -> hydroxo

O^2-, oxide -> oxo

O₂^– peroxide, -> peroxo

CN^–, cyanide -> cyano

N₃^–, azide -> axido

N^3-, nitride -> nitrido

NH₂^–, amide -> amido

CO₃^2-, carbonate -> carbonato

-ONO₂^–, nitrate -> nitrato (when bonded through O)

-NO₃^–, nitrate -> nitro (when bonded through N)

S^2-, sulfide -> sulfide

SCN^–, thiocyanate -> thiocyanato-S

NCS^–, thiocyanate -> thiocyanato-N

-(CH₂-N(CH₂COO^–)₂)₂, ethylenediaminetetraacetato (EDTA)

The names of complexes start with the ligands, the anionic ones first, followed with neutral ligands and the metal. If the complex is negative, the name ends with “ate”. At the very end are some Romaumerals representing the oxidation state of the metal.

To give and remember all rules of nomenclature are hard to do. Pay attention to the names whenever you encounter any complexes is the way to learn.

[Co(NH₃)₅Cl]Cl₂, Chloropentaamminecobalt(III) chloride

[Cr(H₂O)₄Cl₂]Cl, Dichlorotetraaquochromium(III) chloride

K[PtCl₃NH₃], Potassiumtrichloroammineplatinate(II)

PtCl₂(NH₃)₂, Dichlorodiammineplatinum

Co(en)₃Cl₃, tris(ethylenediamine)cobalt(III)chloride

Ni(PF₃)₄, tetrakis(phosphorus(III)fluoride)nickel(0)

A bridgin ligand is indicated by placing a m- before its name. The m- should be repeated for every bridging ligand. For example,

(H₃N)₃Co(OH)₃Co(NH₃)₃, Triamminecobalt(III)-m-trihydroxotriamminecobalt(III)

Example 2

Give the structural formula for chlorotriphenylphosphinepalladium(II)- m-dichlorochlorotriphenylphosphinepalladium(II).

Solution

The structure is

(C6H5)3P    Cl    Cl

        \  /  \  /

         Pd    Pd

        /  \  /  \

      Cl    Cl    P(C6H5)3

Example 3

Name the complex:

        NH        

(en)2Co<  >Co(en)2 Cl3

        OH        

Solution

The name is Bis(ethylenediamine)cobalt(III)-m- imido-m-hydroxobis(ethylenediamine)cobalt(III) ion.

 

Cisplatin, PtCl₂(NH₃)₂

A platinum atom with four ligands

In chemistry, a coordination complex or metal complex, consists of an atom or ion (usually metallic), and a surrounding array of bound molecules or anions, that are in turn known as ligands or complexing agents. Many metal-containing compounds consist of coordination complexes.

Nomenclature and terminology

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. A typical complex 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. 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.

1. Order of naming ions

The positive ion is named first followed by the name of the negative ion. In the case of ionic compounds, the name of the cation is mentioned first followed by the name of the anion, e.g., K₂PtCl₆ is named as potassiumhexchloroplatinate (IV) in which cation potassium is written first followed by the anion Chloro. In the case of non ionic compound the name is written in one word.

For example, [Ni (NH₃)₄Cl₂] is named as tetraamminedichloronickel (II).

 

2. Naming of coordination sphere

The name of the legends are written first followed by the name of the central ion and theoxidatioumber of the central metal atom is expressed in romaumerals just after the name of the central atom. Or the sequence can be written as [Name of legend] [Name of central metal atom] [Oxidatioumber of metal in roman].

For example, [Co(NH₃)₆]³⁺, the coordination sphere is named as Hexaamminecobalt (III) ion.

 

3. Naming of Ligand

The central metal ion is surrounded by positive or negative or neutral ligands. If a ligand carries a negative charge the name has a characteristic ending word. For example, SCN – Thiocynate, HS – hydrogen sulphido etc. For a ligand carrying a positive charge, the name has the characteristic ending words. For example, NO^—Nitrosonium, NO^{2 +} Nitronium etc. For neutral ligands no characteristic ending is used.

For example, H₂O Aqua, etc.

 

4. Order of naming ligands

According to the IUPAC conventions the names of ligands surrounding the central metal atom are written in alphabetical order of preference irrespective of whether they are negative or neutral. For example, in the complex [Co(NH₃)₄Cl(NO₂)], the ligands are named in the order  ammine, chloro and nitro. The prefixes di, tri, etc., are not considered while determining the alphabetical order.

 

5. Numerical prefixes to indicate number of ligands

The number of each kind of ligand is specified by the prefixes di, tri, tetra, etc., i.e., CO (en)₃]Cl₃ is named as Tris (ethylendiamine) cobalt (III) chloride.

Numerical Prefixes

 

 

6. Ending of name of central atom

If the complex ion is a cation, the metal is named the same as the element. For example, Co in a complex cation is called cobalt and Pt is called platinum. If the complex ion is an anion, the name of the metal ends with the suffix rate. So, Co in a complex anion will be cobaltate and Pt is called platinate.

 

7. Oxidation state of central atom or ion

The oxidation state is designated by Romaumerals (I, II, III, etc.) in the bracket at the end of the name of the complex without a gap between the two.

E.g. Pt(NH₃)₂Cl₄  diamminetetrachloro platinum (IV), Fe(CO)^5- pentacarbonyliron(0), etc.

 

8. Naming of ambidentate ligands

The ligands which can coordinate with more than one atom are called ambi-dentate ligands, and in these ligands the point of attachment is indicated by putting the symbol of an atom through which it is coordinating with the rest. Example of this is thiocyanate, SCN^−which can get attached to either the sulfur atom or the nitrogen atom, similar NO^2- ion.

 

9. Naming of Geometrical isomers and optical isomers

In these isomers, the term cis is used to show similar groups at adjacent position & trans is for similar groups at opposite positions.

In the case of optical isomers Dextro & laevo rotatory optically active compounds are designated either by (+) & (-) or by d- & l- respectively.

 

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. Examples are

 

Werner’s Theory of Coordination Compounds

The coordination chemistry was pioneered by Nobel Prize winner Alfred Werner (1866-1919). He received the Nobel Prize in 1913 for his coordination theory of transition metal-amine complexes. At the start of the 20th century, inorganic chemistry was not a prominant field until Werner studied the metal-amine complexes such as [Co(NH₃)₆Cl₃].

Alfred Werner,  (born Dec. 12, 1866, Mulhouse, France—died Nov. 15, 1919, Zürich, Switz.), Swiss chemist and winner of the Nobel Prize for Chemistry in 1913 for his research into the structure of coordination compounds.

 

Education

Werner was the fourth and last child of Jean-Adam Werner, a foundry worker and former locksmith, and his second wife, Salomé Jeanette Werner, who was a member of a wealthy family. Alsace had become part of the second German Empire in 1871, but French continued to be spoken by the family. Although most of Werner’s articles were published in German in German journals, his cultural and political sympathies remained with France.

Although Werner’s later interest in religion was minimal, his family was Roman Catholic, and he attended the École Libre des Frères (1872–78), followed by the École Professionelle, a technical school where he studied chemistry (1878–85). He spent one year (1885–86) of compulsory military service in the German army at Karlsruhe, where he audited chemistry lectures at the Technische Hochschule. In 1886 he enrolled in the Eidgenössisches Polytechnikum (now the Eidgenössische Technische Hochschule [ETH], or Swiss Federal Institute of Technology) in Zürich, from which he received a technical chemical degree (1889). Because the Polytechnikum was not empowered to grant the doctorate until 1909, Werner received a doctorate formally from the University of Zürich in 1890.

Early research

Werner’s first publication, a cornerstone of stereochemistry, based on his doctoral dissertation and written with his research supervisor, Arthur Hantzsch, applied Joseph-Achille Le Bel and Jacobus Henricus van ’t Hoff’s concept of the tetrahedral carbon atom (1874) to the nitrogen atom. It explained numerous cases of cis-trans isomerism among trivalent nitrogen compounds such as the oximes, led to the discovery of new isomers, and placed the stereochemistry of nitrogen on a consistent theoretical foundation. During the winter semester of 1891–92, Werner worked on thermochemical studies at the Collège de France in Paris with Marcellin Berthelot.

Assessment

Werner not only explained known coordination compounds but also predicted the existence of numerous series of unknown compounds, which were discovered by him and his students during a quarter-century tour de force of synthetic activity that confirmed his theory in almost every particular. His concepts of ionogenic and nonionogenic bonding adumbrated the current distinction between electrostatic and covalent bonding by a full generation. His ideas soon encompassed almost the entire field of inorganic chemistry and even found application in organic, analytical, and physical chemistry, as well as biochemistry, geochemistry, and mineralogy. He was one of the first to show that stereochemistry is not limited to organic chemistry but is a general phenomenon. His coordination theory has had an effect on inorganic chemistry comparable to that exerted on organic chemistry by the ideas of Kekule, Archibald Scott Couper, Le Bel, and van ’t Hoff. Consequently, he is sometimes called “the inorganic Kekule.”

Following his resolution of series after series of coordination compounds beginning in 1911, Werner became the first Swiss chemist to win the Nobel Prize for Chemistry, “in recognition of his work on the linkage of atoms in molecules, by which he has thrown fresh light on old problems and opened new fields of research, particularly in inorganic chemistry.” Shortly thereafter he began to suffer from a general, progressive, degenerative arteriosclerosis, especially of the brain, aggravated by years of excessive drinking and overwork. He died in Burghölzli, a psychiatric hospital. He was not only the founder of modern inorganic stereochemistry but also one of the major chemists of all time.

In 1892 Werner became a Privatdozent (unsalaried lecturer) at the Polytechnikum upon acceptance of his Habilitationsschrift (an original research paper required in order to teach at a university). In this work, which elicited little notice because it was published (1891) in an obscure local journal, he proposed replacing August Kekule’s rigidly directed valence bonds in organic compounds with a more flexible approach of viewing affinity as a variously divisible force acting equally in all directions from the atom’s centre.

Major theoretical work

In 1893 Werner published his third major article on stereochemistry, setting forth his controversial theory of coordination compounds, which had occurred to him in a dream. Although his knowledge of inorganic chemistry was extremely limited, he awoke one night in 1892 at 2:00 am with the solution to the puzzle of what were then called “molecular compounds.” He wrote his most important theoretical paper by 5:00 pm. It brought him almost instant fame and an appointment as extraordinarius (associate) professor at the University of Zürich, where he spent the rest of his career. In 1894 he became a Swiss citizen and married Emma Wilhelmina Giesker, with whom he had two children, Alfred and Charlotte. An enthralling lecturer and prolific researcher, he was promoted to full professor in 1895.

At the time of its inception, Werner’s theory was largely without experimental verification. He had done no work in the field, and the data that he cited in support of his ideas had been obtained by others, especially by his primary scientific adversary, the Danish chemist Sophus Mads Jørgensen. Jørgensen adhered to the rival Blomstrand-Jørgensen “chain theory,” which was eventually superseded by Werner’s theory, the basis for modern coordination chemistry.

Werner discarded Kekule’s artificial distinction between “valence compounds,” amenable to classical valence theory, and “molecular compounds,” those not explainable by this theory. Among the latter were the metal-ammines, which contain a metal salt as well as ammonia (a neutral molecule), both of which were capable of independent existence. The basic property of the ammonia was “masked” in that it did not react with acids. Also, the nature of the strong bond between the metal salt and the ammonia was unexplained.

Werner proposed a revolutionary approach in which the constitution and configuration of metal-ammines (now colloquially called “Werner complexes”), double salts, and metal salt hydrates were logical consequences of a new concept, the coordinatioumber. He divided metal-ammines into two classes—those with coordinatioumber six, for which he postulated an octahedral configuration, and those with coordinatioumber four, for which he proposed a square planar or tetrahedral configuration. He also postulated two types of valence—primary valence, which bonded the anion to the metal atom, and secondary valence, which bonded the ammonia to the metal atom.

Werner demonstrated the validity of his views by citing numerous reactions, transformations, and cases of isomerism. He showed that loss of ammonia from metal-ammines was not a simple loss but a substitution in which a change in function of the anions occurred simultaneously, resulting in a complete transition from cationic compounds through nonelectrolytes to anionic compounds. He also showed how ammonia could be replaced by water or other groups, and he demonstrated the existence of transition series between ammines, double salts, and metal hydrates. In addition, he speculated on other subjects such as the state of salts in solution and the polarization effects involved in chemical bonding.

Werner recognized the existence of several forms of cobalt-ammonia chloride. These compounds have different color and other characteristics. The chemical formula has three chloride ions per mole, but the number of chloride ions that precipitate with Ag⁺ ions per formula is not always three. He thought only ionized chloride ions will form precipitate with silver ion. In the following table, the number below the Ionized Cl^– is the number of ionized chloride ions per formula. To distinguish ionized chloride from the coordinated chloride, Werner formulated the Complex formula and explained structure of the cobalt complexes (See page 241 ofInorganic Chemistry by Swaddle).

Alfred Werner (1866-1919), a Swiss chemist was the first to formulate his ideas about the structures of coordination compounds. He prepared and characterised a large number of coordination compounds and studied their physical and chemical behaviour by simple experimental techniques. Werner proposed the concept of a primary valence and a secondary valence for a metal ion. Binary compounds such as CrCl₃, CoCl₂ or PdCl₂ have primary valence of 3, 2 and 2 respectively. In a series of compounds of cobalt(III) chloride with ammonia, it was found that some of the chloride ions could be precipitated as AgCl on adding excess silver nitrate solution in cold but some remained in solution.

1 mol CoCl₃.6NH₃ (Yellow) gave 3 mol AgCl

1 mol CoCl₃.5NH₃ (Purple) gave 2 mol AgCl

1 mol CoCl₃.4NH₃ (Green) gave 1 mol AgCl

1 mol CoCl₃.4NH₃ (Violet) gave 1 mol AgCl

These observations, together with the results of conductivity measurements in solution can be explained if (i) six groups in all, either chloride ions or ammonia molecules or both, remain bonded to the cobalt ion during the reaction and (ii) the compounds are formulated as shown in Table 9.1, where the atoms within the square brackets form a single entity which does not dissociate under the reaction conditions. Werner proposed the term secondary valence for the number of groups bound directly to the metal ion; in each of these examples the secondary valences are six.

Note that the last two compounds in Table 9.1 have identical empirical formula, CoCl₃.4NH₃, but distinct properties. Such compounds are termed as isomers. Werner in 1898, propounded his theory of coordination compounds. The main postulates are:

1. In coordination compounds metals show two types of linkages (valences)-primary and secondary.

2. The primary valences are normally ionisable and are satisfied by negative ions.

3. The secondary valences are non ionisable. These are satisfied by neutral molecules or negative ions. The secondary valence is equal to the coordinatioumber and is fixed for a metal.

4. The ions/groups bound by the secondary linkages to the metal have characteristic spatial arrangements corresponding to different coordinatioumbers.

In modern formulations, such spatial arrangements are called coordination polyhedra. The species within the square bracket are coordination entities or complexes and the ions outside the square bracket are called counter ions.

He further postulated that octahedral, tetrahedral and square planar geometrical shapes are more common in coordination compounds of transition metals. Thus, [Co(NH₃)₆]³⁺, [CoCl(NH₃)₅]²⁺ and [CoCl₂(NH₃)₄]⁺ are octahedral entities, while [Ni(CO)₄] and [PtCl₄]²⁻ are tetrahedral and square planar, respectively.

Structure of hexol

Coordination complexes were known – although not understood in any sense – since the beginning of chemistry, e.g. Prussian blue and copper vitriol. The key breakthrough occurred when Alfred Werner proposed in 1893 that Co(III) bears six ligands in an octahedral geometry. His theory allows one to understand the difference between coordinated and ionic in a compound, for example chloride in the cobalt ammine chlorides and to explain many of the previously inexplicable isomers.

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

 

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)]⁻.

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

In coordination chemistry, a structure is first described by its coordinatioumber, 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. Coordinatioumbers 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 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 coordinatioumber 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 coordinatioumbers, e.g. [Mo(CN)₈]⁴⁻. Small metals with large ligands lead to low coordinatioumbers, e.g. Pt[P(CMe₃)]₂. Due to their large size, lanthanides, actinides, and early transition metals tend to have high coordinatioumbers.

Different ligand structural arrangements result from the coordinatioumber. 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 indistinct geometrically 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.

·         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 coordinatioumber 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

·         Coordinatioumbers 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, andhexacarbonylchromium (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. 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 coordination number 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 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.

 

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. 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.

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:

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

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 and anions 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+),

 

 

STEREOISOMERISM

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

Cis–trans isomerism occurs in octahedral and square planar complexes (but not tetrahedral). When two ligands are mutually 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

Optical isomerism occurs when a molecule is not superposable 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.

Λ-[Fe(ox)₃]³⁻

· 

Δ-[Fe(ox)₃]³⁻

· 

Λ-cis-[CoCl₂(en)₂]⁺

· 

Δ-cis-[CoCl₂(en)₂]⁺

Structural isomerism

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 pentaaminebromidocobalt(III)sulfate [Co(NH₃)₅Br]SO₄ is red violet and in solution gives a precipitate with barium chloride, confirming the presence of sulfate ion, while pentaaminesulfatecobalt(III)bromide [Co(NH₃)₅SO₄]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, [Cr(H₂O)₅Cl]Cl₂·H₂O is blue-green, and [Cr(H₂O)₄Cl₂]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]

ENANTIOMERS 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, 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 methyl (CH₃) group.

 

Electronic properties

Many of the properties of 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.

 

Color

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

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. 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.

Magnetism

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 TiX₂[(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.

The magnetism discussed in this article is paramagnetism. Paramagnetism occurs when there are one or more unpaired electrons in a compound. (The opposite, when all electrons are paired, is called diamagnetism). Di- and para-magnetism are often affected by the presence of coordination complexes, which the transition metals (d-block) readily form.

Singular electrons have a spin, denoted by the quantum number m_s as +(1/2) or –(1/2). This spin is negated when the electron is paired with another, but creates a slight magnetic field when the electron is unpaired. The more unpaired electrons, the more likely paramagnetic a material is. The electron configuration of the transition metals (d-block) changes when in a compound. This is due to the repulsive forces between electrons in the ligands and electrons in the compound. Depending on the strength of the ligand, the compound may become paramagnetic or diamagnetic.

Ferromagnetism

Some paramagnetic compounds are capable of becoming ferromagnetic. This means that the compound shows permanent magnetic properties rather than exhibiting them only in the presence of a magnetic field. In a ferromagnetic element, electrons of atoms are grouped into domains, where each domain has the same charge. In the presence of a magnetic field, these domains line up so that charges are parallel throughout the entire compound. Whether a compound can be ferromagnetic or not depends on how many unpaired electrons it has and on its atomic size.

·        Small atoms pair up too easily and their charges cancel.

·        Large atoms are difficult to keep together, their charge interaction is too weak.

Therefore, only the right sized atoms will work together to group themselves into domains. Elements with the right size include: Fe, Co, Ni. That means that Fe, Co and Ni are paramagnetic with the capability of permanent magnetism; they are also ferromagnetic.

Ligand Field Theory Background

For a full explanation, please see the article on Ligand Field Theory. An element can have up to 10 d electrons in 5 d-orbitals, d_xy, d_xz, d_yz, d_z², and d_x²-y². During the formation of a complex, the degeneracy (equal energy) of these orbitals is broken and the orbitals are at different energy levels.

In an octahedral complex, the ligands approach along the x, y, and z axes, so the repulsion is strongest in the orbitals along these axes (dz² and dx²-y²). As a result, the dz² and d_x²-y² orbitals are higher in energy than the dxy, dxz, and dyz orbitals. In a tetrahedral complex, the splitting is opposite, with the dxy, dxz, and dyz orbitals higher in energy to avoid the ligands approaching between the axes. The splitting in a square planar complex has four levels (lowest to highest): d_yz and d_xz, d_xy, dz², dx²-y².

Depending on the strength of the ligand, the splitting energy between the different d-orbitals may be large or small. Ligands producing a smaller splitting energy are called ‘weak field’ ligands, and those with a larger splitting energy are called ‘strong field’ ligands.

 

Filling of d-orbitals in a complex

Hunds’ Rule states that electrons will fill all available orbitals with single electrons before pairing up, while maintaining parallel spins (paired electrons have opposing spins). For a set of degenerated d-orbitals (not in a complex), electrons fill all orbitals before pairing to conserve the pairing energy, otherwise needed. With the addition of ligands, the situation becomes more complicated. The splitting energy between the d-orbitals means that additional energy is required to place single electrons into the higher-energy orbitals. Once the lower-energy orbitals have been half-filled (one electron per orbital), an electron can either be placed in a higher-energy orbital (preserving Hund’s rule) or pair up with an electron in a lower-energy orbital (when the splitting energy is greater than the pairing energy). The strength of the ligands determine which option is chosen.

With a strong-field ligand, the splitting energy is very large and low-spin complexes are usually formed. With a weak-field ligand, the electrons can easily enter the higher-energy orbitals before pairing (high-spin).

How does this relate to magnetism?

Low-spin complexes contain more paired electrons since the splitting energy is larger than the pairing energy. These complexes, such as [Fe(CN)₆]^3-, are more often diamagnetic or weakly paramagnetic. High-spin complexes usually contain more unpaired electrons since the pairing energy is larger than the splitting energy. With more unpaired electrons, high-spin complexes  are often paramagnetic.

The unpaired electrons in paramagnetic compounds create tiny magnetic fields, similar to the domains in ferromagnetic materials (see above or the related article). The higher the number of unpaired electrons (often the higher-spin the complex), the stronger the paramagnetism of a coordination complex. We can predict paramagnetiism and its relative strength by determining whether a compound is a weak field ligand or a strong field ligand. Once we have determined whether a compound has a weak or a strong ligand, we can predict its magnetic properties:

 

How can we measure magnetism in a compound?

The Gouy balance is used to measure paramagnetism by suspending the complex in question against an equivalent weight with access to a magnetic field. We first weigh the complex without a magnetic field in its presence, then, we weigh it again in the presence of a magnetic field. If the compound is paramagnetic, it will be pulled visibly towards the electromagnet, which is the distance proportional to the magnitude of the compound’s paramagnetism. If the compound, however, is diamagnetic, it will not be pulled towards the electromagnet, instead, it might even slightly be repelled by it. This will be proven by the decreased weight or the no change in weight. The change in weight directly corresponds to the amount of unpaired electrons in the compound.

How Magnetic Properties relate to the “Real World”

Ferromagnetism, the permanent magnetism associated with nickel, cobalt, and iron, appears throughout everyday life, from Aristotle’s discussion in 625 BC, through the use of the compass in 1187, up to the modern-day refrigerator. Einstein declared that electricity and magnetism are inextricably linked in his theory of “special relativity.” He also showed examples that a magnet can be disturbed by electricity.

Paramagnetism and diamagnetism explain and describe some of the properties of certain elements and complexes, which we work with on a regular basis. In the early days of complex-compound chemistry, paramagnetism was often used to help identify the shape of complexes. A technique known as electron paramagnetic resonance has been used in systems with certain para- and dia- magnetic properties to distinguish between bond types and identify the probable location of an individual element within a compound.

 

 

Reactivity

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(H₂O)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

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.

Coordination compounds are the compounds in which the central metal atom is linked to a number of positive or negative ions or neutral molecules by coordinate bonds and the donor atoms, molecules or ions which donate a pair of electrons to the central metal atom or ion and form a coordinate bond called ligand.

For example nickel tetra carbonyl, [Ni(CO)4] in which CO molecules are linked to the central nickel atom by coordinate bonds by donating lone pair of electrons is a coordinate compound or [Fe(CN)₆]^4-,[Cu(NH₃)⁴]²⁺, etc.

Types of complexes

 

THERE ARE THREE TYPES OF COMPLEXES:

1. Cationic complex

A complex in which the complex carries a net positive charge is called cationic complex. For example, [Co (NH₃)₆]³⁺

2. Anionic complex

 

A complex in which the complex carries a net negative charge is called anionic complex. For example, [Fe (CN)₆]^4-

3. Neutral Complex

 

A complex carrying no net charge is called a neutral complex or simply a complex. For exmaple, [Mn₂(CO)₁₀]

 

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

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: natural hemoglobin

Many 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

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

The basic procedure for naming a complex:

1.            Wheaming a complex ion, the ligands are named before the metal ion.

2.            Write the names of the ligands in the order,-neutral, negative, positive. If there are multiple ligands of the same charge type, they are named 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: NH₃ 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 Latiame will be used if available (except for mercury).

4.            If the central atom’s oxidation state needs to be specified (when it is one of several possible, or zero), write it as a Romaumeral (or 0) in parentheses.

5.            Name cation then anion as separate words (if applicable, as in last example)

Examples:

[NiCl₄]²⁻ → tetrachloridonickelate(II) ion

[CuNH₃Cl₅]³⁻ → amminepentachloridocuprate(II) ion

[Cd(en)₂(CN)₂] → dicyanidobis(ethylenediamine)cadmium(II)

[Co(NH₃)₅Cl]SO₄ → pentaamminechloridocobalt(III) sulfate

The coordinatioumber 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.

Application of coordination compounds

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 for electroplating 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⁺ in hard water.

5.                     Silver and gold are extracted by treating zinc with their cyanide complexes.

A complex is a substance in which a metal atom or ion is associated with a group of neutral molecules or anions called ligands. Coordination compounds are neutral substances (i.e. uncharged) in which at least one ion is present as a complex. You will learn more about coordination compounds in the lab lectures of experiment 4 in this course.

The coordination compounds are named in the following way. (At the end of this tutorial we have some examples to show you how coordination compounds are named.)

A. To name a coordination compound, no matter whether the complex ion is the cation or the anion, always name the cation before the anion. (This is just like naming an ionic compound.)

B. Iaming the complex ion:

1. Name the ligands first, in alphabetical order, then the metal atom or ion. Note: The metal atom or ion is written before the ligands in the chemical formula.

2. The names of some common ligands are listed in Table 1.

� For anionic ligands end in “-o”; for anions that end in “-ide”(e.g. chloride), “-ate” (e.g. sulfate, nitrate), and “-ite” (e.g. nirite), change the endings as follows: -ide -o; -ate -ato; -ite -ito

� For neutral ligands, the commoame of the molecule is used e.g. H₂NCH₂CH₂NH₂ (ethylenediamine). Important exceptions: water is called ‘aqua’, ammonia is called ‘ammine’, carbon monoxide is called ‘carbonyl’, and the N₂ and O₂ are called ‘dinitrogen’ and ‘dioxygen’.

Names of Some Common Ligands

3. Greek prefixes are used to designate the number of each type of ligand in the complex ion, e.g. di-, tri- and tetra-. If the ligand already contains a Greek prefix (e.g. ethylenediamine) or if it is polydentate ligands (ie. can attach at more than one binding site) the prefixes bis-, tris-, tetrakis-, pentakis-, are used instead. (See examples 3 and 4.) The numerical prefixes are listed in Table 2.

Numerical Prefixes

 

4. After naming the ligands, name the central metal. If the complex ion is a cation, the metal is named same as the element. For example, Co in a complex cation is call cobalt and Pt is called platinum. (See examples 1-4). If the complex ion is an anion, the name of the metal ends with the suffix –ate. (See examples 5 and 6.). For example, Co in a complex anion is called cobaltate and Pt is called platinate. For some metals, the Latiames are used in the complex anions e.g. Fe is called ferrate (not ironate).

Name of Metals in Anionic Complexes

  

5. Following the name of the metal, the oxidation state of the metal in the complex is given as a Romaumeral in parentheses.

C. To name a neutral complex molecule, follow the rules of naming a complex cation. Remember: Name the (possibly complex) cation BEFORE the (possibly complex) anion.See examples 7 and 8.

For historic reasons, some coordination compounds are called by their commoames. For example, Fe(CN)₆^3- and Fe(CN)₆^4- are named ferricyanide and ferrocyanide respectively, and Fe(CO)₅ is called iron carbonyl.

Examples of Common Coordination Numbers

Note that the charge on the complex is always the sum of the charges on the ions or molecules that form the complex.

Cu²⁺ + 4 NH₃ = Cu(NH₃)₄²⁺

Pb²⁺ + 2 OAc^– = Pb(OAc)₂

Fe²⁺ + 6 CN^– = Fe(CN)₆^4-

Note also that the coordinatioumber of a complex often increases as the charge on the metal ion becomes larger.

Cu⁺ + 2 NH₃ = Cu(NH₃)₂⁺

Cu²⁺ + 4 NH₃ = Cu(NH₃)₄²⁺

 

Examples Give the systematic names for the following coordination compounds:

1. [Cr(NH₃)₃(H₂O)₃]Cl₃

Answer: triamminetriaquachromium(III) chloride

Solution: The complex ion is inside the parentheses, which is a cation.

The ammine ligands are named before the aqua ligands according to alphabetical order.

Since there are three chlorides binding with the complex ion, the charge on the complex ion must be +3 ( since the compound is electrically neutral).

From the charge on the complex ion and the charge on the ligands, we can calculate the oxidatioumber of the metal. In this example, all the ligands are neutral molecules. Therefore, the oxidatioumber of chromium must be same as the charge of the complex ion, +3.

2. [Pt(NH₃)₅Cl]Br₃

Answer: pentaamminechloroplatinum(IV) bromide

Solution: The complex ion is a cation, the counter anion is the 3 bromides.

The charge of the complex ion must be +3 since it bonds with 3 bromides.

The NH₃ are neutral molecules while the chloride carries – 1 charge. Therefore, the oxidatioumber of platinum must be +4.

3. [Pt(H₂NCH₂CH₂NH₂)₂Cl₂]Cl₂

Answer: dichlorobis(ethylenediamine)platinum(IV) chloride

Solution: ethylenediamine is a bidentate ligand, the bis- prefix is used instead of di-

4. [Co(H₂NCH₂CH₂NH₂)₃]₂(SO₄)₃

Answer: tris(ethylenediamine)cobalt(III) sulfate

Solution: The sulfate is the counter anion in this molecule. Since it takes 3 sulfates to bond with two complex cations, the charge on each complex cation must be +3.

Since ethylenediamine is a neutral molecule, the oxidatioumber of cobalt in the complex ion must be +3.

Again, remember that you never have to indicate the number of cations and anions in the name of an ionic compound.

5. K₄[Fe(CN)₆]

Answer: potassium hexacyanoferrate(II)

Solution: potassium is the cation and the complex ion is the anion.

Since there are 4 K⁺ binding with a complex ion, the charge on the complex ion must be – 4.

Since each ligand carries –1 charge, the oxidation number of Fe must be +2.

The commoame of this compound is potassium ferrocyanide.

6. Na₂[NiCl₄]

Answer: sodium tetrachloronickelate(II)

Solution: The complex ion is the anion so we have to add the suffix –ate in the name of the metal.

7. Pt(NH₃)₂Cl₄

Answer: diamminetetrachloroplatinum(IV)

Solution: This is a neutral molecule because the charge on Pt⁺⁴ equals the negative charges on the four chloro ligands.

If the compound is [Pt(NH₃)₂Cl₂]Cl₂, eventhough the number of ions and atoms in the molecule are identical to the example, it should be named: diamminedichloroplatinum(II) chloride, a big difference.

8. Fe(CO)₅

Answer: pentacarbonyliron(0)

Solution: Since it is a neutral complex, it is named in the same way as a complex cation. The commoame of this compound, iron carbonyl, is used more often.

9. (NH₄)₂[Ni(C₂O₄)₂(H₂O)₂]

Answer: ammonium diaquabis(oxalato)nickelate(II)

Solution: The oxalate ion is a bidentate ligand.

10. [Ag(NH₃)₂][Ag(CN)₂]

Answer: diamminesilver(I) dicyanoargentate(I)

You can have a compound where both the cation and the anion are complex ions. Notice how the name of the metal differs even though they are the same metal ions.

Can you give the molecular formulas of the following coordination compounds?

1. hexaammineiron(III) nitrate

2. ammonium tetrachlorocuprate(II)

3. sodium monochloropentacyanoferrate(III)

4. potassium hexafluorocobaltate(III)

Can you give the name of the following coordination compounds?

5. [CoBr(NH₃)₅]SO₄

6. [Fe(NH₃)₆][Cr(CN)₆]

7. [Co(SO₄)(NH₃)₅]⁺

8. [Fe(OH)(H₂O)₅]²⁺

Answers:

1. [Fe(NH₃)₆](NO₃)₃

2. (NH₄)₂[CuCl₄]

3. Na₃[FeCl₁(CN)₅]

4. K₃[CoF₆]

5. pentaamminebromocobalt(III) sulfate

6. hexaammineiron(III) hexacyanochromate (III)

7. pentaamminesulfatocobalt(III) ion

8. pentaaquahydroxoiron(III) ion

A characteristic feature of the transition metals is their ability to form a group of compounds called coordination compounds, complex compounds, or sometimes simply complexes. A typical coordination compound is the intensely blue solid substance Cu(NH₃)₄SO₄ which can be crystallized from solutions of CuSO₄ to which a very large excess of concentrated NH3 has been added. These crystals contain two polyatomic ions, one of which is the sulfate ion, SO4²–, and the other of which is the complex ion Cu(NH3)4²⁺ which is responsible for the blue color.

We can regard a complex ion such as Cu(NH₃)₄²⁺ as being the result of the interaction of :NH₃ acting as a Lewis base with the Cu²⁺ ion acting as a Lewis acid. Each NH3 molecule can be considered as donating a pair of electrons to a central Cu²⁺, thus forming four coordinate covalent bonds to it:

Most coordination compounds contain a complex ion similar to Cu(NH₃)₄²⁺. This ion can be either positively charged like Cr(H₃O)₆³⁺, or it can be negatively charged like CoCl₆^3–. Neutral complexes like Pt(NH₃)₂Cl₂ are also known. All these species contain a central metal ion attached by coordinate covalent bonds to several ligands. These ligands are invariably Lewis bases. Some typical examples of ligands are H₂O, NH₃, Cl^–, OH^–, CN^–, Br^–, and SCN^–. The number of ligands attached to the central metal ion is said to be its coordination number and is usually 2, 4, or 6. The group of ligands bonded to the metal taken collectively is said to constitute the metal’s coordination sphere.

When writing the formula of a coordination compound containing complex ions, square brackets are usually used to enclose the coordination sphere. Examples are

 

[Cu(NH₃)₄]SO₄      [Cr(H₂O)₆]Cl₃       [Pt(NH₃)₂Cl₂]

[Cu(NH₃)₄](NO₃)₂      K₃[Fe(CN)₆]      [Pt(NH₃)₄][PtCl₄]

 

When such compounds are dissolved in H₂O, each of the ions present in the solid becomes an independent species with its own chemical and physical properties. Thus, when 1 mol [Cr(H₂O)₆]Cl₃ crystal is dissolved in H₂O the solution contains 1 mol Cr(H₂O)₆³⁺ ion which can be recognized by its characteristic grayish-violet color and 3 mol Cl^– which can be detected by the precipitate of AgCl which forms when AgNO₃ is added to the solution.

 

Observations on Complex Compounds Containing PtCl₂, NH₃, and KCl.

 

 

An even better example of how the various ions in a coordination compound can behave independently when dissolved in water is provided by the set of Pt(II) complexes shown in the table. The first of these compounds contains the complex ion [Pt(NH₃)₄]²⁺ and in each subsequent compound one of the NH₃ ligands in the coordination sphere of the Pt is replaced by a Cl^– ligand. As a result each compound contains a Pt complex of different composition and also of different charge, and when dissolved in H₂O, it shows just the conductivity and other properties we would expect from the given formula. When 1 mol [Pt(NH₃)₃Cl]Cl is dissolved in H₂O, it furnishes 1 mol Pt(NH₃)₃Cl⁺ ions and 1 mol Cl^– ions. The strongest evidence for this is the molar conductivity of the salt (1.2 A V^–1 dm² mol^–1), which is very similar to that of other electrolytes like NaCl (1.3 A V^–1 dm² mol^–1) which also yield a +1 ion and a –1 ion in solution, but very different from that of electrolytes like MgCl₂ (2.5 A V^–1 dm² mol^–1) which yield one + 2 ion and two –1 ions in solution. The conductivity behavior also suggests that the Pt atom is part of a cation, since the Pt moves toward the cathode during electrolysis. The addition of AgNO₃ to the solution serves to confirm this picture. One mol AgCl is precipitated immediately, showing 1 mol free Cl^– ions. After a few hours a further mole of AgCl is precipitated, the Cl^– this time originating from the coordination sphere of the Pt atom due to the slow reaction

 

[Pt(NH₃)₃Cl]⁺(aq) + Ag⁺(aq) + H₂O → [Pt(NH₃)₃H₂O]²⁺(aq) + AgCl(s)

 

It is worth noting that in all these compounds, Pt has an oxidatioumber of + 2. Thus the combination of Pt with one NH₃ ligand and three Cl^– ligands yields an overall charge of 2(for Pt) – 3(for Cl) + 0(for NH₃) = –1. The ion is thus the anion [PtNH₃Cl₃]^– found in compound 4.

 

EXAMPLE

What is the oxidation state of Pt in the compound Ca[Pt(NH₃)Cl₅]₂?

Solution Since there are two complex ions for each Ca²⁺ ion, the charge on each must be –1. Adding the charge on each ligand, we obtain –5(for Cl^–) + 0(for NH₃) = –5. If the oxidatioumber of Pt is x, then x – 5 must equal the total charge on the complex ion:

x – 5 = –1

or     x = +4

The geometry of a complex is governed almost entirely by the coordinatioumber. We will consider only the most common coordinatioumbers, namely, 2, 4, and 6.

 

Coordinatioumber = 2     Complexes with two ligands are invariably linear. The best-known examples of such compounds are Ag(I) and Au(I) complexes such as

and     

 

Both of these complexes are important. The Au(CN)₂^– complex is used to extract minute gold particles from the rock in which they occur. The crushed ore is treated with KCN solution and air is blown through it:

4Au(s) + 8CN^–(aq) + O₂(g) + 2H₂O(l) → 4[Au(CN)₂]^–(aq) + 4OH^–(aq)

The resultant complex is water soluble. The silver complex is also water soluble and affords a method for dissolving AgCl, which is otherwise very insoluble.

AgCl(s) + 2NH₃(aq) → [Ag(NH₃)₂]⁺(aq) + Cl^–(aq)

This reaction is often used in the laboratory to be sure a precipitate is AgCl(s).

 

Figure 1 Structure of (a) Pt(NH₃)₄²⁺ (square planar) and (b) Zn(NH₃)₄²⁺ (tetrahedral). Metal ions are dark gray, nitrogen atoms are light gray, and hydrogen atoms are light color. Geometrical isomers can exist for square planar complexes but not for tetrahedral complexes.

Coordinatioumber = 4. Two geometries are possible for this coordinatioumber. Some complexes, like the [Pt(NH₃)₄]²⁺ ion shown in Fig. 1, are square planar, while others, like Cd(NH₃)₄^2–, are tetrahedral. Most of the four-coordinated complexes of Zn(II), Cd(II), and Hg(II) are tetrahedral, while the square planar arrangement is preferred by Pd(II), Pt(II), and Cu(II) complexes.

Because the square planar geometry is less symmetrical than the tetrahedral geometry, it offers more possibilities for isomerism. A well-known example of such isomerism is given by the two square planar complexes

 

These two isomers are called geometrical isomers. That isomer in which two identical ligands are next to each other is called the cis isomer, while that in which they are on opposite sides is called the trans isomer. Though these two isomers have some properties which are similar, no properties are identical and some are very different. For example, the cis isomer of the above complex is used as an anti-tumor drug to treat cancerous cells. The trans form, by contrast, shows no similar biological activity.

It is worth noting that cis-trans isomerism is not possible in the case of tetrahedral complexes. As you can quickly verify by examining any three-dimensional tetrahedral shape, any given corner of a tetrahedron is adjacent to the other three. Since all the corners are cis to each other, none are trans.

 

Figure 2 (a) Cis isomer and (b) trans isomer of [Co(NH3)4Cl2]+. On the macroscopic level the cis form is violet while the trans form is green.

Coordinatioumber = 6     When there are six ligands, the geometry of the complex is almost always octahedral, like the geometry of SF₆, or of [Cr(H₂O)₆]³⁺. All ligands are equidistant from the central atom, and all ligand-metal-ligand angles are 90°. An octahedral complex may also be thought of as being derived from a square planar structure by adding a fifth ligand above and a sixth below on a line through the central metal ion and perpendicular to the plane.

The octahedral structure also gives rise to geometrical isomerism. For example, two different compounds, one violet and one green, have the formula [Co(NH₃)₄Cl₂]Cl. The violet complex turns out to have the cis structure and the green one trans.

Although we have confined our discussion so far to simple ligands such as Cl^–, NH₃, or H₂O, much larger and more complicated molecules can also donate electron pairs to metal ion. An important and interesting example of this is the chelating agents—ligands which are able to form two or more coordinate covalent bonds with a metal ion. One of the most common of these is 1,2-diaminoethane (usually called ethylenediamine and abbreviated en.)

When both nitrogens coordinate to a metal ion, a stable five-member ring is formed. The word chelating, derived from the Greek chele, “claw,” describes the pincerlike way in which such a ligand can grab a metal ion.

A chelating agent which forms several bonds to a metal without unduly straining its own structure is usually able to replace a similar simpler ligand. For example, although both form coordinate covalent bonds via groups, ethylenediamine can readily replace ammonia from most complexes:

 

 

For metals which display a coordinatioumber of 6, an especially potent ligand is ethylenediaminetetraacetate ion (abbreviated EDTA):

 

 

All six electron pairs marked in color are capable of coordinating to a metal ion, in which case the EDTA ion wraps completely around the metal and is very difficult to dislodge. EDTA is used to treat lead and mercury poisoning because of its ability to chelate these metals and aid their removal from the body.

Chelate complexes are often important in living systems. The coordination of iron in proteins such as myoglobinor hemoglobin involves four nitrogen of the heme group and one from a histidine side chain. Since iroormally has a coordinatioumber of 6, this leaves one open site, to which oxygen can bond. The presence of carbon monoxide, a stronger ligand than oxygen, causes displacement of oxygen from hemoglobin. This prevents transport of oxygen from the lungs to the brain, causing drowsiness, loss of consciousness, and even death upon long exposure to carbon monoxide. Consequently operating an automobile in a closed garage, a cookstove in a tent, or burning any fossil fuel incompletely in an enclosed space may be hazardous to one’s health.

Another important application of chelates is transport of metal ions across membranes. The interior of a biological membranes contain the nonpolar, hydrophobic tails of lipid molecules. This makes it quite difficult for ionic species such as K⁺ and Na⁺ to travel from one side of a membrane to the other. One way in which this barrier may be circumvented is by carrier molecules, called ionophores. Ionophores are able to chelate an ion, but also have a hydrophobic exterior.

One such ionophore is the antibiotic nonactin, a medium-sized organic molecule with the formula

 

This molecule is able to transport K⁺ ions but not Na⁺ ions. Apparently the Na⁺ ion is too small to fit in among the eight coordinating O’s, while the K⁺ ion can (see Fig. 1). Other than these O’s, most of the nonactin molecule is a hydrocarbon chain. Therefore once K⁺ is chelated, the outer part of the complex is quite hydrophobic. It can easily pass through the interior of a membrane, releasing K⁺ on the other side. The toxic effect of nonactin and several related antibiotics is the result of their ability to transport alkali-metal ions to regions of a cell where they should not be. This breaks down ion gradients the cell has created to perform tasks and store energy. Consequently the cell wastes energy pumping K⁺ and other ions out again.

 

Figure 1 The K+ complex of nonactin (K+ is light gray). Coordinate covalent bonds are indicated as thin solid lines. Hydrogen atoms have been omitted for clarity.

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 emptyd 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 d subshells (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 thenitrogen 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 constantsof 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).

Equilibrium constants for the formation of various nickel-amine complexes

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.

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, (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, and (4) mixed complexes containing other ligands in addition to CO: pentacarbonylchloromanganese, tetracarbonylhydridocobalt, and tricarbonylnitrosylcobalt.

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.

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 for a variety of reactions. More recently, the research field has been expanded to include organometallic species.

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. Such similarities may be deceptive, however, for the additional electron ieutral 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)}₂(μ₂−NH₂)(μ₂−NO)]) or three (e.g., [Mn₃(η⁵−C₅H₅)₃(μ₂−NO)₃(μ₃−NO)]) metal atoms. (The η⁵ indicates that five carbon atoms of the C₅H₅⁻ group are bonded to the chromium atom.)

Cyano 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 andcyanide 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 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, andmercury, usually by substitution in aqueous solution with no change in oxidation state. The most important complexes are anionic with the formula [Mⁿ⁺(CN)_x]^{(x− n)−}, 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 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 d¹⁰transition metal ions, the octa-, hepta-, hexa-, penta-, and tetracoordinate complexes, respectively, are species with maximum coordinatioumber.

Mixed complexes of type [M(CN)₅X]ⁿ⁻ (where X = H₂O, NH₃, CO, NO, H, or a halogen) 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. Cyanide complexes have figured prominently iumerous kinetic 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; 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) and organic isocyanides (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 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—that is, compounds with multiple bonds between carbon atoms.

Isopoly and heteropoly anions

The amphoteric 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 acidanhydride, 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 elementscan 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 theheteropolymolybdates and heteropolytungstates are the best known, have very high molecular weights (some above 4,000) as compared with other inorganic electrolytes, 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), 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 chemistJöns Jacob Berzelius 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, 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.

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 PO₄tetrahedron, 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

Acid-base

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

[M(H₂O)_x]ⁿ⁺⇌ [M(OH)(H₂O)_{x− 1}]^{(n− 1)+}+ H⁺,

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.

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 metalcation 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.

Lewis Acid-Lewis base Approach to Bonding in Complexes

G. N. Lewis was the first to recognize that the reaction between a transition-metal ion and ligands to form a coordination complex was analogous to the reaction between the H⁺ and OH^– ions to form water. The reaction between H⁺ and OH^– ions involves the donation of a pair of electrons from the OH^– ion to the H⁺ ion to form a covalent bond.

The H⁺ ion can be described as an electron-pair acceptor. The OH^– ion, on the other hand, is an electron-pair donor. Lewis argued that any ion or molecule that behaves like the H⁺ ion should be an acid. Conversely, any ion or molecule that behaves like the OH^– ion should be a base. A Lewis acid is therefore any ion or molecule that can accept a pair of electrons. A Lewis base is an ion or molecule that can donate a pair of electrons.

When Co³⁺ ions react with ammonia, the Co³⁺ ion accepts pairs of nonbonding electrons from six NH₃ ligands to form covalent cobalt-nitrogen bonds as shown in the figure below.

The metal ion is therefore a Lewis acid, and the ligands coordinated to this metal ion are Lewis bases.

The Co³⁺ ion is an electron-pair acceptor, or Lewis acid, because it has empty valence-shell orbitals that can be used to hold pairs of electrons. To emphasize these empty valence orbitals we can write the configuration of the Co³⁺ ion as follows.

Co³⁺: [Ar] 3d⁶ 4s⁰ 4p⁰

There is room in the valence shell of this ion for 12 more electrons. (Four electrons can be added to the 3d subshell, two to the 4sorbital, and six to the 4p subshell.) The NH₃ molecule is an electron-pair donor, or Lewis base, because it has a pair of nonbonding electrons on the nitrogen atom.

According to this model, transition-metal ions form coordination complexes because they have empty valence-shell orbitals that can accept pairs of electrons from a Lewis base. Ligands must therefore be Lewis bases: They must contain at least one pair of nonbonding electrons that can be donated to a metal ion.

 

Substitution

One of the most general reactions exhibited by coordination compounds is that of substitution, 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 coordinatioumber 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 coordinatioumber. 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 coordinatioumber; one of the original ligands is then lost to restore the initial coordinatioumber. 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:

CO, CN⁻, C₂H₄> PR₃, H⁻> NO₂⁻, I⁻, SCN⁻> Br⁻, Cl⁻> NH₃, H₂O.

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 transto ammonia.

Lability and inertness

In considering the mechanisms of substitution (exchange) reactions, Canadian-born American chemist Henry Taube 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 termsunstable 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 of reaction (ΔG), whereas inert complexes merely possess large positive free energies of activation (ΔG^*).

While the existence of geometric or optical isomers 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 oflability.

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).

Henry Taube

Henry Taube,  (born Nov. 30, 1915, Neudorf, Sask., Can.—died Nov. 16, 2005, Stanford, Calif., U.S.), Canadian-born American chemist, who won the Nobel Prize for Chemistry in 1983 for his extensive research into the properties and reactions of dissolved inorganic substances, particularly oxidation-reduction processes involving the ions of metallic elements (see oxidation-reduction reaction).

Taube was educated at the University of Saskatchewan (B.S., 1935; M.S., 1937) and the University of California, Berkeley (Ph.D., 1940). He later taught at Cornell University (1941–46) and the University of Chicago (1946–61) before joining the faculty of Stanford University in 1962; he was named professor emeritus in 1986. Taube became a U.S. citizen in 1942.

In the late 1940s Taube carried out experiments with isotopes to show that in water solution the ions of metals form chemical bonds with several molecules of water and that the stability and geometric arrangement of the resulting hydrates, or coordination compounds, vary widely, depending on the identity and oxidation state of the ion. He also helped develop other techniques for studying such substances, and he devised an interpretation of their properties in terms of their electronic configurations. Analogous coordination compounds form in the presence of ammonia, chloride ions, or numerous other chemical species, which are called ligands when they engage in these reactions.

The oxidation or reduction of one metal ion by another involves their exchange of one or more electrons. Many such reactions occur rapidly in aqueous solution despite the fact that the stable shells of water molecules or other ligands should keep the two ions from getting close enough for electron exchange to occur directly. Taube showed that, in an intermediate stage of the reaction, a chemical bond must form between one of the ions and a ligand that is still bonded to the other. This ligand acts as a temporary bridge between the two ions, and its bond to the original ion can later break in such a way as to effect—indirectly—the electron transfer that completes the reaction. Taube’s findings have been applied in selecting metallic compounds for use as catalysts, pigments, and superconductors and in understanding the function of metal ions as constituents of certain enzymes.

Taube was the recipient of numerous honours, including two Guggenheim fellowships (1949, 1955) and the National Medal of Science (1976). In 1959 he became a member of the National Academy of Sciences.

Isomerization

Coordination compounds that exist in two or more isomeric forms may undergo reactions that convert one isomer to another. Examples are the linkage isomerization and cis-transisomerization reactions depicted below.

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

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

 

Oxidation-reduction

Transition metals commonly exhibit two or more stable oxidation states, and their complexes accordingly are able to undergooxidation-reduction reactions. The simplest such reactions involveelectron transfer between two complexes, with little if any accompanying rearrangement or chemical change. An example is shown below:

Two limiting mechanisms of electron transfer, commonly designated outer-sphere and inner-spheremechanisms, 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, 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).

 

Transitional Metal Ions in Aqueous Solutions

We often write transition-metal ions in aqueous solution with symbols such as Cr³⁺, Cu²⁺, and Fe³⁺ as though they were monatomic, but this is far from being the case. These ions are actually hydrated in solution and can be regarded as complex ions. Thus, for example, the grayish-violet color of many chromium(III) salts when dissolved in H₂O is due to the species [Cr(H₂O)₆]³⁺ rather than to a bare Cr³⁺ ion. The same color is evident in many crystalline solids such as [Cr(H₂O)₆]Cl₃ which are known to contain the Cr³⁺ ion surrounded octahedrally by six H₂O molecules. In much the same way the blue color of many solutions of copper(II) salts can be attributed to the species [Cu(H₂O)₄]²⁺ and the pale violet color of some solutions of iron(III) salts to the [Fe(H₂O)₆]³⁺ ion. Because [Fe(H₂O)₆]³⁺ is capable of donating a proton, the conjugate base, [Fe(H₂O)₅OH]²⁺ is generally present when Fe³⁺ is dissolved in water. This imparts a yellow color to the solution. Fig. 1 shows examples of colored ion complexes in aqueous solution.

Not all salts of transition-metal ions yield the hydrated ion when dissolved in H₂O. Thus when CuCl₂ is dissolved in H₂O, a beautiful green color due mainly to the complex [CuCl₂(H₂O)₂] is produced. This is obviously different from the sky-blue color of [Cu(H₂O)₄]²⁺ which is obtained when Copper(II) sulfate or copper(II) nitrate are dissolved. This is because the Cl^– ion is a stronger Lewis base with respect to the Cu²⁺ ion than is H₂O. Thus, if there is a competition between H₂O and Cl^– to bond as a ligand to Cu²⁺, the Cl^– ion will usually win out over the H₂O.

The superior strength of the Cl^– as a Lewis base is easily demonstrated by adding Cl^– ions to a sky-blue solution of copper(II) sulfate. A green color immediately appears due to the formation of chloro complexes:

 

[Cu(H₂O)₄]²⁺ + Cl^– [Cu(H₂O)₃Cl]⁺ + H₂O

 

[Cu(H₂O)₃Cl]⁺ + Cl^– [Cu(H₂O)₂Cl₂] + H₂O

                                                                     Green

If a large excess of Cl^– ion is added, the solution changes color again from green to yellow. This is because of even further displacement of H₂O ligands by Cl^– ligands:

 

[Cu(H₂O)₂Cl₂] + Cl^– [Cu(H₂O)₂Cl₃] ^– + H₂O

[Cu(H₂O)₂Cl₃] ^– + Cl^– [CuCl₄]^2– + H₂O

                  Yellow

Figure 2 compares these three aqueous copper complexes.

 

Under favorable circumstances yellow crystals of salts like Cs₂[CuCl₄], containing the complex ion CuCl₄^2– can be obtained from these solutions.

Coordination compounds, such as the FeCl₄^– ion and CrCl₃ 6 NH₃, are called such because they contain ions or molecules linked, or coordinated, to a transition metal. They are also known as complex ions or coordination complexes because they are Lewis acid-base complexes. The ions or molecules that bind to transition-metal ions to form these complexes are called ligands (from Latin, “to tie or bind”). The number of ligands bound to the transition metal ion is called the coordination number.

Although coordination complexes are particularly important in the chemistry of the transition metals, some main group elements also form complexes. Aluminum, tin, and lead, for example, form complexes such as the AlF₆^3-, SnCl₄^2- and PbI₄^2- ions.

Alfred Werner developed a model of coordination complexs which explains the following observations.

·  At least three different cobalt(III) complexes can be isolated when CoCl₂ is dissolved in aqueous ammonia and then oxidized by air to the +3 oxidation state. A fourth complex can be made by slightly different techniques. These complexes have different colors and different empirical formulas.

·  The reactivity of the ammonia in these complexes has been drastically reduced. By itself, ammonia reacts rapidly with hydrochloric acid to form ammonium chloride.

NH₃(aq) + HCl(aq)   NH₄⁺(aq) + Cl^–(aq)

These complexes don’t react with hydrochloric acid, even at 100^oC.

·  Solutions of the Cl^– ion react with Ag⁺ ion to form a white precipitate of AgCl.

Ag⁺(aq) + Cl^–(aq)   AgCl(s)

When excess Ag⁺ ion is added to solutions of the CoCl₃ 6 NH₃ and CoCl₃ 5 NH₃ H₂O complexes, three moles of AgCl are formed for each mole of complex in solution, as might be expected. However, only two of the Cl^– ions in the CoCl₃ 5 NH₃ complex and only one of the Cl^–ions in CoCl₃ 4 NH₃ can be precipitated with Ag⁺ ions.

·  Measurements of the conductivity of aqueous solutions of these complexes suggest that the CoCl₃ 6 NH₃ and CoCl₃ 5 NH₃ H₂O complexes dissociate in water to give a total of four ions. CoCl₃ 5 NH₃ dissociates to give three ions, and CoCl₃ 4 NH₃ dissociates to give only two ions.

Werner explained these observations by suggesting that transition-metal ions such as the Co³⁺ ion have a primary valence and a secondary valence. The primary valence is the number of negative ions needed to satisfy the charge on the metal ion. In each of the cobalt(III) complexes previously described, three Cl^– ions are needed to satisfy the primary valence of the Co³⁺ ion.

The secondary valence is the number of ions of molecules that are coordinated to the metal ion. Werner assumed that the secondary valence of the transition metal in these cobalt(III) complexes is six. The formulas of these compounds can therefore be written as follows.

The cobalt ion is coordinated to a total of six ligands in each complex, which satisfies the secondary valence of this ion. Each complex also has a total of three chloride ions that satisfy the primary valence. Some of the Cl^– ions are free to dissociate when the complex dissolves in water. Others are bound to the Co³⁺ ion and neither dissociate nor react with Ag⁺.

The [Co(NH₃)₆]Cl₃ complex dissociates in water to give a total of four ions, and all three Cl^– ions are free to react with Ag⁺ ion.

One of the chloride ions is bound to the cobalt in the [Co(NH₃)₅Cl]Cl₂ complex. Only three ions are formed when this compound dissolves in water, and only two Cl^– ions are free to precipitate with Ag⁺ ions.

Once again, the three Cl^– ions are free to dissociate when [Co(NH₃)₅(H₂O)]Cl₃ dissolves in water, and they precipitate when Ag⁺ ions are added to the solution.

Two of the chloride ions are bound to the cobalt in [Co(NH₃)₄Cl₂]Cl. Only two ions are formed when this compound dissolves in water, and only one Cl^– ion is free to precipitate with Ag⁺ ions.

Werner assumed that transition-metal complexes had definite shapes. According to his theory, the ligands in six-coordinate cobalt(III) complexes are oriented toward the corners of an octahedron, as shown in the figure below.

Because they might very possibly form complexes with it, one must be careful about what ions are added to a solution containing hydrated transition-metal ions. Not only the chloride ions, but the other halide ions are liable to complex, and the same is true of species like NH₃ and CN^–. These ligands differ quite a lot in their affinity for a particular metal ion, but the rules governing this situation are not simple. One finds, for instance, that although NH₃ will complex very readily with Cu²⁺ it has little or no affinity for Fe³⁺. In other words, a ligand which is a strong Lewis base with respect to one metal ion is not necessarily a strong base with respect to another. There are some ions, however, which almost always function as very weak Lewis bases. The perchlorate ion, ClO₄^– in particular, forms almost no complexes. The nitrate ion, NO₃^–, and sulfate ion, SO₄^2–, only occasionally form complexes.

The addition of ligands to a solution in order to form a highly colored complex is often used to detect the presence or absence of a given metal in solution. The deep blue color of [Cu(NH₃)₄]²⁺ produced when excess NH₃ is added to solution of Cu(II) salts is a case in point. This can be seen in the following video, where a aqueous solution of ammonia is added to a copper sulfate solution.

The initial copper sulfate solution is sky blue, due to the [Cu(H₂O)₄]²⁺ complex. When ammonia is added, a precipitate of Cu(OH)₂(s) is formed. as it settles to the bottom, it can be seen that the remaining solution is a dark blue, due to the [Cu(NH₃)₄]²⁺ complex formed by copper with ammonia.

Other well-known color reactions are the blood-red complex formed between Fe(III) ions and the thiocyanate ion, SCN^–, as well as the pink-red complex of Ni(II) with dimethylglyoxime.

While most of the reactions we have been describing are very fast and occur just as quickly as the solutions are mixed, this is not always the case. With certain types of complexes, ligand substitution is quite a slow process. For example, if Cl^– ions are added to a solution containing [Cr(H₂O)₆]³⁺ ions, it is a few days before the grayish-violet color of the original ion is replaced by the green color of the chloro complexes [Cr(H₂O₅) Cl]²⁺ and [Cr(H₂O)₄ Cl]⁺. Alternatively the solution may he heated, in which case the green color will usually appear within 10 min. The reaction

 

[Cr(H₂O)₆]³⁺ + Cl^– → [Cr(H₂O)₅Cl]³⁺ + H₂O

 

is thus a slow reaction with a high activation energy. Ligand substitution reactions of other Cr(III) complexes behave similarly. In consequence Cr(III) complexes are said to be inert, as opposed to a complex like Fe(H₂O)₆³⁺ which swaps ligands very quickly and is said to be labile. Other examples of inert complexes are those of Co(III), Pt(IV), and Pt(II). Almost all the compounds which were used to establish the nature and the geometry of coordination compounds were inert rather than labile. There is very little point in trying to prepare cis and trans isomers of a labile complex, for example, because either will quickly react to form an equilibrium mixture of the cis and trans forms.

A final complication in dealing with aqueous solutions of transition-metal complexes is their acid-base behavior. Hydrated metal ions like [Cr(H₂O)₆]³⁺ are capable of donating protons to water and acting as weak acids. Most hydrated ions with a charge of + 3, like Al³⁺ and Fe³⁺ behave similarly and are about as strong as acetic acid. The hydrated Hg(II) ion is also noticeably acidic in this way. Perhaps the most obvious of these cationic acids is the hydrated Fe(III) ion. When most Fe(III) salts are dissolved in water, the color of the solution is yellow or brown, though the Fe(H2O)₆³⁺ ion itself is pale violet. The yellow color is due to the conjugate base produced by the loss of a proton. The equilibrium involved is

 

[Fe(H₂O)₆]³⁺ + H₂O [Fe(H₂O)₅OH]²⁺ + H₃O⁺

Pale violet                         Brown

 

If solutions of Fe(III) salts are acidified with perchloric acid or nitric acid, the brown base is protonated and the yellow color disappears from the solution entirely.

Two or more different compounds having the same formula are called isomers. Two principal types of isomerism are known among coordination compounds. Each of which can be further subdivided.

1. Stereoisomerism.

a) Geometrical isomerism

b) Optical isomerism

2. Structural Isomerism.

a) Coordination isomerism

b) Ionisation isomerism

c) Hydrate isomerism

d) Linkage isomerism

1. Stereoisomers

Stereoisomers have the same atoms, same sets of bonds, but differ in the relative orientation of these bonds.

Geometric isomers are possible for both square planar and octahedral complexes, but not tetrahedral.

Optical isomers are possible for both tetrahedral and octahedral complexes, but not square planar.

The earliest examples of stereoisomerism involve complexes of Co(III). In 1889, Jorgensen observed purple and green salts of [CoCl₂(en)₂]+, which Werner later correctly identified as the cis- and trans- geometric isomers. In 1911, the first resolution of optical isomers was reported by Werner and King for the complexes cis-[CoX(NH₃)(en)₂]²⁺, where X=Cl- or Br-.

Geometric Isomers

The number of geometric isomers expected for common stereochemistries are as follows:

Square Planar:

Compound type       No. of isomers

Ma₂b₂                    2 (cis- and trans-)

Mabcd                   3 (use cis- and trans- relations)

here a, b, c, and d refer to monodentate ligands.

A number of examples of these types have been isolated and characterised and they show very different chemical and biological properties. Thus for example, cis-PtCl₂(NH₃)₂ is an anti-cancer agent (cisplatin) whereas the trans- isomer is inactive against cancer (it is toxic), and so not useful in Chemotherapy.

cis- and trans- isomers of [PtCl₂(NH₃)₂]

cis- and trans- refer to the position of 2 groups relative to each other. In the cis- isomer they are “next to each other” i.e. at 90 degrees in relation to the central metal ion, whereas in the trans- isomer they are “opposite each other”, i.e. at 180 degrees relative to the central metal ion.

 

 

 

3 geometric isomers of a square planar complex [PtBrClNH₃(pyr)].

The first report of the three geometric isomers being isolated and characterised for complexes of the type [Mabcd] was by Il’ya Chernyaev in 1928. The example above was reported by Anna Gel’man in 1948.

Question. Does cis-amminebromo-cis-chloropyridineplatinum(II) uniquely identify isomer 9(ii)??

Octahedral:

Compound type No. of isomers

Ma₄b₂ 2 (cis- and trans-)

Ma₃b₃ 2 (fac- and mer-)

MAA₂b₂ 3 (2*cis- and 1 trans-)

here a, and b, represent monodentate ligands and AA is a bidentate ligand.

In the second example, new labels are introduced to reflect the relative positions of the ligands around the octahedral structure. Thus; placing the 3 groups on one face of the octahedral gives rise to the facial isomer and placing the 3 groups around the centre gives rise to the meridional isomer.

fac- and mer- isomers of [RhCl₃(pyr)₃].

 

[Mabcdef] is expected to give 15 geometric isomers. In the case of [PtBrClI(NH₃)(pyr)NO₂], several of these were isolated and characterised by Anna Gel’man and reported in 1956. Optical isomers are possible for each of these 15 forms, making a total of 30 isomers.

The cis- isomer of MAA₂b₂ may also exhibit optical isomerism although we will concentrate largely on optical isomers of the type M(AA)₃ (see below).

Optical Isomers

Optical isomers are related as non-superimposable mirror images and differ in the direction with which they rotate plane-polarised light. These isomers are referred to as enantiomers or enantiomorphs of each other and their non-superimposable structures are described as being asymmetric.

Various methods have been used to denote the absolute configuration of optical isomers such as R or S, Λ or Δ or C and A. The IUPAC rules suggest that for general octahedral complexes C/A scheme is convenient to use and that for bis and tris bidentate complexes the absolute configuration be designated Lambda Λ (left-handed) and Delta Δ (right-handed).

Priorities are assigned for mononuclear coordination systems based on the standard sequence rules developed for enantiomeric carbon compounds by Cahn, Ingold and Prelog (CIP rules). These rules use the coordinating atom to arrange the ligands into a priority order such that the highest atomic number gives the highest priority number (smallest CIP number). For example the hypothetical complex [Co Cl Br I NH₃ NO₂ SCN]^2- would assign the I- as 1, Br as 2, Cl as 3, SCN as 4, NO₂ as 5 and NH₃ as 6.

Here is one isomer where the I and Cl, and Br and NO₂ were found to be trans- to each other.

The reference axis for an octahedral centre is that axis containing the ligating atom of CIP priority 1 and the trans ligating atom of lowest possible priority (highest numerical value). The atoms in the coordination plane perpendicular to the reference axis are viewed from the ligand having that highest priority (CIP priority 1) and the clockwise and anticlockwise sequences of priority numbers are compared. The structure is assigned the symbol C or A, according to whether the clockwise (C) or anticlockwise (A) sequence is lower at the first point of difference. In the example shown above this would be C.

The two optical isomers of [Co(en)₃]³⁺ have identical chemical properties and just denoting their absolute configuration does NOT give any information regarding the direction in which they rotate plane-polarised light. This can ONLY be determined from measurement and then the isomers are further distinguished by using the prefixes (-) and (+) depending on whether they rotate left or right.

To add to the confusion, when measured at the sodium D line (589nm), the tris(1,2-diaminoethane)M(III) complexes (M= Rh(III) and Co(III)) with IDENTICAL absolute configuration, rotate plane polarised light in OPPOSITE directions!

The left-handed (Λ)-[Co(en)₃]³⁺ isomer gives a rotation to the right and therefore corresponds to the (+) isomer.

Since the successful resolution of an entirely inorganic ion (containing no C atoms) (hexol) only a handful of truly inorganic complexes have been isolated as their optical isomers e.g. (NH₄)₂Pt(S₅)₃.2H₂O.

For tetrahedral complexes, R and S would be used in a similar method to tetrahedral Carbon species and although it is predicted that tetrahedral complexes with 4 different ligands should be able to give rise to optical isomers, in general they are too labile and caot be isolated.

2. Structural Isomers

There are several types of this isomerism frequently encountered in coordination chemistry and the following represents some of them.

·  a) Coordination isomerism: where compounds containing complex anionic and cationic parts can be thought of as occurring by interchange of some ligands from the cationic part to the anionic part.

one isomer [Co(NH₃)₆] [Cr(C₂O₄)₃]

another isomer [Co(C₂O₄)₃] [Cr(NH₃)₆]

·  b) Ionisation isomers: where the isomers can be thought of as occurring because of the formation of different ions in solution.

one isomer [PtBr(NH₃)₃]NO₂ -> NO₂– anions in solution

another isomer [Pt(NO₂)(NH₃)₃]Br -> Br- anions in solution

Notice that both anions are necessary to balance the charge of the complex, and that they differ in that one ion is directly attached to the central metal but the other is not. A very similar type of isomerism results from replacement of a coordinated group by a solvent molecule (Solvate Isomerism). In the case of water, this is called Hydrate isomerism.

·  c) Hydrate isomerism: the best known example of this occurs for chromium chloride “CrCl₃.6H₂O” which may contain 4, 5, or 6 coordinated water molecules.

[CrCl₂(H₂O)₄]Cl.2H₂O bright-green

[CrCl(H₂O)₅]Cl₂.H₂O grey-green

[Cr(H₂O)₆]Cl₃ violet

These isomers have very different chemical properties and on reaction with AgNO₃ to test for Cl- ions, would find 1, 2, and 3 Cl- ions in solution respectively.

·  d) Linkage isomerism occurs with ambidentate ligands. These ligands are capable of coordinating in more than one way. The best known cases involve the monodentate ligands SCN- / NCS- and NO₂^– / ONO^–.

For example:

[Co(ONO)(NH₃)₅]Cl the nitrito isomer -O attached

[Co(NO₂)(NH₃)₅]Cl the nitro isomer – N attached.

Inorganic Nomenclature

As part of this course, you are required to make yourselves familiar with the rules related to Inorganic Nomenclature.

Uses of Coordination Compounds

A brief survey of some of the uses of coordination compounds includes:

l. Dyes and Pigments: Coordination compounds have been used from the earliest times as dyes and pigments, for example madder dye which is red, was used by the ancient Greeks and others. It is a complex of Hydroxyanthraquinone. A more modern example is the pigment copper phthalocyanine, which is blue.

2. Analytical Chemistry: You have already encountered many such uses during the laboratory course.

(a) Colour Tests: Since many complexes are highly coloured they can be used as colourimetric reagents e.g. formation of red 2,2′-bipyridyl and l,l0-phenanthroline complexes as a test for Fe(II)

(b) Gravimetric Analysis: Here chelating ligands are often used to form insoluble complexes e.g. Ni(DMG)₂ and Al(oxine)₃ (see laboratory manual).

(c) Complexometric Titrations and Masking Agents: An example of this is the use of EDTA in the volumetric determination of a wide variety of metal ions in solution, e.g. Zn²⁺, Pb²⁺, Ca²⁺,Co²⁺, Ni²⁺, Cu²⁺, etc. By careful adjustment of the pH and using suitable indicators, mixtures of metals can be analysed, e.g. Bi³⁺ in the presence of Pb²⁺ (see laboratory manual). Alternatively, EDTA may be used as a masking agent to remove a metal ion which would interfere with the analysis of a second metal ion present.

3. Sequestering Agents: Related to their use as masking agents is the use of ligands for “sequestering” i.e. for the effective removal of objectionable ions from solution in industrial processing, e.g. EDTA is used to “soften” water. The addition of EDTA to water is used in boilers etc., to prevent “scaling” or build up of insoluble calcium salts.

4. Extraction of Metals: Sometimes certain metals can be leached from their ores by formation of stable complexes e.g. Ag and Au as complexes of cyanide ion.

5. Bio-Inorganic Chemistry: Naturally occurring complexes include haemoglobin, chlorophyll, vitamin B₁₂ etc.

EDTA and other complexing agents have been used to speed the elimination of harmful radioactive and other toxic elements from the body. (e.g. Pb²⁺). In these cases soluble metal chelate complexes are formed.

6. Chemo-therapy: an example here is the use of cis-PtCl₂(NH₃)₂ as an anti-tumour drug.

 

Example

For example, Tetraammineaquachloridecobalt (III) chloride- In this the anion is=Cl^–

• Cation [Co(NH₃)₄(H₂O)Cl]^x where x is the charge on the complex ion.

• So, x = oxidation state of cobalt + charge on the ligand

• x = +3+0+0+(-1) = +2 as the H₂O & NH₃ are neutral group so carry no charge & chloride group carry -1 charge.

• Charge on anion = -1 & on the complex ion = +2. Since the overall complex molecule should be neutral, two anions caeutralize the charge of the cation.

• Hence the formula of given compound is [Co(NH₃)₄(H₂O)Cl]Cl₂.

Example

Calculate the charge on the transition-metal ion in the following complexes.

(a) Na₂Co(SCN)₄

(b) Ni(NH₃)₆(NO₃)₂

(c) K₂PtCl₆

Molecular Examples

Coordination Number 2

The linear [Ag(NH₃)₂]⁺ ion

Although [Ag(en)]ClO₄ involves a normally bidentate ligand, in this case the structure is polymeric and the silver ion still retains a CN=2 with the N atoms (from different ligands) at ~180 degrees to each other.

 

Coordination Number 3

Trigonal planar – D_3h

[Cu(CN)₃]^2-

[Cu(PPh₃)₂Br]

To help view more easily, the H atoms are turned off.

Trigonal pyramid

T-shaped

[Rh(PPh₃)₃]⁺

To help view more easily, the H atoms are turned off.

 

Coordination Number 4

Tetrahedral

Copyr₂Cl₂

Square Planar

cisplatin – cis-Pt(NH₃)₂Cl₂

The cis- isomer is a powerful anti-cancer drug whereas the trans- is inactive.

 

Coordination Number 5

Square pyramid

Trigonal Bipyramid

[Ni(CN)₅]^3-

 

Coordination Number 6

Hexagonal planar

Trigonal prism

tris(cis-1,2-diphenylethene-1,2-dithiolato)rhenium

The ReS6 geometry is perfectly trigonal prismatic.

Octahedral

Hexol

The first ‘truly’ inorganic complex to be resolved into its optical isomers.

[Co(en)₃]Cl₃

The classic example of optical isomerism in octahedral coordination complexes (H atoms not shown).

[Co(NH₃)₅CO₃]⁺

 

Coordination Number 7

Capped octahedron (C_3v)

K₃[NbOF₆]

Capped trigonal prism (C_2v)

[V(III)(Hedta)(H₂O)]H₂O

Pentagonal Bipyramid (D_5h)

bis-(tert-butylacac)₂(DMSO)di-oxoUranium

The UO7 geometry fits a pentagonal bipyramid.

 

Coordination Number 8

Dodecahedron (D_2d)

Zr(acac)₂(NO₃)₂

[Zr(C₂O₄)₄]^4- is reported to have this shape as well.

Cube (O_h)

Square antiprism (D_4d)

U(acac)₄

Hexagonal bipyramid (D_6h)

UO₂(OAc)₃

 

Coordination Number 9

Three-face centred trigonal prism (D_3h)

Hydrated salts of the lanthanide elements eg Eu(H₂O)₉]³⁺

 

Coordination Number 10

Bicapped square antiprism (D_4d)

Tetrakis(nitrato-O,O’)-bis(triphenylphosphine oxide) cerium(IV)

Another example is [Ce(NO₃)₅]^2-

 

Coordination Number 11

All-faced capped trigonal prism (D_3h)

This is not a common stereochemistry. In aqua-(12-crown-4)-tris(nitrato-O,O’)-cerium(III) (12-crown-4) solvate and (15-crown-5)-tris(nitrato-O,O’)-cerium(III) the Cerium ion is 11 coordinate.

 

Coordination Number 12

cuboctahedron (Oh)

Ceric ammonium nitrate -(NH₄)₂Ce(NO₃)₆

References:

1. The abstract of the lecture.

2. intranet.tdmu.edu.ua/auth.php

3. Atkins P. W. Physical chemistry / P.W. Atkins. – New York, 1994. – P.299‑307.

4. Cotton F. A. Chemical Applications of Group Theory / F. A. Cotton. ‑ John Wiley & Sons : New York, 1990.

5. Girolami G. S. Synthesis and Technique in Inorganic Chemistry / G. S. Girolami, T. B. Rauchfuss, R. J. Angelici. ‑ University Science Books : Mill Valley, CA, 1999.

6. Russell J. B. General chemistry / J B. Russell. New York.1992. – P. 550‑599.

7. Lawrence D. D. Analytical chemistry / D. D. Lawrence. –New York, 1992. – P. 218–224.

8. http://www.lsbu.ac.uk/water/ionish.html

 

Prepared by PhD Falfushynska H.