Revision Notes: Coordination Compounds

Werner'S Theory

Werner explained the nature of bonding in complexes and he concluded that in complexes the metal shows two different types of valency.

Primary Valency

The primary valency is equal to the oxidation state of the central metal atom or ion. These are non-directional.

Secondary Valency

The number of ligand atoms co-ordinated to the central metal atom is called secondary valency. These are directional and so a complex ion has a particular shape.

The ligand which satisfy the secondary valencies are directed toward fixed positions in space. The geometry of the complex ion depends on the coordination number. If the metal has coodination number 6, the complex is octahedral, i.e. six positions around the metal are occupied by six donor atoms of the ligands octahedrally. On the other hand, if the coordination number is 4, the geometry of the complex may be tetrahedral or square planar. This postulate predicted the existence of different types of isomerism in coordination compounds.

Effective Atomic Number (Ean)

Sidgwick proposed effective atomic number abbreviated as EAN, which is defined as the resultant number of electrons with the metal atom or ion after gaining electrons from the donor atoms of the ligands. The effective atomic number (EAN) generally coincides with the atomic number of next inert gas in some cases. EAN is calculated by the following relation :

EAN = Atomic number of the metal - number of electrons lost in ion formation + number of electrons gained from the donor atoms of the ligands. (2 × CN)

The EAN values of various metals in their respective complexes are tabulated below:

Valence Bond Theory

The bonding in coordination compounds can be explained by Valence Bond Theory (VBT) since majority of the complexes formed by the transition metals have their d- orbitals incomplete. Valence bond takes into account the hybridisation of orbitals since penultimate d-orbitals are near in energy to s and p-orbitals of the outer most shell, various kinds of hybridization is possible.
VBT makes the following assumption

1. A number of empty orbitals are available on the central metal ion which can accomodate electrons donated by the ligands. The number of empty d-orbitals is equal to the coordination number of the metal ion for the particular complex.
2. The metal orbitals and ligand orbitals overlap to form strong bonds. Greater the extent of overlapping, more is the stability of the complex. Different orbitals (s, p or d) hydridize to give a set of equivalent hybridized orbital which take part in bonding with the ligands.
3. Each ligand donates a pair of electrons to the central metal ion/atom.
4. The non-bonding metal electrons present in the inner orbitals do not take part in chemical bonding.
5. If the complex contains unpaired electrons, the complex is paramagnetic. If it does not contain unpaired electron, the complex is diamagnetic in nature.
6. Under the influence of strong ligand (CN, CO) the electrons can be forced to pair up against the Hund's rule of multiplicity.

Common Types of Hybridisation

1. The change in the properties of the ligands and the metal ions could not be explained.
2. The valence bond theory does not explain why certain complexes are more labile than other.
3. The VBT does not provide satisfactory explanation for the existence of inner orbital and outer orbital complexes.
4. The VBT could not explain the colour of complexes

Crystal Field Theory

The Crystal Field Theory is more widely accepted than the valence bond theory. It assumes that the attraction between the central metal and the ligands in a complex is purely electrostatic. In the crystal field the following assumptions are made.

1. Ligands are treated as point charges.
2. There is no interaction between metal orbitals and ligand orbitals.
3. The d orbitals on the metal all have the same energy (that is degenerate) in the free atom. However, when a complex is formed the ligands destroy the degeneracy of these orbitals, i.e. the orbitals now have different energies.
   

Octahedral complexes

In an octahedral complex, the metal is at the centre of the octahedron, and the ligands are at the six corners. The directions x, y and z point to three adjacent corners of the octahedron as shown.
The lobes of the e_g orbitals point along the x, y and z axes. the lobes of the t_2g orbitals (d_xy, d_xz and d_yz) point in between the axes. The approach of six ligands along the x, y, z, -x, - y and - z directions will increase the energy of the orbitals (which point along the axes) much more than it increases the energy of the d_xy, d_xz and d_yz orbitals (which point between the axes). Thus under the influence of an octahedral ligand field the d orbitals split into two groups of different energies.

Ligands which cause only a small degree of crystal field splitting are termed weak field ligands. Ligands which cause a large splitting are called strong field ligands. The common ligands can be arranged in ascending order of crystal field splitting Δ.

Spectrochemical Series

The total crystal field stabilization energy is given by 

where are the number of electrons occupying the t_2g and e_g orbitals respectively. The CFSE is zero for ions with d⁰ and d¹⁰ configurations in both strong and weak ligand field. The CFSE is also zero for d⁵ configurations in a weak field.

Tetrahedral Complexes

A regular tetrahedron is related to a cube. One atom is at the centre of the cube, and four of the eight corners of the cube are occupied by ligands as shown.

The directions x, y and z point to the centres of the faces of the cube. The e orbitals point along x, y and z axes (that is to the centres of the faces). The t₂ orbitals point between x, y and z axes (that is towards the centres of the edges of the cube). The direction of approach of the ligands does not coincide exactly with either the e or the t₂ orbitals.
Thus the t₂ orbitals are nearer to the direction of the ligands than the e orbitals. The approach of the ligands raises the energy of both sets of orbitals. The energy of the t₂ orbitals is raised most because they are closest to the ligands. The crystal field splitting is the opposite way round to that in octahedral complexes.
The t₂ orbitals are 0.4Δ_t above weighted average energy of the two groups (the Bari centre) and the e orbitals are 0.6Δ_t below the average.
The magnitude of the crystal field splitting Δ_t in tetrahedral complexes is considerably less than in octahedral fields. There are two reasons for this :

1. There are only four ligands instead of six, so the ligand field is only two third the size ; hence the ligand field splitting is also two third the size.
2. The direction of the orbitals does not coincide with the direction of the ligands. This reduces the crystal field splitting by roughly a further two third.

Thus the tetrahedral crystal field splitting Δ_t is roughly 2/3 × 2/3 = 4/9 of the octahedral crystal field splitting Δ_o.

Organometallic Compounds

Compounds that contain at least one carbon-metal bond are called organometallic compounds.

Grignard reagent, RMgX is a familiar example of organometallic compounds where R is an alkyl group. Diethyl zinc [Zn(C₂H₅)₂], lead tetraethyl [Pb(C₂H₅)₄], ferrocene [Fe(C₅H₅)₂], dibenzene chromium [Cr(C₆H₆)₂], metal carbonyls are other examples of organometallic compounds.

Organometallic compounds may be classified in three classes:-

1. Sigma (σ) bonded complexes,
2. Pi (π) bonded complexes,
3. Complexes containing both σ - and π - bonding characteristics.

Sigma bonded complexes

In these complexes, the metal atom and carbon atom of the ligand are joined together with a sigma bond, i.e., the ligand contributes one electron and is, therefore, called one electron donor. Examples are:
(i) Grignard reagent, R-Mg-X where R is an alkyl or aryl group and X is halogen.
(ii) Zinc compounds of the formula R₂Zn such as (C₂H₅)₂Zn. This was first isolated by Frankland in 1849. Other similar compounds are (CH₃)₄Sn, (C₂H₅)₄Pb, Al₂(CH₃)₆, Al₂(C₂H₅)₆ and Pb(CH₃)₄, etc.

π-bonded organometallic compounds

These are the compounds of metals with alkenes, alkynes, benzene and other ring compounds. In these complexes, the metal and ligand form a bond that involves the π electrons of the ligand. Three common examples are Zeise's salt, ferrocene and dibenzene chromium. These are shown here:

σ- and π-bonded organometallic compounds

Metal carbonyls, compounds formed between metal and carbon monoxide belong to this class. These compounds posses both σ- and π- bonding. The oxidation state of metal atoms in these compounds is zero. Carbonyls may be monomeric, bridged or polynuclear.

In a metal carbonyl, the metal-carbon bond possesses both the σ- and π-character. A σ-bond between metal and carbon atom is formed when a vacant hybrid orbitals of the metal atom overlaps with an orbital on C atom of carbon monoxide containing a lone pair of electrons.
Formation of π-bond is caused when a filled orbital of the metal atom overlaps with a vacant antibonding π* orbital of C atom of carbon monoxide. This overlap is also called back donation of electrons by metal atom to carbon. It has been shown below:
The π-overlap is perpendicular to the nodal plane of σ-bond. In olefinic complexes, the bonding π-orbital electrons are donated to the empty orbital of the metal atom and at the same time back bonding occurs from filled orbital of the metal atom to the antibonding π-orbital of the olefin.

Isomerism

The compounds having same molecular formula but different structural formula are called isomers.

Structural Isomerism

Ionisation Isomerism: This type of isomerism arises when the coordination compounds give different ions in solution. For example, there are two isomers of the formula Co (NH₃)₅ BrSO₄.

(Violet) [Co(NH₃)₅ Br] SO₄ Pentaammine bromido cobalt(III) ion

This isomer gives a white percipitate of BaSO₄ with BaCl₂ solution.

Above isomer gives light yellow precipitate with AgNO₃ solution.

Hydrate Isomerism: This type of isomerism arises when different number of water molecules are present inside and outside the coordination sphere. This isomerism is best illustrated by the three isomers that have the formula CrCl₃.6H₂O.
[Cr(H₂O)₆]Cl₃ , [Cr(H₂O)₅Cl]Cl₂.H₂O, and [Cr(H₂O)₄Cl₂] Cl.2H₂O are its Hydrate Isomers.

Cordination Isomerism: This type of isomerism is observed in the coordination compounds having both cationic and anionic complex ions. The ligands are interchanged in both the cationic and anionic ions to form isomers. An examples is:

Linkage Isomerism: This type of isomerism occurs in complex compounds which contain ambidentate ligands like

For example, [Co(NH₃)₅NO₂] Cl₂ and [Co(NH₃)₅ONO] Cl₂ are linkage isomers as NO₂^- is linked through N or through O.

Ligand Isomerism: Some ligands themselves are capable of existing as isomers, e.g., diamino propane can exist both as 1, 2-diamino propane (pn) and 1, 3-diamino propane, also called trimethylene diamine (tn).

When these ligands (i.e., pn and tn) are associated into complexes the complexes are isomers of each other. One example of isomeric complexes having this ligand is: [Co(pn)₂ Cl₂]⁺ and [Co(tn)₂ Cl₂]⁺ ions.

Coordination position isomerism: This type of isomerism is exhibited by polynuclear complexes by changing the position of ligands with respect to different metal atoms present in the complex. For example,

Stereo Isomerism

Stereo isomerism is exhibited by those compounds which have the same position of atoms or groups but these atoms or groups have different arrangement around the central atom.

Geometrical Isomerism:
The complex compounds which have the same ligands in the co-ordination sphere but the relative position of the ligands around the central metal atom is different are called geometrical isomers and the phenomenon is called geometrical isomerism.

Geometrical Isomerism in square planar complexes:
A square planar complexe having similar ligands at adjacent positions (90º a part) is called cis - isomer while a square planar complex having two similar ligands at opposite positions (180º a part) is called trans-isomer.

1. Ma₂b₂

Example 1: Draw the geometrical isomers of [PtCl₂(NH₃)₂]
Sol.

2. Ma₂bc

Example 2: Draw the geometrical isomers of [PtCl₂(NH₃)py]
Sol.

3. Mabcd

Example 3: Draw the geometrical isomers of [PtClBrpy (NH₃)]
Sol.

4. M(AB)₂
Example 4: Draw the geometrical isomers of [Pt(gly)₂]
Sol.

Geomertical Isomerism in octahedral complexes:

1. Ma₄b₂

Example 5: Draw the geometrical isomers of [CrCl₂(NH₃)₄]⁺
Sol.

5. M(AA)₂b₂
Example 8: Draw the geometrical isomers of [CoCl₂(en)₂]⁺
Sol.

6. M(AA)₂bc
Example 9: Draw the geometrical isomers of [Co^III (en₂) (NH₃) (Cl)]²⁺
_Sol.

7. M(AA)a₂b₂
Example 10: Draw the geometrical isomers of [Co^III(en) (NH₃)₂ Cl₂]⁺
_Sol.

8. Ma₂b₂c₂
Example 11: Draw the geometrical isomers of [PtCl₂Br₂(NH₃)₂]
_Sol.

Optical Isomerism in octahedral complexes:

1. Mabcdef
Example 12: Draw the optical isomers of [Pt(Cl)(Br)(I)(py)(NO₂)(NH₃)]
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Optical d and l-forms [Pt^IV(py) (NH₃) (NO₂) (Cl) (Br) (I)]⁰

2. M(AA)₃
Example 13: Draw the optical isomers of [Co(en)₃]³⁺
_Sol.

The two optical isomeric forms of the complex [Co(en)₃]³⁺

3. M(AB)₃
Example 14: Draw the optical isomers of [Cr(gly)₃]
_Sol.

4. cis M(AA)₂b₂
Example 15: Draw the optical isomers of RhCl₂(en)₂]⁺
_Sol.

5. cis Ma₂b₂c₂
Example 16: Draw the optical isomers of [PtCl₂Br₂(NH₃)₂]
_Sol.

_{6. cis M(AA)b2}c₂
Example 17: Draw the optical isomers of [CoCl₂ (en) (NH₃)₂]⁺
_Sol.

7. cis M(AA)₂bc
Example 18: Draw the optical isomers of [CoCl (en)₂ Br]²⁺
_Sol.

Stability of Coordination Compounds

The stability of a complex in solution refers to the degree of association between the two species involved in the state of equilibrium. If we have a reaction of the type:
M + 4L ⇌ ML₄
then the larger the stability constant, the higher the proportion of ML₄ that exists in solution. Free metal ions rarely exist in the solution so that M will usually be surrounded by solvent molecules which will compete with the ligand molecules, L, and be successively replaced by them. For simplicity, we generally ignore these solvent molecules and write four stability constants as follows :

where K₁, K₂, etc., are referred to as stepwise stability constants. Alternatively, we can write the overall stability constant thus :

M + 4L ⇌ ML₄     
β₄ = [ML₄]/[M][L]⁴

The stepwise and overall stability constant are therefore related as follows:
β₄ = K₁ × K₂ × K₃ × K₄   or more generally,
β_n = K₁ × K₂ × K₃ × K₄ ......... K_n

The instability constant or the dissociation constant of coordination compounds is defined as the reciprocal of the formation constant.

Importance And Applications Of Coordination Compounds

1. Analytical chemistry :The analytical applications of coordination chemistry are in
(a) Qualitative and quantitative analysis :Metals ions form colored coordination compounds on reaction with a number of ligands. These reactions are used for detection of the metal ions. The colored complexes formed can be used for the estimation of metals by classical or instrumental methods such as gravimetry or colorimetry. Some examples are given as follows :
The presence of iron ions (Fe³⁺) can be detected by the addition potassium ferrocyanide solution, which results in formation of Prussian blue complex.
Fe³⁺ + K₄[Fe(CN)₆] → KFe[Fe(CN)₆] + 2K⁺

(b) Volumetric analysis : Hardness of water can be estimated by titration with EDTA. The metal ions causing hardness, that is Ca²⁺ and Mg²⁺, form stable complexes with EDTA.

2. Metal extraction and purification : Extraction of metals, such as silver and gold, is carried out by forming their water soluble cyanide complexes with the ore. Pure gold can then be obtained from the solution by addition of zinc. Similarly, metals can be purified by formation and then decomposition of their coordination compounds. For example, impure nickel obtained after extraction may be converted into pure nickel by first converting it to nickel carbonyl and then decomposing.

3. Catalysis :Coordination compounds are used as catalysts in important commercial processes. For example,

1. The Zeigler-Natta catalyst (TiCl₄ and trialkyl aluminium) is used as a catalyst in the formation of polyethene.
2. The Wilkinson catalyst - RhCl(PPh₃)₃ is used in the hydrogenation of alkenes.
3. In the Monsanto acetic acid process, various rhodium complexes, such as [Rh(CO)₂I₂], [Rh(Cl)(CO)(PPh₃)₂] or [Rh(Cl)(CO)₂]₂, are used as catalyst in the presence of CH₃I, I₂ or HI as activator.

4. Electroplating : Coordination compounds of gold, silver and copper are used as components in the baths used for electroplating articles of other metals with these metals. For example, in silver plating, K[Ag(CN)₂] is used as an electrolyte; in gold plating, K[Au(CN)₂] is used as an electrolyte; and in copper plating, K₃[Cu(CN)₄] is used as an electrolyte.

5. Biological importance : Some important biological compounds are coordination complexes. For example, chlorophyll is a complex of Mg²⁺. This green pigment plays a vital role in photosynthesis in plants. Similarly, haemoglobin, the red pigment present in blood, is a coordination complex of Fe²⁺ and vitamin B₁₂, an essential nutrient, is a complex compound of Co³⁺.

6. Medicinal uses : Complexing or chelating agents are used in treating metal poisoning, wherein, the coordination complex is formed between toxic metal in excess metal and the complexing agent. For example, EDTA is used in lead poisoning. EDTA, when injected intravenously into the bloodstream, traps lead forming a compound that is flushed out of the body with the urine. Other heavy metal poisonings that can be treated similarly with chelation therapy are mercury, arsenic, aluminium, chromium, cobalt, manganese, nickel, selenium, zinc, tin and thallium. Similarly, chelating ligands D-penicillamine and desferrioxime B are used for removal of excess copper and iron, respectively. New potent drugs are being created using various derivatives of metallocene. A platinum complex [PtCl₂(NH₃)₂] called cis-platin is used in treatment of cancer.