Coordination compounds

 1. Double Salts: Double salts are the addition compounds which are stable in solid state but break up into constituent ions when dissolved in water or any other solvent. In these compounds the individual properties of constituent are not lost. Some common examples are:

Mohr’s salt	FeSO4.(NH4)2SO4.6H2O
Potash alum	K2SO4.Al2(SO4)3.24H2O
Carnallite	KCl.MgCl2.6H2O

2. Complex Ion: An electrically charged species, carrying positive or negative charge, in which the central metal atom or ion is surrounded by fixed number of ions or molecules, e.g., [Co(NH3)4]2+, [Fe(CN)6]4–, etc. The complex ion does not dissociate into simple ions in aqueous solutions.

3. Coordination Compounds: Coordination compounds are the compounds which contain complex ions, e.g., [Co(NH3)6]Cl3, K4[Fe(CN)6], etc. These compounds contain a central metal atom or cation which is attached with a fixed number of anions or molecules (called ligands) through coordinate bonds.

4. Coordination Entity: A coordination entity constitutes a central metal atom or ion bonded to a fixed number of molecules or ions (ligands), e.g., [Co(NH3)3Cl3], [Fe(CN)6]4–, [Cu(NH3)4]2+, etc.

5. Central Atom or Ion: In a coordination entity, the atom or ion to which a fixed number of ions or molecules are bound in a definite geometrical arrangement. For example, in the complex ion [CoCl(NH3)5]2+, the
Co2+ ion is the central ion.

6. Ligands: Ligands are the atoms, molecules or ions which donate a pair of electrons to central metal atom or ion and form a coordinate bond with it. Depending upon the number of donor atoms available for coordination, the ligands may be classified as:

n Unidentate ligands: Contain one donor atom, e.g.,  etc.

n Bidentate ligands: Contain two donor atoms, e.g.,

n Polydentate ligands: Contain several donor atoms, e.g.,

n Ambidentate ligand: A ligand which contains two donor atoms but only one of them forms a coordinate bond at a time with central metal atom/ion is called an ambidentate ligand. Some common examples are given below:

n Chelating ligand: When a bidentate or a polydentate ligand uses its two or more donor atoms to bind a single metal ion, then a ring-like structure is obtained. It is called chelate and the ligand is known as chelating ligand. The chelating ligands form more stable complexes than the unidentate ligands. This is because when chelation occurs entropy increases and the process becomes more favourable.

7. Coordination Number (CN): The coordination number of a metal ion in a complex may be defined as the total number of ligand donor atoms to which the metal ion is directly bonded. For example, in the complex ions, [Co(NH3)6]3+ and [Fe(C2O4)3]3–, the coordination numbers of both Co and Fe is 6.

8. Coordination Sphere and Counter Ions: The central metal atom or ion and the ligands directly attached to it are enclosed in a square bracket and is collectively termed as the coordination sphere. The ionisable groups are written outside the bracket and are called counter ions. For example, in the complex [Cr(NH3)3(H2O)3]Cl3, the coordination sphere is [Cr(NH3)3(H2O)3]3+ and the counter ion is Cl–.

9. Coordination Polyhedron: The spatial arrangement of the ligand atoms which are directly attached to the central atom/ion is known as the coordination polyhedron around the central atom/ion. Tetrahedral, square planar, octahedral, square pyramidal and trigonal bipyramidal are common shapes of coordination polyhedra.

10. Oxidation Number of Central Atom: The oxidation number of the central atom in a complex is defined as the charge it would carry if all the ligands are removed along with the electron pairs that are shared with the central atom.

11. Homoleptic Complex: The complex in which metal atom is bound to only one kind of donor groups, e.g., [Cu(CN)4]3–.

12. Heteroleptic Complex: The complex in which metal atom is bound to more than one kind of donor groups, e.g., [Co(NH3)4Cl2]+.

13. Table 9.1: Nomenclature of Some Coordination Compounds

S. No.	Formula	Name
(i)	[Pt(NH3)2ClNO2]	diamminechloridonitrito–N–platinum(II)
(ii)	[CoCl2(en)2]Cl	dichloridobis(ethane-1, 2–diamine) cobalt(III) chloride
(iii)	K3[Fe(C2O4)3]	potassium trioxalatoferrate(III)
(iv)	[Ag(NH3)2] [Ag(CN)2]	diamminesilver(I) dicyanoargentate(I)

14. Isomerism: Two or more compounds having the same molecular formula but different arrangement of atoms and hence differ in one or more physical or chemical properties are called isomers and the phenomenon is called isomerism.

Structural isomers have different bonds. Stereoisomers have the same chemical formula and chemical bonds but differ in their spatial arrangements.

(a) Structural isomerism:

(i) Linkage isomerism: This type of isomerism arises due to the presence of an ambidentate ligand in a coordination compound. Some examples of linkage isomers are:

[Co(NH3)5NO2]Cl2 and [Co(NH3)5ONO]Cl2

[Mn(CO)5SCN] and [Mn(CO5)NCS]

(ii) Coordination isomerism: This type of isomerism arises due to interchange of ligands between the cationic and anionic entities of different metal ions present in a complex. Some examples are

[Co(NH3)6] [Cr(CN)6] and [Cr(NH3)6][Co(CN)6]

[Pt(NH3)4] [PtCl4] and [PtCl(NH3)3] [PtCl3(NH3)]

(iii) Ionisation isomerism: Ionisation isomerism arises when the counter ion in a complex salt is itself a potential ligand and can displace a ligand which can then become the counter ion. Some examples are:

[Co(NH3)5SO4]Br and [Co(NH3)5Br]SO4

[Pt(NH3)4Cl2]Br2 and [Pt(NH3)4Br2]Cl2

[Co(NH3)5NO2]SO4 and [Co(NH3)5SO4]NO2

(iv) Solvate isomerism: Solvate isomers differ by whether or not a solvent molecule is directly bonded to metal ion or merely present as free solvent molecule in the crystal lattice. If water is the solvent then this form of isomerism is known as “hydrate isomerism.

Some examples in which hydrate isomerism is observed are:

[Cr(H2O)6]Cl3 and [Cr(H2O)5Cl]Cl2.H2O

[Co(NH3)4(H2O)Cl]Cl2 and [Co(NH3)4Cl2]Cl.H2O

[Co(NH3)4(H2O)Cl]Br2 and [Co(NH3)4Br2]Cl.H2O

(b) Stereoisomerism: Stereoisomerism is of two types:

(i) Geometrical isomerism: This type of isomerism arises in heteroleptic complexes due to difference in geometrical arrangement of the ligand around the central metal ion. If the same kind of ligand occupy adjacent positions, the isomer is called cis, and if these are opposite to each other, the isomer is called trans.

This type of isomerism is very common in complexes with coordination number 4 and 6.

l Isomerism in complexes with coordination number 4.

These complexes can either have a tetrahedral or square planar geometry. Tetrahedral complexes do not show geometrical isomerism as relative position of the ligands attached to central metal atom is same with respect to each other (adjacent). The square planar complexes on the other hand show this type of isomerism as given below:

n Tetra coordinated square planar complexes of the type [MA4], [MA3X], [MAX3] are incapable of showing geometrical isomerism because all possible arrangements of ligands in each of these complexes are exactly the same.

n Type [MA2X2]: [Pt(NH3)2Cl2], [Pd(NH3)2(NO2)2], [Pt(Py)2Cl2]

n Type [MA2XY]: [Co(NH3)2ClBr], [Pt(Py)2(NH3)Cl]

n Type [MABXY]: [Pt(NH3)Py(NH2OH)(NO2)]+

n Type [M(AB)2]: [Pt(gly)2] where gly = H2NCH2COO–

l Isomerism in complexes with coordination number 6.

n Octahedral complexes of the type [MA6] and [MA5X] are incapable of showing geometrical isomerism.

n Type [MA4X2] or [MA2X4]: [Co(NH3)4Cl2]+, [Cr(NH3)4Cl2]+ ion

n Type [MA3X3]: [Rh(Py)3Cl3], [Co(NH3)3Cl3], [Co(NH3)3(NO2)3]

n Type [M(AA)2X2] or [M(AA)2XY]: [Co(en)2Cl2]+, [Ni(ox)2Cl2]+

(ii) Optical isomerism: This type of isomerism is exhibited by chiral molecules. Optical isomers are mirror images that cannot be superimposed on one another. These are called as enantiomers and rotate the plane of polarised light equally but in opposite directions. The isomer which rotates the plane of polarised light towards left is called laevorotatory (l) while which rotates the plane towards right is called dextrorotatory (d).

Optical isomerism is common in octahedral complexes involving bidentate ligands. Some examples of different types of octahedral complexes showing optical isomerism are given below:

n Type [M(AA)3]: [Cr(ox)3]3–, [Co(en)3]3+

n Type [M(AA)2X2] or [M(AA)2XY]: cis-[Co(en)2Cl2]+, cis-[Pt(en)2Cl2]+, cis-[Cr(ox)2Cl2]3–, etc.

n Type [M(AA)X2Y2]: [Co(en)(NH3)2Cl2]

15. Werner’s Theory of Coordination Compounds: The main postulates of this theory are:

(i) In coordination compounds, metals show two types of valencies: primary and secondary.

(ii) The primary valencies are normally ionisable and are satisfied by negative ions.

(iii) The secondary valencies are non ionisable and are satisfied by neutral molecules or negative ions. The secondary valency is equal to the coordination number and is fixed for a metal.

(iv) The ions or groups bound by secondary linkage to the metal have characteristic spatial arrangements corresponding to different coordination number.

16. Bonding in Coordination Compounds

(a) Valence bond theory: The main features of the valence bond theory as applied to coordination compounds are as follows:

(i) The number of metal–ligand coordinate bonds which can be formed in case of given metal ion depends upon the number of vacant orbitals for bonding in metal ion and is known as the coordination number of metal ion.

(ii) The metal atom or ion under the influence of ligands uses its (n – 1) d, ns, np or ns, np, nd orbitals for hybridisation to yield a set of equivalent orbitals of definite geometry.

Table 9.2: Number of Orbitals and Types of Hybridisations

Coordination Number	Type of Hybridisation	Distribution of Hybrid Orbitals in Space
4	sp3	Tetrahedral
4	dsp2	Square planar
5	sp3d	Trigonal bipyramidal
6	sp3d 2	Octahedral
6	d 2sp3	Octahedral

(iii) The empty hybrid orbitals of the central metal atom or ion overlap with the filled orbitals of the ligand containing the lone pair of electrons. In this way a metal–ligand coordinate bond is formed.

(iv) The inner orbital (low spin) or the outer orbital (high spin) complexes are formed depending upon whether the d-orbitals of inner shell or d-orbitals of outer shell are used in hybridisation.

(v) The complex will be diamagnetic if all electrons are paired. If unpaired electrons are present then the complex will be paramagnetic.

Table 9.3: Application of Valence Bond Theory to Some Complexes

Ion/Complex	Central Metal Ion	Configuration of Metal Ion	Hybridisation of Metal Ion Involved	Geometry of the Complex	Number of Unpaired Electrons	Magnetic Behaviour
[Ti(H2O)6]3+	Ti3+	d 1	d 2 sp3	Octahedral	1	Paramagnetic
[V(H2O)6]3+	V3+	d 2	d 2 sp3	Octahedral	2	Paramagnetic
[Cr(H2O)6]3+	Cr3+	d 3	d 2 sp3	Octahedral	3	Paramagnetic
[Cr (NH3)6]3+	Cr3+	d 3	d 2 sp3	Octahedral	3	Paramagnetic
[MnF6]3–	Mn3+	d 4	sp3 d 2	Octahedral	4	Paramagnetic
[Mn(CN)6]3–	Mn3+	d 4	d 2 sp3	Octahedral	2	Paramagnetic
[MnCl4]2–	Mn2+	d 5	sp3	Tetrahedral	5	Paramagnetic
[FeF6]3–	Fe3+	d 5	sp3 d 2	Octahedral	5	Paramagnetic
[Fe(H2O)6]3+	Fe3+	d 5	sp3 d 2	Octahedral	5	Paramagnetic
[Fe(CN)6]3–	Fe3+	d 5	d 2 sp3	Octahedral	1	Paramagnetic
[Fe(CN)6]4–	Fe2+	d 6	d 2 sp3	Octahedral	0	Diamagnetic
[FeCl4 ]2–	Fe2+	d 6	sp3	Tetrahedral	4	Paramagnetic
[Co(NH3)6]3+	Co3+	d 6	d 2 sp3	Octahedral	0	Diamagnetic
[CoF6]3–	Co3+	d 6	sp3 d 2	Octahedral	4	Paramagnetic
[Ni(CO)4]	Ni	3d 8 4s2	sp3	Tetrahedral	0	Diamagnetic
[Ni(CN)4 ]2–	Ni2+	d 8	dsp2	Square planar	0	Diamagnetic
[NiCl4 ]2–	Ni2+	d 8	sp3	Tetrahedral	2	Paramagnetic
[Ni(H2O)6]2+	Ni2+	d 8	sp3d 2	Octahedral	2	Paramagnetic
[CuCl4 ]2–	Cu2+	d 9	sp3	Tetrahedral	1	Paramagnetic
[Zn(NH3)4 ]2+	Zn2+	d 10	sp3	Tetrahedral	0	Diamagnetic
[Pt(NH3)Cl2]	Pt2+	d 8	dsp2	Square planar	0	Diamagnetic

Limitations of valence bond theory

(i) It involves a number of assumptions.

(ii) It does not give quantitative interpretation of magnetic data.

(iii) It does not distinguish between weak and strong ligands.

(iv) It does not explain the colour exhibited by complexes.

(v) It does not give an exact explanation of thermodynamic or kinetic stabilities of coordination compounds.

(vi) It does not make exact predictions regarding the tetrahedral and square planar structures of 4 – coordinated complexes.

(b) Crystal field theory

n According to crystal field theory, the bonding between central metal ion and ligand is purely electrostatic. Ligands are considered as point charges in case of anions or dipoles in case of neutral molecules.

n The five d-orbitals in an isolated gaseous metal atom/ion have same energy, i.e., they are degenerate orbitals. However, in a complex due to asymmetrical negative field of ligands, the d-orbitals are no more degenerate and split into two sets of orbitals. The pattern of splitting depends upon the nature of the crystal field.

n In an octahedral environment, the dx2 – y2 and dz2 orbitals which point towards the axes along the direction of ligand will experience more repulsion and will be raised in energy; and the dxy, dyz and dxz orbitals which are directed between the axes will be lowered in energy as compared to average energy in the spherical crystal field. Thus, in an octahedral complex, the degeneracy of the five d-orbitals is partially removed due to ligand electron–metal electron repulsions to yield three orbitals of lower energy, t2g set and two orbitals of higher energy, eg set.

n The splitting of the degenerate orbitals due to the presence of ligands in a definite geometry is known as crystal field splitting and the difference of energy between two sets of degenerate orbitals as a result of crystal field splitting is known as crystal field stabilisation energy (CFSE). It is usually denoted by symbol Do (the subscript o stands for octahedral).

It is found that eg orbitals are 35Δ0 above the average energy level and t2g orbitals are 25Δ0 below the average energy level.

n The magnitude of ∆0 depends upon the field produced by ligand and charge on the metal ion. The arrangement of ligands in a series in the order of increasing field strength is called spectrochemical series.

I – < Br– < SCN– < Cl– < S2– < F – < OH – < C2O42– < H2O < NCS– < EDTA4 – < NH3 < en < CN– < CO

n In d 2 and d 3 coordination entities, the d-electrons occupy the t2g orbitals singly in accordance with Hund’s rule. For d 4 ions, the electronic distribution depends on crystal field stabilisation energy (D0) and pairing energy (P). The two options are:

(i) If D0 < P, the fourth electron enters one of the eg orbitals giving the configuration t32g e1g. Ligands for which D0 < P are known as weak field ligands and form high spin complexes.

(ii) If D0 > P, it becomes more energetically favourable for the fourth electron to occupy a t2g orbital with configuration t42g e0g. Thus, ligands for which Do > P, are known as strong field ligands and form low spin complexes.

n In tetrahedral coordination entities, the d-orbitals splitting is inverted and is smaller as compared to the octahedral field splitting.

For the same metal, the same ligand and metal ligand distances, Δt=49Δ0. Consequently, the orbital splitting energies are not sufficiently large for forcing pairing, therefore, low spin configuration are rarely observed.

l Colour in coordination compounds: The crystal field theory attributes the colour of the coordination compounds to d–d transition of the electron from t2g to eg orbitals in octahedral complexes and from eg to t2g orbitals in tetrahedral complexes. In the absence of ligand, crystal field splitting does not occur and hence the substance is colourless. For example, removal of water from [Ti(H2O)6]Cl3 on heating renders it colourless. Similarly, anhydrous CuSO4 is white, but CuSO4.5H2O is blue in colour.

l Limitations of crystal field theory:

(i) As the ligands are considered as point charges, the anionic ligands should exert greater splitting effect. However, the anionic ligands are found at the low end of the spectrochemical series.

(ii) It does not take into account the covalent character of metal ligand bond.

17. Metal Carbonyls: Metal carbonyls are the organometallic compounds in which carbon monoxide acts as the ligand. These compounds are formed by most of the transition metals. Structures of some metal carbonyls are given below:

18. Bonding in Metal Carbonyls: The metal–carbon bond in metal carbonyls have both s and p character. The metal–carbon s-bond is formed by the donation of lone pair of electrons from the carbonyl carbon into a vacant orbital of the metal. The metal–carbon p-bond is formed by the donation of a pair of electrons from a filled d-orbital of metal into the vacant anti-bonding pi-molecular orbital (p*) of carbon monoxide. The metal to ligand bonding creates a synergic effect which strengthens the bond between CO and the metal.

19. 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. The stability of a complex MLn is measured in terms of magnitude of (stability or formation) equilibrium constant. For a reaction of the type

We can write stability constant as follows:

Where K1, K2, K3 etc., are referred to as stepwise stability constants.

Alternatively, we can write the overall stability constant (βn) as:

M + nL MLn ; βn = [MLn]/[M][L]n

The stepwise and overall stability constant are therefore as follows:

βn = K1 × K2 × K3 × .... × Kn

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

20. Factors Affecting the Stability of Complexes

The stability of a complex depends upon

(i) The nature of the central ion: Greater the charge density (i.e., charge/radius ratio) on the central metal ion, greater is the stability of the complex. For example, complexes of Fe3+ are more stable than Fe2+. This is supported by the fact that the stability constant of [Fe(CN)6]3– is 1.21 × 1031 and that of [Fe(CN)6]4– is only 1.8 × 106.

(ii) Nature of the ligand: In general, more basic ligands have a tendency to donate the electron pairs to central metal ion more easily and hence the resulting complex is very stable. For example, the complexes involving F – ions are more stable than those involving Cl – ions or Br– ions.

(iii) Chelate effect: When chelation occurs, entropy increases and therefore, the formation of the complex becomes more favourable. For example, chelate [Cd(en)4]2+ is 10,000 times more stable than the simple complex [Cd(CH3NH2)2]2+.

21. Applications of Coordination Compounds

 In metallurgical separation: Extraction of several metals from their ores involves complex formation. For example, silver and gold are extracted from their ores as cyanide complex.

4Au+8KCN+2H2O+O2→4K[Au(CN)2]Potassiumdicyanoaurate(I)+4KOH

2K[Au(CN)2]+Zn→K2[Zn(CN)4]Potassiumtetracyanozincate(II)+2Au⏐↓⏐⏐

Purification of some metals can be achieved through complex formation. For example, in Mond’s process, impure nickel is converted to [Ni(CO)4], which is decomposed to yield pure nickel.

 In analytical chemistry: Complex formation is frequently encountered in qualitative and quantitative chemical analysis.

(i) Qualitative Analysis:

Detection of Cu2+ is based on the formation of a blue tetraammine copper (II) ion.

Cu2++4NH3→[Cu(NH3)4]2+DeepBlue

Ni2+ is detected by the formation of a red complex with dimethyl glyoxime (DMG).

NiCl2+2DMG+2NH4OH→[Ni(DMG)2]Nickelbisdimethylglyoximate+2NH4Cl+2H2O

The separation of Ag+ and Hg2+ in group I is based on the fact that while AgCl dissolves in NH3, forming a soluble complex, Hg2Cl2 forms an insoluble black substance.

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

Hg2Cl2(s)+NH3→Hg(NH2)Cl.HgBlackinsolublesubstance+HCl

(ii) Quantitative Analysis

Gravimetric estimation of Ni2+ is carried out by precipitating Ni2+ as red nickel dimethyl oxime complex in the presence of ammonia.

Ni2++2DMG−→NH3[Ni(DMG)2]Redppt.

EDTA is used in the complexometric determination of several metal ions such as Ca2+, Zn2+, Fe2+, Co2+, Ni2+, etc.

l In Medicinal Chemistry:

n The platinum complex, cis-[Pt(NH3)2Cl2], known as cisplatin is used in the treatment of cancer.

n EDTA complex of calcium is used in the treatment of lead poisoning. Ca-EDTA is a weak complex; when it is administered, calcium in the complex is replaced by the lead present in the body and is eliminated with the urine.

n The excess of copper and iron present in animal system are removed by the chelating ligands, D-penicillamine and desferrioxime B via the formation of complexes.

l In Biological System:

n Haemoglobin, the red pigment of blood which acts as oxygen carrier is a complex of Fe2+ with porphyrin.

n The pigment responsible for photosynthesis, chlorophyll, is a complex of Mg2+ with porphyrin.

n Vitamin B12 (Cyanocobalamine), the antipernicious anaemia factor, is a complex of cobalt.

 In the Estimation of Hardness of Water: The hardness of water is estimated by simple titration with Na2EDTA. The Ca2+ and Mg2+ ions form stable complexes with EDTA. The selective estimation of these ions can be done due to difference in the stability constants of calcium and magnesium complexes.

 In Photography: In black and white photography, the developed film is fixed by washing with hypo solution which dissolves the undecomposed AgBr to form a complex ion, [Ag(S2O3)2]3–.

 In Catalysis: Coordination compounds are used as catalysts for many industrial processes. For example, Wilkinson’s catalyst, a rhodium complex having formula [(Ph3P)3RhCl] is used for the selective hydrogenation of alkenes.