Transition metals have one of the important abilities to form complex compounds in which the metal atoms are bound to a number of ions or neutral molecules by coordinate bonds. Such compounds are called coordinate compounds. In this unit, you will study different aspects of coordination compounds including their nomenclature and applications. You will also appreciate the postulates of Werner's theory of coordination compounds. You will also be able to understand the nature of bonding in coordination compounds in terms of the Valence Bond and Crystal Field theories.
(i) Hexaamminecobalt (III) chloride
(ii) Pentaamminechloridocobalt (III) chloride
(iii) Potassium hexacyanoferrate (III)
(iv) Potassium trioxalatoferrate (III)
(v) Potassium tetrachloridopalladate (II)
(vi) Diamminechlorido(methylamine)platinum (II) chloride
When ionization isomers are dissolved in water, they ionize to give different ions. These ions then react differently with different reagents to give different products.
Ni is in the +2 oxidation state i.e., in d8 configuration.
There are 4 CN- ions. Thus, it can either have a tetrahedral geometry or square planar geometry. Since CN- ion is a strong field ligand, it causes the pairing of unpaired 3d electrons.
It now undergoes dsp2 hybridization. Since all electrons are paired, it is diamagnetic.
In case of [Ni(CN)4]2- , Cl- ion is a weak field ligand. Therefore, it does not lead to the pairing of unpaired 3d electrons. Therefore, it undergoes sp3 hybridization.
Since there are 2 unpaired electrons in this case, it is paramagnetic in nature.
Though both [NiCl4]2- and [Ni(CO)4] are tetrahedral, their magnetic characters are different. This is due to a difference in the nature of ligands. Cl- is a weak field ligand and it does not cause the pairing of unpaired 3d electrons. Hence, [NiCl4]2- is paramagnetic.
In Ni(CO)4, Ni is in the zero oxidation state i.e., it has a configuration of 3d8 4s2.
But CO is a strong field ligand. Therefore, it causes the pairing of unpaired 3d electrons. Also, it causes the 4s electrons to shift to the 3d orbital, thereby giving rise to sp3 hybridization. Since no unpaired electrons are present in this case, [Ni(CO)4] is diamagnetic.
| [Co(NH3)6]3+ | [Ni(NH3)6]2+ |
| Oxidation state of cobalt = +3 | Oxidation state of Ni = +2 |
| Electronic configuration of cobalt = d6 | Electronic configuration of nickel = d8 |
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[Pt(CN)4]2-
In this complex, Pt is in the +2 state. It forms a square planar structure. This means that it undergoes dsp2 hybridization. Now, the electronic configuration of Pd(+2) is 5d8.
CN - being a strong field ligand causes the pairing of unpaired electrons. Hence, there are no unpaired electrons in [Pt(CN)4]2-
| [Mn(H2O)6]2+ | [Mn(CN)6]4- |
| Mn is in the +2 oxidation state. | Mn is in the +2 oxidation state. |
| The electronic configuration is d5. | The electronic configuration is d5. |
| The crystal field is octahedral. Water is a weak field ligand. Therefore, the arrangement of the electrons in [Mn(H2O)6]2+ is t2g3eg2. |
The crystal field is octahedral. Cyanide is a strong field ligand. Therefore, the arrangement of the electrons in [Mn(CN)6]4- is T2g5eg0. |
| Hence, hexaaquo manganese (II) ion has five unpaired electrons | while hexacyano ion has only one unpaired electron. |
β4 = 2.1 × 1013
The overall complex dissociation equilibrium constant is the reciprocal of the overall stability constant, β4.
1/β4 = 1 / 2.1 × 1013
= 4.7 x 10-14
In order to explain the nature of bonding and structure of coordination compounds, Werner in 1893 proposed a theory called as Werner’s Coordination Theory. The postulates of theory are as follows:
1. In coordination compounds, metal atoms exhibit 2 types of valencies namely, primary and secondary valency. The primary valency is ionizable whereas the secondary valency is non- ionizable, i.e. primary valency corresponds to oxidation state and secondary valency corresponds to coordination number.
2. Every metal atom has a fixed number of secondary valencies i.e. fixed coordination number.
3. The metal atom tends to satisfy both its primary as well as secondary valencies. Primary valencies are satisfied by negative ions whereas secondary valencies are satisfied by negative or by neutral molecules.
4. The secondary valencies are always directed towards the fixed positions in space and this leads to definite geometry of the coordination compound. For example: if the metal ion has four secondary valencies, these are arranged in either tetrahedral or square planar arrangement around the central metal ion. The secondary valencies, thus, determine the stereochemistry of the complex.
Both the compounds i.e., FeSO4 .(NH4)2 SO4 .6H2O and [Cu(NH3)4] SO4 .5H2O fall under the category of addition compounds with only one major difference i.e., the former is an example of a double salt, while the latter is a coordination compound.
A double salt is an addition compound that is stable in the solid state but that which breaks up into its constituent ions in the dissolved state. These compounds exhibit individual properties of their constituents. For e.g. FeSO4 .(NH4)2 SO4 .6H2O breaks into Fe2+, NH4+, and SO42 - ions. Hence, it gives a positive test for Fe2+ ions.
A coordination compound is an addition compound which retains its identity in the solid as well as in the dissolved state. However, the individual properties of the constituents are lost. This happens because [Cu(NH3)4] SO4 .5H2O does not show the test for Cu2+. The ions present in the solution of [Cu(NH3)4] SO4 .5H2O are [Cu(NH3)4]2+ and SO42-.
(i) Coordination entity:
A coordination entity is an electrically charged radical or species carrying a positive or negative charge. In a coordination entity, the central atom or ion is surrounded by a suitable number of neutral molecules or negative ions ( called ligands). For example:
[Ni(NH3)6]2+ , [Fe(CN6)]4+ = cationic complex
[PtCl4]2- , [Ag(CN)2]- = anionic complex
[Ni(CO)4] , [Co(NH3)4 Cl2] = neutral complex
(ii) Ligands The neutral molecules or negatively charged ions that surround the metal atom in a coordination entity or a coordinal complex are known as ligands. For example,
, Cl - , - OH. Ligands are usually polar in nature and possess at least one unshared pair of valence electrons.
(iii) Coordination number: The total number of ligands (either neutral molecules or negative ions) that get attached to the central metal atom in the coordination sphere is called the coordination number of the central metal atom. It is also referred to as its ligancy.
For example:
(a) In the complex, K2[PtCl6], there as six chloride ions attached to Pt in the coordinate sphere. Therefore, the coordination number of Pt is 6.
(b) Similarly, in the complex [Ni(NH3)4]Cl2, the coordination number of the central atom (Ni) is 4.
(iv) Coordination polyhedron: Coordination polyhedrons about the central atom can be defined as the spatial arrangement of the ligands that are directly attached to the central metal ion in the coordination sphere. For example: (a) (b) Tetrahedral
(v) Homoleptic complexes: These are those complexes in which the metal ion is bound to only one kind of a donor group. For eg: [Co(NH3)6]3+ , [PtCl4]2- etc.
(vi) Heteroleptic complexes: Heteroleptic complexes are those complexes where the central metal ion is bound to more than one type of a donor group. For e.g.: [Co(NH3)4 Cl2]+ , [Co(NH3)5 Cl]2+
A ligand may contain one or more unshared pairs of electrons which are called the donor sites of ligands. Now, depending on the number of these donor sites, ligands can be classified as follows:
(a) Unidentate ligands: Ligands with only one donor sites are called unidentate ligands. For e.g., 
, Cl - etc.
(b) Didentate ligands: Ligands that have two donor sites are called didentate ligands. For e.g.,
(The donar atom is N)
(i) [Zn(OH)4]2-
(ii) K2[PdCl4]
(iii) [Pt(NH3)2Cl2]
(iv) K2[Ni(CN)4]
(v) [Co(ONO) (NH3)5]2+
(vi) [Co(NH3)6]2 (SO4)3
(vii) K3[Cr(C2O4)3]
(viii) [Pt(NH3)6]4+
(ix) [Cu(Br)4]2-
(x) [Co[NO2)(NH3)5]2+
(i) Hexaamminecobalt (III) chloride
(ii) Diamminechlorido (methylamine) platinum (II) chloride
(iii) Hexaquatitanium (III) ion
(iv) Tetraamminichloridonitrito-N-Cobalt (III) chloride
(v) Hexaquamanganese (II) ion
(vi) Tetrachloridonickelate (II) ion
(vii) Hexaamminenickel (II) chloride
(viii) Tris (ethane-1, 2-diammine) cobalt (III) ion
(ix) Tetracarbonylnickel (0)
(a) Geometric isomerism:
This type of isomerism is common in heteroleptic complexes. It arises due to the different possible geometric arrangements of the ligands. For example:
(b) Optical isomerism:
This type of isomerism arises in chiral molecules. Isomers are mirror images of each other and are non-superimposable.
(c) Linkage isomerism:
This type of isomerism is found in complexes that contain ambidentate ligands. For example:
[Co(NH3)5 (NO2)]Cl2 and [Co(NH3)5 (ONO)Cl2
Yellow form Red form
(d) Coordination isomerism:
This type of isomerism arises when the ligands are interchanged between cationic and anionic entities of differnet metal ions present in the complex.
[Co(NH3)6] [Cr(CN)6] and [Cr(NH3)6] [Co(CN)6]
(e) Ionization isomerism:
This type of isomerism arises when a counter ion replaces a ligand within the coordination sphere. Thus, complexes that have the same composition, but furnish different ions when dissolved in water are called ionization isomers. For e.g., Co(NH3)5SO4)Br and Co(NH3)5Br]SO4.
(f) Solvate isomerism: Solvate isomers differ by whether or not the solvent molecule is directly bonded to the metal ion or merely present as a free solvent molecule in the crystal lattice.
[Cr[H2O)6]Cl3 [Cr(H2O)5Cl]Cl2 . H2O [Cr(H2O)5Cl2]Cl . 2H2O
Violet Blue-green Dark green
(i) For[Cr(C2O4)3]3-, no geometric isomer is possible as it is a bidentate ligand.
(ii) [Co(NH3)3Cl3]
Two geometrical isomers are possible.
(i) [Cr(C2O4)3]3-
Mirror
(ii) [Co(NH3)Cl(en)2]2+
[Pt(NH3)(Br)(Cl)(py)
From the above isomers, none will exhibit optical isomers. Tetrahedral complexes rarely show optical isomerization. They do so only in the presence of unsymmetrical chelating agents.
Aqueous CuSO4 exists as [Cu(H2O)4]SO4. It is blue in colour due to the presence of
[Cu[H2O)4]2+ ions.
(i) When KF is added:
(ii) When KCl is added:
In both these cases, the weak field ligand water is replaced by the F - and Cl - ions.
CuSO4(aq) + 4KCN(aq) → K2[Cu(CN)4](aq) + K2SO4(aq)
i.e., [Cu(H2O)4]2+ + 4CN- → [Cu(CN)4]2- + 4H2O
Thus, the coordination entity formed in the process is K2[Cu(CN)4]. K2[Cu(CN)4 is a very stable complex, which does not ionize to give Cu2+ ions when added to water. Hence, Cu2+ ions are not precipitated when H2S(g) is passed through the solution.
(i) [Fe(CN)6]4-
In the above coordination complex, iron exists in the +II oxidation state.
Fe2+ : Electronic configuration is 3d6
Orbitals of Fe2+ ion:
Hence, the geometry of the complex is octahedral and the complex is diamagnetic (as there are no unpaired electrons).
(ii) [FeF6]3-
In this complex, the oxidation state of Fe is +3.
Orbitals of Fe+3 ion:
Hence, the geometry of the complex is found to be octahedral.
(iii) [Co(C2O4)3]3-
Cobalt exists in the +3 oxidation state in the given complex.
Orbitals of Co3+ ion:
Hence, the geometry of the complex is found to be octahedral.
(iv) [CoF6]3- Cobalt exists in the +3 oxidation state.
Orbitals of Co3+ ion:
Again, fluoride ion is a weak field ligand. It cannot cause the pairing of the 3d electrons. As a result, the Co3+ ion will undergo sp3d2 hybridization. sp3d2 hybridized orbitals of Co3+ ion are:
The splitting of the d orbitals in an octahedral field takes palce in such a way that dx2y2 , dz2 experience a rise in energy and form the eg level, while dxy, dyzand dzx experience a fall in energy and form the t2g level.
A spectrochemical series is the arrangement of common ligands in the increasing order of their crystal-field splitting energy (CFSE) values. The ligands present on the R.H.S of the series are strong field ligands while that on the L.H.S are weak field ligands. Also, strong field ligands cause higher splitting in the d orbitals than weak field ligands.
I- < Br - < S2- < SCN- < Cl- < N3 < F- < OH- < C2O42- ~ H2O < NCS- ~ H- < CN- < NH3 < en ~ SO32- < NO2- < phen < CO
The degenerate d-orbitals (in a spherical field environment) split into two levels i.e., eg and t2g in the presence of ligands. The splitting of the degenerate levels due to the presence of ligands is called the crystal-field splitting while the energy difference between the two levels (eg and t2g) is called the crystal-field splitting energy. It is denoted by Δo.
After the orbitals have split, the filling of the electrons takes place. After 1 electron (each) has been filled in the three t2g orbitals, the filling of the fourth electron takes place in two ways. It can enter the eg orbital (giving rise to t2g3 eg1 like electronic configuration) or the pairing of the electrons can take place in the t2g orbitals (giving rise to t2g4 eg0 like electronic configuration). If the Δo value of a ligand is less than the pairing energy (P), then the electrons enter the eg orbital. On the other hand, if the Δo value of a ligand is more than the pairing energy (P), then the electrons enter the t2g orbital.
Cr is in the +3 oxidation state i.e., d3 configuration. Also, NH3 is a weak field ligand that does not cause the pairing of the electrons in the 3d orbital.
Cr3+
Therefore, it undergoes d2sp3 hybridization and the electrons in the 3d orbitals remain unpaired. Hence, it is paramagnetic in nature.
In [Ni(CN)4]2-, Ni exists in the +2 oxidation state i.e., d8 configuration.
Ni2+:
CN- is a strong field ligand. It causes the pairing of the 3d orbital electrons. Then, Ni2+ undergoes dsp2 hybridization.
As there are no unpaired electrons, it is diamagnetic.
In [Ni(H2O)6]2+, H2Ö is a weak field ligand. Therefore, there are unpaired electrons in Ni2+. In this complex, the d electrons from the lower energy level can be excited to the higher energy level i.e., the possibility of d - d transition is present. Hence, [Ni(H2O)6]2+ is coloured.
In [Ni(CN)4]2 - , the electrons are all paired as CN- is a strong field ligand. Therefore, d-d transition is not possible in [Ni(CN)4]2 - . Hence, it is colourless.
The colour of a particular coordination compound depends on the magnitude of the crystal-field splitting energy, Δ. This CFSE in turn depends on the nature of the ligand. In case of [Fe(CN)6]4- and [Fe(H2O)6]2+, the colour differs because there is a difference in the CFSE. Now, CN- is a strong field ligand having a higher CFSE value as compared to the CFSE value of water. This means that the absorption of energy for the intra d-d transition also differs. Hence, the transmitted colour also differs.
The metal-carbon bonds in metal carbonyls have both σ and π characters. A σ bond is formed when the carbonyl carbon donates a lone pair of electrons to the vacant orbital of the metal. A π bond is formed by the donation of a pair of electrons from the filled metal d orbital into the vacant anti-bonding π orbital (also known as back bonding of the carbonyl group). The σ bond strengthens the π bond and vice-versa. Thus, a synergic effect is created due to this metal-ligand bonding. This synergic effect strengthens the bond between CO and the metal.
(i) K3[Co(C2O4)3]
The central metal ion is Co.
Its coordination number is 6.
The oxidation state can be given as:
x - 6 = -3
x = + 3
The d orbital occupation for Co3+ is t2g6eg0.
(ii) cis-[Cr(en)2Cl2]Cl
The central metal ion is Cr.
The coordination number is 6.
The oxidation state can be given as:
x + 2(0) + 2(-1) = +1
x - 2 = +1
x = +3
The d orbital occupation for Cr3+ is t2g3.
(iii) (NH4)2[CoF4]
The central metal ion is Co.
The coordination number is 4.
The oxidation state can be given as:
x - 4 = -2
x = + 2
The d orbital occupation for Co2+ is eg4 t2g3.
(iv) [Mn(H2O)6]SO4
The central metal ion is Mn.
The coordination number is 6.
The oxidation state can be given as:
x + 0 = +2
x = +2
The d orbital occupation for Mn is t2g3 eg2.
(i) Potassium diaquadioxalatochromate (III) trihydrate.
Oxidation state of chromium = 3
Electronic configuration: 3d3 : t2g3
Coordination number = 6
Shape: octahedral
Stereochemistry:
(ii) [Co(NH3)5Cl]Cl2
IUPAC name: Pentaamminechloridocobalt(III) chloride
Oxidation state of Co = +3
Coordination number = 6
Shape: octahedral.
Electronic configuration: d6: t2g6.
Stereochemistry:
Magnetic Moment = 0
(iii) CrCl3(py)3
IUPAC name: Trichloridotripyridinechromium (III)
Oxidation state of chromium = +3
Electronic configuration for d3 = t2g3
Coordination number = 6
Shape: octahedral.
Stereochemistry:
Both isomers are optically active. Therefore, a total of 4 isomers exist.
(iv) Cs[FeCl4]
IUPAC name: Caesium tetrachloroferrate (III)
Oxidation state of Fe = +3
Electronic configuration of d6 = eg2t2g3
Coordination number = 4
Shape: tetrahedral
Stereochemistry: optically inactive
Magnetic moment:
(v) K4[Mn(CN)6]
Potassium hexacyanomanganate(II)
Oxidation state of manganese = +2
Electronic configuration: d5+: t2g5
Coordination number = 6
Shape: octahedral.
Streochemistry: optically inactive
The stability of a complex in a solution refers to the degree of association between the two species involved in a state of equilibrium. Stability can be expressed quantitatively in terms of stability constant or formation constant.
M + 3L ↔ ML3
Stability Constant, β = [ML3] / [M][L3]
For this reaction, the greater the value of the stability constant, the greater is the proportion of ML3 in the solution.
Stability can be of two types:
(a) Thermodynamic stability:
The extent to which the complex will be formed or will be transformed into another species at the point of equilibrium is determined by thermodynamic stability.
(b) Kinetic stability:
This helps in determining the speed with which the transformation will occur to attain the state of equilibrium.
Factors that affect the stability of a complex are:
(a) Charge on the central metal ion: Thegreater the charge on the central metal ion, the greater is the stability of the complex.
(b) Basic nature of the ligand: A more basic ligand will form a more stable complex.
(c) Presence of chelate rings: Chelation increases the stability of complexes.
When a ligand attaches to the metal ion in a manner that forms a ring, then the metal- ligand association is found to be more stable. In other words, we can say that complexes containing chelate rings are more stable than complexes without rings. This is known as the chelate effect.
For example:
(i) Role of coordination compounds in biological systems:
We know that photosynthesis is made possible by the presence of the chlorophyll pigment. This pigment is a coordination compound of magnesium. In the human biological system, several coordination compounds play important roles. For example, the oxygen-carrier of blood, i.e., haemoglobin, is a coordination compound of iron.
(ii) Role of coordination compounds in medicinal chemistry:
Certain coordination compounds of platinum (for example, cis- platin) are used for inhibiting the growth of tumours.
(iii) Role of coordination compounds in analytical chemistry:
During salt analysis, a number of basic radicals are detected with the help of the colour changes they exhibit with different reagents. These colour changes are a result of the coordination compounds or complexes that the basic radicals form with different ligands.
(iv) Role of coordination compounds in extraction or metallurgy of metals:
The process of extraction of some of the metals from their ores involves the formation of complexes. For example, in aqueous solution, gold combines with cyanide ions to form [Au(CN)2]. From this solution, gold is later extracted by the addition of zinc metal.
(iii) The given complex can be written as Co(NH3)6Cl2.
Thus, [Co(NH3)6]+ along with two Cl- ions are produced.
(i) No. of unpaired electrons in [Cr(H2O)6]3+ = 3
(ii) No. of unpaired electrons in[Fe(H2O)6]2+ = 4
Then, μ
(iii) No. of unpaired electrons in [Zn(H2O)6]2+ = 0
Hence, [Fe(H2O)6]2+ has the highest magnetic moment value.
We know that CO is a neutral ligand and K carries a charge of +1.
Therefore, the complex can be written as K+ [Co(CO)4]-. Therefore, the oxidation number of Co in the given complex is -1. Hence, option (iii) is correct.
We know that the stability of a complex increases by chelation. Therefore, the most stable complex is [Fe(C2O4)3]3-.
The central metal ion in all the three complexes is the same. Therefore, absorption in the visible region depends on the ligands. The order in which the CFSE values of the ligands increases in the spectrochemical series is as follows:
H2O < NH3 < NO2 -
Thus, the amount of crystal-field splitting observed will be in the following order:
Hence, the wavelengths of absorption in the visible region will be in the order:
[Ni(H2O)6]2+ > [Ni(NH3)6]2+ > [Ni(NO2)6]4-
