Comprehensive theory, key formulas, diagrams, and memory aids for Coordination Compounds.
Coordination compounds are molecular compounds in which a central metal atom or ion is surrounded by a number of ions or molecules called ligands. These compounds, also called complex compounds, are distinguished from simple addition compounds by their ability to exist independently in solution and by their distinctive properties. The study of coordination compounds is central to inorganic chemistry and has applications in analytical chemistry, metallurgy, dyeing, and biological systems such as haemoglobin, chlorophyll, and vitamin B12.
The chapter begins with the basic terminology of coordination chemistry: the central metal atom, the ligands, the coordination number, the coordination sphere, and the charge on the complex. Nomenclature rules are established for naming coordination compounds systematically, and the concept of isomerism, including geometrical and optical isomerism, is explained. These fundamental ideas allow chemists to describe the structure and bonding of complexes precisely.
The bonding in coordination compounds is explained by the valence bond theory and, more quantitatively, by crystal field theory. These theories account for the magnetic properties, colour, geometry, and stability of complexes. The chapter concludes with the applications of coordination compounds in quantitative analysis, extraction of metals, and biological processes, demonstrating the practical importance of this branch of chemistry.
Werner's theory, proposed in 1898, was the first to explain the bonding in coordination compounds. Werner proposed that metals have two types of valency: - Primary valency: The oxidation state of the metal, satisfied by negative ions, shown outside the coordination sphere. - Secondary valency: The coordination number, satisfied by ligands, shown inside the coordination sphere (brackets).
For example, in [Co(NH3)6]Cl3, the primary valency is +3 and the secondary valency is 6.
The rules for naming coordination compounds are:
Examples: - [Co(NH3)6]Cl3: Hexaamminecobalt(III) chloride - K3[Fe(CN)6]: Potassium hexacyanoferrate(III) - [Ni(CO)4]: Tetracarbonylnickel(0) - [Pt(NH3)2Cl2]: Diamminedichloridoplatinum(II)
Isomers are compounds with the same molecular formula but different arrangements of atoms, and they have different properties.
The valence bond theory (VBT) explains the bonding in coordination compounds in terms of hybridisation of the metal orbitals. The inner orbital complexes involve (n-1)d orbitals while outer orbital complexes involve nd orbitals.
For [Co(NH3)6]3+, cobalt is in +3 state with configuration 3d6. The six NH3 ligands require six hybrid orbitals formed by d2sp3 hybridisation, giving an inner orbital octahedral complex that is diamagnetic.
For [CoF6]3-, cobalt is also 3d6, but the weak field ligand F- uses sp3d2 hybridisation involving the outer d-orbitals, giving an outer orbital octahedral complex that is paramagnetic with four unpaired electrons.
VBT explains magnetic behaviour and geometry but fails to explain the colour of complexes and the detailed spectrochemical properties, which are handled by crystal field theory.
Crystal field theory (CFT) treats the ligand-metal interaction as purely electrostatic. The ligands, which may be negative ions or neutral molecules with lone pairs, generate a crystal field that splits the degeneracy of the d-orbitals.
In an octahedral field, the d-orbitals split into two sets: - The eg orbitals (dz2 and dx2-y2) point directly at the ligands and have higher energy. - The t2g orbitals (dxy, dyz, dxz) point between the ligands and have lower energy.
The energy difference between the two sets is the crystal field splitting energy, Îo.
The splitting energy depends on the nature of the ligand, giving the spectrochemical series: I- < Br- < SCN- < Cl- < F- < OH- < H2O < NH3 < en < NO2- < CN- < CO
Strong field ligands like CN- and CO produce large splitting and tend to give low spin complexes; weak field ligands like Cl- and I- give high spin complexes.
The colour arises from d-d transitions between the split levels, and the number of unpaired electrons determines the magnetic behaviour. Crystal field theory also explains why [Co(NH3)6]3+ is diamagnetic (low spin, all electrons paired) while [CoF6]3- is paramagnetic (high spin).
CFT ignores the covalent character of metal-ligand bonds and does not account for the spectrochemical series order fully; ligand field theory addresses these shortcomings.
The stability of a complex in solution is expressed by the stability constant. The overall formation constant, Îēn, is the equilibrium constant for the stepwise formation of the complex. A higher stability constant means a more stable complex. Chelating ligands form more stable complexes because of the entropically favourable chelate effect. The stability of complexes is important in analytical chemistry, where complexes are used to mask interfering ions.
| Ligand | Type | Donor Atom(s) |
|---|---|---|
| NH3 | Monodentate | N |
| Cl- | Monodentate | Cl |
| H2O | Monodentate | O |
| en (ethylenediamine) | Bidentate | 2 N |
| oxalate (ox) | Bidentate | 2 O |
| EDTA | Hexadentate | 6 (N, O) |
| NO2- / ONO- | Ambidentate | N or O |
| Type | Origin | Example |
|---|---|---|
| Geometrical | Different spatial arrangement | [Pt(NH3)2Cl2] cis/trans |
| Optical | Non-superimposable mirror image | [Co(en)3]3+ |
| Linkage | Ambidentate ligand | [Co(NH3)5(NO2)]2+ / (ONO) |
| Coordination | Ligand exchange between ions | [Co(NH3)6][Cr(CN)6] |
| Ionisation | Different ions in solution | [Co(NH3)5Br]SO4 |
| Solvate | Water distribution | [Cr(H2O)6]Cl3 violet/green |
graph TD
A["Coordination Compounds"] --> B["Terminology and Werner's Theory"]
A --> C["Nomenclature"]
A --> D["Isomerism"]
A --> E["Valence Bond Theory"]
A --> F["Crystal Field Theory"]
A --> G["Applications"]
B --> B1["Primary and secondary valency"]
B --> B2["Ligands, coordination number, chelates"]
D --> D1["Geometrical and optical"]
D --> D2["Linkage, coordination, ionisation, solvate"]
E --> E1["Inner orbital: d2sp3"]
E --> E2["Outer orbital: sp3d2"]
F --> F1["Octahedral splitting: t2g and eg"]
F --> F2["Spectrochemical series"]
F --> F3["High spin vs low spin"]
G --> G1["Analytical chemistry"]
G --> G2["Metallurgy, medicine, biology"]
Coordination compounds reveal a rich chemistry of metal-ligand interactions that extends far beyond the simple salts studied earlier. Werner's theory laid the conceptual foundation, and modern theories such as valence bond theory and crystal field theory provide detailed explanations of structure, colour, and magnetism. The systematic nomenclature and the study of isomerism equip chemists to describe and distinguish the many possible structures of complexes. The stability of complexes and the chelate effect are exploited in analysis, metallurgy, and medicine, while biological systems depend on coordination compounds such as haemoglobin and chlorophyll for their function. This chapter connects theoretical inorganic chemistry with practical applications, making it one of the most important and rewarding topics in the Class 12 curriculum.