Coordination Compounds — Study Notes

1. Introduction

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.

2. Werner's Theory and Basic Terminology

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.

Key Terms

• Central metal atom/ion: The metal atom to which ligands are bound, e.g., Co in [Co(NH3)6]3+.
• Ligand: An ion or molecule that donates a pair of electrons to the metal. A ligand can be monodentate (NH3, Cl-), bidentate (ethylenediamine, en), or polydentate (EDTA).
• Coordination number: The number of ligand donor atoms directly bonded to the central metal atom.
• Coordination sphere: The metal ion and its ligands enclosed in square brackets.
• Oxidation state: The charge that the metal would carry if all the ligands were removed along with their electron pairs.
• Chelate: A complex formed by polydentate ligands that bind through more than one donor atom; chelates are more stable because of the chelate effect.
• Ambidentate ligands: Ligands that can bind through more than one donor atom, e.g., NO2- (nitro) and ONO- (nitrito).

3. Nomenclature of Coordination Compounds

The rules for naming coordination compounds are:

1. In ionic complexes, the cation is named first followed by the anion.
2. The ligands are named before the metal, in alphabetical order of the ligand name (not the number prefix).
3. The number of each ligand is indicated by prefixes: di, tri, tetra, penta, hexa. For complex ligand names, bis, tris, and tetrakis are used.
4. The oxidation state of the metal is written in Roman numerals in parentheses after the metal name.
5. In coordination cations, the metal is named as such; in coordination anions, the metal name ends in -ate. For example, Co becomes cobaltate, Fe becomes ferrate, and Cu becomes cuprate.
6. Anionic ligands end in -o (chloro, hydroxo, cyano), while neutral ligands retain their names, with the exceptions aqua (H2O), ammine (NH3), carbonyl (CO), and nitrosyl (NO).

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)

4. Isomerism in Coordination Compounds

Isomers are compounds with the same molecular formula but different arrangements of atoms, and they have different properties.

Stereoisomerism

• Geometrical isomerism: Arises when ligands are arranged differently in space around the metal. In square planar complexes of formula MA2B2, cis and trans isomers exist. In octahedral complexes MA2B4 and MA3B3, cis and trans isomers are also possible. For example, [Pt(NH3)2Cl2] exists as cis (both Cl adjacent) and trans (both Cl opposite) isomers; the cis isomer is the anti-cancer drug cisplatin.
• Optical isomerism: A complex is optically active if it has a non-superimposable mirror image. The two forms are called enantiomers, and they rotate plane polarised light in opposite directions. Octahedral complexes with bidentate ligands, such as [Co(en)3]3+, show optical isomerism.

Structural Isomerism

• Linkage isomerism: Arises when an ambidentate ligand can bind through different atoms, e.g., [Co(NH3)5(NO2)]2+ (nitro) and [Co(NH3)5(ONO)]2+ (nitrito).
• Coordination isomerism: Arises when the distribution of ligands between the cation and anion differs, e.g., [Co(NH3)6][Cr(CN)6] and [Cr(NH3)6][Co(CN)6].
• Ionisation isomerism: Different ions are produced in solution, e.g., [Co(NH3)5Br]SO4 (gives SO4 2- ions) and [Co(NH3)5SO4]Br (gives Br- ions).
• Solvate (hydrate) isomerism: Water is differently distributed between coordination sphere and lattice, e.g., [Cr(H2O)6]Cl3 is violet, [Cr(H2O)5Cl]Cl2.H2O is green, and [Cr(H2O)4Cl2]Cl.2H2O is dark green.

5. Valence Bond Theory

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.

Limitations of VBT

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.

6. 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.

Octahedral Splitting

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, Δo.

Spectrochemical Series

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.

Colour and Magnetic Properties

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).

Limitations of CFT

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.

7. Stability of Coordination Compounds

The stability of a complex in solution is expressed by the stability constant. The overall formation constant, βn, 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.

8. Applications of Coordination Compounds

1. In analytical chemistry: Estimation of hardness of water uses EDTA; qualitative tests for metal ions use complex formation.
2. In metallurgy: Extraction of silver and gold uses the cyanide complex; extraction of nickel uses the carbonyl complex.
3. In biological systems: Haemoglobin contains an iron porphyrin complex and chlorophyll contains a magnesium porphyrin; vitamin B12 contains cobalt.
4. In medicine: Cisplatin is used as an anti-cancer drug; chelate therapy uses EDTA to remove toxic metal ions such as lead and mercury from the body.
5. In photography: Silver halide complexes in fixing solutions remove unexposed silver bromide.

Quick Revision Tables

Table 1: Ligands and Their Types

Table 2: Isomerism Types

Mind Map

Important Diagrams (SVG)

Diagram 1: Geometrical Isomerism in Square Planar Complexes

Diagram 2: Crystal Field Splitting in Octahedral Field

Common Mistakes

1. Confusing primary and secondary valency in Werner's theory; primary is the oxidation state and secondary is the coordination number.
2. Writing the oxidation state of the metal incorrectly; always compute it from the net charge of the complex and the charges of the ligands.
3. Naming the ligands in the wrong order; the alphabet is ordered by the ligand name, not the metal.
4. Believing that all octahedral complexes show geometrical isomerism; MA6, MA5B, and MA4B2 complexes do not show cis-trans isomerism in the same way.
5. Using the wrong hybridisation; d2sp3 is inner orbital and sp3d2 is outer orbital, and they have different magnetic properties.
6. Placing ligands in the spectrochemical series incorrectly; CN- and CO are strong field while Cl- and I- are weak field.
7. Assuming that optical isomerism requires no symmetry at all; it specifically requires a non-superimposable mirror image, as in [Co(en)3]3+.

Exam Tips

1. Memorise the naming rules, especially the suffixes: -o for anionic ligands, -ate for anionic complexes, and Roman numerals for oxidation states.
2. Learn the spectrochemical series in order and associate strong field with low spin and weak field with high spin.
3. Practise determining oxidation state and coordination number from the formula of the complex.
4. Remember one example for each isomerism type; cisplatin (geometrical) and [Co(en)3]3+ (optical) are the most frequently asked.
5. For magnetic moment problems, count unpaired electrons using mu = sqrt(n(n+2)) BM.
6. Know the applications: EDTA in water hardness, cisplatin in cancer therapy, and the cyanide process for gold and silver.

Conclusion

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.

Test Your Understanding

1. Define coordination number, ligand, and coordination sphere with an example.
2. Write the IUPAC names of [Co(NH3)6]Cl3 and K3[Fe(CN)6].
3. Distinguish between inner orbital and outer orbital octahedral complexes.
4. What is the spectrochemical series and how does it relate to high spin and low spin complexes?
5. Give two applications of coordination compounds in medicine and two in analysis.