Revision: Std XII >> Coordination Compounds MAH-MHT CET (PCM/PCB) Coordination Compounds

Coordination compounds are those molecular compounds which retain their identity in solid as well as in aqueous solution. In these compounds, metals or atoms are bonded to a number of anions or neutral molecules by a coordinate bond.

Lewis bases: The ligands being electron pair donors are Lewis bases.

Lewis acids: The central metal ion, as an electron-pair acceptor, serves as a Lewis acid. For example, in the coordination compound [Cu(NH₃)₄]²⁺, NH₃ is a Lewis base, and Cu²⁺ is a Lewis acid.

In the coordination compound, the species surrounding the central metal atom or ion are called ligands.

A ligand is a molecule, ion or group that is bonded to the metal atom or ion in a complex or coordination compound by a coordinate bond.

A monodentate ligand is one in which a single donor atom shares an electron pair with the centre metal ion to create a coordinate bond.

The central metal ion and ligands linked to it are enclosed in a square bracket. This is called a coordination sphere. This is a discrete structural unit. The ionisable groups shown outside the bracket are the counter ions. For example, the compound K₄[Fe(CN)₆] has [Fe(CN)₆]^4- coordination sphere with the ionisable K^⊕ ions representing counter ions.

A negatively charged coordination sphere or a coordination compound having a negatively charged coordination sphere is called anionic complex or anionic sphere complex.

Coordination number is the number of ligand donor atoms directly bonded to the central metal atom or ion in a complex.

Example: In [Co(NH₃)₅Cl]²⁺, the coordination number of cobalt (Co) is 6 because 5 ammonia molecules and 1 chloride ion are attached to it.

The total number of electrons present on the central metal atom/ion, including those gained by it in bonding, is called EAN.

Two or more coordination compounds which contain the same number and types of atoms, and bonds (i.e., the connectivity between atoms is the same), but which have different spatial arrangements of the atoms are called distereoisomers.

Isomers in which there is exchange of solvent (water) ligands between coordination and ionization spheres are called hydrate isomers.

Isomers which show interchange of ligands between cationic and anionic spheres of different metal ions are called co-ordination isomers.

Isomerism is the phenomenon in which compounds have the same molecular formula but differ in their physical or chemical properties due to a different arrangement of atoms or groups in space or structure.

The stability of a coordination complex refers to the extent to which it exists in a solution as a coordination sphere.

In coordination compounds, the central metal exhibits two types of valencies:

• Primary valency
• Secondary valency

Primary valency:

• Satisfied by anions only
• Depends on the oxidation state of the metal
• Ionisable and non-directiona

Secondary valency:

• Satisfied by ligands
• Represents the coordination number
• Non-ionisable and directional
• Determines the geometry of the complex

Basic Terms:

• Coordination sphere: Metal + ligands inside brackets
• Charge number: Net charge on complex ion
• Oxidation state: Charge on central metal ion
• Coordination number: Number of donor atoms attached to metal

Double Salts vs Coordination Compounds:

1. On the basis of ligands

• Homoleptic complex:
  Contains only one type of ligand
  Example: [Co(NH₃)₆]³⁺
• Heteroleptic complex:
  Contains two or more types of ligands
  Example: [Co(NH₃)₄Cl₂]⁺

2. On the basis of charge

• Cationic complex:
  Positively charged coordination sphere
  Example: [Zn(NH₃)₄]²⁺
• Anionic complex:
  Negatively charged coordination sphere
  Example: [Fe(CN)₆]³⁻
• Neutral complex:
  No overall charge
  Example: [Ni(CO)₄]

1. Order

• Ligand → Metal
• Cation first, then anion

2. Ligand Names

• Cl⁻ → chloro
• CN⁻ → cyano
• OH⁻ → hydroxo
• NH₃ → ammine
• H₂O → aqua
• CO → carbonyl

3. Number Prefix

• di, tri, tetra, penta, hexa
• Special: bis, tris (if ligand has number)

4. Order of Ligands

• Alphabetical order

5. Metal Name

• Neutral/cation → normal name
• Anionic complex → ends with “-ate”
  Fe → ferrate
  Cu → cuprate
  Co → cobaltate

6. Oxidation State

• Write in Roman (II), (III)

7. Important

• Counter ions not named
• Complex always in [ ]

8. Examples

Neutral complexes

• [Co(NO₂)₃(NH₃)₃] → Triamminetrinitrocobalt(III)
• Fe(CO)₅ → Pentacarbonyliron(0)

Cationic complexes

• [Cu(NH₃)₄]²⁺ → Tetraamminecopper(II) ion
• [Fe(H₂O)₅(NCS)]²⁺ → Pentaaquathiocyanatoiron(III) ion

Anionic complexes

• [Ni(CN)₄]²⁻ → Tetracyanonickelate(II) ion
• [Fe(CN)₆]⁴⁻ → Hexacyanoferrate(II) ion

Compounds (Very Important)

• [Co(NH₃)₅Cl]Cl₂ → Pentaamminechlorocobalt(III) chloride
• K₃[Al(C₂O₄)₃] → Potassium trioxalatoaluminate(III)
• Na₃[Co(NO₂)₆] → Sodium hexanitrocobaltate(III)

Factors affecting the stability of a complex:

1. Charge on the central metal ion 
2. Nature of the metal ion

For the same ligand, the stability of complexes formed by divalent metal ions follows the order:

Cu²⁺ > Ni²⁺ > Co²⁺ > Fe²⁺ > Mn²⁺ > Cd²⁺

• Proposed by Heitler and London (1927), further developed by Pauling and Slater.
• A covalent bond is formed when half-filled valence atomic orbitals of similar energies overlap, each containing one unpaired electron.
• Greater the overlap → stronger the bond.

Types of Orbital Overlap:

Hybridisation & Shapes:

• d-orbitals of metal ion split into lower and higher energy levels in the presence of ligands
• In a free ion, d-orbitals are degenerate (same energy)
• Spectrochemical Series

I⁻ < Br⁻ < Cl⁻ < S²⁻ < F⁻ < OH⁻ < C₂O₄²⁻ < H₂O < NCS⁻ < EDTA < NH₃ < en < CN⁻ < CO

• d-orbitals split into t₂g (low) and e_g (high) energy levels
• The energy difference lies in the visible region
• Electron jumps from t₂g to e_g energy level, which is called d–d transition.
• d¹ to d⁹ ions → coloured
• d⁰ and d¹⁰ ions → colourless