Coordination Compounds

Werner's theory (1893) — Nobel Prize 1913. The first coherent model of coordination compounds.

Two types of valence:

1. Primary (ionizable, ionic): the oxidation state of the metal. Satisfied by negative ions.
2. Secondary (non-ionizable, covalent): the coordination number. Satisfied by ligands (neutral or anionic, donating lone pairs).

Example: CoCl₃ · 6NH₃ is actually [Co(NH₃)₆]Cl₃. The Cl⁻ outside the sphere are ionizable; the NH₃ are coordinated.

Vocabulary:

• Central metal atom/ion (CMA): at the centre, accepts lone pairs.
• Ligand: group donating electrons to CMA.
• Coordination number: number of bonds from metal to ligands. Typical: 4 or 6.
• Coordination sphere: the metal + its ligands, written in [ ].
• Counter ions: outside [ ], balance charge.

Types of ligands:

By number of donor atoms (denticity):

• Monodentate (unidentate): 1 donor — H₂O, NH₃, Cl⁻, CN⁻, CO.
• Bidentate: 2 donors — ethylenediamine (en), oxalate (ox), glycinate.
• Polydentate: more — EDTA (hexadentate).

Chelate: complex with a cyclic ring formed by a polydentate ligand. Especially stable due to entropy (chelate effect).

By charge:

• Neutral: H₂O (aqua), NH₃ (ammine), CO (carbonyl), NO (nitrosyl).
• Anionic: Cl⁻ (chloro), CN⁻ (cyano), OH⁻ (hydroxo), CH₃COO⁻ (acetato), NO₂⁻ (nitro or nitrito).

IUPAC NOMENCLATURE RULES:

1. Cation named before anion (just like NaCl is "sodium chloride").

2. Within a complex, ligands named alphabetically before the metal. Ignore prefixes (di-, tri-) when alphabetizing.

3. Anion ligands end in -o (chloro, cyano, hydroxo, nitro).
   Neutral ligands keep their name with some exceptions: aqua (H₂O), ammine (NH₃ — note double m), carbonyl (CO), nitrosyl (NO).

4. Number of ligands: use Greek prefixes di-, tri-, tetra-, penta-, hexa-. For ligand names that already contain Greek prefixes or are complex, use bis-, tris-, tetrakis- with parentheses.

5. Oxidation state of metal in Roman numerals in parentheses.

6. If the complex is an anion, metal name ends in -ate (and Latin name often used: ferrate from iron, cuprate from copper, plumbate from lead, stannate from tin, aurate from gold, argentate from silver).

7. No space between ligand names; space between last ligand and metal.

Worked examples:

(Use chlorido in latest IUPAC; chloro in older but JEE accepts both.)

EAN (Effective Atomic Number) rule:

EAN = (atomic number of metal) − (oxidation state) + (electrons donated by ligands).

Stable complexes often have EAN equal to that of nearest noble gas. Example: Ni(CO)₄: EAN = 28 − 0 + 4×2 = 36 (Kr). [Fe(CN)₆]⁴⁻: 26 − 2 + 6×2 = 36.

Notable exception: many 4d and 5d complexes don't obey 18-electron rule.

Crystal Field Theory (CFT) treats ligands as point negative charges that split the degenerate d-orbitals of a transition metal.

Octahedral splitting. In an octahedral field, the 5 d-orbitals split into:

• t₂g (lower energy): d_xy, d_yz, d_xz — three orbitals between the axes
• e_g (higher energy): d_x²-y², d_z² — two orbitals along the axes

Energy gap = Δ_o ("crystal field splitting energy").

Filling the d-electrons:

High-spin (weak-field ligands like F⁻, H₂O, OH⁻): Δ_o is small. Electrons spread out via Hund's rule before pairing. Maximum unpaired electrons.

Low-spin (strong-field ligands like CN⁻, CO, NO₂⁻): Δ_o is large. Pairing in t₂g is energetically favored over occupying e_g. Fewer unpaired electrons.

Spectrochemical series (weak → strong):
I⁻ < Br⁻ < S²⁻ < SCN⁻ < Cl⁻ < NO₃⁻ < F⁻ < OH⁻ < H₂O < NH₃ < en < CN⁻ ≈ CO

Crystal field stabilization energy (CFSE):
CFSE = (number in t₂g) × (−0.4 Δ_o) + (number in e_g) × (+0.6 Δ_o)

Worked example: [Fe(H₂O)₆]³⁺ vs [Fe(CN)₆]³⁻ (both d⁵).

H₂O is weak field → high-spin. Fe³⁺ d⁵: t₂g³ e_g². Unpaired electrons = 5. Magnetic moment μ = √(5(5+2)) = √35 ≈ 5.92 BM. Paramagnetic.

CN⁻ is strong field → low-spin. Fe³⁺ d⁵: t₂g⁵ e_g⁰. Unpaired electrons = 1. μ = √3 ≈ 1.73 BM. Still paramagnetic but less so.

Tetrahedral splitting is reversed (e below t₂) and smaller (Δ_t = 4/9 Δ_o). Tetrahedral complexes are nearly always high-spin because Δ_t < pairing energy.

Why are coordination compounds coloured? d-d transitions in the visible spectrum. Ti(H₂O)₆³⁺ is purple (absorbs ~500 nm green). Empty (Sc³⁺ d⁰) and full (Zn²⁺ d¹⁰) d-shells → no d-d transitions → colorless.