Coordination Chemistry

Coordination Compounds: Definition & Basic Terminology

• Compounds in which a central metal atom/ion is directly bonded (via coordinate bonds) to a fixed number of negative or neutral molecules/ions.
• Entire metal–ligand aggregate enclosed in square brackets is called the coordination sphere / coordination entity / complex.
  • Outside-bracket ion(s) = counter-ions.
• Charge on the sphere may be positive (cationic complex), negative (anionic complex) or zero (neutral complex).
  • Examples
    • Cationic: $[Co(NH<em>3)</em>6]^{3+},\;[Fe(H<em>2O)</em>6]^{2+}$
    • Anionic: $[Fe(CN)<em>6]^{4-},\;[Ag(CN)</em>2]^-$
    • Neutral: $[Ni(CO)<em>4],\;[Fe(CO)</em>5]$

Central Metal Atom / Ion

• Must possess vacant orbitals to accept lone pairs from ligands.
• Usually a transition-metal (small size, high nuclear charge, variable oxidation states, vacant d-orbitals).
• Acts as a Lewis acid.

Ligands

• Species that donate lone pair(s) to the central metal – behave as Lewis bases.
• Classified by charge
  • Negative: $F^- ,\;Cl^- ,\;Br^- ,\;I^- ,\;OH^- ,\;CN^- ,\;NO<em>2^- ,\;SCN^-$ • Neutral: $H</em>2O,\;NH<em>3,\;CO,\;NO,\;N</em>2$
  • Positive: $NO^+,\;NH_2^+$
• Denticity = number of donor atoms actually attached to metal.
  • Monodentate (one site): $H<em>2O,\;NH</em>3,\;Cl^- ,\;CN^-$
  • Bidentate (two sites): oxalate $(C<em>2O</em>4)^{2-},$ ethane-1,2-diamine (en), glycinato (gly)
  • Tridentate: diethylenetriamine (dien)
  • Tetradentate: triethylenetetramine (trien)
  • Hexadentate: EDTA$^{4-}$
• Polydentate = denticity ≥3; create chelate rings → enhanced stability (chelate effect).
• Ambidentate: two possible donor atoms; can bind through either but not both at once.
  • $CN^-$ (C-donor = cyano, N-donor = isocyano)
  • $SCN^-$ (S-donor = thiocyanato-S, N-donor = isothiocyanato-N)
  • $NO_2^-$ (N-donor = nitro, O-donor = nitrito-O)
• Flexidentate: variable denticity in different complexes (e.g., sulphato).

Coordinate (Dative) Bond

• Shared electron pair supplied by ligand only.
  • Notation: $M{:}L$ or $M\leftarrow L$
• Bridged to Lewis concept: metal = Lewis acid (acceptor), ligand = Lewis base (donor).

Classification of Complexes

• Homoleptic: only one kind of ligand, e.g. $[Fe(CN)_6]^{3-}$.
• Heteroleptic: ≥2 different ligands, e.g. $[Co(NH<em>3)</em>4Cl_2]^+$.

Coordination Number (CN) & Oxidation Number (ON)

• $CN = \text{number of coordinate bonds} = \Sigma(\text{number of ligands} \times \text{denticity})$.
• ON: formal charge on metal if all ligands are removed with electron pairs.
  • Example $[Fe(CN)_6]^{3-}:\;x + 6(-1) = -3 \Rightarrow x = +3$.

Common Coordination Polyhedra

• $ML_4$ tetrahedral or square-planar.
• $ML_5$ trigonal-bipyramidal or square-pyramidal.
• $ML_6$ octahedral (most common).

IUPAC Nomenclature Rules (highlights)

1. Cationic part named before anionic part.
2. Inside a sphere: (a) numerical prefix, (b) ligand names (alphabetical, ignore prefixes), (c) metal name + oxidation state in Roman numerals.
3. Prefixes: mono- (rarely used), di, tri, tetra, penta, hexa. If ligand already contains di/tri etc., use bis, tris, tetrakis.
4. For anionic complexes, metal ends with “-ate”; Latin roots for some (ferrate, cuprate, argentate…).

• Sample names
  • $[Co(NH<em>3)</em>5(ONO)]^{2+}$ → pentaammine​nitrito-O-cobalt(III) ion.
  • $K<em>2[NiCl</em>4]$ → potassium tetrachloridonickelate(II).

Werner’s Theory (Historical)

• Metals show primary valency (ionisable; oxidation state) and secondary valency (coordinate; CN).
• Secondary valencies are directional → geometry.
• Negative ligands can satisfy both valencies (e.g., $[CoCl<em>2(NH</em>3)_4]^+$).

Difference: Double Salts vs Complex Salts

• Double salts dissociate completely to give constituent ions in solution (e.g., Mohr’s salt $FeSO<em>4(NH</em>4)<em>2SO</em>4·6H_2O$).
• Complex salts retain the complex ion intact; properties differ markedly from simple ions (e.g., $K<em>4[Fe(CN)</em>6]$).

Isomerism in Coordination Compounds

Structural Isomerism

1. Linkage: ambidentate ligand binds through different donor atoms.
   • $[Co(NH<em>3)</em>5(NO<em>2)]^{2+}$ (nitro-N, red) vs $[Co(NH</em>3)_5(ONO)]^{2+}$ (nitrito-O, yellow).
2. Coordination (interchange of ligands between two metal centres in bimetallic salts).
   • $[Cr(NH<em>3)</em>6][Co(CN)<em>6]$ vs $[Cr(CN)</em>6][Co(NH<em>3)</em>6]$.
3. Ionisation: exchange between sphere ligand and counter-ion → different ions in solution.
   • $[Co(NH<em>3)</em>5SO<em>4]Br$ vs $[Co(NH</em>3)<em>5Br]SO</em>4$.
4. Hydrate (solvate): different number of coordinated vs lattice H₂O.
   • $[Cr(H<em>2O)</em>6]Cl<em>3$ (violet), $[Cr(H</em>2O)<em>5Cl]Cl</em>2·H<em>2O$ (green), $[Cr(H</em>2O)<em>4Cl</em>2]Cl·2H_2O$ (blue-green).

Stereoisomerism

1. Geometrical (cis–trans; facial (fac)/meridional (mer)).
   • Square planar $[PtCl<em>2(NH</em>3)<em>2]$: cis (optically active when using bidentate ligands) & trans. • Octahedral $[Co(NH</em>3)<em>4Cl</em>2]^+$: cis vs trans.
   • $[MA<em>3B</em>3]$ gives fac (three identical ligands on one face) vs mer (around meridian).
2. Optical: non-superimposable mirror images (enantiomers). Requirements
   • No plane of symmetry; typical in octahedral complexes with one or more bidentate ligands (e.g., $[Co(en)_3]^{3+}$).
   • Tetrahedral rarely show optical; square planar never (has symmetry plane).

Valence Bond Theory (VBT) Synopsis

• Metal provides empty orbitals; hybridises to form equivalent set → coordinate bonds.
• Inner-orbital (d²sp³) vs outer-orbital (sp³d²) octahedral depending on whether inner 3d or outer 4d/5d used.
• Magnetic behaviour predicted from number of unpaired electrons.
  • Example $[CoF<em>6]^{3-}$ (weak-field): outer-orbital $sp^3d^2$, high-spin, 4 unpaired e⁻, $\mu=\sqrt{n(n+2)}=\sqrt{24}\,\text{BM}$. • $[Co(CN)</em>6]^{3-}$ (strong-field): inner-orbital $d^2sp^3$, low-spin, diamagnetic.
• Hybridisation summary
  • $sp^3$ (tetrahedral), $dsp^2$ (square planar), $sp^3d$ (trigonal-bipyramidal), $d^2sp^3 / sp^3d^2$ (octahedral).

Crystal Field Theory (CFT) Essentials

• Treat metal ion as point positive; ligands as point charges/dipoles; bonding purely electrostatic.
• Degenerate d-orbitals split when ligands approach:
  • Octahedral: $e<em>g$ (higher, $d</em>{x^2-y^2}, d<em>{z^2}$) up by $+0.6\Delta</em>o$; $t<em>{2g}$ (lower) down by $-0.4\Delta</em>o$.
  • Tetrahedral: reverse ordering; magnitude $\Delta<em>t \approx \dfrac{4}{9}\Delta</em>o$.
• Electron placement depends on comparison of $\Delta<em>o$ with pairing energy P. • $\Delta</em>o < P$ → high-spin (weak-field). • $\Delta_o > P$ → low-spin (strong-field).
• Explains colour (d–d transitions) and magnetism.

Spectrochemical Series

I^- < Br^- < SCN^- < Cl^- < F^- < OH^- < C2O4^{2-} < H2O < edta^{4-} < NH3 < en < NO_2^- < CN^- < CO

• Left side = weak field, small $\Delta$, high-spin; right side = strong field, large $\Delta$, low-spin.

Stability of Complexes

• Expressed by stability (formation) constant $K<em>f$ : $M^{n+} + L \rightleftharpoons [ML]^{n+}$, $K</em>f = \frac{[ML]}{[M][L]}$.
• Factors enhancing stability
  • Higher metal charge.
  • Smaller ionic radius.
  • 3d < 4d < 5d (due to greater covalency) – but 4f lanthanides behave differently. • Strong-field ligands. • Chelate effect (five- or six-membered rings most stable). • Macrocyclic ligands > open-chain polydentate.

Metal Carbonyls

• Compounds where CO acts as neutral ligand; examples & geometries
  • $Ni(CO)<em>4$ tetrahedral, $Fe(CO)</em>5$ trigonal-bipyramidal, $Cr(CO)_6$ octahedral.
• Bonding (synergic $\sigma$-donation + $\pi$-back donation)
  • $\sigma$: CO donates lone pair on C to vacant metal orbital.
  • $\pi$: filled metal d-orbital back-donates to vacant $\pi^*$ of CO.
  • Strengthens M–C and weakens C–O; accounts for stability & low CO stretching frequency.

Applications & Biological Relevance

• Analytical chemistry: hardness of water by EDTA titration; detection of metal ions via complexation.
• Catalysts: $[PtCl<em>2(PPh</em>3)<em>2]$ (Wilkinson), $[Ni(CN)</em>4]^{2-}$ in Kapler process, metal carbonyls in hydroformylation.
• Metallurgy: Mond process uses $Ni(CO)_4$ for nickel purification.
• Medicine: chelation therapy for heavy-metal poisoning (EDTA, BAL), cisplatin $[PtCl<em>2(NH</em>3)_2]$ as anticancer drug.
• Bio-molecules: chlorophyll (Mg-porphyrin complex), haemoglobin (Fe-porphyrin), vitamin B$_{12}$ (Co-corrin).

Quick Reference: Magnetic Moment Formula

$\mu_{\text{spin only}} = \sqrt{n(n+2)}\,\text{B.M.}$ where $n =$ number of unpaired d-electrons.

Essential Numerical / Symbolic Facts

• $CN =\text{ No. ligands}\times\text{denticity}$.
• $ON(M) = x$ obtained from $x + \sum(\text{ligand charges}) = \text{overall charge}$.
• $\Delta<em>t \approx \dfrac{4}{9}\,\Delta</em>o$.
• Chelate effect: $K<em>f(\text{chelated}) \gg K</em>f(\text{monodentate})$.