Coordination Compounds

Introduction

• Transition metals form numerous stable compounds in which the metal shares electron pairs with anions/neutral molecules → modern term: coordination compounds.

• Coordination chemistry underpins modern inorganic, bio–inorganic chemistry & many industrial processes.

  • Biological: $\text{Chlorophyll (Mg)},\ \text{Haemoglobin (Fe)},\ \text{Vitamin B}_{12}\,(\text{Co})$ are classic coordination entities.

  • Practical: metallurgical extractions, analytical titrations, homogeneous/heterogeneous catalysis, electroplating, dyeing, medicines.

Werner’s Theory of Coordination Compounds

• Alfred Werner (1866–1919) prepared/characterised many complexes; proposed two valences for metal ions:

  1. Primary valence = usual ionic valence; ionisable; satisfied by anions (e.g., $\text{CrCl}_3$ has primary valence $+3$).

  2. Secondary valence = coordination number (CN); non-ionisable; satisfied by ligands directly bound to metal.

• Key postulates

  • Secondary valence is fixed for a given metal and dictates a definite spatial arrangement (coordination polyhedron).

  • Common geometries: octahedral, tetrahedral, square-planar.

• Evidence: Cobalt(III) chloride–ammonia series; AgNO$_3$ precipitation & conductance showed $6$ groups remain bound (Table sample):

  • $[\text{Co(NH}<em>3)</em>6]Cl_3$ (yellow) → 3 AgCl, 1:3 electrolyte (secondary valence 6)

  • $[\text{CoCl(NH}<em>3)</em>5]Cl_2$ (purple) → 2 AgCl, 1:2 electrolyte

  • $[\text{CoCl}<em>2\text{(NH}</em>3)_4]Cl$ (green/violet) → 1 AgCl, 1:1 electrolyte

  • Demonstrated isomerism (same empirical formula, different properties).

Fundamental Terminology

• Coordination entity: central atom/ion + ligands inside $[\,]$ (e.g., $[\text{Ni(CO)}_4]$).

• Central atom/ion: Lewis-acidic metal site (e.g., $\text{Fe}^{3+}$ in $[\text{Fe(CN)}_6]^{3-}$).

• Ligand: ion/molecule donating lone pair(s).

  • Denticity: unidentate, didentate (e.g., $\text{en}=H<em>2NCH</em>2CH<em>2NH</em>2$), polydentate (e.g., $\text{EDTA}^{4-}$-hexadentate).

  • Chelate ligand: uses ≥2 donor atoms simultaneously → chelate complexes more stable.

  • Ambidentate: two different donor atoms; binds through either, not both (e.g., $\text{NO}_2^-$, $\text{SCN}^-$).

• Coordination number (CN): number of σ-donor atoms attached (π bonds not counted); examples:

  • $[\text{PtCl}<em>6]^{2-}$ (CN=6), $[\text{Ni(NH}</em>3)_4]^{2+}$ (CN=4).

• Coordination sphere: full entity in brackets; outside ions = counter-ions.

• Coordination polyhedron: spatial arrangement of donors (octahedral, tetrahedral, square-planar, trigonal-bipyramidal etc.).

• Oxidation number: hypothetical charge on metal after heterolytic removal of ligands; given in Roman numerals e.g., Cu(I).

• Homoleptic vs heteroleptic: single ligand type (e.g., $[\text{Co(NH}<em>3)</em>6]^{3+}$) vs multiple types (e.g., $[\text{Co(NH}<em>3)</em>4Cl_2]^+$).

• Homoleptic Complex:A coordination complex in which only one type of ligand is attached to the central metal atom/ion.Example: [Co(NH3)6]3+[Co(NH_3)_6]^{3+}[Co(NH3​)6​]3+ → All ligands are NH3

• Heteroleptic complex: A coordination complex with multiple types of ligands attached to the central metal atom/ion (e.g., [ ext{Co(NH}3)4Cl_2]^+}).

Nomenclature Rules (IUPAC)

Writing Formulae

1. Metal symbol first, then ligands (alphabetically by ligand symbol/abbreviation, independent of charge).

2. Polyatomic ligand formulas in parentheses; whole entity in $[\,]$.

3. Overall charge outside as superscript (e.g., $[\text{Cr(H}<em>2\text{O})</em>6]^{3+}$).

4. Counter-ions written left (cations) or right (anions) of bracket to achieve electroneutrality.

Naming

1. Name cation before anion in full compound.

2. Within entity list ligands alphabetically, then metal.

   • Anionic ligands end in -o (IUPAC 2004 draft: -ido; e.g., chloro→chlorido).

   • Neutral: usual names except aqua (H$2$O), ammine (NH$3$), carbonyl (CO), nitrosyl (NO).

3. Prefixes mono-, di-, tri- etc.; if ligand already contains prefix, use bis, tris, tetrakis [ligand] .

4. Oxidation state of metal in Roman numerals ( ) after the entity.

5. For anionic complexes, metal name ends in -ate (sometimes Latin root, e.g., ferrate).

6. Neutral complex named like cationic.

Illustrative examples

• $[\text{Cr(NH}<em>3)</em>3(H<em>2O)</em>3]Cl_3$ → triamminetriaquachromium(III) chloride.

• $[\text{Ag(NH}<em>3)</em>2][\text{Ag(CN)}_2]$ → diamminesilver(I) dicyanidoargentate(I).

• $\text{K}<em>2[\text{Zn(OH)}</em>4]$ → potassium tetrahydroxidozincate(II).

Isomerism in Coordination Compounds

• Stereoisomerism

  • Geometrical (cis/trans, fac/mer) – Isomers have the same empirical formula but differ in the spatial arrangement of ligands around the central metal atom. This can manifest as cis/trans isomerism e.g., square-planar $[\text{Pt(NH}<em>3)</em>2Cl<em>2]$; octahedral $[\text{Co(NH}</em>3)<em>4Cl</em>2]^+$; fac/mer in $[\text{Ma}<em>3b</em>3]$.

  • Optical (enantiomers) – Optical isomerism (enantiomerism) occurs in chiral coordination complexes that are non-superimposable on their mirror images and lack a plane of symmetry. Examples include [Co(en)3]3+[Co(en)3]3+ and cis-−[PtCl2(en)2]2+−[PtCl2(en)2​]2+.; e.g., $[\text{Co(en)}<em>3]^{3+}$, cis-$[\text{PtCl}</em>2(en)_2]^{2+}$.

• Structural isomerism

  1. Linkage – ambidentate ligand binds through different atoms: $[\text{Co(NH}<em>3)</em>5ONO]^{2+}$ vs $[\text{Co(NH}<em>3)</em>5NO_2]^{2+}$.

  2. Coordination – exchange of ligands between cat./an. entities: $[\text{Co(NH}<em>3)</em>6][\text{Cr(CN)}_6]$ vs vice-versa.

  3. Ionisation – interchange of ligands & counter-ions: $[\text{Co(NH}<em>3)</em>5SO<em>4]Br$ vs $[\text{Co(NH}</em>3)<em>5Br]SO</em>4$.

  4. Solvate (hydrate) – solvent in/out of sphere: $[\text{Cr(H}<em>2O)</em>6]Cl<em>3$ (violet) vs $[\text{Cr(H}</em>2O)<em>5Cl]Cl</em>2·H_2O$ (grey-green).

Valence Bond Theory (VBT)

• Metal ion hybridises its orbitals (d, s, p) to accept σ-donor lone pairs.

  • Common schemes & geometries:
    • $sp^3$ → tetrahedral (CN 4);
    • $dsp^2$ → square-planar (CN 4);
    • $sp^3d$ → trigonal bipyramidal (CN 5);
    • $sp^3d^2$ (outer) or $d^2sp^3$ (inner) → octahedral (CN 6).

• Examples

  • Inner-orbital low-spin $[\text{Co(NH}<em>3)</em>6]^{3+}$ uses $d^2sp^3$ (3d pair up; diamagnetic).

  • Outer-orbital high-spin $[\text{CoF}_6]^{3-}$ uses $sp^3d^2$ (4d involved; paramagnetic).

  • $[\text{NiCl}_4]^{2-}$: $sp^3$ (tetrahedral, 2 unpaired e⁻, paramagnetic).

  • $[\text{Ni(CN)}_4]^{2-}$: $dsp^2$ (square-planar, all paired, diamagnetic).

• Limitations: qualitative only; cannot predict colours, energies, ligand strength, etc.

Crystal Field Theory (CFT)

• Treat ligands as point charges/dipoles; purely electrostatic.

• Octahedral field: degeneracy of $d$ orbitals splits into lower $t<em>{2g}(d</em>{xy},d<em>{yz},d</em>{xz})$ and higher $e<em>g(d</em>{x^2−y^2},d_{z^2})$.

  • Energy gap $\Delta<em>o$; $e</em>g$ raised $+\tfrac{3}{5}\Delta<em>o$, $t</em>{2g}$ lowered $-\tfrac{2}{5}\Delta_o$.

• Tetrahedral field: inverse splitting; $\Delta<em>t = \tfrac{4}{9}\Delta</em>o$; usually high-spin.

• Spectrochemical series (increasing field strength $\Delta<em>o$): I^- < Br^- < SCN^- < Cl^- < S^{2-} < F^- < OH^- < C2O4^{2-} < H2O < NCS^- < edta^{4-} < NH_3 < en < CN^- < CO

• Electron configurations depend on rivalry between $\Delta_o$ and pairing energy $P$:

  • If \Delta_o < P → high-spin (weak-field ligands).

  • If \Delta_o > P → low-spin (strong-field ligands).

• Colour originates from $d\rightarrow d$ transitions absorbing specific visible wavelengths; complementary colour observed.

  • $[\text{Ti(H}<em>2O)</em>6]^{3+}$ (d$^1$) absorbs blue-green (~498 nm) → appears violet.

  • Changes in ligand alter $\Delta$ → colour change e.g., Ni(II) series with en (green → violet).

• Magnetic moments (spin-only): $\mu = \sqrt{n(n+2)}$ BM; support high vs low spin assignments.

• Limitations: neglects covalency, incorrectly predicts field order of anionic ligands; improved by Ligand Field & MO theories.

Metal Carbonyls

• Homoleptic carbonyls: only CO ligands; zero-oxidation metals.

  • Geometry examples: $\text{Ni(CO)}<em>4$ tetrahedral; $\text{Fe(CO)}</em>5$ trigonal-bipyramidal; $\text{Cr(CO)}_6$ octahedral.

  • Bridged dimers: $[\text{Mn}<em>2(CO)</em>{10}]$ (Mn–Mn bond); $[\text{Co}<em>2(CO)</em>8]$ (2 μ-CO bridges).

• Synergic M–CO bond: dual interaction

  1. σ-donation: CO lone pair (C) → empty metal orbital.

  2. π-back-donation: filled metal $d_{\pi}$ → CO antibonding $\pi^*$.

  • Reinforces both M–C and weakens C–O (observed IR red-shift).

Importance & Applications

• Analytical chemistry: colourimetric detection (e.g., DMG for Ni), complexometric EDTA titrations (water hardness).

• Metallurgy: cyanide leaching of Au/Ag via $[\text{Au(CN)}<em>2]^-$; purification via carbonyls: Mond process $\text{Ni + 4CO} \rightarrow \text{Ni(CO)}</em>4$.

• Biological: chlorophyll (Mg-porphyrin), haemoglobin (Fe-porphyrin), vitamin $B_{12}$ (Co-corrin), metallo-enzymes (carbonic anhydrase, carboxypeptidase A).

• Industrial catalysis: Wilkinson’s catalyst $[(Ph<em>3P)</em>3RhCl]$ (alkene hydrogenation), many others.

• Electroplating/Photography: smoother deposits from $[\text{Ag(CN)}<em>2]^- , [\text{Au(CN)}</em>2]^-$; fixer hypo dissolves AgBr via [\text{Ag(S2O3)_2}]^{3-}.

• Medicinal: chelate therapy (EDTA for Pb, D-penicillamine for Cu); anti-cancer drugs e.g., cis-platin $\text{[PtCl}<em>2(NH</em>3)_2]$.

Tables & Key Data

• Hybridisation vs Geometry

  • $sp^3$ → tetrahedral (CN 4)

  • $dsp^2$ → square-planar (CN 4)

  • $sp^3d$ → trigonal-bipyramidal (CN 5)

  • $sp^3d^2$ / $d^2sp^3$ → octahedral (CN 6)

• Example oxidation-state calculations

  • $[Co(H<em>2O)(CN)(en)</em>2]^{2+}$: let x=OS(Co); $x+0+(-1)+0=+2$ → $x=+3$.

  • $[PtCl_4]^{2-}$: $x+4(-1)=-2$ → $x=+2$.

Representative Worked Illustrations

• Magnetic moment & geometry: $[\text{MnBr}_4]^{2-}$ has $\mu=5.9$ BM ⇒ $n=5$ unpaired e⁻ ⇒ high-spin d$^5$ tetrahedral ($sp^3$), not square-planar.

• Predicting spin state: $[\text{Fe(H<em>2O)}</em>6]^{3+}$ (weak field) → d$^5$ high-spin, 5 unpaired; $[\text{Fe(CN)}_6]^{3-}$ (strong field) → low-spin, 1 unpaired.

• Square-planar diamagnetism: $[\text{Ni(CN)}_4]^{2-}$ requires dsp$^2$; d$^8$ electrons pair up, no unpaired e⁻, hence diamagnetic.

  Created by:MAYA D N ,HSST CHEMISTRY

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