Coordination Compounds

In this chapter: Werner's theory, double salts and complexes, key terms (ligand, denticity, chelate, coordination number, oxidation number), IUPAC nomenclature, structural and stereoisomerism, valence bond theory, crystal field theory in octahedral and tetrahedral fields, colour, bonding in metal carbonyls, and uses of coordination compounds.

Werner's theory

Werner prepared a series of cobalt(III) chloride-ammonia compounds and counted how many chloride ions each gave as AgCl with silver nitrate in the cold:

• A metal shows two kinds of valence. Primary valence is ionisable and corresponds to the oxidation state; it is satisfied by negative ions. Secondary valence is non-ionisable and equals the coordination number; it is satisfied by neutral molecules or negative ions.
• Groups satisfying the secondary valence are arranged in definite directions in space, giving geometries such as octahedral, square planar and tetrahedral.

A double salt such as Mohr's salt, FeSO₄·(NH₄)₂SO₄·6H₂O, or potash alum dissociates completely into simple ions in water. A complex such as K₄[Fe(CN)₆] does not: its solution gives no test for Fe²⁺ or CN⁻.

Key terms

• Coordination entity: a central metal atom or ion bonded to a fixed number of ions or molecules, e.g. [CoCl₃(NH₃)₃] or [Fe(CN)₆]⁴⁻. The part in square brackets is the coordination sphere; ions outside it are counter ions.
• Ligands: ions or molecules bonded to the central atom through donor atoms. Unidentate: Cl⁻, H₂O, NH₃. Didentate: ethane-1,2-diamine (en, H₂NCH₂CH₂NH₂) and oxalate (C₂O₄²⁻). Hexadentate: ethylenediaminetetraacetate (EDTA⁴⁻), which binds through two N and four O atoms.
• Chelate: a di- or polydentate ligand that binds through two or more donor atoms to the same metal ion, forming a ring. Chelate complexes are more stable than similar complexes of unidentate ligands.
• Ambidentate ligand: can bind through two different atoms, e.g. NO₂⁻ (through N as nitrito-N, or through O as nitrito-O) and SCN⁻ (through S as thiocyanato-S, or N as thiocyanato-N).
• Coordination number (CN): the number of ligand donor atoms directly bonded to the metal. In [Co(en)₃]³⁺ it is 6, not 3.
• Oxidation number of the central atom: the charge it would carry if all ligands were removed along with their electron pairs.
• Homoleptic complexes have only one kind of ligand ([Co(NH₃)₆]³⁺); heteroleptic ones have more than one ([Co(NH₃)₄Cl₂]⁺).

Worked example: Find the coordination number and oxidation state of cobalt in [Co(en)₂Cl₂]Cl.
Solution: each en is didentate, so the donor atoms are 2 × 2 + 2 Cl = 6. The complex ion carries +1 (one Cl⁻ outside). en is neutral and each Cl is −1, so x + 0 − 2 = +1, giving x = +3.

IUPAC nomenclature

1. Name the cation first, then the anion, whether or not the complex is the cation.
2. Inside the coordination entity, name ligands in alphabetical order before the metal. Multiplying prefixes (di, tri) are ignored for alphabetical order.
3. Anionic ligands end in -o: chlorido, cyanido, hydroxido, oxido, oxalato, nitrito-N. Neutral ligands keep their names except aqua (H₂O), ammine (NH₃), carbonyl (CO) and nitrosyl (NO).
4. Use di, tri, tetra for simple ligands; bis, tris, tetrakis for ligands whose names already contain a number, with the ligand name in brackets, e.g. tris(ethane-1,2-diamine).
5. Give the oxidation state of the metal in Roman numerals in brackets after it.
6. If the complex ion is an anion, the metal name ends in -ate: cobaltate, ferrate (Fe), cuprate (Cu), argentate (Ag), aurate (Au), plumbate (Pb), stannate (Sn).

In the formula, the central atom is written first, followed by the ligands in alphabetical order of their symbols, whatever their charge. So the formula is written [CoCl(NH₃)₅]Cl₂, with Cl before NH₃.

Isomerism

Structural isomerism

Geometrical isomerism

Geometrical isomers differ in the relative position of ligands. It occurs in square planar complexes of type [MX₂L₂] and in octahedral complexes of types [MX₂L₄] and [MX₂(L-L)₂] (cis and trans), and [MA₃B₃] (fac, with the three A on one face, and mer, with them around a meridian). Tetrahedral complexes do not show geometrical isomerism, because all four positions are adjacent to each other.

Optical isomerism

Optical isomers are non-superimposable mirror images (enantiomers), called dextro (d) and laevo (l) forms. They are common in octahedral complexes with didentate ligands: [Co(en)₃]³⁺ is optically active, and so is cis-[PtCl₂(en)₂]²⁺. The trans isomer of [PtCl₂(en)₂]²⁺ has a mirror plane and is optically inactive.

Valence bond theory

The metal uses hybridised orbitals to accept lone pairs from the ligands. The hybridisation fixes the shape, and the number of unpaired electrons fixes the magnetic behaviour.

VBT has clear limits: it involves several assumptions, does not explain colour, gives no quantitative account of magnetic data, cannot distinguish weak and strong field ligands, and cannot say exactly why some 4-coordinate complexes are tetrahedral and others square planar.

Crystal field theory

CFT treats the metal-ligand bond as purely ionic, with ligands acting as point charges (or dipoles). In a free ion the five d orbitals have the same energy. The approaching ligands raise their energy and, because the ligands come from particular directions, split them into groups. This is crystal field splitting.

Octahedral complexes

Six ligands approach along the x, y and z axes. The d_x²−y² and d_z² orbitals point straight at them and are raised (e_g set, +0.6Δ_o); d_xy, d_yz and d_xz point between the axes and are lowered (t_2g set, −0.4Δ_o). The size of Δ_o depends on the ligand. The spectrochemical series, found from experiment, lists ligands in increasing order of field strength:

For d¹ to d³ the electrons go singly into t_2g. From d⁴ onwards there is a choice. If Δ_o is smaller than the pairing energy P (weak field ligands), the fourth electron enters e_g: t_2g³e_g¹, a high spin complex. If Δ_o > P (strong field ligands), it pairs in t_2g: t_2g⁴e_g⁰, a low spin complex.

Tetrahedral complexes

The splitting is inverted (e below, t₂ above) and smaller, Δ_t = (4/9)Δ_o. The energy needed for pairing is generally larger than Δ_t, so low spin tetrahedral complexes are rarely observed. The subscript g is not used because a tetrahedron has no centre of symmetry.

Worked example: Predict the d electron arrangement and magnetic behaviour of [Fe(CN)₆]⁴⁻ and [CoF₆]³⁻.
Solution: in [Fe(CN)₆]⁴⁻, Fe is +2 (3d⁶) and CN⁻ is a strong field ligand, so the arrangement is t_2g⁶e_g⁰: no unpaired electrons, diamagnetic, CFSE = 6 × (−0.4Δ_o) = −2.4Δ_o (plus pairing energy). In [CoF₆]³⁻, Co is +3 (3d⁶) and F⁻ is a weak field ligand, so it is t_2g⁴e_g²: 4 unpaired electrons, μ = √24 = 4.90 BM.

Colour

When white light falls on a complex, an electron absorbs light of energy Δ and moves from t_2g to e_g (a d-d transition). The colour seen is complementary to the colour absorbed. [Ti(H₂O)₆]³⁺ (d¹) absorbs in the blue-green region and looks violet. Without ligands there is no splitting, so anhydrous CuSO₄ is white while CuSO₄·5H₂O is blue. Changing the ligand changes Δ and therefore the colour.

CFT also has limits: treating ligands as point charges means anionic ligands should give the largest splitting, yet they sit at the weak end of the series; and it ignores the covalent character of metal-ligand bonds.

Bonding in metal carbonyls

Homoleptic carbonyls: [Ni(CO)₄] is tetrahedral, [Fe(CO)₅] trigonal bipyramidal and [Cr(CO)₆] octahedral. The M-C bond has both σ and π character. CO donates its carbon lone pair into a vacant metal orbital (σ bond), and the metal donates electrons from a filled d orbital into the empty antibonding π* orbital of CO (π back bonding). This synergic bonding strengthens the M-C bond and weakens the C-O bond.

Importance and applications

• Analysis: hardness of water is estimated by titration with Na₂EDTA, which forms stable complexes with Ca²⁺ and Mg²⁺. Ni²⁺ is detected with dimethylglyoxime.
• Biology: chlorophyll is a Mg complex, haemoglobin an Fe complex and vitamin B₁₂ (cyanocobalamin) a Co complex.
• Metallurgy: gold and silver are extracted as cyanide complexes such as [Au(CN)₂]⁻; nickel is purified through [Ni(CO)₄].
• Catalysis: Wilkinson's catalyst, [(Ph₃P)₃RhCl], is used for hydrogenation of alkenes.
• Medicine: EDTA removes excess lead from the body in lead poisoning; cis-platin inhibits the growth of tumours.
• Photography: undecomposed AgBr on a developed film dissolves in hypo as [Ag(S₂O₃)₂]³⁻.

Common mistakes: (1) Counting ligands instead of donor atoms for the coordination number. (2) Alphabetising by prefix ("dichlorido" under d). Use the ligand name, so ammine comes before chlorido. (3) Writing "ammine" for an amine or "amine" for NH₃. (4) Forgetting that a trans isomer with didentate ligands is usually optically inactive. (5) Assuming every Co³⁺ complex is diamagnetic. With weak field F⁻ it is high spin.

JEE and NEET focus

• Werner's theory problems: number of ions in solution, moles of AgCl precipitated, conductivity comparisons.
• IUPAC names both ways, including anionic complexes (-ate names) and ambidentate ligands.
• Counting geometrical and optical isomers for [MX₂L₂], [MX₂L₄], [MA₃B₃] and [M(AA)₃] types.
• Hybridisation, geometry and spin-only magnetic moment from VBT and from CFT with the spectrochemical series.
• Octahedral and tetrahedral splitting diagrams, CFSE, the reason for colour, and synergic bonding in carbonyls.

Practice questions