What are coordination compounds?
Haemoglobin, chlorophyll and vitamin B₁₂ all contain a metal ion bound to surrounding molecules or ions. Such compounds, in which a central metal atom or ion is attached to a number of ions or molecules by coordinate bonds, are coordination compounds.
Werner's theory
Alfred Werner (Nobel Prize 1913) studied complexes of cobalt(III) chloride with ammonia. He proposed that metals show two kinds of valence:
Primary valence: the ionisable valence, equal to the oxidation state, satisfied by negative ions.
Secondary valence: non-ionisable, equal to the coordination number, satisfied by ligands. It is directed in space, giving complexes definite shapes.
Example. In CoCl₃·6NH₃, written as [Co(NH₃)₆]Cl₃, all three Cl⁻ are ionisable: adding AgNO₃ precipitates 3 mol AgCl. In CoCl₃·5NH₃, written [CoCl(NH₃)₅]Cl₂, only 2 mol AgCl precipitate, because one Cl is bound in the coordination sphere.
Key terms
Coordination entity: the central metal and its attached ligands, written in square brackets: [Co(NH₃)₆]³⁺.
Central atom/ion: the metal, a Lewis acid.
Ligands: ions or molecules bound to the metal, which donate electron pairs (Lewis bases).
Coordination number: the number of donor atoms bound to the metal. In [PtCl₆]²⁻ it is 6; in [Ni(NH₃)₄]²⁺ it is 4.
Counter ions: the ions outside the square brackets.
Ligands by number of donor atoms (denticity):
Type | Donor atoms | Examples |
|---|---|---|
Unidentate | 1 | Cl⁻, H₂O, NH₃, CN⁻ |
Didentate | 2 | ethane-1,2-diamine (en), oxalate (C₂O₄²⁻) |
Polydentate | several | EDTA⁴⁻ (hexadentate: 6 donor atoms) |
A di- or polydentate ligand that binds through two or more atoms forms a ring and is called a chelating ligand. Chelate complexes are more stable.
Ambidentate ligands can bind through either of two different atoms: NO₂⁻ (through N, nitrito-N, or through O, nitrito-O) and SCN⁻ (through S or N).
IUPAC nomenclature
Name the cation first, then the anion.
Within the complex, name the ligands in alphabetical order, then the metal.
Anionic ligands end in -o: chlorido, cyanido, hydroxido, oxalato. Neutral ligands keep their names, except aqua (H₂O), ammine (NH₃), carbonyl (CO) and nitrosyl (NO).
Prefixes di-, tri-, tetra- give the number of each ligand (ignored for alphabetical order). For complex ligand names, use bis-, tris-, tetrakis-.
Give the metal's oxidation state in Roman numerals in brackets.
If the complex is an anion, the metal's name ends in -ate: cobaltate, nickelate; some use Latin roots: ferrate (Fe), cuprate (Cu), argentate (Ag), plumbate (Pb).
Examples:
[Co(NH₃)₆]Cl₃: hexaamminecobalt(III) chloride
[CoCl(NH₃)₅]Cl₂: pentaamminechloridocobalt(III) chloride
K₃[Fe(CN)₆]: potassium hexacyanidoferrate(III)
[Pt(NH₃)₂Cl₂]: diamminedichloridoplatinum(II)
Isomerism
Structural isomers have different bonds:
Type | Arises from | Example |
|---|---|---|
Linkage | an ambidentate ligand binding through different atoms | [Co(NH₃)₅(NO₂)]Cl₂ vs [Co(NH₃)₅(ONO)]Cl₂ |
Coordination | exchange of ligands between cationic and anionic complexes | [Co(NH₃)₆][Cr(CN)₆] vs [Cr(NH₃)₆][Co(CN)₆] |
Ionisation | exchange of a ligand and a counter ion | [Co(NH₃)₅(SO₄)]Br vs [Co(NH₃)₅Br]SO₄ |
Solvate (hydrate) | water inside or outside the coordination sphere | [Cr(H₂O)₆]Cl₃ (violet) vs [CrCl(H₂O)₅]Cl₂·H₂O (green) |
Stereoisomers have the same bonds but different arrangements in space:
Geometrical (cis-trans) isomerism occurs in square planar [MX₂L₂] complexes, such as [Pt(NH₃)₂Cl₂], which exists as cis and trans forms, and in octahedral [MX₂L₄] complexes. The fac and mer isomers of octahedral [MA₃B₃] complexes are another case.
Optical isomerism occurs when the complex and its mirror image can't be superimposed (chiral). Example: cis-[Co(en)₂Cl₂]⁺ and [Co(en)₃]³⁺. The two forms (d and l) rotate plane-polarised light in opposite directions.
Valence bond theory
The metal uses hybridised orbitals to accept electron pairs from ligands.
Coordination number | Hybridisation | Shape | Example |
|---|---|---|---|
4 | sp³ | tetrahedral | [NiCl₄]²⁻ |
4 | dsp² | square planar | [Ni(CN)₄]²⁻ |
6 | d²sp³ (inner orbital, low spin) | octahedral | [Co(NH₃)₆]³⁺ |
6 | sp³d² (outer orbital, high spin) | octahedral | [CoF₆]³⁻ |
[Co(NH₃)₆]³⁺ is diamagnetic (all 3d electrons paired, using inner d orbitals), while [CoF₆]³⁻ is paramagnetic with 4 unpaired electrons.
Limitations: VBT can't explain colour, doesn't distinguish weak and strong ligands quantitatively, and makes several assumptions.
Crystal field theory
CFT treats ligands as point charges. In an octahedral field, the five d orbitals split into two sets:
t₂g (three orbitals, lower energy)
e_g (two orbitals, higher energy)
The energy difference is the crystal field splitting energy, Δₒ.
Whether electrons pair up in t₂g or go into e_g depends on Δₒ compared with the pairing energy P:
Strong-field ligands (large Δₒ > P): electrons pair in t₂g, giving low-spin complexes.
Weak-field ligands (small Δₒ < P): electrons spread into e_g, giving high-spin complexes.
Spectrochemical series (increasing field strength):
I⁻ < Br⁻ < SCN⁻ < Cl⁻ < S²⁻ < F⁻ < OH⁻ < C₂O₄²⁻ < H₂O < NCS⁻ < edta⁴⁻ < NH₃ < en < CN⁻ < CO
In a tetrahedral field, the splitting is reversed and smaller: Δₜ ≈ (4/9) Δₒ, so tetrahedral complexes are almost always high spin.
Colour: complexes absorb light that promotes an electron from t₂g to e_g, and show the complementary colour. [Ti(H₂O)₆]³⁺ (d¹) absorbs green-yellow light and looks violet. Anhydrous CuSO₄ is white, but with water ligands it becomes blue.
Uses
Analysis: EDTA titrations measure the hardness of water (Ca²⁺ and Mg²⁺); dimethylglyoxime detects Ni²⁺.
Metallurgy: gold and silver are extracted as cyanide complexes, [Au(CN)₂]⁻.
Biology: chlorophyll (Mg), haemoglobin (Fe) and vitamin B₁₂ (Co) are coordination compounds.
Medicine: cisplatin, cis-[Pt(NH₃)₂Cl₂], is used against cancer; EDTA removes excess lead from the body.
Catalysis: Wilkinson's catalyst, [(Ph₃P)₃RhCl], hydrogenates alkenes.
Summary
Werner: primary (ionisable) and secondary (coordination number) valences.
Name cations first, ligands alphabetically, metal with oxidation state; anionic complexes end in -ate.
Isomerism: linkage, coordination, ionisation, solvate; geometrical and optical.
VBT gives hybridisation and shape; CFT explains colour and magnetism via Δₒ.