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Must-know facts

20 facts

1Werner: 1 mol each of CoCl₃·6NH₃, ·5NH₃ and ·4NH₃ give 3, 2 and 1 mol AgCl; they are 1:3, 1:2 and 1:1 electrolytes.
2Primary valence = ionisable = oxidation state; secondary valence = non-ionisable = coordination number.
3Double salts (Mohr's salt, carnallite, potash alum) give all their simple ions in water; complexes such as K₄[Fe(CN)₆] do not.
4EDTA⁴⁻ is hexadentate (2 N + 4 O); en and oxalate are didentate; NO₂⁻ and SCN⁻ are ambidentate.
5Coordination number counts σ-bonded donor atoms: [Co(en)₃]³⁺ and [Fe(C₂O₄)₃]³⁻ both have CN 6.
6Formula: metal first, ligands alphabetical. Name: ligands alphabetical, then metal; anionic complex ends in -ate (ferrate, argentate, cobaltate).
7aqua = H₂O, ammine = NH₃, carbonyl = CO, nitrosyl = NO; anionic ligands end in -ido in the 2004 IUPAC draft (chlorido, cyanido).
8Tetrahedral complexes show no geometrical isomerism; square planar [MABXL] has three geometrical isomers.
9[Co(NH₃)₃(NO₂)₃] has fac and mer isomers; [Co(en)₃]³⁺ is optically active; only cis-[PtCl₂(en)₂]²⁺ is chiral.
10Linkage: [Co(NH₃)₅(ONO)]²⁺ red, [Co(NH₃)₅(NO₂)]²⁺ yellow. Hydrate: [Cr(H₂O)₆]Cl₃ violet, [Cr(H₂O)₅Cl]Cl₂·H₂O grey-green.
11[Co(NH₃)₆]³⁺: d²sp³, inner orbital, diamagnetic. [CoF₆]³⁻: sp³d², outer orbital, 4 unpaired.
12[NiCl₄]²⁻: sp³, tetrahedral, 2 unpaired. [Ni(CN)₄]²⁻: dsp², square planar, diamagnetic. [Ni(CO)₄]: tetrahedral, diamagnetic, Ni(0).
13[Fe(CN)₆]³⁻ has 1 unpaired electron, [FeF₆]³⁻ has 5; [Mn(CN)₆]³⁻ has 2, [MnCl₆]³⁻ has 4.
14Octahedral splitting: eg +(3/5)Δo, t₂g −(2/5)Δo; tetrahedral Δt = (4/9)Δo, inverted, nearly always high spin.
15Spectrochemical series ends: I⁻ weakest ... en < CN⁻ < CO strongest; H₂O < NH₃ < en.
16Δo < P: high spin t₂g³eg¹ (weak field); Δo > P: low spin t₂g⁴eg⁰ (strong field).
17[Ti(H₂O)₆]³⁺ absorbs 498 nm (blue-green) and looks violet; [Cu(H₂O)₄]²⁺ absorbs 600 nm (red) and looks blue.
18Ruby: Al₂O₃ with 0.5-1% Cr³⁺; emerald: Cr³⁺ in beryl, Be₃Al₂Si₆O₁₈.
19Carbonyls: Ni(CO)₄ tetrahedral, Fe(CO)₅ trigonal bipyramidal, Cr(CO)₆ octahedral; bonding is synergic σ donation + π back donation.
20Wilkinson's catalyst [(Ph₃P)₃RhCl] hydrogenates alkenes; cis-platin fights tumours; EDTA treats lead poisoning; hypo gives [Ag(S₂O₃)₂]³⁻.

Common traps

Where marks are lost

Counting every chloride in CoCl₃·5NH₃ as precipitable by AgNO₃.

Only chlorides outside the bracket are free: [CoCl(NH₃)₅]Cl₂ gives 2 mol AgCl, not 3.

Giving [Co(en)₃]³⁺ a coordination number of 3.

Count donor atoms, not ligands. Each en binds through two N atoms, so CN = 6.

Writing ligands in a formula with anions first, as in older books.

In the formula all ligands go in alphabetical order irrespective of charge, as the 2004 IUPAC draft recommends; in the name they are alphabetical too.

Naming the metal in an anionic complex without -ate, e.g. 'potassium hexacyanidoiron(III)'.

Anionic complexes end in -ate: potassium hexacyanidoferrate(III), with the Latin root for iron.

Expecting cis and trans forms for tetrahedral [MA₂B₂].

In a tetrahedron every corner touches every other, so there is only one arrangement. Geometrical isomers need square planar or octahedral geometry.

Calling trans-[PtCl₂(en)₂]²⁺ optically active.

Only the cis form is chiral; the trans form has a mirror plane and is superimposable on its image.

Assuming that since [NiCl₄]²⁻ and [Ni(CO)₄] are both tetrahedral they are both paramagnetic.

[NiCl₄]²⁻ is Ni²⁺, 3d⁸, with 2 unpaired electrons; [Ni(CO)₄] is Ni(0) and CO pairs everything, so it is diamagnetic.

Thinking anionic ligands must give the largest splitting because they are charged.

That is exactly where point-charge CFT fails: I⁻, Br⁻ and Cl⁻ are among the weakest ligands; CN⁻ and CO are the strongest.

Reporting the colour absorbed as the colour seen, e.g. saying [Ti(H₂O)₆]³⁺ looks blue-green.

We see the complement: it absorbs blue-green (498 nm) and looks violet.

Using the eg/t₂g labels with a g for tetrahedral complexes and drawing the octahedral order.

Tetrahedral splitting is inverted (e below t₂) and written without g, because a tetrahedron has no centre of symmetry. Δt = (4/9)Δo.

Formulas

5 to know

Oxidation state of the central atom

charge on entity = x + Σ(ligand charges)

K₃[Fe(CN)₆]: −3 = x − 6, so x = +3.

Octahedral splitting

E(eg) = +(3/5)Δo; E(t₂g) = −(2/5)Δo

Measured from the average (spherical-field) energy of the d orbitals.

Tetrahedral splitting

Δt = (4/9)Δo

Same metal, ligands and metal-ligand distance; order of levels is inverted.

Spin state

Δo < P → high spin (t₂g³eg¹ for d⁴); Δo > P → low spin (t₂g⁴eg⁰)

P is the energy needed to pair two electrons in one orbital.

Spin-only magnetic moment

μ = √[n(n + 2)] BM

[MnBr₄]²⁻: 5.9 BM means n = 5, so the ion is tetrahedral.

Key terms

17 terms

Coordination entity: A central metal atom or ion together with the fixed set of ligands bonded to it.
Ligand: An ion or molecule that gives an electron pair to the central atom in a complex.
Denticity: The number of donor atoms through which one ligand binds a single metal ion.
Chelate: A complex in which one ligand grips the same metal through two or more donor atoms, forming a ring.
Ambidentate ligand: A ligand with two different donor atoms that can bind through either, such as NO₂⁻ or SCN⁻.
Coordination number: The number of ligand donor atoms σ-bonded directly to the central atom.
Coordination sphere: The central atom and its ligands, written inside square brackets.
Counter ion: An ionisable ion written outside the square bracket that balances the complex's charge.
Coordination polyhedron: The shape traced by the donor atoms around the central atom, such as an octahedron.
Homoleptic / heteroleptic: Having only one kind of ligand / having more than one kind.
Primary valence: Werner's ionisable valence, satisfied by anions; today the oxidation state.
Secondary valence: Werner's non-ionisable valence; today the coordination number.
Inner / outer orbital complex: A complex that hybridises (n−1)d orbitals (d²sp³, low spin) / nd orbitals (sp³d², high spin).
Crystal field splitting (Δo): The energy gap between the eg and t₂g orbitals created by six ligands in an octahedron.
Pairing energy (P): The energy cost of putting a second electron into an orbital that already holds one.
Spectrochemical series: Ligands ranked by the size of the d-orbital splitting they cause, found from absorption spectra.
Synergic bonding: σ donation from CO to the metal and π back donation from metal to CO, each strengthening the other.

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