What this chapter is about
Coordination compounds are substances in which a central metal atom or ion is surrounded by molecules or ions called ligands, which donate electron pairs to the metal. These compounds are essential in chemistry because they explain the colours of many transition-metal solutions, the functioning of haemoglobin and chlorophyll, and the extraction of metals like silver and gold. You meet this chapter now because it builds directly on your knowledge of d-block elements, chemical bonding and electron configurations from earlier study.
After working through this chapter, you should be able to write IUPAC names and formulas for coordination compounds, identify the oxidation state and coordination number of the central metal, distinguish types of ligands, apply Werner's theory and the valence bond approach, and predict geometries, magnetic behaviour and isomerism in complexes. You will also understand how crystal field theory explains the colour and magnetism of these compounds.
Key ideas
- A coordination entity consists of a central metal atom or ion bonded to a fixed number of ligands; the charge on the whole entity is the sum of the metal's oxidation state and the charges on the ligands.
- Ligands are classified by the number of donor atoms: monodentate (one, e.g. Cl⁻, NH₃), bidentate (two, e.g. ethylenediamine, en), polydentate (many, e.g. EDTA⁴⁻ with six donor atoms).
- Werner's coordination theory distinguishes primary valency (oxidation state of metal, satisfied by anions) from secondary valency (coordination number, satisfied by ligands in a definite geometry).
- Coordination number is the total number of donor atoms directly attached to the metal; common values are 4 (tetrahedral or square planar) and 6 (octahedral).
- Isomerism in coordination compounds includes structural types (ionisation, linkage, coordination) and stereoisomers (geometrical cis-trans, optical mirror-image forms).
- Valence bond theory uses hybridisation (sp³, dsp², sp³d², d²sp³) to predict geometry and magnetic properties based on whether electrons are paired or unpaired.
- Crystal field theory treats ligands as point charges that split the d-orbitals of the metal; in an octahedral field the five d-orbitals split into a lower t₂g set and a higher eg set, with splitting energy Δₒ.
- Colour arises when a complex absorbs visible light whose energy matches Δₒ (or Δₜ for tetrahedral), promoting a d-electron to a higher orbital; the transmitted light is the complementary colour.
Formulas and facts to remember
- Oxidation state of metal = total charge on complex − sum of charges on all ligands.
- Coordination number = number of ligand donor atoms bonded to the metal (not the number of ligand molecules).
- IUPAC naming order: cation first, then anion; within the coordination sphere, ligands are named alphabetically, with anionic ligands ending in -o (e.g. chlorido), neutral ligands by their molecule name (aqua for H₂O, ammine for NH₃).
- Effective atomic number (EAN) = atomic number of metal − oxidation state + 2 × coordination number; a value of 36, 54 or 86 (noble-gas configuration) often indicates extra stability.
- Crystal field splitting energy Δₒ for octahedral complexes is about 10 Dq; for tetrahedral, Δₜ ≈ (4/9)Δₒ.
- Spectrochemical series (increasing Δ): I⁻ < Br⁻ < Cl⁻ < F⁻ < OH⁻ < H₂O < NH₃ < en < NO₂⁻ < CN⁻ < CO.
- High-spin complexes form when Δ is small (weak-field ligands); low-spin complexes form when Δ is large (strong-field ligands), causing electrons to pair in the lower set.
- Chelate effect: polydentate ligands form more stable complexes than equivalent monodentate ligands because the entropy change on complex formation is more favourable.
Worked examples
Example 1 – Finding oxidation state and coordination number
Problem: In the complex ion [Co(NH₃)₄Cl₂]⁺, determine the oxidation state and coordination number of cobalt.
Solution:
- NH₃ is a neutral ligand, so its contribution to charge is 0.
- Each Cl⁻ carries a charge of −1; two chloride ions give −2.
- Let the oxidation state of Co be x. The overall charge is +1.
- Equation: x + 4(0) + 2(−1) = +1 → x − 2 = +1 → x = +3.
- Co is in the +3 oxidation state.
- Coordination number = number of donor atoms = 4 (from NH₃) + 2 (from Cl) = 6.
Example 2 – IUPAC naming
Problem: Write the IUPAC name of K₃[Fe(CN)₆].
Solution:
- K⁺ is the cation; the anionic complex is [Fe(CN)₆]³⁻.
- CN⁻ (cyanido) is an anionic ligand; six of them give a total ligand charge of −6.
- For the complex charge to be −3: x + (−6) = −3 → x = +3, so Fe is in +3 oxidation state.
- Anionic complexes have the metal name ending in -ate; iron becomes ferrate.
- Name: potassium hexacyanidoferrate(III).
Example 3 – Predicting geometry and magnetism using crystal field theory
Problem: Predict the geometry, number of unpaired electrons and magnetic behaviour of [Ni(CN)₄]²⁻.
Solution:
- Ni in +2 state has electronic configuration [Ar] 3d⁸.
- CN⁻ is a strong-field ligand (high in the spectrochemical series), so pairing occurs.
- For coordination number 4 with strong-field ligands and a d⁸ metal, the favoured geometry is square planar (dsp² hybridisation).
- In square planar d⁸ complexes, all eight d-electrons occupy the four lower-energy orbitals and pair up completely.
- Unpaired electrons = 0; the complex is diamagnetic.
Common mistakes
- Adding the number of ligand molecules instead of donor atoms when finding coordination number → count each donor atom, not each ligand molecule (en has 2 donor atoms).
- Forgetting that neutral ligands contribute zero charge when calculating oxidation state → always assign charge only to ionic ligands.
- Naming ligands in the order they appear in the formula rather than alphabetically → list ligand names in strict alphabetical order, ignoring prefixes like di-, tri-.
- Assuming all four-coordinate complexes are tetrahedral → with d⁸ metals and strong-field ligands, square planar geometry is common.
- Using the crystal field splitting value Δₒ for a tetrahedral complex → use Δₜ ≈ (4/9)Δₒ; tetrahedral splitting is smaller.
Quick revision
- Coordination number = total donor atoms bonded to the metal.
- Oxidation state = complex charge minus sum of ligand charges.
- Spectrochemical series ranks ligands by their ability to cause d-orbital splitting (I⁻ weakest, CO strongest).
- Strong-field ligands produce low-spin, often diamagnetic complexes; weak-field ligands produce high-spin, paramagnetic complexes.
- Colour of a complex depends on the energy gap Δ; absorbed light's complementary colour is observed.
- Chelating ligands form extra-stable complexes due to the chelate effect.