Coordination Compounds: An Introduction to Complex Ions

Add a few drops of ammonia solution to pale blue copper(II) sulfate solution, and a pale blue precipitate forms. Add more ammonia and the precipitate dissolves, leaving a deep, intense royal-blue solution. Nothing has been oxidised or reduced — the copper is still Cu²⁺. What has changed is the set of molecules bonded directly to the copper ion. This is the chemistry of coordination compounds, also called complexes. It explains the colours of gemstones, how haemoglobin carries oxygen, how some cancer drugs work, and how water hardness is measured.

Key vocabulary

A complex (or coordination compound) consists of a central metal atom or ion surrounded by molecules or ions called ligands, each bonded to the metal by a dative (coordinate) covalent bond.

• Ligand: a molecule or ion with at least one lone pair that it donates to the metal. Common ligands: H₂O, NH₃, Cl⁻, OH⁻, CN⁻, CO.
• Donor atom: the atom in the ligand that provides the lone pair (O in water, N in ammonia).
• Coordination number: the number of dative bonds to the central metal ion.
• Complex ion: a complex with an overall charge, written in square brackets, e.g. [Cu(H₂O)₆]²⁺.

The metal ion acts as a Lewis acid (electron-pair acceptor), and the ligands act as Lewis bases (electron-pair donors) (see Lewis acids and bases and dative covalent bonds).

Why transition metals form complexes

Complexes can form with many metal ions, but they’re especially characteristic of transition metals (see transition metals). Transition metal ions:

• are relatively small and often highly charged, so they attract lone pairs strongly;
• have empty orbitals (3d, 4s, 4p) available to accept electron pairs.

The copper–ammonia example explained

1. In water, Cu²⁺ exists as the hexaaquacopper(II) ion, [Cu(H₂O)₆]²⁺ — six water molecules each donating a lone pair from oxygen. This ion is pale blue.
2. A little ammonia solution provides OH⁻ ions, which remove H⁺ from coordinated water molecules, giving a neutral complex that precipitates: copper(II) hydroxide, pale blue.
3. Excess ammonia replaces four of the water ligands (a ligand substitution reaction), forming a soluble, deep blue complex:

[Cu(H₂O)₆]²⁺ + 4NH₃ ⇌ [Cu(NH₃)₄(H₂O)₂]²⁺ + 4H₂O

The deep blue colour is used as a test for Cu²⁺ ions.

Coordination number and shape

The number and size of ligands decides the shape of the complex (see VSEPR and molecular geometry):

Coordination number	Shape	Bond angles	Examples
2	Linear	180°	[Ag(NH₃)₂]⁺ (Tollens’ reagent)
4	Tetrahedral	109.5°	[CuCl₄]²⁻, [CoCl₄]²⁻
4	Square planar	90°	[Pt(NH₃)₂Cl₂] (cisplatin), [Ni(CN)₄]²⁻
6	Octahedral	90°	[Cu(H₂O)₆]²⁺, [Fe(CN)₆]⁴⁻, [Co(NH₃)₆]³⁺

Six is the most common coordination number, with small ligands like water and ammonia. Large ligands such as Cl⁻ often give four, because fewer fit around the metal. Adding concentrated hydrochloric acid to pale blue copper(II) sulfate solution replaces the six water molecules with four chloride ions, forming the yellow-green tetrahedral [CuCl₄]²⁻ ion.

Unlike VSEPR shapes for main-group molecules, square planar and tetrahedral shapes for four-coordinate complexes depend on the metal’s d electrons, not just on counting electron pairs.

Types of ligand: denticity

Ligands are classified by how many dative bonds each one forms:

• Monodentate (“one-toothed”): one bond per ligand — H₂O, NH₃, Cl⁻, CN⁻.
• Bidentate: two bonds — 1,2-diaminoethane (en, H₂NCH₂CH₂NH₂) donates through both nitrogens; ethanedioate (oxalate, C₂O₄²⁻) through two oxygens.
• Multidentate (polydentate): several bonds — EDTA⁴⁻ can form six dative bonds, wrapping around a metal ion completely.

Complexes with multidentate ligands are called chelates (from the Greek for “claw”). They’re usually much more stable than complexes with monodentate ligands. Replacing six water molecules with one EDTA ion increases the number of free particles in solution (from 2 to 7), a large increase in entropy that strongly favours the chelate. This is the chelate effect.

Naming complex ions

Complex names follow IUPAC rules:

1. Name the ligands first, in alphabetical order, with Greek prefixes for numbers (di-, tri-, tetra-, hexa-).
   • H₂O → aqua
   • NH₃ → ammine (double m)
   • Cl⁻ → chlorido (older: chloro)
   • OH⁻ → hydroxido (older: hydroxo)
   • CN⁻ → cyanido (older: cyano)
   • CO → carbonyl
2. Name the metal, with its oxidation state in Roman numerals.
3. If the complex ion is negative, the metal name ends in -ate, often using the Latin root: ferrate (Fe), cuprate (Cu), argentate (Ag), plumbate (Pb), stannate (Sn), aurate (Au).

Formula	Name
[Cu(H₂O)₆]²⁺	hexaaquacopper(II) ion
[Cu(NH₃)₄(H₂O)₂]²⁺	tetraamminediaquacopper(II) ion
[Ag(NH₃)₂]⁺	diamminesilver(I) ion
[CuCl₄]²⁻	tetrachloridocuprate(II) ion
[Fe(CN)₆]³⁻	hexacyanidoferrate(III) ion
[Co(NH₃)₆]Cl₃	hexaamminecobalt(III) chloride

For the metal’s oxidation state, see naming compounds with transition metals; it’s found by subtracting the ligand charges from the complex’s overall charge.

Example: in [Fe(CN)₆]³⁻, six CN⁻ ligands give −6; overall charge −3; so Fe is +3.

Why complexes are coloured

In an isolated transition metal ion, the five d orbitals have the same energy. When ligands surround the ion, they split the d orbitals into groups of slightly different energies. An electron can absorb a photon of visible light and jump from a lower to a higher d orbital. The energy gap, ΔE, determines which wavelength is absorbed:

ΔE = hf = hc/λ

We see the complementary colour of the light absorbed. [Cu(H₂O)₆]²⁺ absorbs red-orange light and appears blue.

The size of the splitting — and so the colour — depends on:

• the metal and its oxidation state;
• the ligands (ammonia splits more than water, which is why the copper complex deepens to royal blue);
• the coordination number and shape.

This is why the same metal ion can show different colours, as in the copper complexes above, and why iron(II) and iron(III) solutions look different. Ions with empty (d⁰) or completely full (d¹⁰) d subshells, such as Sc³⁺ and Zn²⁺, have no possible d–d transitions and are usually colourless (see transition metal electron configurations).

Gemstones owe their colours to the same effect: rubies and emeralds both contain small amounts of Cr³⁺ ions, but in different crystal environments, which split the d orbitals differently — red in ruby, green in emerald.

Isomerism in complexes

Complexes can form isomers:

• Cis–trans (geometric) isomerism in square planar and octahedral complexes. Cisplatin, cis-[Pt(NH₃)₂Cl₂], has the two chloride ligands next to each other and is an important anticancer drug; its trans isomer is far less active. Cisplatin works by binding to DNA, forming links between nearby guanine bases that disrupt replication (see DNA structure).
• Optical isomerism in octahedral complexes with three bidentate ligands, such as [Ni(en)₃]²⁺, which exist as non-superimposable mirror images.

Complexes in biology and everyday life

• Haemoglobin: Fe²⁺ bonded to four nitrogen atoms of a haem ring and a histidine from the protein; O₂ binds at the sixth position. Carbon monoxide binds much more strongly, which is why it’s poisonous (see haemoglobin and carbon monoxide poisoning).
• Chlorophyll: a Mg²⁺ ion held by four nitrogen atoms (see chlorophyll).
• Vitamin B₁₂: a cobalt complex.
• Water hardness: EDTA is used to titrate Ca²⁺ and Mg²⁺ (see complexometric titration with EDTA).
• Chelation therapy: EDTA and similar ligands bind toxic metal ions such as lead so the body can excrete them.
• Tollens’ reagent: [Ag(NH₃)₂]⁺ distinguishes aldehydes from ketones by forming a silver mirror.
• Food preservation: EDTA is added to some foods to bind metal ions that would otherwise catalyse spoilage.

Common misconceptions

• “Ligands form ionic bonds to the metal.” They form dative covalent bonds (often with ionic character).
• “Coordination number equals the number of ligands.” Only for monodentate ligands; one EDTA gives six bonds.
• “Colour changes mean the metal was oxidised.” Ligand substitution changes colour without any redox change.
• “Ammine” is a spelling mistake. It’s the correct ligand name for NH₃ (amine, with one m, is an organic compound).

Key takeaways

• A complex has a central metal ion bonded to ligands through dative covalent bonds.
• Coordination number = number of dative bonds; common shapes are linear (2), tetrahedral or square planar (4), and octahedral (6).
• Multidentate ligands form stable chelates, driven partly by entropy.
• Names list ligands (aqua, ammine, chlorido, cyanido…) then the metal with its oxidation state; negative complexes end in -ate.
• d-orbital splitting by ligands explains the colours of complexes; complexes are central to biology, medicine and analysis.