Why Some Salts Keep Their Ions Locked Inside a Complex

The red of blood, the green of plant leaves and the deep blue of copper sulphate in ammonia all come from metal ions surrounded by attached molecules or ions. These coordination compounds follow their own rules of bonding, naming and shape.

This part covers Werner's theory, the key terms, IUPAC naming, and isomerism in coordination compounds.

Werner's theory says a metal in a coordination compound shows two kinds of valence — ionisable primary valence, equal to its oxidation state, and non-ionisable secondary valence, equal to its coordination number and directed in space to give a definite shape; a double salt breaks completely into simple ions in water, while a coordination compound keeps its complex ion intact.

Postulates:

- Primary valence — ionisable, satisfied by negative ions
- Secondary valence — non-ionisable, satisfied by neutral molecules or negative ions, and fixed for a given metal
- Secondary valences point in fixed directions, giving shapes such as octahedral, tetrahedral or square planar

Worked example. Excess AgNO${}_3$ precipitates different amounts of AgCl per mole of three cobalt(III) ammine chlorides:

- CoCl${}_3\cdot 6$NH${}_3$ gives $3$ mol, so it is $[{\text{Co(NH}}_3{)}_6]{\text{Cl}}_3$
- CoCl${}_3\cdot 5$NH${}_3$ gives $2$ mol, so it is $[{\text{CoCl(NH}}_3{)}_5]{\text{Cl}}_2$
- CoCl${}_3\cdot 4$NH${}_3$ gives $1$ mol, so it is $[{\text{CoCl}}_2({\text{NH}}_3{)}_4]\text{Cl}$

Only chloride ions outside the square brackets are free to precipitate.

Double salts versus coordination compounds:

- Mohr's salt, FeSO${}_4\cdot ({\text{NH}}_4{)}_2{\text{SO}}_4\cdot 6{\text{H}}_2\text{O}$, and potash alum dissociate fully into simple ions
- Potassium ferrocyanide, K${}_4[{\text{Fe(CN)}}_6]$, gives K${}^{+}$ and $[{\text{Fe(CN)}}_6{]}^{4-}$ but no free Fe${}^{2+}$ or CN${}^{-}$ ions

An everyday example. **Potash alum, sold as phitkari and used to settle muddy water, is a double salt that releases all its ions in solution.

The substance. Precipitation and conductivity tests reveal how many ions a complex really releases** — exactly how formulas like those above are fixed.

A central metal atom or ion is bonded to ligands that donate electron pairs; the number of donor atoms attached is the coordination number, the metal and ligands in square brackets form the coordination sphere, their arrangement in space is the coordination polyhedron, and the oxidation number is the charge the metal would carry if every ligand were removed with its electron pair.

Types of ligand:

- Unidentate — one donor atom, such as Cl${}^{-}$, H${}_2$O or NH${}_3$
- Didentate — two donor atoms, such as ethane-1,2-diamine (en) or oxalate
- Polydentate — several donor atoms, such as EDTA${}^{4-}$, which binds through six
- Ambidentate — can bind through either of two atoms, such as NO${}_2^{-}$ through N or O, and SCN${}^{-}$ through S or N
- Chelating — a di- or polydentate ligand that forms a ring with the metal, giving extra stability

Homoleptic complexes contain one kind of ligand, as in $[{\text{Co(NH}}_3{)}_6{]}^{3+}$; heteroleptic complexes contain more than one, as in $[{\text{Co(NH}}_3{)}_4{\text{Cl}}_2{]}^{+}$.

Worked example. In K${}_3[{\text{Fe(C}}_2{\text{O}}_4{)}_3]$, each oxalate is didentate with charge $-2$:

$$\text{coordination number}=3\times 2=6,x+3(-2)=-3\Rightarrow x=+3$$

An everyday example. EDTA is added to some packaged foods and shampoos because it chelates metal ions that would otherwise spoil the product.

The substance. Coordination number counts donor atoms, not ligands — three didentate oxalates give a coordination number of $6$.

Name the cation before the anion; inside the complex, list ligands alphabetically before the metal, give the metal's oxidation state in Roman numerals, and end the metal's name in -ate when the complex is an anion.

Rules:

- Ligands first, alphabetically, ignoring multiplying prefixes such as di- and tri-
- Anionic ligands end in -o — chlorido, cyanido, oxalato, hydroxido; neutral ones keep their names except aqua, ammine, carbonyl and nitrosyl
- Prefixes di-, tri- and tetra- for simple ligands; bis- and tris- for ligands with complex names such as ethane-1,2-diamine
- Oxidation state in Roman numerals after the metal
- Anionic complexes take -ate endings: cobaltate, ferrate for iron, cuprate for copper, argentate for silver

Worked example 1 — formula to name:

- $[{\text{Co(NH}}_3{)}_5\text{Cl}]{\text{Cl}}_2$ — pentaamminechloridocobalt(III) chloride
- K${}_3[{\text{Fe(CN)}}_6]$ — potassium hexacyanidoferrate(III)
- $[{\text{Pt(NH}}_3{)}_2{\text{Cl(NO}}_2)]$ — diamminechloridonitrito-N-platinum(II)

Worked example 2 — name to formula. Tetraamminediaquacobalt(III) chloride has four NH${}_3$, two H${}_2$O and Co${}^{3+}$, so the complex carries $+3$ and needs three chloride ions: $[{\text{Co(NH}}_3{)}_4({\text{H}}_2\text{O}{)}_2]{\text{Cl}}_3$.

An everyday example. Prussian blue, a pigment used in inks and paints, is iron(III) hexacyanidoferrate(II) — a name that shows both iron oxidation states at once.

The substance. Alphabetical order uses the ligand's name, not its prefix — so ammine comes before chlorido even in tetraamminedichlorido.

Coordination compounds show structural isomerism — ionisation, linkage, coordination and solvate isomers, which differ in their bonds — and stereoisomerism — geometrical and optical isomers, which have the same bonds arranged differently in space.

Structural isomers:

- Ionisation — ions swap between the sphere and outside: $[{\text{Co(NH}}_3{)}_5{\text{SO}}_4]\text{Br}$ and $[{\text{Co(NH}}_3{)}_5\text{Br}]{\text{SO}}_4$
- Linkage — an ambidentate ligand binds through different atoms, as with nitrito-N and nitrito-O
- Coordination — ligands swap between cationic and anionic complexes: $[{\text{Co(NH}}_3{)}_6][{\text{Cr(CN)}}_6]$ and $[{\text{Cr(NH}}_3{)}_6][{\text{Co(CN)}}_6]$
- Solvate (hydrate) — water inside or outside the sphere: violet $[{\text{Cr(H}}_2{\text{O)}}_6]{\text{Cl}}_3$ and grey-green $[{\text{CrCl(H}}_2{\text{O)}}_5]{\text{Cl}}_2\cdot {\text{H}}_2\text{O}$

Stereoisomers:

- Geometrical — in square planar $[{\text{Pt(NH}}_3{)}_2{\text{Cl}}_2]$ the chlorides can be adjacent (cis) or opposite (trans); octahedral $[{\text{MA}}_3{\text{B}}_3]$ complexes show fac and mer forms
- Optical — non-superimposable mirror images, common in octahedral complexes with didentate ligands such as $[{\text{Co(en)}}_3{]}^{3+}$

Worked example. $[{\text{Co(NH}}_3{)}_4{\text{Cl}}_2{]}^{+}$ has two geometrical isomers, cis and trans. In $[{\text{CoCl}}_2(\text{en}{)}_2{]}^{+}$ only the cis form is optically active, so it has three stereoisomers in all: trans, and a pair of cis enantiomers.

An everyday example. Cisplatin, the cis isomer of $[{\text{Pt(NH}}_3{)}_2{\text{Cl}}_2]$, is used as an anticancer medicine, while its trans isomer is not effective.

The substance. Tetrahedral complexes show no geometrical isomerism, because every position in a tetrahedron is adjacent to every other.