Complex Ions: Ligands, Shapes, Names and Isomers

1Blue that becomes deeper blue

Copper(II) sulfate dissolves in water to give a pale blue solution. Add dilute ammonia solution a drop at a time: first a pale blue solid appears, and then, with more ammonia, it dissolves to give a solution of an intense, deep blue. The copper is still copper(II) throughout. What changes is the set of molecules attached to the copper ion. In water it is surrounded by six water molecules; in excess ammonia four of them are replaced by ammonia molecules. These are complex ions, and they are one of the defining features of transition-metal chemistry.

2The shape follows the coordination number

The lone pairs from the ligands repel one another, so the ligands take up positions as far apart as possible, as in the VSEPR theory you used in S4. Small ligands such as water and ammonia usually give six-coordinate, octahedral complexes. Large ligands such as $\mathrm{C}{\mathrm{l}}^{-}\backslash mathrm\{Cl^-\}$ fit only four around a first-row ion, giving tetrahedral complexes. A few $d^{8}d^8$ ions, such as $\mathrm{N}{\mathrm{i}}^{2+}\backslash mathrm\{Ni^\{2+\}\}$ with cyanide and $\mathrm{P}{\mathrm{t}}^{2+}\backslash mathrm\{Pt^\{2+\}\}$, form square planar complexes, and $\mathrm{A}{\mathrm{g}}^{+}\backslash mathrm\{Ag^\{+\}\}$ forms linear ones.

Figure 3.1: The four common shapes of complex ions, with one example of each.

Figure 3.2: Bidentate ligands grip the metal ion with two donor atoms, forming a ring called a chelate.

3Ligand exchange and the chelate effect

One ligand can replace another. Two important exchanges change the colour:

$$[\mathrm{C}\mathrm{u}({\mathrm{H}}_2\mathrm{O}{)}_6{]}^{2+}(\mathrm{a}\mathrm{q})+4\mathrm{N}{\mathrm{H}}_3(\mathrm{a}\mathrm{q})\rightleftharpoons [\mathrm{C}\mathrm{u}(\mathrm{N}{\mathrm{H}}_3{)}_4({\mathrm{H}}_2\mathrm{O}{)}_2{]}^{2+}(\mathrm{a}\mathrm{q})+4{\mathrm{H}}_2\mathrm{O}(\mathrm{l})\backslash mathrm\{[Cu(H\_2O)\_6]^\{2+\}(aq)\; +\; 4NH\_3(aq)\; \backslash rightleftharpoons\; [Cu(NH\_3)\_4(H\_2O)\_2]^\{2+\}(aq)\; +\; 4H\_2O(l)\}$$

Pale blue becomes deep blue; the coordination number stays 6 and the shape stays octahedral.

$$[\mathrm{C}\mathrm{u}({\mathrm{H}}_2\mathrm{O}{)}_6{]}^{2+}(\mathrm{a}\mathrm{q})+4\mathrm{C}{\mathrm{l}}^{-}(\mathrm{a}\mathrm{q})\rightleftharpoons [\mathrm{C}\mathrm{u}\mathrm{C}{\mathrm{l}}_4{]}^{2-}(\mathrm{a}\mathrm{q})+6{\mathrm{H}}_2\mathrm{O}(\mathrm{l})\backslash mathrm\{[Cu(H\_2O)\_6]^\{2+\}(aq)\; +\; 4Cl^-(aq)\; \backslash rightleftharpoons\; [CuCl\_4]^\{2-\}(aq)\; +\; 6H\_2O(l)\}$$

With concentrated hydrochloric acid, blue becomes yellow-green: the chloride ions are larger, so only four fit, and the shape changes from octahedral to tetrahedral. Adding water pushes the equilibrium back.

Polydentate ligands replace monodentate ones very readily. In $[\mathrm{C}\mathrm{u}({\mathrm{H}}_2\mathrm{O}{)}_6{]}^{2+}+\mathrm{E}\mathrm{D}\mathrm{T}{\mathrm{A}}^{4-}\to [\mathrm{C}\mathrm{u}(\mathrm{E}\mathrm{D}\mathrm{T}\mathrm{A}){]}^{2-}+6{\mathrm{H}}_2\mathrm{O}\backslash mathrm\{[Cu(H\_2O)\_6]^\{2+\}\; +\; EDTA^\{4-\}\; \backslash rightarrow\; [Cu(EDTA)]^\{2-\}\; +\; 6H\_2O\}$, two particles become seven. The entropy of the system increases greatly, so $\mathrm{\Delta }G\backslash Delta\; G$ is strongly negative. This is the chelate effect. It is why EDTA is used to measure water hardness, to keep traces of metal ions from spoiling food, and in medicine to remove lead from the body. In your blood, each haem group holds an $\mathrm{F}{\mathrm{e}}^{2+}\backslash mathrm\{Fe^\{2+\}\}$ ion with four nitrogen atoms of a ring-shaped ligand; oxygen binds at another position, and carbon monoxide binds there more than 200 times more strongly, which is why it is so poisonous.

4Naming complex ions

1. In a compound, name the cation first, then the anion, as for any salt.
2. Inside the complex, name the ligands first, in alphabetical order, then the metal. Ignore the number prefixes when you alphabetise.
3. Show the number of each ligand with di-, tri-, tetra-, penta-, hexa-.
4. Neutral ligands: aqua (${\mathrm{H}}_2\mathrm{O}\backslash mathrm\{H\_2O\}$), ammine ($\mathrm{N}{\mathrm{H}}_3\backslash mathrm\{NH\_3\}$). Negative ligands end in -ido: chlorido ($\mathrm{C}{\mathrm{l}}^{-}\backslash mathrm\{Cl^-\}$), cyanido ($\mathrm{C}{\mathrm{N}}^{-}\backslash mathrm\{CN^-\}$), hydroxido ($\mathrm{O}{\mathrm{H}}^{-}\backslash mathrm\{OH^-\}$). Older books write chloro, cyano and hydroxo.
5. Give the metal's oxidation state in Roman numerals in brackets.
6. If the complex is an anion, the metal name ends in -ate, using the Latin stem for some metals: iron → ferrate, copper → cuprate, silver → argentate, lead → plumbate, tin → stannate. Otherwise: cobaltate, nickelate, chromate, manganate, zincate.

Table 3.1: Examples of names

5Isomerism in complexes

Table 3.2: Types of isomerism in transition-metal complexes

Figure 3.3: cis–trans isomers of a square planar complex. Only the cis isomer is an anticancer drug.

The cis isomer, cisplatin, has been used since the 1970s to treat several cancers. Its two chloride ligands are replaced by two nitrogen atoms on the same strand of a cancer cell's DNA, which stops the cell dividing. Transplatin cannot bind in this way and is not an effective drug: the same atoms in a different arrangement give a different biological effect.

Figure 3.4: Optical isomers of $[\mathrm{N}\mathrm{i}(\mathrm{e}\mathrm{n}{)}_3{]}^{2+}\backslash mathrm\{[Ni(en)\_3]^\{2+\}\}$: non-superimposable mirror images, like a left and a right hand.