Transition metals as d-block elements with variable oxidation states; ligands and complexes with coordinate bonds, coordination number and shape; the origin of colour in d-d transitions and the splitting of d orbitals; and the catalytic properties of transition metals.

What this key area is asking

The SQA wants you to describe transition metals as d-block elements with variable oxidation states, to define ligands and complexes and assign coordination number and shape, to explain the origin of colour in terms of d-orbital splitting and d-d transitions, and to explain why transition metals act as catalysts. Assigning oxidation numbers, naming the bond and shape, and the colour explanation are reliable exam earners.

d-block elements and oxidation states

The oxidation number of a metal in a compound or complex is found from the rule that the sum of oxidation numbers equals the overall charge, with neutral ligands counted as zero and charged ligands at their ionic charge. For example, in $[\text{Fe(CN)}_6]^{3-}$ the six cyanide ligands carry $-1$ each, so iron must be $+3$.

Ligands and complexes

Ligands are classified by how many bonds they make: monodentate ligands (such as water, ammonia and chloride) form one bond each, while bidentate ligands (such as ethanediamine) form two and polydentate ligands form several. The coordination number sets the shape: a coordination number of $6$ gives an octahedral complex, while $4$ gives either a tetrahedral or a square planar complex.

The origin of colour

The colour of a transition metal complex comes from the way ligands affect the d orbitals.

The size of the gap, and therefore the colour, depends on the metal, its oxidation state and the nature of the ligands. A complex needs a partly filled d subshell for d-d transitions, which is why ions with empty ($d^0$) or full ($d^{10}$) d subshells, such as $\text{Zn}^{2+}$, are colourless.

Catalytic properties

Transition metals and their compounds are widely used as catalysts for two reasons:

• Variable oxidation states let the metal accept and then donate electrons, acting as a temporary electron store in redox reactions. Iron in the Haber process and vanadium(V) oxide in the Contact process work this way.
• Forming intermediate complexes with reactant molecules provides an alternative reaction pathway with a lower activation energy, which speeds up the reaction without the catalyst being consumed overall.

Examples in context

Transition metal chemistry runs through industry, biology and analysis. The intense colours of $[\text{Cu(H}_2\text{O})_6]^{2+}$ (pale blue) and $[\text{Cu(NH}_3)_4(\text{H}_2\text{O})_2]^{2+}$ (deep blue) show how changing the ligand changes the gap and the colour, the basis of the ammonia test for copper(II). Coloured complexes are quantified by colorimetry, where absorbance is proportional to concentration. In biology, haemoglobin is an iron(II) complex that carries oxygen as a ligand, and many enzymes have transition metal ions at their active sites. Industrially, iron, nickel, platinum and vanadium(V) oxide catalyse the Haber process, hydrogenation, catalytic converters and the Contact process respectively.

Try this

Q1. Define the term ligand. [1 mark]

• Cue. A molecule or ion with a lone pair that forms a coordinate bond to a central metal ion.

Q2. State the shape of a complex with coordination number 6. [1 mark]

• Cue. Octahedral.

Q3. Explain why $\text{Zn}^{2+}$ compounds are colourless. [2 marks]

• Cue. $\text{Zn}^{2+}$ has a full $d^{10}$ subshell, so no d-d transition is possible and no visible light is absorbed.