Transition metal characteristics, variable oxidation states, complex ions and ligands, coloured ions, catalysis, and ligand substitution.

What this dot point is asking

WJEC wants you to describe transition metal characteristics, explain variable oxidation states, define complex ions and ligands, account for colour, describe catalysis, and write ligand-substitution reactions.

The answer

Transition metal characteristics

Complex ions and ligands

The origin of colour

The ligands split the five $d$ orbitals into two energy levels. An electron absorbs visible light to make a $d$-$d$ transition; the colour seen is complementary to the light absorbed. The gap, and so the colour, depends on the metal, its oxidation state and the ligands.

Catalysis

Transition metals and their compounds catalyse reactions because their variable oxidation states let them accept and donate electrons (for example iron in the Haber process, vanadium(V) oxide in the Contact process).

What affects the colour

The size of the energy gap between the split $d$ orbitals, and so the colour seen, depends on three things: the metal and its oxidation state, the type of ligand, and the coordination number (geometry). A stronger-field ligand splits the orbitals more, raising the energy gap and shifting the absorbed wavelength. This is why $[\text{Cu}(\text{H}_2\text{O})_6]^{2+}$ is pale blue but $[\text{Cu}(\text{NH}_3)_4(\text{H}_2\text{O})_2]^{2+}$ is deep blue: substituting ammonia for water changes the gap and so the colour. Colorimetry exploits this, measuring how much light a coloured complex absorbs to find its concentration.

Variable oxidation states and catalysis

Transition metals show variable oxidation states because the $4s$ and $3d$ sub-shells are close in energy, so different numbers of electrons can be lost with similar energy cost. Iron commonly shows $+2$ and $+3$, manganese ranges from $+2$ to $+7$. This ability to switch oxidation states lets transition metals act as catalysts: they can accept electrons from one reactant (being reduced) and pass them to another (being reoxidised), providing a lower-energy electron-transfer pathway. Iron in the Haber process and manganese(IV) oxide in hydrogen peroxide decomposition both work this way.

Examples in context

Haemoglobin. Iron(II) in haemoglobin forms a complex with oxygen as a ligand; carbon monoxide binds more strongly to the same site, which is why it is toxic, a biological complex-ion example. Industrial catalysts. Iron (Haber), vanadium(V) oxide (Contact) and nickel (hydrogenation) all rely on variable oxidation states, central to large-scale chemical manufacture.

Try this

Q1. Define a complex ion. [1 mark]

• Cue. A central metal ion surrounded by ligands bonded by dative covalent bonds.

Q2. State the coordination number and shape of $[\text{Fe}(\text{H}_2\text{O})_6]^{2+}$. [1 mark]

• Cue. Coordination number $6$, octahedral.

Q3. Explain why $\text{Zn}^{2+}$ solutions are colourless. [1 mark]

• Cue. The $d$ sub-shell is full, so no $d$-$d$ transition is possible.

Q4. State two factors that affect the colour of a transition metal complex. [2 marks]

• Cue. Any two of: the metal and its oxidation state, the ligand, and the coordination number (geometry).

Q5. Explain why transition metals can act as catalysts. [1 mark]

• Cue. Their variable oxidation states let them accept and donate electrons, providing a lower-energy pathway.