Ligand Exchange & Coloured Ions (AQA A Level Chemistry): Video

Hi, I'm Eleanor with 3 years of experience teaching A-Level Chemistry, and this video is about ligand exchange and coloured ions.

Swapping the ligands on a metal ion changes the size of the energy gap between its split d orbitals, and that same energy gap is what decides the colour we see.

Changing the ligands surrounding a transition metal ion can change the size of the energy gap between its split d orbitals. The light energy absorbed corresponds to the energy gap between these orbitals. Different ligands produce different energy gaps, so they absorb different wavelengths of visible light. The wavelengths that are not absorbed are transmitted or reflected and are seen as the colour of the complex ion. Therefore, ligand exchange can change the colour of a transition metal complex.

So in this video: we'll look at ligand exchange itself, what happens when substitution is incomplete or the coordination number changes, two real examples — haem and the chelate effect — and then why complexes are coloured in the first place, and how colorimetry uses that colour to find a concentration.

Ligand exchange is when one ligand in a complex is replaced by another, forming a new, more stable complex. If the ligands are a similar size, the coordination number and shape don't change. Take cobalt(II): the pink hexaaqua complex turns pale yellow when ammonia replaces all six of the water ligands. Add excess concentrated ammonia and the cobalt is oxidised to cobalt(III), giving a brown solution.

Substitution isn't always complete. With a copper(II) ion complex with six water ligands, sodium hydroxide added dropwise gives a light blue precipitate where only two of the six water ligands have been replaced. Excess concentrated ammonia then dissolves that precipitate into a deep blue solution.

Ligands can also differ in size. Chloride ions are bigger than water, so when concentrated hydrochloric acid is added to a copper(II) ion complex with six water ligands, all six are replaced by four chloride ions. The coordination number drops from six to four, and the solution turns from blue to yellow. It's reversible — adding water brings the blue complex straight back.

Haem is an iron(II) complex. Oxygen binds weakly to the iron(II), which is what lets it be picked up and released again in the blood. Carbon monoxide binds more strongly and blocks that site, which is why it's toxic.

Separately, bidentate and multidentate ligands can displace monodentate ones — this is the chelate effect. It's driven by entropy: replacing six water ligands with one EDTA turns two particles into seven, producing a large positive entropy change that makes the reaction favourable even with a small enthalpy change.

So why are these complexes coloured at all? The d orbitals in an isolated metal ion are all the same energy. Once ligands bond to the metal, the d orbitals split into two energy levels. An electron can absorb light energy with a wavelength equal to the energy gap and jump from the lower to the higher level which is called electron promotion. The energy needed is delta E equals h multiplied by ν. H is Planck's constant, 6.626 times 10 to the minus 34 joule-second, and v is the wave frequency measured in hertz or per second. The wavelengths that aren't absorbed pass through, so what we see is the complementary colour of whatever was absorbed.

That energy gap, delta E, isn't fixed — it depends on the ligand, the coordination number, and the oxidation state. Different ligands split the d orbitals by different amounts, so the same metal ion can look completely different colours depending on what's bonded to it.

Coordination number matters too, because it usually comes with a change in geometry. An oxidation state change alters the metal ion's nuclear charge — which is why iron(III), with a stronger pull on its ligands, has a bigger delta E than manganese(II).

That colour change is also useful — it's how a colorimeter measures concentration. Light passes through a filter that lets through the colour the solution absorbs most, then through the sample to a detector. You build a calibration curve from solutions of known concentration, then read an unknown solution's absorbance off that line to find its concentration. If a solution's colour is too pale to measure accurately, you can intensify it by adding thiocyanate ions to iron(III), for example, forms a much more strongly coloured blood-red complex.

When explaining why transition metal complexes are coloured explain that electrons in the lower-energy d-orbitals absorb a specific wavelength or frequency of visible light or white light to be promoted to the higher-energy d-orbital. Just writing absorbs light without specifying visible or white is a very common mistake that will cost you marks

The remaining wavelengths of visible light are transmitted or reflected to give the complementary colour seen. Never say that the colour seen is emitted as examiners explicitly penalise this.

Ligand exchange is when one ligand is replaced by another, which can change the coordination number, shape, stability, and colour of a complex. The chelate effect explains why bidentate and multidentate ligands form more stable complexes.

Transition metal complexes are coloured because their d orbitals split into different energy levels. Electrons absorb specific wavelengths of visible light to move between these levels, and the remaining wavelengths give the complementary colour we see. Changing the ligand, coordination number, or oxidation state changes the energy gap and therefore the colour. Colorimetry uses this absorbance to determine the concentration of a coloured solution.