Multidentate Ligands and the Chelate Effect Revision Notes | A Level AQA Chemistry

1. From Monodentate to Multidentate: Why Ligands Can Grip More Than Once

You already know that a ligand is a molecule or ion that donates a lone pair to a transition metal ion, forming a coordinate bond. Ammonia (NH3) and water (H2O) each donate one lone pair — they are monodentate ligands, from the Latin for 'one tooth'. The question worth asking is: what happens when a ligand has two or more lone pairs positioned so that it can bite onto the metal at two points simultaneously?

That is precisely what bidentate and multidentate ligands do. Rather than one lone pair reaching the metal, these ligands wrap around it, forming multiple coordinate bonds in a single act of coordination. The resulting ring-shaped structures — called chelate rings — are far more stable than their monodentate equivalents, and understanding why is the heart of this topic.

2. Bidentate Ligands: Two Points of Attachment

A bidentate ligand donates two lone pairs to the same metal ion, forming two coordinate bonds and creating a five- or six-membered chelate ring. The geometry of the ligand must allow both donor atoms to reach the metal simultaneously — which is why the spacing between donor atoms matters.

The ethanedioate ion, $C_2O_4^{2-}$ (oxalate), is the other key AQA bidentate example. Here, two oxygen atoms — one on each carboxylate group — each donate a lone pair to the metal. Because the two donor oxygens are on adjacent carbon atoms, the geometry again allows a five-membered chelate ring to form. Notice that ethanedioate carries a 2– charge, so it also contributes to reducing the overall charge on the complex.

A typical octahedral complex with three bidentate ligands — for example $[Co(en)_3]^{3+}$ — uses all six coordination sites, with each ethane-1,2-diamine molecule occupying two adjacent sites. Three ligands replace what would have been six separate monodentate ligands. That replacement is the starting point for understanding why chelate complexes are so stable.

3. Multidentate Ligands: EDTA4– and Six-Point Coordination

If bidentate ligands are impressive, multidentate ligands take coordination chemistry to another level. The most important multidentate ligand at A-level is EDTA4– (the fully deprotonated form of ethylenediaminetetraacetic acid). Its structure contains two nitrogen donor atoms and four carboxylate oxygen donor atoms — six lone pairs in total, all capable of coordinating to a single metal ion simultaneously.

When EDTA4– coordinates to a metal ion such as $Fe^{3+}$ or $Ca^{2+}$, it forms the complex $[Fe(EDTA)]^-$ or $[Ca(EDTA)]^{2-}$. One EDTA4– molecule completely envelops the metal, replacing all six monodentate ligands that would otherwise be needed. The resulting complex is extraordinarily stable — a fact that has direct practical consequences in medicine, analytical chemistry, and water treatment.

4. The Chelate Effect: Why Polydentate Complexes Are More Stable

Here is where the chemistry becomes genuinely interesting. Compare two reactions: replacing six ammonia ligands in $[Ni(NH_3)_6]^{2+}$ with either six separate monodentate ligands, or with three bidentate ethane-1,2-diamine molecules. Both reactions swap out the same number of coordinate bonds, and the enthalpy changes are similar. Yet the bidentate reaction is overwhelmingly more favourable. The reason is entropy.

Consider the substitution reaction below. When three en molecules replace six NH3 ligands, the total number of independent species in solution increases from four (one complex + three en) to seven (one new complex + six NH3). More particles in solution means a larger number of possible arrangements — a greater entropy ($\Delta S > 0$). This positive entropy change makes the reaction thermodynamically favourable even when the enthalpy change is only modestly negative.

Worked example: Counting Particles to Understand the Entropy Change

Compare the particle count before and after the substitution: $[Ni(NH_3)_6]^{2+}(aq) + 3en(aq) \rightleftharpoons [Ni(en)_3]^{2+}(aq) + 6NH_3(aq)$

1. 1 complex ion $[Ni(NH_3)_6]^{2+}$ + 3 en molecules = 4 independent species — Each dissolved species contributes to the disorder of the solution. We start with 4 independent particles.
2. 1 complex ion $[Ni(en)_3]^{2+}$ + 6 NH3 molecules = 7 independent species — After substitution, 7 independent particles are present in solution — a net gain of 3 particles.
3. $\Delta n_{\text{particles}} = 7 - 4 = +3$ — More independent particles in solution means more possible arrangements, so entropy increases: $\Delta S > 0$.
4. $\Delta G = \Delta H - T\Delta S$ — With $\Delta S > 0$ and T > 0, the $-T\Delta S$ term is negative. Even if $\Delta H$ is only slightly negative (or close to zero), $\Delta G$ becomes negative, making the reaction spontaneous.
5. $\Delta G < 0$ because the large positive $\Delta S$ dominates — The chelate complex $[Ni(en)_3]^{2+}$ is thermodynamically more stable than $[Ni(NH_3)_6]^{2+}$ primarily because of the entropy gain on releasing more free particles into solution.

Answer: The substitution releases a net gain of 3 free particles into solution, giving $\Delta S > 0$. Combined with a modestly negative $\Delta H$, this makes $\Delta G < 0$, so the chelate complex is thermodynamically favoured. This is the chelate effect.

The enthalpy change ($\Delta H$) for chelate substitution is typically small and negative — the new M–N or M–O bonds formed are of similar strength to the M–NH3 bonds broken. So enthalpy alone cannot explain the dramatic stability difference. It is the entropy term ($T\Delta S$) that drives the equilibrium strongly towards the chelate complex, particularly at higher temperatures where the $T\Delta S$ contribution grows larger.

5. EDTA4– and the Extreme Case of the Chelate Effect

The chelate effect becomes even more pronounced with EDTA4–. When EDTA4– replaces six water molecules coordinated to a metal ion — for example in $[Fe(H_2O)_6]^{3+}$ — the particle count changes dramatically. One EDTA4– replaces six water molecules, changing the particle count from 2 species (the original complex plus EDTA4-) to 7 species (the new complex plus 6 water molecules), a net gain of 5 particles.

This very large positive entropy change, combined with a small negative enthalpy change — since the six coordinate bonds formed are of similar strength to the six bonds broken — gives an exceptionally negative $\Delta G$. The stability constant ($K_{stab}$) for EDTA4– complexes is therefore enormous — often $10^{20}$ or greater — reflecting how strongly the equilibrium lies towards the chelate complex. This is why EDTA is used clinically to remove toxic metal ions such as lead from the bloodstream: once bound, the metal is effectively locked in place.

6. Exam Technique: Writing High-Mark Answers on the Chelate Effect

AQA questions on the chelate effect typically carry 3–6 marks and use the command words 'explain' or 'explain in terms of'. A common student error is to write only about entropy and forget to address enthalpy — or to state that enthalpy is the main driving force when it is not. The mark scheme expects you to address both thermodynamic quantities and to link them explicitly to the Gibbs equation.

Another frequent trap is describing the entropy change vaguely as 'disorder increases'. Examiners want a mechanistic reason: the number of independent particles in solution increases, so there are more possible arrangements of those particles, giving a positive $\Delta S$. Quoting the particle count change (for example, 4 particles become 7 in the en substitution) earns the mark that a vague statement about 'disorder' does not.

The key exam idea is not just that EDTA forms six bonds, but that replacing several separate ligands with one multidentate ligand changes the particle count in solution. That entropy gain is what students must connect to the Gibbs free energy equation.