Ligand substitution reactions

In a nutshell

​Ligand exchange involves a ligand being swapped for another ligand. A colour change is observed when ligand exchange occurs. Ligand exchange reactions are driven by positive entropy change.

Ligand exchange reactions

Ligand exchange involves a ligand being swapped for another ligand. Ligand exchange usually results in a colour change.

When a large, charged ligand is swapped for a small, uncharged ligand, the coordination number and the shape of the complex will change.

Example 

​$$[Cu(H_2O)_6]^{2+} \, (aq) \, + \, \, 4Cl^- \, (aq) \leftrightharpoons [CuCl_4]^{2-} \, (aq) + 6H_2O \, (l)$$​​

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In this ligand exchange, the relatively small, uncharged water ligands are being substituted for large, charged ligands. The complex ion changes from an octahedral shape to a tetrahedral shape.

The $$[Cu(H_20)_6]^{2+}$$ solution is pale blue. The $$[CuCl_4]^{2-}$$ solution is yellow. When the ligand exchange occurs, a colour change is observed.

If the ligands are similar in size, the coordination number and the shape of the complex ion will not change. For instance, $$CN^-$$​and $$OH^-$$ ligands are similar in size and have the same charge, so when forming a complex ion will adopt the same shape.

Example

​$$H_2O$$​ and $$NH_3$$​ are similar in size. In the ligand exchange below, the shape of the complex ion does not change; both $$[Cr(H_2O)_6]^{3+}$$ and $$[Cr(NH_3)_6]^{3+}$$​ complex ions are octahedral in shape.​

$$[Cr(H_2O)_6]^{3+} \, (aq) \space + \space 6NH_3 \, (aq) \rightleftharpoons [Cr(NH_3)_6]^{3+} \, (aq) \space + \space 6H_2O \, (l)$$​​

Sometimes ligand exchange involves partial substitution. 

​​Example

The reaction below only happens when an excess ammonia is added. 

​$$[Cu(H_2O)_6]^{2+}(aq) \space + \space 4NH_3(aq) \rightleftharpoons [Cu(NH_3)_4(H_2O)_2]^{2+}(aq) \space + \space 4H_2O(l)$$​​

Positive entropy change

Ligand exchange reactions involve the breaking and formation of new bonds. Ligand exchange reactions have very small enthalpy changes as the bonds broken and the new bonds formed have very similar strengths. 

Ligand exchange reactions are driven by positive entropy change. Substituting monodentate ligands with bidentate or multidentate ligands increases the number of particles in the solution and therefore increases the entropy. 

Example

​$$[Cr(NH_3)_6]^{3+}(aq) \space + \space EDTA^{4-}(aq) \rightarrow [Cr(EDTA)]^-(aq) \space + \space 6NH_3(aq)$$

In this reaction the monodentate ligands are replaced by EDTA, a multidentate ligand. A more stable complex forms. This reaction has a very small enthalpy change. 

This reaction has a large positive entropy change as it goes from two particles to seven particles; this gives rise to a negative free energy change and the reaction is feasible.

Coloured precipitates 

When transition elements are added to water, metal-aqua ions form. The metal-aqua ions have the formula $$[M(H_2O)_6]^{n+}$$ but can be written as $$M^{n+}$$ if there are only water ligands present.​

When an aqueous solution of a transition metal ion is mixed with aqueous ammonia or sodium hydroxide an acid-base reaction takes place. The water ligands of the complex ions are deprotonated and a coloured hydroxide precipitate forms.

To reverse the reaction, an acid can be added to the hydroxide precipitate to protonate the hydroxide ligands. The metal-aqua ions formed are soluble and the precipitate will eventually dissolve.

Iron($$II$$​)

$$[Fe(H_2O)_6]^{2+}$$ solution is pale green. The $$[Fe(OH)_2(H_2O)_4]$$ precipitate formed is green. The precipitate darkens over time; this is because the precipitate gets oxidised by oxygen and water in the air which forms iron($$III$$​) hydroxide.

​$$[Fe(H_2O)_6] ^{2+}(aq) \space + \space 2OH^-(aq) \rightarrow [Fe(OH)_2(H_2O)_4](s) \space + \space 2H_2O(l)$$

$$[Fe(H_2O)_6] ^{2+}(aq) \space + \space 2NH_3(aq) \rightarrow [Fe(OH)_2(H_2O)_4](s) \space + \space 2NH_4^+ (aq)$$

Iron($$III$$​)

$$[Fe(H_2O)_6] ^{3+}$$ solution is yellow. The $$[Fe(OH)_3(H_2O)_3]$$ precipitate formed is orange and darkens over time.​

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$$[Fe(H_2O)_6] ^{3+}(aq) \space + \space 3OH^-(aq) \rightarrow [Fe(OH)_3(H_2O)_3](s) \space + \space 3H_2O(l)$$

​$$[Fe(H_2O)_6] ^{3+}(aq) \space + \space 3NH_3(aq) \rightarrow [Fe(OH)_3(H_2O)_3](s) \space + \space 3NH_4^+(aq)$$​​

Copper($$II$$​)

$$[Cu(H_2O)_6] ^{2+}$$ solution is pale blue. The $$[Cu(OH)_2(H_2O)_4]$$ precipitate formed is blue.​

​$$[Cu(H_2O)_6] ^{2+}(aq) \space + \space 2OH^-(aq) \rightarrow [Cu(OH)_2(H_2O)_4](s) \space + \space 2H_2O(l)$$

$$[Cu(H_2O)_6] ^{2+}(aq) \space + \space 2NH_3(aq) \rightarrow [Cu(OH)_2(H_2O)_4](s) \space + \space 2NH_4^+(aq)$$​​​

When excess ammonia is added to copper($$II$$​) hydroxide, a ligand exchange reaction takes place. The blue precipitate changes to a deep blue solution.

​$$[Cu(OH)_2(H_2O)_4](s) \space + \space 4NH_3(aq) \rightarrow [Cu(NH_3)_4(H_2O)_2] ^{2+}(aq) \space + \space 2OH^-(aq) \space + \space 2H_2O(l)$$

Cobalt($$II$$​)

​$$[Co(H_2O)_6] ^{2+}$$ solution is pale pink. The $$[Co(OH)_2(H_2O)_4]$$ precipitate formed is blue but turns brown over time.​

​$$[Co(H_2O)_6] ^{2+}(aq) \space + \space 2OH^-(aq) \rightarrow [Co(OH)_2(H_2O)_4](s) \space + \space 2H_2O(l)$$

$$[Co(H_2O)_6] ^{2+}(aq) \space + \space 2NH_3(aq) \rightarrow [Co(OH)_2(H_2O)_4](s) \space + \space 2NH_4^+(aq)$$​​​

​When an excess ammonia is added to cobalt($$II$$) hydroxide complex, a ligand exchange reaction takes place. The blue precipitate dissolves to give a yellow-brown solution, containing $$[Co(NH_3)_6]^{2+}$$. 

$$[Co(OH)_2(H_2O)_4](s) \space + \space 6NH_3(aq) \rightarrow [Co(NH_3)_6] ^{2+}(aq) \space + \space 2OH^-(aq) \space + \space 4H_2O(l)$$​​

On standing, oxidation takes place and a brown solution forms, containing the complex $$[Co(NH_3)_6]^{3+}$$ .