Ligand exchange and redox
Ligand exchange and redox
- In ligand exchange, one ligand replaces another — often with a colour change.
- Copper(II) is the classic example.
- Transition ions also do useful redox titrations.
Practice
Ligand exchange is when:
Swapping a ligand for another changes the complex and usually its colour.
Ligand exchange (copper II)
- $[\text{Cu}(\text{H}_2\text{O})_6]^{2+}$ is pale blue.
- Add ammonia → $[\text{Cu}(\text{NH}_3)_4(\text{H}_2\text{O})_2]^{2+}$, deep blue.
- Add concentrated HCl → $[\text{CuCl}_4]^{2-}$, yellow.
Practice
Adding concentrated HCl to pale-blue [Cu(H₂O)₆]²⁺ gives:
Chloride ligands replace water, giving the yellow tetrachlorocuprate(II) ion.
Redox titrations
- Use $E^{\ominus}$ values to predict whether a redox reaction is feasible.
- Classic titrations: $\text{MnO}_4^-$ with $\text{Fe}^{2+}$ (purple → colourless), and $\text{Cu}^{2+}$ with $\text{I}^-$ (makes iodine, then titrated with thiosulfate).
Practice
In a MnO₄⁻/Fe²⁺ titration in acid, the colour change at the end point is:
MnO₄⁻ (purple) is reduced to colourless Mn²⁺; the first excess drop gives a faint pink.
Practice
To predict whether a redox reaction of transition ions happens, you use:
A positive E°cell (from the E° values) tells you the redox reaction is feasible.
You've got it
Key idea
- ligand exchange swaps one ligand for another (colour change)
- copper(II): pale blue (water) → deep blue (ammonia) → yellow ($[\text{CuCl}_4]^{2-}$)
- use $E^{\ominus}$ for feasibility; classic titrations: $\text{MnO}_4^-/\text{Fe}^{2+}$ (purple → colourless), $\text{Cu}^{2+}/\text{I}^-$