Compare ligand exchange colours in copper(II) and cobalt(II)
Use the syllabus examples to explain how changing ligands changes d-orbital splitting and observed complex colour.
A Level; Cambridge International Chemistry 9701, 2028–2030 syllabus, Topic 28. Use the first-row transition elements titanium to copper and the species explicitly named in the syllabus. Where redox potentials or stability constants are needed, use the values supplied in the question. Colour descriptions can depend on conditions and concentration. Written practice is self-checked
Before you begin
- Use electron configurations, oxidation numbers, ionic equations, equilibrium expressions and the AS redox and bonding models.
Work on paper. Try each question before revealing help, and explain why your method works.
Quick readiness check
Why can Cu²⁺ complexes with different ligands have different colours?
What you will learn
- Use Cu²⁺ aqua/ammine/chloro examples to connect ligand exchange and colour
- Use Co²⁺ aqua/chloro examples to connect ligand exchange and colour
- Explain that colour differences reflect changed ligand fields rather than necessarily a redox change
Copper(II) changes colour when its ligand set changes
The pale-blue Cu²⁺ aqua complex and the deep-blue ammine complex have different ligands around Cu²⁺. In chloride-rich solution, a chloro complex can also form. These observations illustrate ligand exchange and different d-orbital energy gaps.
The ligand change can alter ΔE, which changes the frequency absorbed and the complementary colour seen. Keep the copper oxidation state at +2 unless a redox equation gives evidence otherwise.
Cobalt(II) aqua/chloro exchange changes the observed colour
Pink [Co(H₂O)₆]²⁺ and blue [CoCl₄]²⁻ illustrate the effect of replacing water ligands with chloride ligands. Their geometries and ligand fields differ, so the splitting and absorbed light differ.
A colour change supports a change in coordination environment; it does not by itself prove a change in oxidation state. For a mixture at equilibrium, observed colour can be intermediate and depends on composition.
Pause and explain
What remains the same in the pink-aqua to blue-chloro cobalt(II) ligand exchange?
Worked example
Explain why pink cobalt(II) aqua solution can become blue in excess chloride without invoking redox.
Show the worked solution
- Write the aqua/chloro ligand exchange.
- The cobalt remains +2 on both sides.
- Different ligands give a different splitting and absorbed frequency.
Answer Chloride replaces water to form [CoCl₄]²⁻ from [Co(H₂O)₆]²⁺. Co remains +2; the ligand environment changes ΔE and the absorbed frequency, giving a different complementary colour.
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Write your answer and reasoning first. Use a hint only if you are stuck. The model solution is for self-checking; your written work is not automatically marked.
Identify the ligand and colour change in the Cu²⁺ aqua-to-excess-ammonia example.
Give me a hint
Name the starting and final ligand environments.
Compare my reasoning
- Starting species is pale-blue [Cu(H₂O)₆]²⁺.
- Excess NH₃ forms the deep-blue ammine complex.
- Cu remains +2.
Water ligands are partly replaced by NH₃ to form deep-blue [Cu(NH₃)₄(H₂O)₂]²⁺ from the pale-blue aqua complex; copper remains Cu(II).
Look for these in your work
- Starting complex identified
- Final ligand set identified
- Oxidation state retained
Compare pink [Co(H₂O)₆]²⁺ and blue [CoCl₄]²⁻ in terms of what is replaced and why colour differs.
Give me a hint
Use ligand exchange followed by ligand-field absorption.
Compare my reasoning
- Cl⁻ replaces coordinated water.
- The coordination environment and splitting change.
- The absorbed and complementary observed colours can change.
Chloride replaces water around Co(II). The ligand field changes ΔE and the frequency absorbed, so a different complementary colour is observed.
Look for these in your work
- Ligand substitution stated
- Energy splitting linked
- Colour mechanism stated
A student reports that the colour change proves Cu²⁺ became Cu⁺. What additional evidence would be needed to make a redox claim?
Give me a hint
Distinguish a change in ligands from transfer of electrons.
Compare my reasoning
- Ligand exchange alone changes the coordination environment.
- A redox claim needs oxidation-state/electron-transfer evidence.
- The colour observation alone is not sufficient.
A colour change alone is not proof of reduction. The complexes can have different colours because their ligand fields differ while Cu remains +2; a redox claim needs an equation or other evidence of electron transfer and a changed oxidation state.
Look for these in your work
- Coordination explanation considered
- Electron transfer required
- Colour evidence qualified
Common mistakes
- Attributing every colour change to oxidation-state change
- Describing absorbed light rather than observed colour
- Ignoring geometry when comparing complexes
- Assuming a mixture has the exact colour of one pure complex
What can you explain now?
Why can Cu²⁺ complexes with different ligands have different colours?
Compare with the explanation
Different ligand fields change ΔE and the light absorbed
The observed colour follows the complementary transmitted/reflected light.
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Make it stick
How do the Cu(II) and Co(II) examples connect ligand exchange, ΔE and observed colour?
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