Explain copper(II) complex formation and ligand exchange
Use equations, ligand donor pairs and observations to explain copper(II) reactions with water, ammonia, hydroxide and chloride.
Section 28 · Chemistry of transition elements
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
What is the central reason excess NH₃ changes the copper(II) complex?
What you will learn
- Describe selected Cu²⁺ complexes with water, ammonia, hydroxide and chloride
- Explain ligand exchange as replacement of coordinated ligands
- Use colour or precipitate evidence while keeping conditions explicit
Water and ammonia form different copper(II) complexes
In water, Cu²⁺ is represented by the pale-blue hexaaquacopper(II) ion, [Cu(H₂O)₆]²⁺. Adding a small amount of aqueous ammonia can first produce a pale-blue Cu(OH)₂ precipitate because ammonia acts as a base. In excess ammonia, the precipitate dissolves and a deep-blue ammine complex forms, commonly represented as [Cu(NH₃)₄(H₂O)₂]²⁺.
The excess-ammonia change replaces coordinated water with ammonia ligands. The visible colour change is evidence of a different complex environment; it does not mean the copper oxidation state changed.
Hydroxide and chloride give other coordination environments
Hydroxide ions form pale-blue Cu(OH)₂(s) from aqueous Cu²⁺. In a sufficiently chloride-rich solution, ligand exchange can produce [CuCl₄]²⁻, changing the observed colour from the aqua-complex colour toward yellow; mixtures can look green.
Concentration matters: a few drops and a large excess of a ligand may give different observations. Use the equation and the stated conditions instead of naming a single colour without context.
Pause and explain
What can a pale-blue precipitate after adding a small amount of NH₃(aq) to Cu²⁺ indicate?
Worked example
Explain the change when excess aqueous ammonia is added to a pale-blue Cu²⁺ solution.
Show the worked solution
- A small amount may first give Cu(OH)₂ because ammonia is basic.
- Excess NH₃ can replace coordinated water around Cu²⁺.
- The deep-blue product is an ammine complex, still containing Cu(II).
Answer In excess NH₃, ligand exchange forms deep-blue [Cu(NH₃)₄(H₂O)₂]²⁺ from the pale-blue aqua complex. The copper remains in oxidation state +2; the colour changes because the ligand environment changes.
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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.
Describe the expected sequence when a small amount, then an excess, of NH₃(aq) is added to a copper(II) solution.
Give me a hint
Separate the initial base reaction from later ligand exchange.
Compare my reasoning
- A small amount gives pale-blue Cu(OH)₂ precipitate.
- Excess ammonia dissolves the precipitate.
- A deep-blue Cu(II) ammine complex forms.
A small addition produces pale-blue Cu(OH)₂(s); in excess NH₃ the precipitate dissolves to form deep-blue [Cu(NH₃)₄(H₂O)₂]²⁺.
Look for these in your work
- Small-addition observation stated
- Excess condition stated
- Final complex identified
Balance the precipitation equation for Cu²⁺(aq) reacting with OH⁻(aq) to form Cu(OH)₂(s).
Give me a hint
Use two hydroxide ions per Cu²⁺ and conserve charge.
Compare my reasoning
- Write one Cu²⁺ ion on the left.
- Add two OH⁻ ions.
- Check total charge is zero on both sides.
Cu²⁺(aq) + 2OH⁻(aq) → Cu(OH)₂(s).
Look for these in your work
- Hydroxide ratio correct
- Atoms balanced
- Charge balanced
A chloride-rich Cu(II) solution changes from pale blue toward yellow or green. Explain the inference and one limitation.
Give me a hint
Connect ligand exchange to a new complex, and account for mixed colours.
Compare my reasoning
- Chloride can replace water ligands to form [CuCl₄]²⁻.
- A new ligand environment changes the absorbed and observed colour.
- Mixtures and concentration can affect the apparent colour.
The colour is consistent with chloride ligand exchange to form [CuCl₄]²⁻. A green appearance can result from overlapping aqua- and chloro-complex colours, so colour alone is not a unique quantitative measure of composition.
Look for these in your work
- Complex identified
- Colour linked to ligand change
- Limitation stated
Common mistakes
- Calling the deep-blue product Cu(I)
- Ignoring the initial hydroxide precipitate when only a small amount of ammonia is added
- Treating hydroxide precipitation as ligand exchange to an aqueous complex
- Giving a colour without stating whether ligand is in excess
What can you explain now?
What is the central reason excess NH₃ changes the copper(II) complex?
Compare with the explanation
NH₃ replaces some coordinated water ligands
The ligand sphere changes while copper remains Cu(II).
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Make it stick
How do ligand identity and reagent amount affect the observed copper(II) species?
If you needed the solution, close it and solve the problem again from a blank page. A correct answer is stronger when you can explain the reason for each step.