The Many Colors of Copper
Summary
A single element gives sky-blue, teal, powder-blue, malachite-green, white, brick-red and black in an afternoon, because the colour of a copper complex1 depends on which ligands2 surround the metal and on its oxidation state. You’ll be able to explain how ligand strength shifts the colour, why anhydrous copper sulfate is white, and how glucose reduces blue Cu²⁺ to brick-red Cu⁺ (Fehling’s test).
History
No metal has left a more colorful mark on human history than copper. The brilliant blue of Egyptian faience beads, the vivid green of malachite used as eye paint in ancient Kohl, the turquoise patina of bronze statues—all are copper compounds. Medieval painters ground malachite and azurite (both natural copper carbonates) into pigments. Verdigris, made by exposing copper to vinegar fumes, was the brightest green available to European artists for centuries. The Statue of Liberty’s famous blue-green skin is nothing more than copper reacting slowly with air, rain, and sea salt.
In 1848, Hermann von Fehling3 developed his glucose test around copper chemistry: an alkaline copper solution turns brick-red when it reacts with glucose. Generations of chemists and doctors used Fehling’s test to diagnose diabetes before modern blood-sugar meters existed.
All these colors come from the same element. This experiment recreates eight of them in an afternoon.
Hazards & preparation
PPE: safety glasses and gloves; an adult supervises the heating steps.
- Copper sulfate and copper salts are toxic if swallowed and harmful to aquatic life — keep off skin and out of drains.
- Heat with care (Samples 6–8): hot foil, candle/hotplate — work on a heatproof surface, tie back hair.
Disposal: copper must not go down the drain — collect solutions and precipitates, let solids settle, bin them (or hold for hazmat), and dilute the remaining liquid heavily. See the Safety page.
Materials
| What | How much |
|---|---|
| Copper sulfate (CuSO₄·5H₂O) | 15 g |
| Cupric chloride (CuCl₂) | 3 g |
| Sodium carbonate | 4 g |
| Dextrose (glucose) | 3 g |
| Copper wire, coins, or scraps | a small piece |
| Acetic acid / white vinegar | 30 mL |
| Water | ~300 mL |
| Test tubes or small glass jars | 6–8 |
| Hotplate or candle | for heating |
| Aluminum foil | small sheet |
| Safety glasses and gloves |
Colour Guide
Here is what you are building toward:
| # | Colour | Compound | Copper oxidation state |
|---|---|---|---|
| 1 | Reddish-orange | Cu metal | Cu⁰ |
| 2 | Sky blue | CuSO₄ solution | Cu²⁺ with water ligands |
| 3 | Blue-green | CuCl₂ solution | Cu²⁺ with chloride ligands |
| 4 | Powder blue | Cu(OH)₂ precipitate | Cu²⁺ |
| 5 | Vivid green | Cu₂(OH)₂CO₃ (synthetic malachite) | Cu²⁺ |
| 6 | White | CuSO₄ anhydrous | Cu²⁺ (no water) |
| 7 | Brick red | Cu₂O (cuprous oxide) | Cu⁺ |
| 8 | Black | CuO (cupric oxide) | Cu²⁺ |
Procedure
Work through the steps in order. Line up your samples as you go—by the end you will have a full colour series on the bench.
Sample 1 — Reddish-orange metal
- Hold a copper coin or a short length of copper wire in your hand. This is Cu⁰. The reddish-orange metallic lustre is copper’s ground state. Set it aside as your reference.
Sample 2 — Sky blue sulfate solution
- Dissolve 5 g of copper sulfate crystals in 50 mL of water. The vivid sky blue comes from Cu²⁺ ions surrounded by six water molecules. Pour into a test tube and label it.
Sample 3 — Blue-green chloride solution
- Dissolve 3 g of cupric chloride in 30 mL of water. Notice the distinctly different blue-green (teal) colour compared to Sample 2—same Cu²⁺ ion, different ligands, different colour. Place beside Sample 2 and compare.
Sample 4 — Powder blue hydroxide
- Take 20 mL of your copper sulfate solution (Sample 2). Add sodium carbonate solution (1 g in 10 mL water) drop by drop with stirring. A powder-blue precipitate forms immediately—this is copper hydroxide, Cu(OH)₂. Stop when roughly half the blue has precipitated. Label and set aside.
Sample 5 — Vivid green malachite
- To a fresh 20 mL portion of copper sulfate solution, add the sodium carbonate solution more generously—2 g in 15 mL water, all at once. The precipitate is now distinctly greener: basic copper carbonate, Cu₂(OH)₂CO₃, the same mineral as malachite. Compare it directly with the powder-blue Sample 4.
Sample 6 — White anhydrous sulfate
- Place a small heap of copper sulfate crystals (about 2 g) on a piece of aluminum foil. Heat gently over a hotplate or candle, moving constantly. The crystals lose their water of crystallisation and turn white or very pale grey. Once cool, add a drop or two of water and watch the blue return—this is one of chemistry’s most satisfying reversible reactions.
Sample 7 — Brick red cuprous oxide (Fehling’s)
- In a test tube, mix 5 mL of copper sulfate solution with 1 g sodium carbonate (powder-blue Cu(OH)₂ forms). Add 1 g of dextrose (glucose). Heat the test tube gently in a water bath at ~80 °C, or hold it carefully over a candle. Within 2–5 minutes, a brick-red precipitate appears—cuprous oxide, Cu₂O. This is Fehling’s test: glucose donates electrons to Cu²⁺, reducing it to Cu⁺. The medical laboratory test for sugar in urine uses exactly this reaction.
Sample 8 — Black copper oxide
- Hold a piece of copper wire or coin in a candle flame or over a hotplate for 30–60 seconds. The surface turns black: Cu²⁺ in CuO. This is the same reaction used to blacken copper jewellery and to make black ceramic glazes. Let it cool—if you then dip it briefly in the copper sulfate solution, the black oxide dissolves back and the copper surface reappears.
What you should see
By the end the bench holds a full spectrum from one element: reddish metal, the sky-blue and teal solutions side by side (same ion, different ligand), powder-blue and grass-green precipitates, a heap of white dehydrated crystals that flush blue again with one drop of water, a brick-red Fehling’s precipitate, and a coin blackened in the flame.
| Symptom | Likely cause | Fix |
|---|---|---|
| Samples 2 and 3 look the same | Too dilute to see the hue difference | Use more salt; compare against a white card |
| Sample 6 won’t go fully white | Under-heated / re-absorbing water | Heat longer, keep moving; add water only once cool |
| No brick-red in Fehling’s (Sample 7) | Not hot enough, or no alkali | Ensure it’s alkaline (carbonate) and heat to ~80 °C for several minutes |
| Precipitates 4 and 5 look identical | Carbonate amount similar | Use a little carbonate for blue Cu(OH)₂, excess for green malachite |
The reactions
Samples 4 & 5 — Precipitation
\[\ce{CuSO4(aq) + Na2CO3(aq) + H2O -> Cu(OH)2(s) + Na2SO4(aq) + CO2(g)}\]
With excess carbonate, the hydroxide converts to basic copper carbonate (synthetic malachite):
\[\ce{2Cu(OH)2(s) + CO2(g) -> Cu2(OH)2CO3(s) + H2O}\]
Sample 6 — Reversible dehydration
\[\ce{CuSO4 * 5H2O(s) <=>[\Delta][\text{+ H}_2\text{O}] CuSO4(s) + 5H2O(g)}\]
Sample 7 — Fehling’s reduction
\[\ce{2Cu^{2+}(aq) + C6H12O6(aq) + 4OH^-(aq) ->[\Delta] Cu2O(s) + C6H12O7(aq) + 2H2O(l)}\]
Copper(II) is reduced to copper(I); glucose is oxidised to gluconic acid.
Sample 8 — Oxidation
\[\ce{2Cu(s) + O2(g) ->[\Delta] 2CuO(s)}\]
the colour of each copper species is set by its ligands and oxidation state; glucose reduces blue Cu²⁺ to brick-red Cu⁺, and heat oxidizes copper metal to black CuO.
The Science
Copper is a transition metal, meaning its electrons fill the d orbitals—the same orbitals responsible for the brilliant colours of most transition metal compounds.
The key idea is that Cu²⁺ absorbs visible light to shuffle electrons between d-orbital energy levels. Which wavelength gets absorbed—and therefore what colour you see—depends on exactly which molecules or ions are bonded to the copper atom. These bonded species are called ligands.
Different ligands split the d-orbital energies by different amounts. Chloride ions (Cl⁻) are weak-field ligands: they cause a small energy split, so the copper absorbs red-orange light and transmits blue-green → teal. Water molecules cause a medium split, absorbing orange light and transmitting blue → sky blue. Hydroxide and carbonate are stronger still, shifting the absorption toward shorter wavelengths → paler blue and green.
This explains the table:
| Ligand | Field strength | Absorption | Transmitted colour |
|---|---|---|---|
| Cl⁻ | Weak | Red-orange | Blue-green (teal) |
| H₂O | Medium | Orange | Sky blue |
| OH⁻ | Stronger | Yellow-orange | Pale blue |
| CO₃²⁻ | Stronger still | Yellow | Green |
The white anhydrous sulfate seems to break the pattern—but it fits perfectly: with no water ligands, the d-orbital geometry collapses and the d-d transition shifts into the ultraviolet, out of visible range altogether. No visible absorption → white.
Cu⁺ (brick red Cu₂O) and Cu⁰ (metal) follow different rules. Cu⁺ is a d¹⁰ ion with fully filled d orbitals—no d-d transitions are possible, so Cu₂O’s red colour comes from a different mechanism (charge-transfer between O²⁻ and Cu⁺). Cu⁰ has delocalised metallic electrons that absorb most light except orange-red wavelengths.
Questions to Explore
Why do copper sulfate and copper chloride solutions look different? Both contain Cu²⁺, and in dilute solution both are surrounded mainly by water — yet concentrated CuCl₂ is a distinct teal. So what role does the anion play?
Hint / answer
In concentrated chloride solution, chloride ions actually do get in close and partly replace water as ligands on the copper, and chloride splits the d-orbitals differently than water — shifting the colour toward teal. It’s not the “spectator” sulfate vs. chloride in bulk, but whether chloride coordinates the metal.
What is happening at the quantum level? Copper absorbs light by electrons jumping between d-orbital levels. What sets the gap, and why does swapping water for hydroxide change the colour?
Hint / answer
The surrounding ligands split the copper’s d-orbitals into higher and lower sets; the size of that split fixes the energy (colour) of light absorbed. Stronger-field ligands like hydroxide split them more, so a different wavelength is absorbed and the transmitted colour shifts.
Why does Cu⁺ not follow the same rules? Cu₂O (brick red) has a full d shell — no d–d transitions possible. What makes it red?
Hint / answer
With completely filled d-orbitals there’s no room for a d–d jump, so the colour can’t come from that. Instead it’s charge transfer — an electron shifting between oxygen and copper — that absorbs part of the visible spectrum and leaves Cu₂O looking red.
Can you reverse the dehydration? Heating to white just removes water; one drop restores blue. What does that say about the change?
Hint / answer
It’s a reversible physical change — only water of crystallization is coming and going, not any breaking of the copper–sulfate bonding. Because nothing is destroyed, adding water lets the blue hydrated complex re-form instantly.
Why is the Statue of Liberty green, not blue? Its patina is a mix of copper carbonates, sulfates, and chlorides. Given ligands and colour, why green?
Hint / answer
The patina is dominated by basic copper carbonates (malachite-type) and related sulfates/chlorides — solid compounds with hydroxide and carbonate ligands, which (as in Sample 5) give green. The bright blue you see in solution needs water ligands, which the dry surface film doesn’t have.
Going further
- A ninth colour. Add ammonia to Sample 2: a pale precipitate forms, then dissolves into a deep royal blue-violet [Cu(NH₃)₄]²⁺ — ammonia is a stronger-field ligand than water.
- Make verdigris. Suspend a copper coin over (not touching) vinegar in a sealed jar for a day or two to grow blue-green copper-acetate crystals, the medieval pigment.
- Next in the Redox track: watch one metal span five oxidation states in The Many Colors of Permanganate.
Footnotes
Complex ion — A central metal ion surrounded by bound molecules or ions (ligands), often intensely coloured.↩︎
Ligand — A molecule or ion that bonds to a central metal atom to form a complex ion.↩︎
Hermann von Fehling — German chemist (1811–1885) whose solution tests for reducing sugars.↩︎