The Many Colors of Permanganate
Difficulty: Easy–Medium | Time: 30–45 min | Visual Impact: Very High
Background
Potassium permanganate is one of chemistry’s most visually striking compounds. Its deep purple color comes from the permanganate ion (MnO₄⁻), in which manganese sits at its highest common oxidation state of +7. From there, it can be reduced stepwise through +6, +4, and +2 — and each state has a distinct, vivid color.
What makes this especially interesting is that the final product of reduction depends on the pH of the solution. Add acid, and permanganate goes all the way to pale, nearly colorless Mn²⁺. Use a neutral solution, and it stops at a deep brown MnO₂ precipitate. In strongly alkaline conditions, it pauses at green MnO₄²⁻ (manganate) before going further. The same chemical reaction, driven to different endpoints by controlling the environment — a clean demonstration of how conditions govern outcomes in chemistry.
The compound was first synthesized in 1659 and was widely used in the 19th century as a disinfectant and oxidizing bleach. Its dramatic purple color made it immediately recognizable, and it became the namesake of the “chameleon mineral” — a term that persists in the classic “chameleon reaction” demonstration today.
Colour Guide
| Sample | Colour | Species | Mn oxidation state | Conditions |
|---|---|---|---|---|
| 1 | Deep purple | MnO₄⁻ (permanganate) | +7 | Starting point, any pH |
| 2 | Green | MnO₄²⁻ (manganate) | +6 | Strong base only |
| 3 | Brown / black | MnO₂ (manganese dioxide) | +4 | Neutral to mild base |
| 4 | Colorless / pale pink | Mn²⁺ | +2 | Acid |
Materials
| What | How much |
|---|---|
| Potassium permanganate | ~0.5 g |
| Sodium hydroxide | 5–8 g |
| Dextrose (glucose) | 1–2 g |
| Ascorbic acid (vitamin C) | 0.5 g |
| Hydrochloric acid, dilute | 10 mL (or a splash of vinegar) |
| Water | ~400 mL total |
| 4 small glass jars or test tubes | |
| Safety glasses and gloves | Permanganate stains skin brown — hard to remove |
Substitutions: Citric acid or white vinegar can replace hydrochloric acid for Sample 4. Baking soda solution can replace sodium hydroxide for a weaker version of Sample 2, though the green state may not be stable.
Procedure
Prepare a stock solution first: dissolve a small pinch (~0.1 g) of potassium permanganate in 200 mL of water. The solution should be a clear, brilliant purple. Use this for all four samples. If it is too dark to see through, dilute further.
Wear gloves throughout. Permanganate stains skin and surfaces an intractable brown — treat every spill immediately with water, then a few drops of ascorbic acid solution to reduce the stain.
Sample 1 — Purple MnO₄⁻
Pour 40 mL of the purple stock solution into a jar. This is permanganate in its fully oxidized form. Set it aside as your color reference for all comparisons. Nothing else is needed.
Sample 2 — Green MnO₄²⁻ (Chameleon Reaction)
This sample demonstrates the fleeting +6 state — only stable in strongly alkaline conditions.
- Dissolve 5 g of sodium hydroxide in 80 mL of water in a sealable jar or bottle. Add 40 mL of the purple stock solution and stir. The solution turns very dark — nearly opaque purple-black.
- Add 0.5 g of dextrose to the jar and seal it tightly.
- Shake the bottle vigorously for a few seconds. The dissolved oxygen reacts with the dextrose, keeping the solution purple.
- Set it down and leave undisturbed for 3–10 minutes. Watch the color shift progressively: the purple will pass through blue and then settle into a deep olive-green as the permanganate is reduced one step to manganate (MnO₄²⁻).
- If you continue watching, the green will slowly turn brown as MnO₂ precipitates, and eventually the liquid above the precipitate may become pale or colorless.
Shaking restores oxygen to the solution, reversing the reaction briefly — the purple flashes back. Setting it down allows the glucose to reduce again. You can cycle this several times before the glucose is exhausted.
Sample 3 — Brown MnO₂
- Pour 40 mL of the purple stock solution into a jar.
- Add 0.2 g of ascorbic acid (or a small pinch of sodium carbonate to make the solution slightly alkaline). Stir.
- Within seconds, the purple fades and a dark brown precipitate forms throughout the solution: manganese dioxide (MnO₂). The liquid above may become yellow-brown or clear.
MnO₂ is a solid, insoluble oxide — in neutral or mildly alkaline conditions, this is as far as the reduction goes. Manganese dioxide is itself a useful oxidizing catalyst (it is what accelerates hydrogen peroxide decomposition in the classic catalase/raw potato demonstration).
Sample 4 — Colorless Mn²⁺
- Pour 40 mL of the purple stock solution into a jar.
- Add 5–10 mL of dilute hydrochloric acid (or enough white vinegar to make the solution clearly acidic). The purple color should remain unchanged for now.
- Add 0.3 g of ascorbic acid and stir. The purple rapidly fades to nearly colorless — a hint of pale pink at most. In acid, the reduction has enough driving force to proceed all the way to Mn²⁺, which has no strong light absorption in the visible range.
Hold Sample 1 and Sample 4 next to each other: same compound, one electron difference in history, entirely different appearance.
Reactions
The three reduction reactions, one for each pH regime:
In acid (Sample 4): \[\text{MnO}_4^- + 8\text{H}^+ + 5e^- \rightarrow \text{Mn}^{2+} + 4\text{H}_2\text{O}\]
Permanganate gains five electrons and requires eight protons from the acid. The large H⁺ requirement is why acid strongly favors this fully reduced endpoint.
In neutral/mild base (Sample 3): \[\text{MnO}_4^- + 2\text{H}_2\text{O} + 3e^- \rightarrow \text{MnO}_2 + 4\text{OH}^-\]
Only three electrons needed; the product is a solid oxide. Hydroxide is generated, which is why MnO₂ formation tends to push the solution more alkaline as it proceeds.
In strong base (Sample 2): \[\text{MnO}_4^- + e^- \rightarrow \text{MnO}_4^{2-}\]
Only one electron transferred. The green manganate ion is only stable in strongly alkaline conditions — in neutral or acid, it disproportionates immediately back to purple permanganate and brown MnO₂.
Why the Colors Differ
Each manganese oxidation state absorbs different wavelengths of visible light, depending on the electronic structure around the metal center. In MnO₄⁻, four oxygen atoms symmetrically surround manganese in a tetrahedral arrangement. The charge transfer between oxygen and manganese occurs at an energy corresponding to green light absorption — so the complement (red + blue = purple) is what we see.
As manganese is reduced and its oxidation state drops, the energy levels shift, the absorption wavelengths change, and different colors emerge. The +6 manganate (MnO₄²⁻) absorbs red light, giving green. MnO₂ is a wide-bandgap semiconductor that absorbs across most of the visible range, giving an opaque dark brown. Mn²⁺ in water absorbs only weakly in the near-ultraviolet, leaving the solution nearly transparent.
Tips and Notes
- Dilution matters: Concentrated permanganate solutions look black rather than purple. If your solution is opaque, dilute by half and try again.
- Green state is fast: The manganate (green) stage in Sample 2 may last only a few minutes before dropping to brown. Watch it closely.
- Temperature speeds things up: Warm solutions (40–50°C) reach the color transitions faster. Cold solutions slow everything down — useful if you want to observe an intermediate state for longer.
- Stain removal: Reduce any permanganate stains with a few drops of ascorbic acid solution (dissolve vitamin C in water). The brown residue from MnO₂ can be dissolved with a few drops of dilute acid.