The Many Colors of Permanganate

A single compound spans five oxidation states — track the color changes as pH and reducing agents drive manganese from deep purple to green, brown, and colorless
Intermediate🕐45 min▲Moderate hazardredox

Summary

Manganese steps down through several oxidation states — each a different colour — as you reduce1 purple permanganate, and which colour it stops at depends on the pH2. You’ll be able to explain how proton availability controls the endpoint, what disproportionation means, and why these colours are so intense.

History

Potassium permanganate’s colour play has fascinated chemists since Johann Rudolf Glauber3 first made it in 1659 by fusing the mineral pyrolusite (MnO₂) with potash. The solution shifted from green to blue to purple as it stood, and the compound earned the nickname mineral chameleon — the very same colour sequence you’ll reproduce in the “chameleon” jar here.

By the 19th century it had found a practical life as Condy’s fluid, a permanganate disinfectant and deodoriser sold by Henry Bollmann Condy4 in the 1850s, and as an oxidising bleach. Its intense purple — visible even in vanishingly dilute solution — made it instantly recognisable, and it remains a workhorse oxidant in water treatment and first-aid kits 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

Hazards & preparation

Warning

PPE: safety glasses and gloves throughout.

  • Potassium permanganate is a strong oxidizer and stains skin and surfaces an intractable brown — treat spills at once with water, then a few drops of vitamin-C solution to reduce the stain.
  • Sodium hydroxide and dilute hydrochloric acid are corrosive — handle carefully and never let permanganate meet concentrated acid.

Disposal: reduce leftover permanganate with vitamin C until brown/colourless, let any MnO₂ solid settle and bin it, then flush the neutralized liquid with water. See the Safety page.

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

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. It should be a clear, brilliant purple you can see through. If it looks black, dilute further.

Sample 1 — Purple MnO₄⁻: pour 40 mL of the purple stock into a jar. This is your fully-oxidized reference. Nothing else needed.

Sample 2 — Green MnO₄²⁻ (chameleon):

  1. Dissolve 5 g sodium hydroxide in 80 mL water in a sealable bottle. Add 40 mL purple stock and stir — it turns nearly opaque purple-black.
  2. Add 0.5 g dextrose and seal tightly.
  3. Shake vigorously for a few seconds — dissolved oxygen keeps it purple.
  4. Set it down, undisturbed, 3–10 minutes. The colour shifts purple → blue → deep olive-green as permanganate is reduced one step to manganate (MnO₄²⁻).
  5. Watch longer and the green turns brown as MnO₂ precipitates. Shaking re-adds oxygen and flashes it back to purple — cycle several times until the glucose is gone.

Sample 3 — Brown MnO₂:

  1. Pour 40 mL purple stock into a jar.
  2. Add 0.2 g vitamin C and stir (a pinch of sodium carbonate can be added to keep the solution slightly alkaline, but the vitamin C is what reduces the permanganate).
  3. Within seconds the purple fades and a dark brown precipitate of manganese dioxide forms throughout.

Sample 4 — Colorless Mn²⁺:

  1. Pour 40 mL purple stock into a jar.
  2. Add 5–10 mL dilute hydrochloric acid (or enough vinegar to make it clearly acidic) — still purple for now.
  3. Add 0.3 g vitamin C and stir. The purple rapidly fades to nearly colourless (a hint of pale pink at most) — in acid the reduction goes all the way to Mn²⁺.

Hold Sample 1 and Sample 4 side by side: the same compound, entirely different appearance.

What you should see

The star is the “chameleon” jar: shaken, a rich royal purple; left standing, it slides through indigo to a murky olive-green, then to muddy brown. Sample 3 drops a dark chocolate-brown cloud within seconds, and Sample 4 simply switches off to water-clear (barely pink) — a striking contrast to the vivid purple beside it.

Symptom Likely cause Fix
Stock looks black, not purple Too concentrated Dilute by half until you can see through it
Green never appears (Sample 2) Not alkaline enough, or reduced too far Use more NaOH; watch early — green is brief
Green vanishes almost instantly Warm / over-reduced Chill the solution; use a touch less glucose
Sample 4 stays purple Not acidic enough Add more acid before the vitamin C

The reactions

In acid (Sample 4): \[\ce{MnO4^- + 8H+ + 5e^- -> Mn^2+ + 4H2O}\]

Permanganate gains five electrons and needs eight protons — the large H⁺ requirement is why acid drives it to the fully reduced Mn²⁺.

In neutral/mild base (Sample 3): \[\ce{MnO4^- + 2H2O + 3e^- -> MnO2 + 4OH^-}\]

Only three electrons; the product is a solid oxide, and hydroxide is generated.

In strong base (Sample 2): \[\ce{MnO4^- + e^- -> MnO4^2-}\]

Only one electron. Green manganate is stable only in strong base; elsewhere it disproportionates back to purple permanganate and brown MnO₂.

permanganate is reduced step by step (+7 → +6 → +4 → +2), and pH sets how far it goes by controlling the supply of protons.

The Science

Each manganese oxidation state absorbs different wavelengths of visible light. In MnO₄⁻, four oxygens symmetrically surround manganese; a charge transfer between oxygen and manganese absorbs green light, so the complement (purple) is what we see. As manganese is reduced, the energy levels shift: +6 manganate absorbs red, giving green; MnO₂ is a wide-bandgap semiconductor absorbing across the visible, giving opaque brown; and Mn²⁺ absorbs only weakly in the near-UV, leaving the solution nearly transparent.

Questions to Explore

  1. Why does pH control which state the reduction stops at? Each equation needs a different number of protons. Why does proton availability set how far the reduction goes?

    Hint / answer

    Going all the way to Mn²⁺ consumes eight H⁺ per permanganate, so it can only happen where protons are plentiful — in acid. Without those protons (neutral or basic), the reaction can only reach a proton-light endpoint like MnO₂ or manganate. The equations literally tell you the pH each needs.

  2. Why is green manganate unstable at neutral pH? It disproportionates to purple + brown. What is disproportionation, and why does a compound split itself into more- and less-oxidized forms?

    Hint / answer

    Disproportionation is when a species in a middle oxidation state converts partly up and partly down at once. Manganate (+6) is only marginally stable, so outside strong base it’s energetically favourable for some to be oxidized to +7 (purple) and the rest reduced to +4 (brown) — two more stable states.

  3. What causes each colour? The intense permanganate and manganate colours come from charge-transfer, not d–d, transitions. What’s the difference, and why is charge transfer more intense?

    Hint / answer

    A d–d transition just shuffles an electron between the metal’s own d-orbitals and is quantum-mechanically weak (pale colours). Charge transfer shifts an electron between oxygen and manganese — a fully allowed transition that absorbs light strongly, so even dilute permanganate is vividly coloured.

  4. Why is permanganate such a strong oxidizer? At +7 it can accept five electrons per ion in acid. How does comparing that to one-electron oxidants explain its power?

    Hint / answer

    Sitting at manganese’s highest common state, permanganate is very “hungry” to be reduced, and in acid each ion can grab five electrons — a large capacity to oxidize other substances. Many common oxidants shuffle only one electron, so permanganate does far more oxidizing per ion.

  5. Where is permanganate used outside the lab? Water treatment, antiseptics, industrial oxidation. Given the pH dependence, what conditions get a specific product?

    Hint / answer

    In water treatment the goal is usually the harmless brown MnO₂ solid that can be filtered out, so plants run near neutral pH where MnO₂ is the endpoint. Choosing pH (and dose) lets operators land on the product — Mn²⁺, MnO₂, or manganate — that suits the job.

Going further

  • Titrate with it. Replace the vitamin C in Sample 4 with measured oxalic acid in dilute acid — the endpoint is a sharp permanent pale pink: 2 MnO₄⁻ + 5 C₂H₂O₄ + 6 H⁺ → 2 Mn²⁺ + 10 CO₂ + 8 H₂O.
  • Slow the chameleon. Chill Sample 2 to hold the green manganate stage longer for a good look at the +6 state.
  • Next in the Redox track: put redox to work dyeing cloth in Indigo Vat Dyeing.

Footnotes

  1. Reduction — The gain of electrons by an atom, ion, or molecule; always paired with oxidation.↩︎

  2. pH — A 0–14 scale measuring how acidic (low) or basic (high) a solution is; 7 is neutral.↩︎

  3. Johann Rudolf Glauber — German-Dutch chemist (1604–1670) who first made sodium sulfate (‘Glauber’s salt’) and potassium permanganate.↩︎

  4. Henry Bollmann Condy — English industrial chemist who marketed ‘Condy’s fluid’, a permanganate disinfectant, in the 1850s.↩︎