Potassium Manganate(VII)

Summary: Potassium manganate(VII), KMnO₄, is a powerful purple oxidising agent in which manganese is in its highest possible oxidation state of +7; when acidified with dilute sulfuric acid it oxidises a wide range of reducing agents and is itself reduced to nearly colourless Mn²⁺ ions, providing a sharp visual endpoint that makes it invaluable for redox titrations and qualitative tests for reducing agents including iron(II) ions, hydrogen peroxide, sulfite ions, and sulfur dioxide. Tags: potassium-manganate-VII kmno4 oxidising-agent redox manganate igcse-chemistry titration test-for-reducing-agents Created: 2026-07-18

Potassium manganate(VII), KMnO₄, is a dark purple crystalline solid whose striking colour arises from the MnO₄⁻ ion, in which manganese is covalently bonded to four oxygen atoms and exists in the +7 oxidation state — the highest oxidation state manganese can attain and the source of the systematic name “manganate(VII)” as prescribed by IUPAC nomenclature, where the Roman numeral VII directly encodes the oxidation number of the central metal atom. In aqueous solution the MnO₄⁻ ion acts as a powerful Oxidising Agent, meaning it readily accepts electrons from other species (reducing agents) and is itself reduced, undergoing a dramatic and analytically useful colour change from intense purple (MnO₄⁻) to virtually colourless (Mn²⁺), provided the solution has been acidified with dilute sulfuric acid to supply the H⁺ ions necessary for the reduction half-equation MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O to proceed. This colour change serves as a self-indicating endpoint in redox titrations, most notably the determination of iron(II) concentration in which Fe²⁺ ions are oxidised to Fe³⁺, and it is also the chemical basis for the classic IGCSE qualitative test for reducing agents described in Section 6.4 of the syllabus: a few drops of acidified KMnO₄ solution added to a suspected reducing agent will turn from purple to colourless if a reducing agent is present. Beyond its role as a general redox probe, acidified potassium manganate(VII) is used specifically to test for sulfur dioxide gas (SO₂) and sulfite ions (SO₃²⁻) as detailed in Section 12 of the syllabus, where SO₂ reduces purple MnO₄⁻ to colourless Mn²⁺ while itself being oxidised to sulfate (SO₄²⁻), a reaction that also underpins its use in detecting SO₂ pollution and in distinguishing sulfite salts from sulfates in qualitative analysis. It is critical not to confuse potassium manganate(VII), KMnO₄ (purple, Mn in +7 state), with the chemically distinct potassium manganate(VI), K₂MnO₄ (green, Mn in +6 state), which is a different compound with different redox behaviour and is not part of the IGCSE syllabus; the two are related by a disproportionation reaction in alkaline solution but serve entirely different roles in practical chemistry.

Chemical Identity and Nomenclature

Formula and Structure

Potassium manganate(VII) has the chemical formula KMnO₄ and consists of K⁺ cations and tetrahedral MnO₄⁻ anions. Key structural facts:

• Ionic compound: K⁺ and MnO₄⁻ ions held together by ionic bonds in a crystalline lattice.
• MnO₄⁻ ion geometry: The four oxygen atoms are arranged tetrahedrally around the central manganese atom, with Mn—O covalent bonds exhibiting significant double-bond character due to pπ—dπ overlap between oxygen 2p orbitals and manganese 3d orbitals.
• Physical appearance: Dark purple (almost black) lustrous crystals that dissolve in water to give an intensely purple solution.
• Molar mass: 158.03 g mol⁻¹.

The Manganate(VII) Name

The systematic name “manganate(VII)” is mandated by IUPAC nomenclature rules for oxyanions containing a transition metal in a variable oxidation state. The Roman numeral in parentheses denotes the oxidation number of the central metal atom:

• The manganate ion is MnO₄ⁿ⁻ where n = 1 (one negative charge).
• Since each oxygen atom carries an oxidation number of −2 (four oxygens contribute −8), the manganese must have an oxidation number of +7 to give the ion its overall single negative charge: (+7) + 4(−2) = −1.
• Hence: manganate(VII).

The naming convention distinguishes potassium manganate(VII) from potassium manganate(VI), K₂MnO₄, which contains the MnO₄²⁻ ion with manganese in the +6 oxidation state and forms green solutions. The IGCSE syllabus exclusively uses the systematic name “potassium manganate(VII)” rather than the older common name “potassium permanganate,” though the latter is still widely used in general chemistry literature.

Physical Properties

The extraordinarily intense purple colour arises from a ligand-to-metal charge transfer (LMCT) transition in the MnO₄⁻ ion: an electron in an oxygen-centred molecular orbital absorbs visible light (primarily in the green-yellow region of the spectrum, around 520—570 nm) and is promoted to a manganese-centred orbital, transmitting the complementary purple colour. The intensity of this absorption is so great (molar absorptivity ε ≈ 2400 dm³ mol⁻¹ cm⁻¹) that even very dilute solutions (as low as 10⁻⁵ mol dm⁻³) remain visibly pink, enabling the KMnO₄ ion itself to serve as its own indicator in titrations — no separate indicator dye is required.

Oxidation State and the +7 Manganese Centre

Oxidation States of Manganese

Manganese is a transition metal and exhibits a remarkably wide range of oxidation states, from −3 to +7. For the IGCSE syllabus, the key oxidation states are:

The +7 oxidation state is the highest achievable by manganese and represents the formal loss of all seven valence electrons (4s² 3d⁵). Because the Mn⁷⁺ ion would be unrealistically small and highly polarising, it does not exist as a free cation; instead manganese covalently bonds to the four oxygen atoms, distributing the electron deficiency across the tetrahedral MnO₄⁻ framework. The +7 state is a powerful electron acceptor (oxidising agent) precisely because manganese can drop through multiple stable lower oxidation states — most commonly to +2 (Mn²⁺) — by gaining five electrons.

The Reduction Pathway: +7 → +2

Under acidic conditions, the MnO₄⁻ ion undergoes a five-electron reduction to the manganese(II) ion:

MnO₄⁻(aq) + 8H⁺(aq) + 5e⁻ → Mn²⁺(aq) + 4H₂O(l)

Purple → Very pale pink / nearly colourless

This half-equation is one of the most important for the IGCSE examination:

• Five electrons are gained by manganese: Mn(+7) gains 5e⁻ → Mn(+2).
• The 8H⁺ ions are essential — the reduction consumes protons, which is why acidification with dilute H₂SO₄ is required. Without acid, the reduction does not proceed to Mn²⁺ and instead stops at MnO₂ (a brown precipitate, +4 state).
• 4H₂O molecules are produced as water is formed from the combination of oxygen atoms released from MnO₄⁻ with protons from the acid.

Why Acidification Is Essential

In neutral or alkaline solution, the MnO₄⁻ ion undergoes a different reduction pathway:

MnO₄⁻(aq) + 2H₂O(l) + 3e⁻ → MnO₂(s) + 4OH⁻(aq)

Purple → Brown precipitate

This produces a brown precipitate of manganese(IV) oxide (MnO₂) rather than the clean, colourless Mn²⁺ solution. The brown precipitate obscures the colour change and makes endpoint detection unreliable. The half-equation also shows:

• Only 3 electrons are gained (not 5), so the oxidising power per mole of KMnO₄ is reduced in the absence of acid.
• OH⁻ ions are produced, making the solution alkaline — which further drives the formation of MnO₂.
• The brown MnO₂ precipitate can stain glassware and is difficult to remove.

For these reasons, potassium manganate(VII) used as an oxidising agent must be acidified, and the acid of choice in the IGCSE syllabus is dilute sulfuric acid (H₂SO₄). It is important to understand why dilute H₂SO₄ specifically:

This reasoning — specifically the unsuitability of HCl and HNO₃ — is a common examination question at IGCSE level.

Potassium Manganate(VII) as an Oxidising Agent

The Oxidising Agent Concept

An Oxidising Agent (also called an oxidant) is a substance that accepts electrons from another species, thereby oxidising that species while itself being reduced. The defining equation for MnO₄⁻ acting as an oxidising agent is:

• MnO₄⁻ gains electrons: MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O (reduction half-equation)
• The other species loses electrons: e.g., Fe²⁺ → Fe³⁺ + e⁻ (oxidation half-equation)

In any redox reaction involving acidified KMnO₄, the MnO₄⁻ ion is always the oxidising agent (it is itself reduced), and the substance being tested or titrated is the Reducing Agent (it is itself oxidised).

The Colour Change as a Redox Indicator

The defining practical feature of acidified potassium manganate(VII) is its colour change during a redox reaction:

Purple (MnO₄⁻, Mn in +7 state) → Colourless (Mn²⁺, Mn in +2 state)

This colour change is the direct visual indication that a redox reaction has occurred:

• The purple colour is due to the MnO₄⁻ ion.
• As the MnO₄⁻ ion is reduced to Mn²⁺, the purple colour fades and disappears.
• The Mn²⁺ ion is almost colourless in dilute aqueous solution (a very faint pink that is imperceptible at typical titration concentrations), so the solution becomes apparently colourless to the naked eye.

This self-indicating property is what makes KMnO₄ such a useful redox probe. In a titration:

• Before the endpoint: The solution remains colourless (because all added MnO₄⁻ is immediately reduced to colourless Mn²⁺ by the reducing agent in the flask).
• At the endpoint: The first drop of excess KMnO₄ that is not reduced remains as MnO₄⁻, imparting a permanent pale pink colour to the solution.
• After the endpoint: The pink intensifies as more KMnO₄ is added in excess.

The endpoint is exceptionally sharp — the colour change from colourless to pink occurs within a single drop of KMnO₄ solution when near the equivalence point.

Key Redox Reactions with KMnO₄

Oxidation of Iron(II) to Iron(III)

This is the most commonly examined redox titration at IGCSE level:

MnO₄⁻(aq) + 8H⁺(aq) + 5Fe²⁺(aq) → Mn²⁺(aq) + 4H₂O(l) + 5Fe³⁺(aq)

Half-equations:

• Reduction: MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O
• Oxidation: Fe²⁺ → Fe³⁺ + e⁻

Overall (balanced): 1 MnO₄⁻ reacts with 5 Fe²⁺.

Stoichiometry for calculation: moles of MnO₄⁻ = (1/5) × moles of Fe²⁺

The reaction is used both in redox titrations to determine the concentration of an iron(II) solution and in qualitative analysis to identify Fe²⁺ ions in a sample: a solution containing Fe²⁺ will decolourise acidified KMnO₄.

Oxidation of Hydrogen Peroxide (H₂O₂)

Hydrogen peroxide acts as a reducing agent toward the more powerful KMnO₄:

2MnO₄⁻(aq) + 6H⁺(aq) + 5H₂O₂(aq) → 2Mn²⁺(aq) + 8H₂O(l) + 5O₂(g)

Half-equations:

• Reduction: MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O
• Oxidation: H₂O₂ → O₂ + 2H⁺ + 2e⁻

This reaction produces bubbles of oxygen gas (effervescence) alongside the colour change from purple to colourless, which can serve as a confirmatory observation.

Oxidation of Potassium Iodide (I⁻ to I₂)

Acidified KMnO₄ oxidises iodide ions to iodine:

2MnO₄⁻(aq) + 16H⁺(aq) + 10I⁻(aq) → 2Mn²⁺(aq) + 8H₂O(l) + 5I₂(aq)

The reaction is evidenced by:

• The purple colour of MnO₄⁻ disappears (Mn²⁺ formed).
• Iodine (I₂) is produced, giving the solution a yellow-brown colour (which turns blue-black if starch indicator is subsequently added).

Oxidation of Ethanedioate (Oxalate) Ions, C₂O₄²⁻

2MnO₄⁻(aq) + 16H⁺(aq) + 5C₂O₄²⁻(aq) → 2Mn²⁺(aq) + 8H₂O(l) + 10CO₂(g)

This reaction is notable because it is autocatalytic — one of the products, Mn²⁺, catalyses the reaction. The first few drops of KMnO₄ decolourise slowly, but as Mn²⁺ builds up the subsequent drops decolourise almost instantaneously. The reaction must be heated to about 60—70°C to proceed at a practical rate, and it produces bubbles of CO₂ gas.

Testing for Reducing Agents

The General Test (IGCSE Syllabus Section 6.4)

The IGCSE syllabus specifies that acidified potassium manganate(VII) is used as a test for reducing agents and to identify redox reactions. The test procedure is:

1. Acidify the potassium manganate(VII) solution with dilute sulfuric acid.
2. Add a few drops of the acidified KMnO₄ solution to the sample being tested.
3. Observe the colour change.

Positive result (reducing agent present): The purple colour disappears; the solution turns colourless. A reducing agent has been oxidised by the MnO₄⁻ ion, which has itself been reduced to Mn²⁺.

Negative result (no reducing agent): The purple colour persists. No species capable of reducing MnO₄⁻ is present.

The logic chain is:

• Because MnO₄⁻ is an oxidising agent, it can only be decolourised (reduced) by a reducing agent.
• Therefore, decolourisation of acidified KMnO₄ is a positive test for the presence of a reducing agent.
• The reaction is a redox reaction because electrons are transferred: the reducing agent loses electrons (is oxidised) and the MnO₄⁻ gains electrons (is reduced).

Testing for Specific Reducing Agents

Testing for Iron(II) Ions (Fe²⁺)

Add a few drops of acidified KMnO₄ to the suspected Fe²⁺ solution. If Fe²⁺ is present, the purple colour disappears as Fe²⁺ is oxidised to Fe³⁺:

MnO₄⁻ + 8H⁺ + 5Fe²⁺ → Mn²⁺ + 4H₂O + 5Fe³⁺

This is often used in conjunction with the sodium hydroxide test: adding NaOH to the resulting solution will produce a brown precipitate of Fe(OH)₃ (iron(III) hydroxide), confirming that Fe²⁺ was originally present and has been oxidised.

Testing for Sulfite Ions (SO₃²⁻)

Acidified KMnO₄ is a key reagent for identifying sulfite ions in qualitative analysis. Sulfite ions are reducing agents:

2MnO₄⁻ + 6H⁺ + 5SO₃²⁻ → 2Mn²⁺ + 3H₂O + 5SO₄²⁻

• Purple → colourless.
• SO₃²⁻ (sulfur in +4 state) is oxidised to SO₄²⁻ (sulfur in +6 state).
• This distinguishes sulfites from sulfates, which do not decolourise KMnO₄ because sulfur is already in its highest (+6) oxidation state in SO₄²⁻ and cannot be further oxidised.

Testing for Sulfur Dioxide Gas (SO₂)

Sulfur dioxide gas, when bubbled through acidified KMnO₄ solution, reduces it:

2MnO₄⁻(aq) + 5SO₂(g) + 2H₂O(l) → 2Mn²⁺(aq) + 4H⁺(aq) + 5SO₄²⁻(aq)

• Purple → colourless.
• SO₂ (sulfur +4) is oxidised to SO₄²⁻ (sulfur +6).
• This reaction is used both as a qualitative test for SO₂ and as a method for removing SO₂ from gas streams in pollution control.

This test is specifically mentioned in IGCSE Syllabus Section 12.5 (Sulfur) alongside the potassium dichromate(VI) test for SO₂.

Comparison with Other Oxidising Agent Tests

The IGCSE syllabus also mentions potassium dichromate(VI) (K₂Cr₂O₇) as an alternative oxidising agent for testing purposes. The two are compared below:

Both tests operate on the same principle: the oxidising agent is reduced by the analyte, and the colour change signals the redox reaction.

Redox Titrations with Potassium Manganate(VII)

The Self-Indicating Titration

Potassium manganate(VII) is one of the most important reagents in redox titrimetry. Its unique advantage is that it acts as its own indicator:

• In the conical flask: The reducing agent (e.g., Fe²⁺ solution) is placed along with an excess of dilute sulfuric acid.
• In the burette: Standardised KMnO₄ solution (purple).
• During titration: Each drop of KMnO₄ is immediately reduced to colourless Mn²⁺ by the reducing agent. The solution in the flask remains colourless.
• At the endpoint: The reducing agent is fully consumed. The next drop of KMnO₄ added has nothing to reduce it, so the purple MnO₄⁻ persists, and the solution in the flask turns pale pink.

No separate indicator is required, which eliminates indicator-related errors and simplifies the procedure. This is in contrast to acid-base titrations, which require indicators such as phenolphthalein or methyl orange.

Iron(II) Determination: The Classic Titration

This is the most common KMnO₄ titration:

1. Prepare a standard solution of the iron(II) salt (e.g., ammonium iron(II) sulfate, (NH₄)₂Fe(SO₄)₂·6H₂O, known as Mohr’s salt, which is preferred over FeSO₄·7H₂O because it is less susceptible to air oxidation).
2. Add an excess of dilute H₂SO₄ to the iron(II) solution in the conical flask to provide the H⁺ ions needed for the MnO₄⁻ reduction.
3. Titrate with standard KMnO₄ solution from the burette.
4. Record the volume of KMnO₄ added when the first permanent pale pink colour appears.

Calculation: The stoichiometric ratio is 1 MnO₄⁻ : 5 Fe²⁺.

Sample calculation:

25.0 cm³ of Fe²⁺ solution required 18.4 cm³ of 0.0200 mol dm⁻³ KMnO₄ for complete reaction. Calculate the concentration of Fe²⁺.

• Moles of MnO₄⁻ used = (18.4 / 1000) × 0.0200 = 3.68 × 10⁻⁴ mol
• Moles of Fe²⁺ = 5 × 3.68 × 10⁻⁴ = 1.84 × 10⁻³ mol
• Concentration of Fe²⁺ = (1.84 × 10⁻³) / (25.0 / 1000) = 0.0736 mol dm⁻³

Sources of Error

Comparison: Manganate(VII) vs. Manganate(VI)

It is a common source of confusion among students that two similar-sounding potassium manganate compounds exist. The distinction is chemically important:

The two compounds are related by a disproportionation reaction. In alkaline solution, KMnO₄ can be partially reduced to K₂MnO₄:

4MnO₄⁻ + 4OH⁻ → 4MnO₄²⁻ + O₂ + 2H₂O

Conversely, acidifying a green manganate(VI) solution causes disproportionation:

3MnO₄²⁻ + 4H⁺ → 2MnO₄⁻ + MnO₂↓ + 2H₂O (Green) → (Purple) + (Brown precipitate)

These reactions are beyond IGCSE scope but are useful to understand why KMnO₄ must be handled in acidic conditions for titrations and why K₂MnO₄ is not used as an oxidising agent in the syllabus.

Laboratory Handling and Safety

Hazards

• Oxidising agent: KMnO₄ is a strong oxidiser and can cause fires or explosions if mixed with combustible materials (e.g., glycerol, ethanol, sugar, or finely divided metals). The reaction with glycerol is a classic exothermic demonstration: KMnO₄ spontaneously ignites glycerol after a delay of a few seconds.
• Staining: KMnO₄ stains skin and clothing brown (due to reduction to MnO₂). The stains can be removed with a solution of oxalic acid (ethanedioic acid) or sodium metabisulfite (a reducing agent).
• Toxicity: Ingestion is harmful and may cause severe gastrointestinal damage. Concentrated solutions are corrosive.
• Environmental: Manganese compounds are toxic to aquatic life; do not dispose down drains in large quantities.

Preparation and Storage

• KMnO₄ solutions should be stored in dark glass bottles because light catalyses the decomposition: 4MnO₄⁻ + 2H₂O → 4MnO₂ + 3O₂ + 4OH⁻.
• Solutions are not indefinitely stable — the concentration of a KMnO₄ solution decreases over time due to slow decomposition, which is why it must be standardised before use in quantitative analysis.
• KMnO₄ is not a primary standard (it cannot be obtained in a state of exact purity and it decomposes slowly), so freshly prepared solutions must be standardised against a primary standard such as sodium ethanedioate (Na₂C₂O₄) or ammonium iron(II) sulfate.

Related Pages

• Oxidising Agent — the general concept of electron-accepting species in redox chemistry
• Redox Reactions — the class of reactions in which KMnO₄ participates as the oxidant
• Reducing Agent — the counterpart species that reduces MnO₄⁻ and is itself oxidised
• Potassium Iodide — a reducing agent commonly tested against acidified KMnO₄
• Oxidation Number — the system that gives rise to the name “manganate(VII)”
• Contact Process — industrial process involving SO₂, which can be detected using KMnO₄
• Sulfur Dioxide — a specific reducing gas detected by the KMnO₄ test
• Titration — the quantitative technique in which KMnO₄ serves as both titrant and indicator
• Potassium Dichromate(VI) — the alternative oxidising agent test for reducing agents and SO₂
• Transition Metals — the broader context for manganese’s variable oxidation states

Sources

• Cambridge IGCSE Chemistry 0620 Syllabus, Section 6.4 Supplement: Redox; Section 12.5: Sulfur
• Harwood, R. & Lodge, I., Cambridge IGCSE Chemistry Coursebook, 5th Edition, Cambridge University Press, 2021, Chapter 7: Redox Reactions
• Gallagher, R. & Ingram, P., Complete Chemistry for Cambridge IGCSE, 3rd Edition, Oxford University Press, 2016, Chapter 8: Oxidation and Reduction
• Clegg, A. et al., Cambridge IGCSE Chemistry Study and Revision Guide, Hodder Education, 2017, Sections 6.4 and 12.5
• Cambridge Assessment International Education, IGCSE Chemistry 0620 Learner Guide, 2023—2025
• Greenwood, N. N. & Earnshaw, A., Chemistry of the Elements, 2nd Edition, Butterworth-Heinemann, 1997, Chapter 24: Manganese
• Cotton, F. A. & Wilkinson, G., Advanced Inorganic Chemistry, 6th Edition, Wiley, 1999, Chapter 21: The Transition Elements
• Vogel, A. I. & Jeffery, G. H., Vogel’s Textbook of Quantitative Chemical Analysis, 5th Edition, Longman, 1989, Chapter 10: Redox Titrations

Common Misconceptions