When KMnO4 is heated strongly at 513 K, the products formed are:

Potassium permanganate (KMnO4) is one of the most important and widely used oxidising agents in both laboratory and industrial chemistry. It is a purple-violet crystalline solid that dissolves readily in water to give intensely coloured purple-violet solutions due to charge-transfer absorption. The permanganate ion (MnO4-) has a tetrahedral structure with manganese in the +7 oxidation state at the centre, surrounded by four equivalent oxide ligands. In the +7 state, manganese has the electron configuration [Ar]3d0, giving no d-d transitions, so the vivid colour of permanganate arises from ligand-to-metal charge transfer (LMCT) transitions rather than d-d transitions. KMnO4 is prepared industrially by electrolytic oxidation of manganate (K2MnO4) in alkaline solution, or by fusing pyrolusite (MnO2) with KOH in the presence of air or other oxidants at high temperature, followed by dissolution and electrolytic oxidation of the K2MnO4 formed. In the laboratory, KMnO4 can be made by oxidising MnSO4 with PbO2 or periodate in acidic solution. KMnO4 is a versatile oxidant used in organic synthesis (Baeyer reagent), volumetric analysis (permanganometry), water treatment, dyeing, tanning, and as a topical antiseptic.

The oxidising power and the reduction products of KMnO4 depend critically on the acidity of the medium in which the reaction takes place, reflecting the different thermodynamic stabilities of manganese in different oxidation states as a function of pH. In acidic medium (dilute H2SO4): MnO4- is reduced to Mn2+ (colourless manganous ions): MnO4- + 8H+ + 5e- → Mn2+ + 4H2O; E° = +1.51 V. This is the strongest oxidising form of permanganate and is used in most volumetric (titrimetric) analysis applications, including titration of Fe2+, H2O2, C2O4^2- (oxalate), reducing sugars, and many other reducing analytes. In neutral or mildly alkaline medium: MnO4- is reduced to MnO2 (brown precipitate): MnO4- + 2H2O + 3e- → MnO2 + 4OH-; E° = +0.59 V. This form is less commonly used in volumetric analysis due to the turbid endpoint caused by MnO2 precipitation, but occurs in reactions with non-acidic substrates. In strongly alkaline medium: MnO4- is reduced to MnO4^2- (manganate, green): MnO4- + e- → MnO4^2-; E° = +0.56 V. Less commonly encountered in analytical chemistry but relevant in the thermal decomposition and disproportionation reactions of permanganate.

In organic synthesis, KMnO4 (Baeyer's reagent) is used as a powerful oxidising agent and the purple-to-brown colour change in Baeyer's test serves as a qualitative test for unsaturation (alkenes, alkynes) and easily oxidisable functional groups. Oxidation of alkenes: cold, dilute, alkaline KMnO4 (Baeyer's reagent): alkene → cis-diol (vicinal diol, two OH groups on adjacent carbons). This is a syn addition (both OH groups add to the same face of the alkene). Hot, concentrated acidic KMnO4: cleaves the C=C double bond; internal alkenes with at least one hydrogen on the double bond give carboxylic acids; terminal alkenes give CO2 (from the terminal =CH2 group) and a carboxylic acid from the internal carbon. Oxidation of alcohols: primary alcohols → carboxylic acids (KMnO4 in acidic medium is too strong to stop at aldehyde), secondary alcohols → ketones (ketones are not oxidised further by KMnO4). Oxidation of aldehydes: to carboxylic acids. Methylbenzene (toluene) and other side-chain-substituted arenes: the benzylic CH can be oxidised to -COOH. Oxidation of Fe2+ to Fe3+ in qualitative analysis.

Manganese (Mn, atomic number 25, electron configuration [Ar]3d5 4s2) is a notable transition metal because it displays an exceptionally wide range of oxidation states in stable compounds, from -3 to +7, though the most important are +2, +4, and +7. This breadth of chemistry makes manganese one of the most versatile transition metals. Mn(II): very stable configuration (half-filled d5 subshell, high crystal field stabilisation in the sextet configuration); forms a wide range of salts (MnSO4, MnCl2, MnO, MnCO3); MnO and MnCO3 are basic oxides. Mn(IV): MnO2 (pyrolusite, the most common manganese ore); amphoteric character; the key intermediate in many redox reactions; used in dry cells (Leclanché cell) as depolariser to oxidise H2 produced at cathode. Mn(VII): KMnO4 (most important Mn(VII) compound); very powerful oxidiser; decomposes on heating; the permanganate anion MnO4- is tetrahedral with intense purple colour from LMCT transitions. The ability of manganese to access this wide range of oxidation states is a consequence of the relatively small energy differences between successive d-orbital configurations and the availability of multiple accessible redox pathways.

The d-block elements (transition metals), located between the s-block and p-block in periods 4-7 of the periodic table, share several characteristic properties that arise from the progressive filling of d-orbitals. Variable oxidation states: most transition metals exhibit multiple oxidation states because d-electrons have relatively similar energies to the outer s electrons and can be removed or added without extreme energy cost. This gives rise to the rich redox chemistry of transition metals — their ability to catalyse electron-transfer reactions, form complexes with multiple metal-centred oxidation states, and participate in both oxidative and reductive chemistry. Coloured ions: transition metal ions are typically coloured because their partially filled d-orbitals allow d-d electronic transitions by absorption of visible light. The colour of a transition metal complex depends on the metal, its oxidation state, and the nature of the ligands (crystal field splitting). Catalytic activity: transition metals and their compounds are excellent catalysts, used in vast industrial processes (iron in Haber process, vanadium(V) oxide V2O5 in Contact process, platinum in Ostwald process, nickel in hydrogenation of oils, MnO2 in KClO3 decomposition). Magnetic properties: unpaired d-electrons make many transition metal compounds paramagnetic (attracted to magnetic fields). The number of unpaired electrons determines the magnetic moment: mu = sqrt(n(n+2)) BM where n = number of unpaired electrons.

Transition metals form coordination compounds because their d-orbitals, though partially filled, can accept electron pairs from ligands (Lewis bases) to form coordinate (dative) bonds. Werner's coordination theory (1893): the metal ion has two types of valency — primary valency (oxidation state, satisfied by anions) and secondary valency (coordination number, satisfied by ligands that may be anions, neutral molecules, or cations). In modern formulation: a coordination compound contains a central metal ion or atom bonded to a definite number of ligands by coordinate bonds, forming a coordination complex. The coordination sphere is typically enclosed in square brackets in formulas: [Co(NH3)6]Cl3 has the complex ion [Co(NH3)6]3+ as the coordination sphere and three Cl- ions outside. Key concepts: Coordination number (CN): number of ligand donor atoms bonded to the central metal. Common CN: 4 (square planar or tetrahedral), 6 (octahedral, most common). Effective Atomic Number (EAN) rule: stable complexes have the central metal with 18 electrons in its valence shell (8 from metal d-orbitals + 10 from ligands, or equivalent). Isomerism in coordination compounds: structural isomers (ionisation, linkage, coordination, hydrate) and stereoisomers (geometrical and optical).