`=>` These oxides are generally formed by the reaction of metals with oxygen at high temperatures. All the metals except scandium form `color{red}(MO)` oxides which are ionic.

`=>` The highest oxidation number in the oxides, coincides with the group number and is attained in `color{red}(Sc_2O_3)` to `color{red}(Mn_2O_7)`.

● Beyond group `7`, no higher oxides of iron above `color{red}(Fe_2O_3)` are known.

● Besides the oxides, the oxocations stabilise `color{red}(V^V)` as `color{red}(VO_2^+, V^(IV))` as `color{red}(VO_2^+)` and `color{red}(Ti^(IV))` as `color{red}(TiO^(2+))`.

`=>` As the oxidation number of a metal increases, ionic character decreases.

● In the case of `color{red}(Mn)`, `color{red}(Mn_2O_7)` is a covalent green oil. Even `color{red}(CrO_3)` and `color{red}(V_2O_5)` have low melting points.

● In these higher oxides, the acidic character is predominant.

● Thus, `color{red}(Mn_2O_7)` gives `color{red}(HMnO_4)` and `color{red}(CrO_3)` gives `color{red}(H_2CrO_4)` and `color{red}(H_2Cr_2O_7)`.

● `color{red}(V_2O_5)` is, however, amphoteric though mainly acidic and it gives `color{red}(VO_3^(4–))` as well as `color{red}(VO_2^+)` salts.

● In vanadium there is gradual change from the basic `color{red}(V_2O_3)` to less basic `color{red}(V_2O_4)` and to amphoteric `color{red}(V_2O_5)`.

● `color{red}(V_2O_4)` dissolves in acids to give `color{red}(VO^(2+))` salts.

● Similarly, `color{red}(V_2O_5)` reacts with alkalies as well as acids to give `color{red}(3 VO )`− and `color{red}(4 VO^+)` respectively.

● The well characterised `color{red}(CrO)` is basic but `color{red}(Cr_2O_3)` is amphoteric.

`=>` Potassium dichromate is a very important chemical used in leather industry and as an oxidant for preparation of many azo compounds.

`color{green}(text(Preparation ))` : Dichromates are generally prepared from chromate, which in turn are obtained by the fusion of chromite ore (`color{red}(FeCr_2O_4)`) with sodium or potassium carbonate in free access of air.

● The reaction with sodium carbonate occurs as follows :

`color{red}(4 FeCr_2O_4 + 8 Na_2CO_3 + 7 O_2 → 8 Na_2 CrO_4 + 2 Fe_2 O_3 + 8 CO_2)`

● The yellow solution of sodium chromate is filtered and acidified with sulphuric acid to give a solution from which orange sodium dichromate, `color{red}(Na_2Cr_2O_7. 2H_2O)` can be crystallised.

`color{red}(2Na_2CrO_4 + 2 H^(+) → Na_2Cr_2O_7 + 2 Na^(+) + H_2O)`

● Sodium dichromate is more soluble than potassium dichromate. The latter is therefore, prepared by treating the solution of sodium
dichromate with potassium chloride.

`color{red}(Na_2Cr_2O_7 + 2 KCl → K_2Cr_2O_7 + 2 NaCl)`

● Orange crystals of potassium dichromate crystallise out.

`=>` The chromates and dichromates are interconvertible in aqueous solution depending upon `pH` of the solution.

`=>` The oxidation state of chromium in chromate and dichromate is the same.

`color{red}(2 CrO_4^(2–) + 2H^(+) → Cr_2O_7^(2–) + H_2O)`

`color{red}(Cr_2O_7^(2–) + 2 OH^(-) → 2 CrO_4^(2–) + H_2O)`

`=>` The structures of chromate ion, `color{red}(CrO_4^(2-))` and the dichromate ion, `color{red}(Cr_2O_7^(2-))` are shown below.

● The chromate ion is tetrahedral whereas the dichromate ion consists of two tetrahedra sharing one corner with `color{red}(Cr–O–Cr)` bond angle of `126°`.

`color{green}(text(Properties ))` : Sodium and potassium dichromates are strong oxidising agents; the sodium salt has a greater solubility in water and is extensively used as an oxidising agent in organic chemistry.

`=>` Potassium dichromate is used as a primary standard in volumetric analysis.

● In acidic solution, its oxidising action can be represented as follows :

`color{red}(Cr_2O_7^(2–) + 14H^(+) + 6e → 2Cr^(3+) + 7H_2O \ \ \ \ (EV = 1.33V))`

● Thus, acidified potassium dichromate will oxidise iodides to iodine, sulphides to sulphur, tin(II) to tin(IV) and iron(II) salts to iron(III). The half-reactions are noted below :

`color{red}(6 I^(–) → 3I_2 + 6 e^(–)`; `3 Sn^(2+) → 3Sn^(4+) + 6 e^–)`

`color{red}(3 H_2 S → 6H^(+) + 3S + 6e^(–)`; `6 Fe^(2+) → 6Fe^(3+) + 6 e^–)`

● The full ionic equation may be obtained by adding the half-reaction for potassium dichromate to the half-reaction for the reducing agent, for e.g.,

`color{red}(Cr_2O_7^(2–) + 14 H^(+) + 6 Fe^(2+) → 2 Cr^(3+) + 6 Fe^(3+) + 7 H_2O)`

`color{green}(text(Preparation ))` : (i) Potassium permanganate is prepared by fusion of `color{red}(MnO_2)` with an alkali metal hydroxide and an oxidising agent like `color{red}(KNO_3)`.

● This produces the dark green `color{red}(K_2MnO_4)` which disproportionates in a neutral or acidic solution to give permanganate.

`color{red}(2MnO_2 + 4KOH + O_2 → 2K_2MnO_4 + 2H_2O)`

`color{red}(3MnO_4^(2–) + 4H^+ → 2MnO_4^(–) + MnO_2 + 2H_2O)`

(ii) Commercially, it is prepared by the alkaline oxidative fusion of `color{red}(MnO_2)` followed by the electrolytic oxidation of manganate (`Vl`).

`color{red}(MnO_2 overset(text{Fused with KOH , oxidised with air or } NNO_3)→undersettext(manganate ion)(MnO_4^(2-)))`;

`color{red}(undersettext(manganate)(MnO_4^(2-)) oversettext(Electrolytic oxidation in alkaline slution) →undersettext(permanganate ion)(MnO_4^(2-)))`

(iii) In the laboratory, a manganese (`II`) ion salt is oxidised by peroxodisulphate to permanganate.

`color{red}(2Mn^(2+) + 5S_2O_8^(2–) + 8H_2O → 2MnO_4^(–) + 10SO_4^(2–) + 16H^+)`

● Potassium permanganate forms dark purple (almost black) crystals which are isostructural with those of `color{red}(KClO_4)`.

● The salt is not very soluble in water (`6.4 g//100 g` of water at `293 K`), but when heated it decomposes at `513 K`.

`color{red}(2KMnO_4 → K_2MnO_4 + MnO_2 + O_2)`

`color{green}(text(Physical Properties ))` : (i) Its intense colour

(ii) It possesses diamagnetism along with weak temperature dependent paramagnetism

● These are explained by the use of molecular orbital theory.

`color{red}(text(Note ))` : The manganate and permanganate ions are tetrahedral; the green manganate is paramagnetic with one unpaired electron but the permanganate is diamagnetic.

The `color{red}(π)`-bonding takes place by overlap of `color{red}(p)` orbitals of oxygen with `color{red}(d)` orbitals of manganese.

`color{green}(text(Chemical properties ))` : Acidified permanganate solution oxidises oxalates to carbon dioxide, iron(II) to iron(III), nitrites to nitrates and iodides to free iodine. The half-reactions of reductants are :

`color{red}(5 Fe^(2+) → 5 Fe^(3+) + 5e^–)`

`color{red}(5NO_2^– + 5H_2O → 5NO_3^(–) + 10H^(+) + 10e^–)`


`color{red}(10I^– → 5I_2 + 10e^–)`


● The full reaction can be written by adding the half-reaction for `color{red}(KMnO_4)` to the half-reaction of the reducing agent, balancing wherever necessary.

● If we represent the reduction of permanganate to manganate, manganese dioxide and manganese(II) salt by half-reactions,

`color{red}(MnO_4^(–) + e^(–) → MnO_4^(2–) \ \ \ \ \ \ \ (E = + 0.56 V))`

`color{red}(MnO_4^(-) + 4H^(+) + 3e^(–) → MnO_2 + 2H_2O \ \ \ \ \ \ (E = + 1.69 V))`

`color{red}(MnO_4^(–) + 8H^(+) + 5e^(–) → Mn^(2+) + 4H_2O \ \ \ \ \ \ \ \ \ (E = + 1.52 V))`

● We see that the hydrogen ion concentration of the solution plays an important part in influencing the reaction.

● Although many reactions can be understood by consideration of redox potential, kinetics of the reaction is also an important factor.

● Permanganate at `color{red}([H^+] = 1)` should oxidise water but in practice the reaction is extremely slow unless either manganese(`ll`) ions are present or the temperature is raised.

A few important oxidising reactions of `color{red}(KMnO_4)` are given below :

1. `color{green}(text(In Acid Solutions ))` :

(a) Iodine is liberated from potassium iodide :

`color{red}(10I^(–) + 2MnO_4^(–) + 16H^(+) → 2Mn^(2+) + 8H_2O + 5I_2)`

(b) `color{red}(Fe^(2+))` ion (green) is converted to `color{red}(Fe^(3+))` (yellow) :

`color{red}(5Fe^(2+) + MnO_4^(–) + 8H^(+) → Mn^(2+) + 4H_2O + 5Fe^(3+))`

(c) Oxalate ion or oxalic acid is oxidised at `333 K` :

`color{red}(5C_2O_4^(2–) + 2MnO_4^(–) + 16H^(+) → 2Mn^(2+) + 8H_2O + 10CO_2)`

(d) Hydrogen sulphide is oxidised, sulphur being precipitated :

`color{red}(H_2S → 2H^(+) + S^(2–))`

`color{red}(5S^(2–) + 2MnO_4^(-) + 16H^(+) → 2Mn^(2+) + 8H_2O + 5S)`

(e) Sulphurous acid or sulphite is oxidised to a sulphate or sulphuric acid :

`color{red}(5SO_3^(2–) + 2MnO_4^(–) + 6H^(+) → 2Mn^(2+) + 3H_2O + 5SO_4^(2–))`

(f) Nitrite is oxidised to nitrate :

`color{red}(5NO_2^(–) + 2MnO_4^(–) + 6H^(+) → 2Mn^(2+) + 5NO_3^(–) + 3H_2O)`

2. `color{green}(text(In Neutral or Faintly Alkaline Solutions ))` :

(a) A notable reaction is the oxidation of iodide to iodate :

`color{red}(2MnO_4^(–) + H_2O + I^– →2MnO_2 + 2OH^(–) + IO_3^–)`

(b) Thiosulphate is oxidised almost quantitatively to sulphate :

`color{red}(8MnO_4^(–) + 3S_2O_3^(2–) + H_2O → 8MnO_2 + 6SO_4^(2–) + 2OH^–)`

(c) Manganous salt is oxidised to `color{red}(MnO_2)`; the presence of zinc sulphate or zinc oxide catalyses the oxidation :

`color{red}(2MnO_4^(–) + 3Mn^(2+) + 2H_2O → 5MnO_2 + 4H^+)`

`color{red}(text(Note ))` : Permanganate titrations in presence of hydrochloric acid are unsatisfactory since hydrochloric acid is oxidised to chlorine.