Important compounds of transition elements

What are the important compounds of the transition elements?

In plain words: the higher the metal’s formal oxidation state, the more it holds on to its oxygens with shared, covalent bonds, and the more its oxide behaves like a non-metal oxide, an acid. Chromium(VI) and manganese(VII) sit at that top end. A metal in so high a state readily takes electrons back, which is why dichromate and permanganate are strong oxidising agents (reasoned).

The oxidation numbers on this page follow the Class 11 rules.

Oxides and oxoanions of metals

• Formation. The oxides are usually made by heating the metal with oxygen. Every metal of the first series except scandium forms an ionic oxide MO (NCERT Section 4.4.1, page 105, repeating page 100).
• Highest oxides. The highest oxidation number in the oxides equals the group number from Sc₂O₃ to Mn₂O₇; beyond group 7, the ordinary textbook series ends with Fe₂O₃ for iron; this is not a catalogue of phases under extreme pressure. Oxocations also stabilise high states: V(V) as VO₂⁺, V(IV) as VO²⁺ and Ti(IV) as TiO²⁺ (NCERT page 105).
• From ionic to covalent, basic to acidic. Mn₂O₇ is a covalent green oil, and even CrO₃ and V₂O₅ have low melting points. In these higher oxides the acidic character dominates: Mn₂O₇ gives HMnO₄, and CrO₃ gives H₂CrO₄ and H₂Cr₂O₇ (NCERT page 105).

Character as NCERT gives it (page 105); oxidation states from the formulas.

• Vanadium’s oxides show the change step by step. V₂O₄ dissolves in acids to give VO²⁺ salts, and V₂O₅, amphoteric though mainly acidic, gives vanadates VO₄³⁻ with alkalis and VO₂⁺ salts with acids (NCERT page 105).
• NCERT slip (page 105): the last vanadium sentence prints “VO₄³⁻ and VO₄⁺”. VO₄⁺ would need vanadium at +9 (x − 8 = +1), more than its five valence electrons allow. The cation is VO₂⁺ (x − 4 = +1, so +5), as NCERT itself writes three sentences earlier (“it gives VO₄³⁻ as well as VO₂⁺ salts”), at the start of the section (V^V as VO₂⁺) and on page 100. In an exam, write VO₂⁺.

Potassium dichromate: preparation, properties and uses

Uses. Potassium dichromate is used in the leather industry and as an oxidant for making many azo compounds. It is a primary standard in volumetric analysis; sodium dichromate, which is more soluble in water, is the one used widely as an oxidising agent in organic chemistry (NCERT pages 105–106).

Preparation from chromite ore, FeCr₂O₄ (NCERT pages 105–106):

1. Fuse the ore with sodium (or potassium) carbonate in free access of air, giving yellow sodium chromate: 4FeCr₂O₄ + 8Na₂CO₃ + 7O₂ → 8Na₂CrO₄ + 2Fe₂O₃ + 8CO₂
2. Filter, and acidify the yellow solution with sulphuric acid; orange sodium dichromate, Na₂Cr₂O₇·2H₂O, crystallises: 2Na₂CrO₄ + 2H⁺ → Na₂Cr₂O₇ + 2Na⁺ + H₂O
3. Add potassium chloride; the less soluble potassium dichromate crystallises out as orange crystals: Na₂Cr₂O₇ + 2KCl → K₂Cr₂O₇ + 2NaCl

Chromate and dichromate interconvert with pH. Acid turns yellow chromate into orange dichromate, and alkali turns dichromate back into chromate; chromium is +6 in both, so this is not a redox change (NCERT page 106). A useful overall equilibrium is 2CrO₄²⁻ + 2H⁺ ⇌ Cr₂O₇²⁻ + H₂O: added H⁺ shifts it right, and OH⁻ removes H⁺ and shifts it left (reasoned from Le Chatelier’s principle). Actual speciation also depends on total chromium concentration and includes hydrogen chromate, HCrO_₄^⁻; the two-ion picture is a school simplification. NCERT writes the two directions as:

• 2CrO₄²⁻ + 2H⁺ → Cr₂O₇²⁻ + H₂O
• Cr₂O₇²⁻ + 2OH⁻ → 2CrO₄²⁻ + H₂O

Structures. The chromate ion is tetrahedral; the dichromate ion is two tetrahedra sharing one corner, with a Cr–O–Cr angle of 126° (NCERT page 106).

Oxidising action. In acid, dichromate is a strong oxidising agent (NCERT page 106):

It oxidises iodide to iodine, sulphides to sulphur, tin(II) to tin(IV) and iron(II) to iron(III). In each, one dichromate ion takes six electrons (NCERT page 106):

• 6I⁻ → 3I₂ + 6e⁻
• 3H₂S → 6H⁺ + 3S + 6e⁻
• 3Sn²⁺ → 3Sn⁴⁺ + 6e⁻
• 6Fe²⁺ → 6Fe³⁺ + 6e⁻

Adding a half-reaction to the dichromate one gives the full equation, for iron(II): Cr₂O₇²⁻ + 14H⁺ + 6Fe²⁺ → 2Cr³⁺ + 6Fe³⁺ + 7H₂O (NCERT page 106). The Class 11 page shows this half-reaction method step by step.

Potassium permanganate: preparation, properties and uses

Preparation (NCERT pages 106–107):

• From MnO₂. Fuse MnO₂ with an alkali metal hydroxide and an oxidising agent such as KNO₃, or air. This gives dark green potassium manganate, which disproportionates in neutral or acid solution to permanganate:
  • 2MnO₂ + 4KOH + O₂ → 2K₂MnO₄ + 2H₂O
  • 3MnO₄²⁻ + 4H⁺ → 2MnO₄⁻ + MnO₂ + 2H₂O
• Commercially, the alkaline oxidative fusion of MnO₂ gives manganate (MnO₄²⁻), which is then oxidised to permanganate (MnO₄⁻) electrolytically in alkaline solution.
• In the laboratory, a manganese(II) salt is oxidised by peroxodisulphate: 2Mn²⁺ + 5S₂O₈²⁻ + 8H₂O → 2MnO₄⁻ + 10SO₄²⁻ + 16H⁺.

Properties (NCERT page 107):

• Dark purple, almost black crystals, isostructural with KClO₄. Not very soluble: 6.4 g per 100 g of water at 293 K.
• On heating it decomposes at 513 K: 2KMnO₄ → K₂MnO₄ + MnO₂ + O₂.
• Its intense colour, and its diamagnetism with a weak temperature-dependent paramagnetism, need molecular orbital theory, beyond NCERT’s scope.
• Manganate and permanganate ions are both tetrahedral, with π bonding from the overlap of oxygen p orbitals with manganese d orbitals. Green manganate is paramagnetic, with one unpaired electron; permanganate has none and is diamagnetic.

Oxidising action. How far permanganate is reduced depends on the hydrogen ion concentration (NCERT page 107):

NCERT’s Table 2.1 (Electrochemistry) lists the last couple as 1.51 V; use the value specified in the question or chosen data table consistently.

• Electrons per MnO₄⁻, by medium. In acid it takes 5, giving Mn²⁺; in neutral or faintly alkaline solution it takes 3, giving MnO₂ (NCERT pages 107–108). The first couple in the table, to green manganate, takes 1. These electron counts are what titration problems use.

Kinetics matters too: at [H⁺] = 1 permanganate should oxidise water, but in practice the reaction is extremely slow unless Mn(II) ions are present or the temperature is raised (NCERT page 107).

In acid solution permanganate is reduced to Mn²⁺ (NCERT pages 107–108):

• iodide to iodine: 10I⁻ + 2MnO₄⁻ + 16H⁺ → 2Mn²⁺ + 8H₂O + 5I₂
• green Fe²⁺ to yellow Fe³⁺: 5Fe²⁺ + MnO₄⁻ + 8H⁺ → Mn²⁺ + 4H₂O + 5Fe³⁺
• oxalate or oxalic acid, at 333 K: 5C₂O₄²⁻ + 2MnO₄⁻ + 16H⁺ → 2Mn²⁺ + 8H₂O + 10CO₂
• hydrogen sulphide, precipitating sulphur: 5S²⁻ + 2MnO₄⁻ + 16H⁺ → 2Mn²⁺ + 8H₂O + 5S (NCERT first writes H₂S → 2H⁺ + S²⁻; with the H₂S molecule itself, 5H₂S + 2MnO₄⁻ + 6H⁺ → 2Mn²⁺ + 5S + 8H₂O, reasoned)
• sulphite to sulphate: 5SO₃²⁻ + 2MnO₄⁻ + 6H⁺ → 2Mn²⁺ + 3H₂O + 5SO₄²⁻
• nitrite to nitrate: 5NO₂⁻ + 2MnO₄⁻ + 6H⁺ → 2Mn²⁺ + 5NO₃⁻ + 3H₂O

In the standard neutral or faintly alkaline reactions listed here it is reduced to MnO₂ (NCERT page 108):

• iodide to iodate: 2MnO₄⁻ + H₂O + I⁻ → 2MnO₂ + 2OH⁻ + IO₃⁻
• thiosulphate almost quantitatively to sulphate: 8MnO₄⁻ + 3S₂O₃²⁻ + H₂O → 8MnO₂ + 6SO₄²⁻ + 2OH⁻
• manganese(II) to MnO₂, catalysed by zinc sulphate or zinc oxide: 2MnO₄⁻ + 3Mn²⁺ + 2H₂O → 5MnO₂ + 4H⁺

Titrations. Permanganate titrations in these school experiments avoid hydrochloric acid, because chloride can be oxidised to chlorine (NCERT page 108).

Uses. Besides analytical chemistry, KMnO₄ is a favourite oxidant in preparative organic chemistry, and its oxidising power is used to bleach wool, cotton, silk and other textile fibres and to decolourise oils (NCERT page 108).

Solved examples

Common mistakes to avoid

Quick revision summary

• Oxides grow less ionic and more acidic as the oxidation state rises: CrO basic, Cr_₂O_₃ amphoteric, CrO_₃ acidic; Mn_₂O_₇ is a covalent green oil.
• K_₂Cr_₂O_₇: chromite → sodium chromate (fusion with Na_₂CO_₃ in air) → sodium dichromate (H_₂SO_₄) → potassium dichromate (KCl). It is a primary standard.
• Chromate (yellow) and dichromate (orange) interconvert with pH; chromium stays +6. Dichromate in acid: E⊖ = 1.33 V.
• KMnO_₄: MnO_₂ fused with KOH and an oxidant → green manganate → permanganate (disproportionation, or electrolytic oxidation).
• Permanganate goes to Mn^²⁺ in acid (E⊖ = 1.52 V) and to MnO_₂ in neutral or faintly alkaline solution; it decomposes at 513 K.

Questions students ask

Why are KMnO₄ and K₂Cr₂O₇ coloured though Mn(VII) and Cr(VI) are d^⁰?

KMnO₄ and K₂Cr₂O₇ owe their colour to charge transfer, not to d–d transitions: Mn(VII) and Cr(VI) are d^⁰, so no d–d transition is possible (reasoned from NCERT pages 103 and 107), but visible light can move an electron from the oxygens into the metal’s empty d orbitals (beyond the NCERT text; Wikipedia, “Charge-transfer band” and “Potassium permanganate”). Such bands absorb far more strongly than d–d bands, so permanganate is intensely purple even when dilute, while molybdate and tungstate, also d^⁰, absorb only in the ultraviolet and are colourless (Wikipedia, “Charge-transfer band”).

Why does acidified potassium dichromate turn green when it oxidises something?

Acidified potassium dichromate turns from orange to green because its chromium is reduced from +6 in Cr₂O₇²⁻ to Cr³⁺ (NCERT page 106), and chromium(III) in the spent solution is green (beyond the NCERT text; Wikipedia, “Potassium dichromate”). The hexaaqua complex [Cr(H₂O)₆]³⁺ is violet, the colour NCERT Table 4.8 gives Cr³⁺ (page 103), but chromium(III) complexes in which anions such as sulphate or chloride replace some water are green (Wikipedia, “Chromium”, “Chromium(III) sulfate”). The colour change tells aldehydes, which it oxidises, from ordinary ketones under mild test conditions; vigorous oxidation can also break down ketones (Wikipedia, “Potassium dichromate”).

Why is potassium dichromate acidified in the usual oxidation reactions?

Potassium dichromate is a strong oxidising agent in acid because its reduction uses up 14 H⁺ ions per Cr₂O₇²⁻, giving 2Cr³⁺ and 7H₂O, with E^⊖ = 1.33 V for standard activities (NCERT page 106). By the Nernst equation, at 298 K the potential falls by 14 × 0.059/6 = 0.138 V for each unit rise in pH, so formally at pH 3 with the other activities held at unity, it is about 0.92 V (reasoned). This illustrates the pH dependence, not the full speciation of a real solution. Alkali shifts the equilibrium towards yellow chromate, which can still oxidise suitable substrates; acid strengthens the usual dichromate oxidising action.

What is the difference between KMnO₄ and K₂Cr₂O₇ as oxidising agents?

KMnO₄ is the stronger oxidant in acid, E^⊖ = 1.52 V against 1.33 V (NCERT pages 106–107), but only K₂Cr₂O₇ is a primary standard (NCERT page 106). MnO₄⁻ takes 5 electrons and Cr₂O₇²⁻ 6, so the equivalent mass of K₂Cr₂O₇ in acid is 294.2/6 = 49.0 g per equivalent, against 31.6 g per equivalent for KMnO₄ (reasoned). Permanganate is its own indicator, while dichromate needs diphenylamine, which it turns blue (NCERT Class 11 Chemistry, page 249).

Why is K₂Cr₂O₇, not the sodium salt, used as a primary standard?

K₂Cr₂O₇ is the primary standard because a weighed sample has a fixed composition: it has no water of crystallisation and, beyond the NCERT text, is not deliquescent, unlike sodium dichromate, Na₂Cr₂O₇·2H₂O (NCERT page 105; Wikipedia, “Potassium dichromate”). Water of crystallisation alone is no bar: the NCERT Laboratory Manual (Chemistry, Class 11) rules out unstable hydrated salts but accepts stable ones such as oxalic acid (page 67). The US National Bureau of Standards recrystallised K₂Cr₂O₇, fused it at 400 °C and weighed it for standard solutions (Fowler and Bright, 1935).

Are K₂Cr₂O₇ and KMnO₄ in the JEE Main and NEET syllabus?

Yes, Unit 11 of both the JEE Main 2027 and the NEET (UG) 2026 syllabus ends its list for the transition elements with the preparation, properties and uses of K₂Cr₂O₇ and KMnO₄, so the preparations, reactions and uses on this page all come under it. The oxides and oxoanions of the 3d metals, NCERT’s Section 4.4.1, are not named as such, though the same unit lists the oxidation states of the first-row transition elements. The CBSE Class XII syllabus for 2026–27 lists the whole section, “Some Important Compounds of Transition Elements”, in Unit 4.

Practice questions

Try each question first, then open the answer to check your working.

1. Q1. Describe the preparation of potassium dichromate from iron chromite ore. What is the effect of increasing the pH of a solution of potassium dichromate? (NCERT Exercise 4.14.)

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   Chromite is fused with sodium carbonate in air, giving sodium chromate; the yellow solution is acidified with sulphuric acid to sodium dichromate; potassium chloride then precipitates the less soluble potassium dichromate (equations as above, NCERT pages 105–106). Raising the pH adds OH⁻, which turns orange dichromate into yellow chromate: Cr₂O₇²⁻ + 2OH⁻ → 2CrO₄²⁻ + H₂O (NCERT page 106).

2. Q2. Write the ionic equation for the reaction of acidified potassium dichromate with iodide. (NCERT Exercise 4.15 (i).)

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   Add the dichromate half-reaction to 6I⁻ → 3I₂ + 6e⁻: Cr₂O₇²⁻ + 14H⁺ + 6I⁻ → 2Cr³⁺ + 3I₂ + 7H₂O (reasoned from NCERT page 106). Check: charge −2 + 14 − 6 = +6 on the left and 2 × 3 = +6 on the right.

3. Q3. Why is V₂O₅ called amphoteric, and what does it give with alkalis and with acids?

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   It reacts with both. With alkalis it gives vanadates, VO₄³⁻; with acids it gives VO₂⁺ salts, vanadium staying +5 in both. It is mainly acidic, like the other high oxides (NCERT page 105).