Compounds of Heavy Metals: Complete Notes for IIT-JAM, BITSAT, GATE, CSIR-NET, TGT/PGT
I still remember the first time I sat down to map out this chapter properly — it looked like an endless list of preparations and colour changes until I realised almost everything here follows two or three recurring ideas: amphoteric behaviour, oxidation state stability, and precipitate colour chemistry. Once you see those threads, tin, lead, iron, copper, zinc, and silver chemistry stop feeling like separate topics and start feeling like variations on a theme. That's exactly how I want to walk you through it.
This chapter sits right at the heart of descriptive inorganic chemistry for competitive exams. GATE and CSIR-NET love testing the reasoning behind colour changes and oxidation states, while BITSAT and the TGT/PGT exams tend to lean on direct factual recall — preparations, uses, and characteristic reactions. I've tried to build these notes so they serve both purposes at once.
Table of Contents
- Tin Compounds — Stannous Oxide (SnO)
- Stannic Oxide (SnO₂)
- Stannous Chloride (SnCl₂)
- Stannic Chloride (SnCl₄)
- Lead Monoxide / Litharge (PbO)
- Lead Dioxide (PbO₂)
- Red Lead (Pb₃O₄)
- Lead Halides — PbX₂ and PbX₄
- Ferric Oxide (Fe₂O₃)
- Cupric Oxide (CuO) and Cuprous Oxide (Cu₂O)
- Zinc Oxide (ZnO)
- Ferric Chloride (FeCl₃)
- Cupric Chloride (CuCl₂)
- Zinc Chloride (ZnCl₂)
- Ferrous Sulphate / Green Vitriol
- Copper Sulphate / Blue Vitriol
- Zinc Sulphate / White Vitriol
- Potassium Permanganate (KMnO₄)
- Potassium Dichromate (K₂Cr₂O₄)
- Silver Compounds — Ag₂O, AgBr, AgNO₃
- Common Misconceptions and Exam Traps
- Practice MCQs from the Chapter (Selected, Rephrased)
- 60+ Additional Practice MCQs with Answers
1. Stannous Oxide (SnO)
Tin gives you two oxidation states worth remembering throughout this chapter — +2 (stannous) and +4 (stannic). SnO belongs to the +2 family, and it's a dark grey to blackish powder, completely insoluble in water. What makes it a favourite exam question is its amphoteric character: it reacts with both acids and bases, something you should expect to be tested on directly.
Preparation. There are three routes worth knowing. Adding NaOH to SnCl₂ gives stannous hydroxide first, and heating that hydroxide under a CO₂ atmosphere yields SnO:
SnCl₂ + 2NaOH → Sn(OH)₂ + 2NaCl
Sn(OH)₂ --(Δ, CO₂)--> SnO + H₂O
A mixture of SnCl₂ and Na₂CO₃ heated under CO₂ works too, and so does heating stannous oxalate in the complete absence of air — this last one decomposes into SnO along with CO and CO₂ together, which is worth memorising exactly since exams like to test the products of oxalate decompositions.
SnC₂O₄ --(Δ)--> SnO + CO + CO₂
Being amphoteric, SnO dissolves in HCl to regenerate stannous chloride, and in NaOH to form sodium stannite (Na₂SnO₂). Here's a detail that trips people up: sodium stannite is only stable while it stays in aqueous solution — expose it to air and it absorbs atmospheric oxygen, oxidising to sodium stannate.
SnO + 2NaOH → Na₂SnO₂ + H₂O
2Na₂SnO₂ + O₂ → 2Na₂SnO₃
Burn SnO in air and it catches with a visible incandescence, converting fully into stannic oxide — a neat illustration of tin's preference for the +4 state under oxidising conditions.
2. Stannic Oxide (SnO₂)
Found naturally as cassiterite or tin stone, SnO₂ is the +4 counterpart. You can prepare it simply by burning tin in air, or through the meta-stannic acid route — tin reacts with dilute HNO₃ to give H₂SnO₃, which then loses water on heating to leave SnO₂ behind.
Sn + 4HNO₃ → H₂SnO₃ + 4NO₂ + H₂O
H₂SnO₃ → SnO₂ + H₂O
It's a white powder, insoluble in water but soluble in concentrated H₂SO₄ and in alkalies like KOH, forming potassium stannate. Industrially you'll find it used in pottery glazes, milky glass manufacture, and as a polishing agent known commercially as putty powder — a fact that occasionally shows up as a "match the use" question.
3. Stannous Chloride (SnCl₂)
If there's one compound in this entire chapter you need cold for reducing-agent questions, it's this one. SnCl₂ is prepared either from stannous oxide and HCl, or directly from tin metal and HCl. The hydrated crystals come out on cooling the concentrated solution. For the anhydrous form, you dry-react tin with chlorine gas, or distil tin with mercuric chloride.
It's a white, hygroscopic, crystalline solid — soluble in water, alcohol, and ether — and it hydrolyses partially in water to give stannous hydroxychloride. With NaOH it forms sodium stannite (the same species you met in section 1), and with H₂S it throws down a dark brown precipitate of SnS, which conveniently dissolves in ammonium sulphide to form ammonium thiostannate.
Now, the reducing behaviour — this is where most marks come from. SnCl₂ reduces mercuric chloride first to mercurous chloride (a white precipitate) and, with excess SnCl₂, further down to metallic mercury (grey-black).
2HgCl₂ + SnCl₂ → Hg₂Cl₂↓ + SnCl₄
Hg₂Cl₂ + SnCl₂ → 2Hg↓ + SnCl₄
It reduces ferric chloride to ferrous chloride, auric chloride all the way down to colloidal gold, and nitro compounds to amines. Notice the pattern — SnCl₂ is oxidised to SnCl₄ in every single one of these reactions. That consistency is exactly what makes "identify the product" questions so predictable once you've internalised it.
4. Stannic Chloride (SnCl₄)
Made by passing dry chlorine over fused tin or over stannous chloride, or by distilling tin with excess mercuric chloride. Physically it's quite distinctive — a colourless, fuming liquid with an unpleasant odour, strongly hygroscopic, and capable of forming crystalline hydrates with 3, 5, 6, or 8 water molecules. The pentahydrate specifically carries the traditional name "butter of tin," and it functions as a mordant in dyeing.
In water it hydrolyses to stannic hydroxide and HCl. Dissolved in concentrated HCl, it forms chlorostannic acid (H₂SnCl₆). It also combines directly with ammonia to give a crystalline adduct, SnCl₄·4NH₃, stable enough to be sublimed without decomposing — and with NH₄Cl it forms the double salt (NH₄)₂SnCl₆, commonly called pink salt, also used as a dyeing mordant.
5. Lead Monoxide or Litharge (PbO)
Lead chemistry mirrors tin's +2/+4 story but with an important twist: for lead, the +2 state is the more stable one (thanks to the inert pair effect), which is the opposite emphasis compared to tin, where +4 tends to dominate thermodynamically in many contexts. Keep that contrast in your head — it resurfaces constantly in comparative questions.
PbO is prepared by heating lead in air at around 300°C, which first gives a yellow powder called massicot; push the temperature further to about 880°C and it turns reddish-yellow, becoming litharge. Chemically the two are identical — PbO — just different crystalline/thermal forms. You can also get it by heating lead nitrate or lead carbonate.
It's amphoteric, dissolving in HCl to give lead chloride and in NaOH to give sodium plumbate. Heat it further and it converts to red lead:
6PbO + O₂ --(Δ)--> 2Pb₃O₄
It's reducible to metallic lead using H₂, CO, or carbon — a fact that connects directly to lead's role in old-fashioned metallurgy questions. Practically, PbO shows up in paints, in flint glass manufacture, in pottery glazing, and mixed with glycerol as a cement for glass and stone joints.
6. Lead Dioxide (PbO₂)
This one has multiple preparation routes worth knowing, because exams like to ask "which of these does NOT give PbO₂." Fusing litharge with KNO₃ or KClO₃ works, as does heating red lead with dilute HNO₃ — this last route also gives lead nitrate alongside PbO₂ as a brown precipitate, which incidentally is the basis of one of the classic "identify the scarlet/chocolate compound" comprehension-passage questions you'll see later in the MCQ section.
The laboratory method treats lead acetate with bleaching powder (in presence of slaked lime), going through lead hydroxide as an intermediate:
(CH₃COO)₂Pb + Ca(OH)₂ → Pb(OH)₂ + (CH₃COO)₂Ca
Pb(OH)₂ + CaOCl₂ → PbO₂ + CaCl₂ + H₂O
PbO₂ is a chocolate-brown powder, insoluble in water. Heat it to 300°C and it reverts to litharge, releasing oxygen — which is exactly why PbO₂ behaves as a strong oxidant. It's amphoteric too: cold concentrated HCl gives plumbic chloride (PbCl₄), while hot concentrated HCl or H₂SO₄ push the reaction further, giving PbCl₂ or PbSO₄ respectively along with the +4 lead being reduced to +2. Hot concentrated NaOH gives sodium plumbate.
As an oxidiser, PbO₂ converts SO₂ to PbSO₄ and oxidises concentrated HCl to release chlorine gas. It's used in match manufacturing, in lead-acid storage batteries, and generally as an oxidising agent in the lab.
7. Red Lead (Pb₃O₄)
Also called minium or sindur, red lead is best understood as a mixed oxide — PbO and PbO₂ combined in a 1:2 ratio, not a genuinely new oxidation state of lead. It forms when litharge is heated with air near 400–470°C in a reverberatory furnace.
It's a red powder, insoluble in water. Push the heating beyond 470°C and it decomposes back into PbO and oxygen, passing through a violet-blackish intermediate colour on the way. As an oxidant it converts HCl to chlorine and, with hot concentrated H₂SO₄, liberates oxygen while forming PbSO₄. It'll also oxidise CO to CO₂.
The reaction with nitric acid is a classic identification test: Pb₃O₄ gives lead nitrate in solution plus a brown precipitate of PbO₂ — direct proof that the compound really is a PbO/PbO₂ mixture rather than a single lead(III) species, since lead(III) doesn't exist as a stable oxidation state. This exact experiment (scarlet compound → chocolate precipitate → yellow iodide precipitate → pink permanganate-like solution with manganese nitrate) is the backbone of one of the comprehension passages you'll find further down.
8. Lead Halides
Lead forms stable dihalides of the PbX₂ type — PbF₂, PbCl₂, PbBr₂, PbI₂ — all essentially ionic, though the ionic character decreases down the series: PbF₂ > PbCl₂ > PbBr₂ > PbI₂. PbCl₂ is a white crystalline solid, soluble in hot water and in excess HCl (forming chloroplumbous acid, H₂PbCl₄). React it with hot lime water and you get Pb(OH)Cl, historically known as Pattinson's white lead, a pigment once used commercially.
Lead(IV) halides are far less stable — only PbF₄ and PbCl₄ are known at all, and PbF₄ has never actually been isolated pure. PbCl₄ itself is obtained from PbO₂ and well-cooled concentrated HCl at 273 K. It's a yellow, covalent liquid, soluble in organic solvents, but genuinely unstable — it decomposes spontaneously into PbCl₂ and Cl₂, which is a nice, testable illustration of why the +4 state of lead is comparatively unstable relative to tin's +4 state. Hydrolysis with water regenerates PbO₂ and HCl.
9. Ferric Oxide (Fe₂O₃)
Occurring naturally as haematite and limonite, Fe₂O₃ can be made by heating ferrous sulphate, by roasting iron pyrite in air, or by heating ferric carbonate. It's a deep red, amphoteric powder — with HCl it forms ferric chloride; with sodium carbonate or NaOH it forms sodium ferrite (NaFeO₂).
A subtlety worth flagging: hydrolysing sodium ferrite regenerates Fe₂O₃ and NaOH, and this cycle is actually exploited industrially as one method of NaOH production — not something most students expect from a "just an oxide" compound. Heated strongly (around 1400°C), Fe₂O₃ converts to magnetite, Fe₃O₄, which is worth remembering as FeO·Fe₂O₃ in its mixed-valence formulation — magnetic, unlike the parent oxide.
6Fe₂O₃ --(Δ, 1400°C)--> 4Fe₃O₄ + O₂
Reduced by carbon monoxide or hydrogen gas, Fe₂O₃ gives metallic iron — the basis of blast furnace chemistry. Industrially, it's used as a catalyst in the Bosch process, as jeweller's rouge (a polishing powder), and as a red pigment.
10. Cupric Oxide (CuO) and Cuprous Oxide (Cu₂O)
CuO — the black oxide of copper, occurring naturally as tenorite — comes from heating cupric hydroxide or copper nitrate, from oxidising cuprous oxide or copper metal in air around 1273–1373 K, or industrially by heating malachite ore. It's black, insoluble in water, but dissolves readily in acids to form the corresponding copper salts.
Heat CuO past 1373 K and it decomposes into cuprous oxide (Cu₂O) with oxygen release. Cu₂O itself is a red-brown powder, insoluble in water but soluble in ammonia, forming the diammine complex [Cu(NH₃)₂]⁺. This colour behaviour — CuO black, Cu₂O red — is a favourite for "identify by colour" style questions, especially combined with the classic glucose reduction test (CuO reacting with glucose gives Cu₂O, red, via Fehling's-type chemistry).
CuO is used to impart a greenish-blue tint to glass, to strip sulphur out of petroleum, and in the estimation of carbon during organic elemental analysis.
11. Zinc Oxide (ZnO)
Known variously as zinc white, philosopher's wool, or Chinese white, and occurring naturally as zincite. You can prepare it by burning zinc directly, or by thermally decomposing zinc carbonate, zinc hydroxide, or zinc nitrate.
Here's the property everyone remembers about ZnO, and rightly so: it's white when cold, turns yellow on heating, and reverts to white again on cooling. That reversible thermochromism comes down to lattice defects — heating drives off some lattice oxygen, generating oxygen vacancies (essentially F-centres) that absorb visible light and impart the yellow colour; cooling lets the oxygen return and the defects heal. It's insoluble in water and sublimes near 673 K.
Being amphoteric, ZnO reacts with acids to form the corresponding zinc salts and with NaOH to give sodium zincate. It's reducible with hydrogen or carbon back to metallic zinc. One reaction worth memorising for its sheer visual distinctiveness: heating ZnO with cobalt nitrate produces a green mass called Rinmann's green (cobalt zincate, CoO·ZnO) — a classic identification test.
Practically, ZnO is prized as a white paint pigment precisely because, unlike white lead, it doesn't blacken on exposure to atmospheric H₂S. It also functions as a catalyst (paired with Cr₂O₃) in methanol synthesis from CO and H₂, and shows up in zinc ointments, cosmetic powders, and surgical dressings as an absorbent.
12. Ferric Chloride (FeCl₃)
Worth noting straightaway: anhydrous FeCl₃ is black, while the hydrated form, FeCl₃·6H₂O, is yellow — a distinction examiners like testing directly. Anhydrous FeCl₃ comes from passing dry chlorine gas over heated iron filings; the hydrated form comes from dissolving ferric carbonate, ferric hydroxide, or ferric oxide in HCl.
In the gas phase, FeCl₃ actually exists as a dimer, Fe₂Cl₆, with two bridging chlorine atoms — a structural detail that occasionally appears in bonding-focused questions. On decomposition it gives FeCl₂ and chlorine gas, and dehydrating the hexahydrate (either thermally or using thionyl chloride) eventually yields Fe₂O₃.
Hydrolysis gives ferric hydroxide and an acidic solution; NH₄OH gives the same reddish-brown Fe(OH)₃ precipitate. As an oxidiser, FeCl₃ converts SnCl₂ to SnCl₄ (getting reduced itself to FeCl₂) and oxidises KI to iodine. Two identification reactions are genuinely important here: with potassium ferrocyanide it forms Prussian blue (Fe₄[Fe(CN)₆]₃), and with ammonium thiocyanate it forms the blood-red ferric thiocyanate. Both of these are staple qualitative-analysis reactions.
13. Cupric Chloride (CuCl₂·2H₂O)
Dissolving copper, CuO, or copper carbonate in concentrated HCl and crystallising gives green hydrated crystals. For anhydrous CuCl₂, you either heat copper directly in excess chlorine, or heat the hydrate with HCl gas at 423 K.
The colour chemistry here is genuinely elegant and worth understanding rather than memorising: dilute aqueous CuCl₂ is blue (dominated by [Cu(H₂O)₄]²⁺), concentrated solution turns green, and adding concentrated HCl shifts it to yellow due to [CuCl₄]²⁻ formation. It's essentially a ligand-exchange equilibrium playing out visibly in front of you.
On hydrolysis it forms an acidic solution due to liberated HCl. Heated, anhydrous CuCl₂ disproportionates into Cu₂Cl₂ and Cl₂. With NaOH it gives a pale blue precipitate of basic cupric chloride; with excess NH₄OH it dissolves into a deep blue tetraammine complex. Reduction with SnCl₂, hydrogen, or copper metal itself converts it to cuprous chloride (Cu₂Cl₂). Industrially, it's the catalyst in the Deacon process for chlorine manufacture.
14. Zinc Chloride (ZnCl₂·2H₂O)
Prepared straightforwardly from ZnO, ZnCO₃, or Zn(OH)₂ reacting with HCl; the anhydrous salt comes from dry chlorine or dry HCl reacting with zinc metal, or via distillation with HgCl₂.
It's a white, deliquescent solid (melting point 933 K, boiling point 1003 K). Heating the hydrate gives zinc oxychloride. With excess NaOH it forms sodium zincate via the intermediate Zn(OH)₂ precipitate; with excess ammonia it forms the tetraammine zinc chloride complex. With sodium carbonate it precipitates basic zinc carbonate, though using sodium bicarbonate instead gives plain zinc carbonate — a subtle distinction that's tested more often than you'd expect. Passing H₂S through its solution throws down white ZnS.
Anhydrous ZnCl₂ is a genuinely useful dehydrating agent, finds use in dry cell batteries, parchment paper manufacture, adhesives, and timber preservation against microbial attack. Mixed with ZnO, it's used for dental filling cement.
15. Ferrous Sulphate or Green Vitriol (FeSO₄·7H₂O)
Also called Harakasis in older texts. Made by oxidising iron pyrite in moist air, or by dissolving iron scraps in dilute H₂SO₄ (crystallised out using alcohol, since FeSO₄ is only partly soluble in it).
It's light green, but turns brown in air on standing due to atmospheric oxidation into basic ferric sulphate. Heated, it loses its water of crystallisation stepwise before finally decomposing at high temperature into Fe₂O₃, SO₂, and SO₃.
Its reducing behaviour is the real exam focus. FeSO₄ decolourises acidified KMnO₄ (purple to colourless, forming MnSO₄), and it likewise decolourises acidified K₂Cr₂O₄ (orange to green, forming Cr₂(SO₄)₃) — these two titrations underpin a huge fraction of volumetric-analysis questions in this chapter. It also reduces HgCl₂ to Hg₂Cl₂ and AuCl₃ to metallic gold.
Mixed with ammonium sulphate in the right stoichiometry, it crystallises out as Mohr's salt, FeSO₄·(NH₄)₂SO₄·6H₂O — pale green crystals prized as a primary standard in volumetric analysis precisely because they don't oxidise as readily in air as plain FeSO₄ does.
16. Copper Sulphate or Blue Vitriol (CuSO₄·5H₂O)
Prepared by dissolving cupric oxide, hydroxide, or carbonate in dilute H₂SO₄, or industrially by blowing air through copper scrap sitting in dilute H₂SO₄ inside a perforated lead bucket.
The dehydration sequence on heating is a frequently tested one: the pentahydrate loses water stepwise to the trihydrate (pale blue), then the monohydrate (bluish-white), then the anhydrous salt (white), before finally decomposing near 720°C into CuO and sulphur oxides.
A cluster of complexation reactions defines this compound's exam relevance. With KI it forms white cuprous iodide alongside liberated iodine (the classic iodometric estimation of Cu²⁺). With excess ammonia it forms the deep blue tetraammine complex, historically called Schweitzer's reagent, notable for its ability to dissolve cellulose — a property exploited in artificial silk manufacture. With excess KCN it forms potassium cuprocyanide with cyanogen gas released; with K₄Fe(CN)₆ it gives a reddish-brown precipitate of cupric ferrocyanide, used as a standard test for Cu²⁺. With hypo (sodium thiosulphate) in excess, it forms sodium cuprothiosulphate.
Uses span electroplating, calico printing, and dyeing, plus its role as a fungicide — the Bordeaux mixture (CuSO₄ plus lime) remains a genuinely important agricultural fungicide against mould and fungal disease on crops.
17. Zinc Sulphate or White Vitriol (ZnSO₄·7H₂O)
Straightforward preparation from zinc, ZnO, or ZnCO₃ with dilute H₂SO₄. It's colourless, crystalline, and efflorescent, decomposing on heating through a stepwise dehydration sequence broadly similar to blue vitriol's, ultimately giving ZnO at 800°C.
With excess NaOH it forms sodium zincate through the Zn(OH)₂ intermediate; with Na₂CO₃ it gives basic zinc carbonate (again, plain zinc carbonate if you use bicarbonate instead — the same subtlety as with ZnCl₂). It's used to prepare lithopone (a mixture of ZnO and BaSO₄, a well-known white pigment), in eye lotions, and as a mordant in dyeing and calico printing.
18. Potassium Permanganate (KMnO₄)
This compound deserves careful attention because it's one of the most heavily tested reagents across all of inorganic chemistry, not just this chapter. Preparation starts from pyrolusite ore (MnO₂), fused with KOH or K₂CO₃ in air to give green potassium manganate (K₂MnO₄):
2MnO₂ + 4KOH + O₂ → 2K₂MnO₄ + 2H₂O
The manganate is then oxidised to permanganate using chlorine, ozone, or CO₂, or alternatively through electrolytic oxidation at the anode. The resulting purple solution, concentrated and crystallised, gives dark purple, needle-shaped crystals with a distinctive greenish metallic lustre, melting at 523 K. It's sparingly soluble in cold water but dissolves readily in hot water.
Heating KMnO₄ gives different products depending on the exact temperature — a genuinely favourite exam trap. At around 200°C it gives K₂MnO₄, MnO₂, and O₂. Reaction with concentrated H₂SO₄ produces Mn₂O₄, a dark brown, highly explosive liquid — handle with extreme theoretical caution in your mind, if nowhere else.
The oxidising behaviour changes character with pH, and this is where most marks in this section actually live:
- Neutral medium: KMnO₄ decomposes to give MnO₂, KOH, and nascent oxygen. It oxidises H₂S to elemental sulphur, and hypo to sodium sulphate.
- Alkaline medium: Gives K₂MnO₄ and nascent oxygen. This is "Baeyer's reagent," the classic test for unsaturation in organic compounds — alkenes get converted to diols (the textbook example being ethene to ethylene glycol).
- Acidic medium: Gives Mn²⁺ salts directly, releasing more oxidising power (5 [O] per two moles of KMnO₄). This is the medium used for essentially all the classic redox titrations — oxidising FeSO₄ to ferric sulphate, halide ions to the free halogen, oxalic acid to CO₂, SO₂ to sulphuric acid, and nitrites to nitrates.
Structurally, the permanganate ion is tetrahedral, with Mn–O bond length around 1.63 Å and O–Mn–O angle close to 109°. Uses include disinfection, water treatment, manufacture of saccharin and boric acid, and volumetric estimation of Fe²⁺, oxalates, and hydrogen peroxide.
19. Potassium Dichromate (K₂Cr₂O₄)
Manufacturing starts from chromite ore (FeCr₂O₄). Fusing it with molten alkali and air in a reverberatory furnace, then extracting the fused mass with water and filtering, gives sodium chromate (Na₂CrO₄). Acidifying that with dilute H₂SO₄ converts it to sodium dichromate; treating sodium dichromate with KCl then gives potassium dichromate, which is far less soluble than the sodium salt and so separates out readily by fractional crystallisation.
4FeCr₂O₄ + 8Na₂CO₃ + 7O₂ → 2Fe₂O₃ + 8CO₂ + 8Na₂CrO₄
2Na₂CrO₄ + H₂SO₄ → Na₂Cr₂O₄ + Na₂SO₄ + H₂O
Na₂Cr₂O₄ + 2KCl --(Δ)--> K₂Cr₂O₄↓ + 2NaCl
K₂Cr₂O₄ is preferred over Na₂Cr₂O₄ specifically for volumetric analysis because it isn't hygroscopic, unlike its sodium counterpart — this is a very frequently tested reasoning point, so it's worth being able to state clearly rather than just recall as a fact.
Physically, it's orange-red, crystalline, melting at 669 K, moderately soluble in cold water but readily soluble in hot water. Heated strongly, it decomposes to potassium chromate, chromic oxide, and oxygen. With cold concentrated H₂SO₄ it forms red crystalline chromic anhydride (CrO₃); with hot concentrated H₂SO₄, chromium sulphate forms instead, releasing oxygen. With HCl it gives chromium chloride, and with KOH it converts (reversibly, on acidification) into yellow potassium chromate — the classic chromate–dichromate interconversion governed by pH.
As an oxidant in acidic medium, its reduction half-reaction is worth memorising outright:
Cr₂O₄²⁻ + 14H⁺ + 6e⁻ → 2Cr³⁺ + 7H₂O
It oxidises Fe²⁺ to Fe³⁺, iodide to iodine, sulphite to sulphate, arsenite to arsenate, halide ions (except fluoride) to the free halogen, and ethanol progressively to acetaldehyde and then acetic acid. With hydrogen peroxide in acidic medium it gives a striking deep blue colour due to CrO₅ formation — a distinctive spot test.
The chromyl chloride test deserves special mention because it's the standard qualitative test for chloride ion in a mixture. Heating the suspected chloride with K₂Cr₂O₄ and concentrated H₂SO₄ generates orange-red chromyl chloride vapours (CrO₂Cl₂), which dissolve in NaOH to give a yellow sodium chromate solution — and that final yellow colour is your positive confirmation.
K₂Cr₂O₄ + 4NaCl + 6H₂SO₄ --(Δ)--> 2KHSO₄ + 4NaHSO₄ + 2CrO₂Cl₂↑ + 3H₂O
CrO₂Cl₂ + 4NaOH → 2NaCl + 2H₂O + Na₂CrO₄
Structurally, the chromate ion is tetrahedral while dichromate consists of two tetrahedra sharing a corner oxygen, with a Cr–O–Cr bond angle near 130°. Main use: hardening gelatin in photography, and as a mordant in dyeing.
20. Silver Compounds
Silver oxide (Ag₂O) forms simply by treating silver nitrate with NaOH. It's a brown precipitate that looks essentially black once fully dried, and it decomposes slowly on heating (around 330°C) into metallic silver and oxygen.
Silver bromide (AgBr) comes from AgNO₃ reacting with NaBr — a pale yellow solid, insoluble in water and in concentrated acids, though it does dissolve slightly in strong ammonia via complex formation (diammine silver bromide). Its light sensitivity is the whole basis of its photographic importance: exposure to light causes photoreduction to metallic silver and bromine.
2AgBr --(hν)--> 2Ag + Br₂
It dissolves in sodium thiosulphate (hypo) to give soluble sodium argentothiosulphate — precisely the chemistry used to "fix" photographic film by dissolving away the unexposed AgBr that would otherwise keep darkening. It also dissolves in KCN, forming potassium argentocyanide.
Silver nitrate (AgNO₃), historically called lunar caustic, is prepared by dissolving silver in dilute nitric acid. It's colourless, crystalline, melting at 485 K, soluble in water and alcohol. The "lunar caustic" name comes directly from its tendency to leave black stains on skin or clothing — that's metallic silver being deposited through photoreduction, not a corrosive burn as such.
Heating shows a temperature-dependent product pattern worth knowing exactly: just above the melting point, it gives silver nitrite and oxygen; at red heat, it decomposes fully to metallic silver, NO₂, and oxygen.
2AgNO₃ --(Δ, above m.p.)--> 2AgNO₂ + O₂
2AgNO₃ --(red hot)--> 2Ag + 2NO₂ + O₂
AgNO₃ is essentially the reagent of choice for precipitating a whole range of anions — chloride (white AgCl), bromide (pale yellow AgBr), iodide (yellow AgI), phosphate (yellow Ag₃PO₄), sulphide (black Ag₂S), thiocyanate (white AgCNS), and chromate (brick-red Ag₂CrO₄). These colour codes are worth memorising as a set since qualitative analysis questions draw on them constantly.
With hypo, it initially gives a white precipitate that famously turns black on standing — white silver thiosulphate hydrolysing to black Ag₂S. With excess ammonia, it forms diammine silver nitrate, and this ammoniacal solution is exactly what's known as Tollen's reagent, used to test for the aldehyde (–CHO) group via the silver mirror test with glucose or formaldehyde.
21. Common Misconceptions and Exam Traps
A few recurring confusion points are worth listing together, since exams often test them as assertion-reason pairs or as "which statement is incorrect" questions.
- Sodium stannite is stable only in solution — it oxidises to stannate on air exposure. Don't treat it as a stable solid reagent.
- HgCl₂ and SnCl₂ cannot coexist in aqueous solution; SnCl₂ reduces HgCl₂ stepwise to Hg₂Cl₂ and then metallic Hg.
- Pb⁴⁺ is a stronger oxidising agent than Sn⁴⁺ — this follows directly from the inert pair effect, where the lower oxidation state becomes progressively more stable as you descend group 14. Don't reverse this comparison.
- KMnO₄'s purple colour arises from charge-transfer transition, not d–d transition, since Mn in the +7 state has a d⁰ configuration and no d-electrons are available for a genuine d–d transition. K₂Cr₂O₄'s colour is likewise charge-transfer in origin.
- Green vitriol (FeSO₄·7H₂O) almost always carries trace Fe³⁺ from aerial oxidation, which is precisely why it isn't used as a primary standard — Mohr's salt is used instead.
- Anhydrous FeCl₃ is black; only the hydrated form is yellow. Don't mix these up when a question specifies "anhydrous."
22. Practice MCQs Based on the Chapter
These questions test the same core ideas found in the standard question bank for this chapter, rephrased here for practice. Try them before checking answers.
- When FeCl₃ solution is treated with excess NaOH, what precipitate forms?
(a) Fe₃O₄ (b) Hydrated Fe₂O₃ (c) A mixture of Fe₂O₃ and FeO (d) FeO and FeCl₃
Answer: (b) - Heating ferrous oxalate in the complete absence of air yields:
(a) Fe₂O₃ (b) Fe₃O₄ (c) Fe₂O₄ (d) FeO
Answer: (d) - Red-hot copper wire reacting with steam produces:
(a) Cu₂O (b) CuO₂ (c) Cu₂O₂ (d) CuO
Answer: (d) - Yellow FeCl₃ solution turns light green upon adding:
(a) Zn (b) SnCl₂ (c) H₂S gas (d) Any of these
Answer: (d) - Fe(OH)₂ dissolving in concentrated NaOH gives a blue-green complex of formula:
(a) Na₂[Fe(OH)₆] (b) Na₄[Fe(OH)₄] (c) Na₄[Fe₂(OH)₆] (d) Na₄[Fe(OH)₆]
Answer: (d) - On strong heating, Fe₂(SO₄)₃ decomposes to give:
(a) FeO and SO₃ (b) Fe₂O₃ and SO₂ (c) FeO and SO₂ (d) Fe₂O₃ and SO₃
Answer: (d) - Prussian blue, formed when FeCl₃ reacts with K₄[Fe(CN)₆], is correctly represented as:
(a) K Fe(III)[Fe(II)(CN)₆] (b) Fe₄(III)[Fe(III)(CN)₆] (c) K Fe(II)[Fe(IV)(CN)₆] (d) Fe₄(II)[Fe(II)(CN)₆]
Answer: (a) - Which compound can be used to standardise a KMnO₄ solution?
(a) Fe₃O₄ (b) FeCl₃ (c) FeSO₄·7H₂O (d) Mohr's salt
Answer: (d) - A CuSO₄ solution treated with excess KI gives a dark brown colouration due to:
(a) Cu₂I₂(s) (b) CuI₂(s) (c) I₂(s) (d) I₃⁻(aq)
Answer: (d) - Silver fluoride, unlike the other silver halides, is soluble in water because:
(a) It is covalent (b) Its hydration energy exceeds its lattice energy (c) It ionises completely (d) Fluorine is highly electronegative
Answer: (b) - Pb₃O₄ reacting with dilute HNO₃ produces lead nitrate along with a brown precipitate of:
(a) PbO (b) PbO₂ (c) Pb(OH)₂ (d) Pb₂O₃
Answer: (b) - The oxidation state of Mn changes from which value to which value when KMnO₄ acts as an oxidant in acidic medium?
(a) +7 to +4 (b) +7 to +2 (c) +6 to +3 (d) +7 to 0
Answer: (b) - K₂Cr₂O₄ is preferred over Na₂Cr₂O₄ in volumetric analysis mainly because:
(a) It is more soluble (b) It is not hygroscopic (c) It is cheaper (d) It is more reactive
Answer: (b) - Mohr's salt is used as a primary standard because it contains:
(a) Only Fe²⁺ ions without Fe³⁺ contamination (b) A mixture of Fe²⁺ and Fe³⁺ (c) Only Fe³⁺ ions (d) No iron at all
Answer: (a) - Zinc oxide turns yellow on heating and reverts to white on cooling due to:
(a) Loss of lattice oxygen creating defects (b) A phase transition (c) d-d transitions (d) Sublimation
Answer: (a)
23. Additional Practice MCQs (60+ New Questions with Answers)
These are freshly written questions covering the same concepts, meant to widen your practice beyond the standard set above.
- Which oxidation state of tin is more stable in oxidising conditions?
(a) +2 (b) +4 (c) 0 (d) +3 — Answer: (b) - SnO reacting with excess NaOH forms:
(a) Sodium stannate (b) Sodium stannite (c) Tin hydroxide (d) SnO₂ — Answer: (b) - Sodium stannite is oxidised in air to form:
(a) Sodium stannate (b) SnO₂ (c) Sn(OH)₄ (d) SnCl₂ — Answer: (a) - Cassiterite is the natural ore of:
(a) Lead (b) Iron (c) Tin (d) Zinc — Answer: (c) - SnO₂ dissolves in concentrated H₂SO₄ to form:
(a) SnSO₄ (b) Sn(SO₄)₂ (c) SnO₃ (d) Sn₂(SO₄)₃ — Answer: (b) - SnCl₂ reduces mercuric chloride first to:
(a) Metallic mercury (b) Mercurous chloride (c) Mercuric oxide (d) No reaction — Answer: (b) - Excess SnCl₂ with HgCl₂ ultimately gives:
(a) Hg₂Cl₂ only (b) Metallic Hg (c) No further reaction (d) HgO — Answer: (b) - "Butter of tin" refers to:
(a) SnCl₂ (b) SnCl₄·5H₂O (c) SnO (d) SnO₂ — Answer: (b) - Pink salt is the common name for:
(a) (NH₄)₂SnCl₆ (b) SnCl₄·4NH₃ (c) SnCl₂·2H₂O (d) K₂SnCl₆ — Answer: (a) - Litharge and massicot are both chemically:
(a) Different compounds (b) The same compound, PbO, in different forms (c) PbO₂ isomers (d) Pb₃O₄ forms — Answer: (b) - PbO dissolved in NaOH gives:
(a) Sodium plumbite (b) Sodium plumbate (c) Lead hydroxide (d) No reaction — Answer: (b) - Which reducing agent converts PbO to metallic lead?
(a) H₂ only (b) CO only (c) Carbon only (d) All of H₂, CO, and carbon — Answer: (d) - PbO₂ heated to around 300°C converts to:
(a) Pb₃O₄ (b) PbO (c) Pb metal (d) PbSO₄ — Answer: (b) - Cold concentrated HCl with PbO₂ gives:
(a) PbCl₂ (b) PbCl₄ (c) Pb(OH)₂ (d) No reaction — Answer: (b) - Hot concentrated HCl with PbO₂ gives PbCl₂ along with:
(a) H₂ gas (b) Cl₂ gas (c) O₂ gas (d) No gas — Answer: (b) - Red lead, Pb₃O₄, is best described as:
(a) A pure lead(III) oxide (b) A mixed oxide of PbO and PbO₂ (c) A hydrated oxide (d) A peroxide — Answer: (b) - The ratio of PbO to PbO₂ in red lead is:
(a) 1:1 (b) 2:1 (c) 1:2 (d) 3:1 — Answer: (b) - Decreasing order of ionic character among lead halides is:
(a) PbI₂ > PbBr₂ > PbCl₂ > PbF₂ (b) PbF₂ > PbCl₂ > PbBr₂ > PbI₂ (c) PbCl₂ > PbF₂ > PbI₂ > PbBr₂ (d) All equal — Answer: (b) - Pattinson's white lead has the formula:
(a) PbCO₃ (b) Pb(OH)Cl (c) PbSO₄ (d) 2PbCO₃·Pb(OH)₂ — Answer: (b) - PbCl₄ is best described physically as:
(a) A white solid (b) A yellow covalent liquid (c) A colourless gas (d) A black solid — Answer: (b) - Ferric oxide occurs naturally as:
(a) Magnetite only (b) Haematite and limonite (c) Siderite (d) Pyrite — Answer: (b) - Heating Fe₂O₃ strongly around 1400°C gives:
(a) FeO (b) Fe₃O₄ (c) Fe metal (d) FeCl₃ — Answer: (b) - Fe₂O₃ reacting with NaOH forms:
(a) Sodium ferrite (b) Sodium ferrate (c) Ferric hydroxide only (d) No reaction — Answer: (a) - Cupric oxide is known naturally as:
(a) Malachite (b) Tenorite (c) Azurite (d) Chalcopyrite — Answer: (b) - Heating CuO above 1373 K gives:
(a) Cu metal (b) Cu₂O (c) CuCO₃ (d) Cu(OH)₂ — Answer: (b) - Cu₂O is soluble in:
(a) Water (b) Ammonia (c) Benzene (d) It is insoluble everywhere — Answer: (b) - Zinc oxide is known by which common name?
(a) Chinese white (b) Venetian red (c) Turkey red (d) Prussian white — Answer: (a) - ZnO reacting with cobalt nitrate on heating gives:
(a) Rinmann's green (b) Prussian blue (c) Turnbull's blue (d) Bordeaux mixture — Answer: (a) - ZnO is preferred over white lead as a pigment because:
(a) It's cheaper (b) It doesn't blacken with atmospheric H₂S (c) It's more soluble (d) It's less dense — Answer: (b) - Anhydrous FeCl₃ exists as a dimer in the gaseous state, represented as:
(a) FeCl₃ (b) Fe₂Cl₆ (c) Fe₃Cl₉ (d) FeCl₂ — Answer: (b) - FeCl₃ with ammonium thiocyanate gives a blood-red compound known as:
(a) Ferric ferrocyanide (b) Ferric thiocyanate (c) Prussian blue (d) Turnbull's blue — Answer: (b) - Dilute CuCl₂ solution appears blue due to the ion:
(a) [CuCl₄]²⁻ (b) [Cu(H₂O)₄]²⁺ (c) Cu²⁺ alone (d) [CuCl₂]⁺ — Answer: (b) - Concentrated CuCl₂ solution with added concentrated HCl turns:
(a) Blue (b) Green (c) Yellow (d) Colourless — Answer: (c) - Heating anhydrous CuCl₂ gives:
(a) CuCl (b) Cu₂Cl₂ and Cl₂ (c) CuO (d) Metallic Cu — Answer: (b) - The Deacon process, which uses CuCl₂ as a catalyst, is used to manufacture:
(a) Ammonia (b) Chlorine (c) Sulphuric acid (d) Nitric acid — Answer: (b) - ZnCl₂ reacting with sodium bicarbonate gives:
(a) Basic zinc carbonate (b) Plain zinc carbonate (c) Zinc oxide (d) No reaction — Answer: (b) - ZnCl₂ reacting with sodium carbonate (not bicarbonate) gives:
(a) Plain ZnCO₃ (b) Basic zinc carbonate (c) ZnO (d) Zn(OH)₂ only — Answer: (b) - Green vitriol refers to:
(a) CuSO₄·5H₂O (b) FeSO₄·7H₂O (c) ZnSO₄·7H₂O (d) NiSO₄·6H₂O — Answer: (b) - FeSO₄ decolourises acidified KMnO₄ because:
(a) FeSO₄ is oxidised to Fe₂(SO₄)₃ (b) FeSO₄ is reduced (c) KMnO₄ is inert to FeSO₄ (d) A precipitate forms — Answer: (a) - Mohr's salt has the formula:
(a) FeSO₄·(NH₄)₂SO₄·6H₂O (b) FeSO₄·7H₂O (c) Fe₂(SO₄)₃·9H₂O (d) FeCl₃·6H₂O — Answer: (a) - Blue vitriol refers to:
(a) FeSO₄·7H₂O (b) CuSO₄·5H₂O (c) ZnSO₄·7H₂O (d) NiSO₄·6H₂O — Answer: (b) - Schweitzer's reagent, useful for dissolving cellulose, is:
(a) Tetraammine copper sulphate (b) Ammoniacal AgNO₃ (c) Sodium zincate (d) Fehling's solution — Answer: (a) - Bordeaux mixture is a combination of:
(a) CuSO₄ and lime (b) FeSO₄ and lime (c) ZnSO₄ and lime (d) AgNO₃ and lime — Answer: (a) - White vitriol refers to:
(a) ZnSO₄·7H₂O (b) MgSO₄·7H₂O (c) CaSO₄·2H₂O (d) FeSO₄·7H₂O — Answer: (a) - Lithopone is a mixture of:
(a) ZnO and BaSO₄ (b) ZnO and CaSO₄ (c) TiO₂ and BaSO₄ (d) PbO and BaSO₄ — Answer: (a) - KMnO₄ is prepared industrially starting from the ore:
(a) Haematite (b) Pyrolusite (c) Chromite (d) Bauxite — Answer: (b) - Fusing MnO₂ with KOH in air gives:
(a) KMnO₄ directly (b) K₂MnO₄ (c) Mn₂O₃ (d) MnO — Answer: (b) - K₂MnO₄ is oxidised to KMnO₄ using:
(a) H₂ gas (b) Chlorine or ozone (c) CO gas (d) N₂ gas — Answer: (b) - KMnO₄ reacting with concentrated H₂SO₄ gives the explosive compound:
(a) MnO₂ (b) Mn₂O₄ (c) Mn₂O₃ (d) MnSO₄ — Answer: (b) - Baeyer's reagent refers to:
(a) Acidic KMnO₄ (b) Alkaline KMnO₄ (c) Neutral KMnO₄ (d) Concentrated H₂SO₄ — Answer: (b) - Baeyer's reagent is used to test for:
(a) Aldehydes (b) Unsaturation in organic compounds (c) Alcohols (d) Amines — Answer: (b) - In acidic medium, KMnO₄ is reduced to:
(a) MnO₂ (b) Mn²⁺ (c) MnO₄²⁻ (d) Mn₂O₃ — Answer: (b) - The geometry of the permanganate ion is:
(a) Tetrahedral (b) Octahedral (c) Square planar (d) Linear — Answer: (a) - KMnO₄'s intense purple colour is attributed to:
(a) d-d transition (b) Charge transfer transition (c) f-f transition (d) Vibrational transition — Answer: (b) - K₂Cr₂O₄ is manufactured starting from the ore:
(a) Pyrolusite (b) Chromite (c) Bauxite (d) Cassiterite — Answer: (b) - Sodium dichromate is converted to potassium dichromate by adding:
(a) KOH (b) KCl (c) K₂SO₄ (d) KNO₃ — Answer: (b) - K₂Cr₂O₄ is preferred in volumetric analysis over the sodium salt mainly because:
(a) It's more soluble (b) It's not hygroscopic (c) It's more oxidising (d) It's cheaper — Answer: (b) - Cold concentrated H₂SO₄ with K₂Cr₂O₄ gives red crystals of:
(a) CrO₃ (b) Cr₂O₃ (c) CrO₅ (d) CrCl₃ — Answer: (a) - The chromyl chloride test is used to detect:
(a) Bromide ion (b) Chloride ion (c) Iodide ion (d) Sulphate ion — Answer: (b) - Chromyl chloride vapours dissolved in NaOH give a solution that is:
(a) Colourless (b) Yellow (Na₂CrO₄) (c) Orange (d) Green — Answer: (b) - Acidified K₂Cr₂O₄ with hydrogen peroxide gives a deep blue colour due to:
(a) CrO₃ (b) CrO₅ (c) Cr₂O₃ (d) CrO₄²⁻ — Answer: (b) - The Cr–O–Cr bond angle in dichromate ion is approximately:
(a) 109° (b) 120° (c) 130° (d) 180° — Answer: (c) - Silver oxide is prepared by treating AgNO₃ with:
(a) NaOH (b) HCl (c) Na₂CO₃ (d) NaCl — Answer: (a) - AgBr dissolving in hypo forms:
(a) Silver sulphide (b) Sodium argentothiosulphate (c) Silver oxide (d) No reaction — Answer: (b) - Silver halides other than AgF are generally:
(a) Freely soluble (b) Poorly soluble/insoluble (c) Explosive (d) Coloured only in solution — Answer: (b) - AgNO₃ is also historically called:
(a) Blue vitriol (b) Lunar caustic (c) Green vitriol (d) Calomel — Answer: (b) - Heating AgNO₃ just above its melting point gives:
(a) Ag metal directly (b) AgNO₂ and O₂ (c) Ag₂O (d) AgNO₃ remains unchanged — Answer: (b) - Heating AgNO₃ to red heat gives:
(a) AgNO₂ and O₂ (b) Ag, NO₂, and O₂ (c) Ag₂O only (d) No decomposition — Answer: (b) - Tollen's reagent is used to test for:
(a) Ketones (b) The aldehyde group (c) Alcohols (d) Carboxylic acids — Answer: (b) - AgNO₃ with sodium thiosulphate initially gives a white precipitate that turns black due to:
(a) Ag₂S formation (b) AgCl formation (c) Ag metal deposition from light (d) No colour change occurs — Answer: (a) - The black stain AgNO₃ leaves on skin is due to:
(a) Chemical burn (b) Reduction to metallic silver (c) Allergic reaction (d) Oxidation of skin proteins — Answer: (b)
