ICSE Class 10 Study of Compounds — HCl, Ammonia, Nitric Acid and Sulphuric Acid
"Study of Compounds" is really four short chapters wearing one name: hydrogen chloride, ammonia, nitric acid and sulphuric acid. Each one is examined in the same shape — how it is prepared in the laboratory, how it is dried and collected, what it reacts with, and how you prove it is that substance. Once you see the shape, the chapter stops being a memory dump. This guide takes the four compounds in that order, with every equation balanced.
Hydrogen chloride
Laboratory preparation. Sodium chloride is warmed with concentrated sulphuric acid below 200 °C:
2NaCl + H2SO4 → Na2SO4 + 2HCl↑ (above 200 °C)
The lower temperature is preferred because the sodium sulphate formed at the higher temperature sticks to the flask and is difficult to remove, and the hard-glass vessel may crack. Concentrated sulphuric acid works here because it is a non-volatile acid — it displaces the volatile acid HCl from its salt.
Drying and collection. The gas is dried by passing it through concentrated sulphuric acid. Quicklime cannot be used, because it is basic and would simply react with the acidic gas: CaO + 2HCl → CaCl2 + H2O. Hydrogen chloride is denser than air and extremely soluble in water, so it is collected by upward displacement of air (downward delivery) and never over water.
Key properties.
- Colourless gas with a sharp, choking smell; fumes strongly in moist air.
- Dry HCl gas does not turn dry blue litmus red. It has to dissolve in water first so that it can ionise into H+ and Cl−. This is one of the most reliable one-mark questions in the chapter.
- Its extreme solubility is demonstrated by the fountain experiment: a jar of the gas inverted over water sucks the water up in a jet as the gas dissolves.
- With ammonia gas it gives dense white fumes: NH3 + HCl → NH4Cl.
Tests for a chloride. Add silver nitrate solution: a curdy white precipitate forms, insoluble in dilute nitric acid but soluble in excess ammonium hydroxide.
2HCl + Pb(NO3)2 → PbCl2↓ + 2HNO3
The lead chloride precipitate is white and dissolves in hot water — a neat way to tell it apart from silver chloride, which does not.
Aqua regia is a mixture of three parts concentrated hydrochloric acid to one part concentrated nitric acid, and it dissolves gold and platinum.
Ammonia
Laboratory preparation. Any ammonium salt is warmed with slaked lime:
Drying and collection. Ammonia is dried over quicklime, CaO. It cannot be dried over concentrated sulphuric acid, which would react with it (2NH3 + H2SO4 → (NH4)2SO4), and it cannot be dried over anhydrous calcium chloride, which forms the addition compound CaCl2·8NH3. Ammonia is lighter than air and very soluble, so it is collected by downward displacement of air (upward delivery).
Manufacture — Haber process.
The conditions are a high pressure, a moderately high temperature of roughly 450–500 °C, and finely divided iron as catalyst with molybdenum as promoter. Different textbook editions print slightly different pressures and temperatures — use the values your own book prints, because that is what your marking scheme will follow.
Key properties.
- It is the only common alkaline gas — it turns moist red litmus blue. That single fact identifies it.
- With hydrogen chloride it gives dense white fumes of ammonium chloride (the "glass rod dipped in concentrated HCl" test).
- It burns in oxygen: 4NH3 + 3O2 → 2N2 + 6H2O.
- Catalytic oxidation over hot platinum gives nitric oxide — the first step of the Ostwald process: 4NH3 + 5O2 → 4NO + 6H2O.
Ammonium hydroxide as a laboratory reagent. Adding it drop by drop, then in excess, identifies metal ions by the colour of the hydroxide precipitate:
| Solution | With a little NH4OH | With excess NH4OH |
|---|---|---|
| FeCl3 (Fe3+) | Reddish-brown precipitate | Insoluble |
| FeSO4 (Fe2+) | Dirty green precipitate | Insoluble |
| CuSO4 (Cu2+) | Pale blue precipitate | Dissolves to a deep inky blue solution |
| ZnSO4 (Zn2+) | White precipitate | Dissolves — colourless solution |
| Pb(NO3)2 (Pb2+) | White precipitate | Insoluble |
Nitric acid
Laboratory preparation. A nitrate is heated with concentrated sulphuric acid below 200 °C in an all-glass apparatus, because hot nitric acid vapour attacks cork and rubber:
Manufacture — Ostwald process, in three stages:
2NO + O2 → 2NO2
4NO2 + O2 + 2H2O → 4HNO3
Why concentrated nitric acid turns yellow on standing. It decomposes slightly in light and heat, and the nitrogen dioxide produced dissolves in the acid:
Bubbling air or carbon dioxide through the acid drives the dissolved NO2 out and restores the colourless liquid.
Nitric acid is both an acid and a strong oxidising agent, which is why it normally does not give hydrogen with metals — the hydrogen produced is immediately oxidised to water. With copper the product depends on concentration:
Cu + 4HNO3(concentrated) → Cu(NO3)2 + 2NO2 + 2H2O
Textbooks note magnesium and manganese with very dilute nitric acid as the exceptions that do release hydrogen. It also oxidises non-metals: C + 4HNO3 → CO2 + 4NO2 + 2H2O and S + 6HNO3 → H2SO4 + 6NO2 + 2H2O.
Test for a nitrate — the brown ring test. Add freshly prepared iron(II) sulphate solution to the nitrate solution, then pour concentrated sulphuric acid gently down the side of the tube. A brown ring appears where the two layers meet.
Sulphuric acid
Manufacture — Contact process, in three stages:
2SO2 + O2 ⇌ 2SO3 (V2O5 catalyst, exothermic)
SO3 + H2SO4 → H2S2O7 (oleum)
H2S2O7 + H2O → 2H2SO4
Why is sulphur trioxide not dissolved directly in water? Because the reaction is violently exothermic and produces a dense mist of acid droplets that is very hard to condense. Absorbing it in concentrated acid to make oleum, then diluting, avoids the mist entirely. Expect this as a two-mark reasoning question.
Concentrated sulphuric acid behaves in three separate roles, and ICSE examiners like to test that you can name which role is at work:
| Role | What happens | Example |
|---|---|---|
| Dehydrating agent | Removes hydrogen and oxygen as water from a compound that never contained liquid water | C12H22O11 → 12C + 11H2O (sugar chars to black carbon) |
| Drying / removing water of crystallisation | Takes away water already present | CuSO4·5H2O → CuSO4 + 5H2O (blue to white) |
| Oxidising agent (hot and concentrated) | Itself reduced to SO2 | Cu + 2H2SO4 → CuSO4 + SO2 + 2H2O |
| Non-volatile acid | Displaces a more volatile acid from its salt | Its use in preparing HCl and HNO3 above |
Test for a sulphate. Barium chloride solution gives a white precipitate of barium sulphate that is insoluble in dilute hydrochloric acid: Na2SO4 + BaCl2 → BaSO4↓ + 2NaCl.
Safety rule you must be able to state: always add the concentrated acid to water slowly with stirring, never water to the acid. Dilution releases a great deal of heat, and adding water to acid can make it spit.
Worked example 1 — mass and volume of ammonia
Q. 10.7 g of ammonium chloride is heated with excess slaked lime. What mass of ammonia is produced, and what volume does it occupy at S.T.P.? (N = 14.007, H = 1.008, Cl = 35.45)
2NH4Cl + Ca(OH)2 → CaCl2 + 2NH3 + 2H2O, so 2 mol NH4Cl give 2 mol NH3 — a 1 : 1 ratio.
M(NH4Cl) = 14.007 + (4 × 1.008) + 35.45 = 14.007 + 4.032 + 35.45 =
53.489 g/mol
Moles of NH4Cl = 10.7 ÷ 53.489 = 0.2000 mol
Moles of NH3 = 0.2000 mol
M(NH3) = 14.007 + (3 × 1.008) = 14.007 + 3.024 = 17.031 g/mol
Mass of NH3 = 0.2000 × 17.031 = 3.41 g
Volume = 0.2000 × 22.4 = 4.48 L using the older S.T.P. molar volume of 22.4 L/mol (273.15 K, 1 atm), which is what ICSE papers use. If your book defines S.T.P. at 1 bar, the molar volume is 22.7 L/mol and the answer is 4.54 L. State which convention you are using and you cannot be marked wrong.
Worked example 2 — hydrogen chloride from common salt
Q. 5.844 g of sodium chloride is warmed with concentrated sulphuric acid below 200 °C. Calculate the mass and the S.T.P. volume of hydrogen chloride formed. (Na = 22.990, Cl = 35.45, H = 1.008)
NaCl + H2SO4 → NaHSO4 + HCl — a 1 : 1 ratio.
M(NaCl) = 22.990 + 35.45 = 58.440 g/mol
Moles of NaCl = 5.844 ÷ 58.440 = 0.1000 mol
M(HCl) = 1.008 + 35.45 = 36.458 g/mol
Mass of HCl = 0.1000 × 36.458 = 3.65 g
Volume at S.T.P. = 0.1000 × 22.4 = 2.24 L (2.27 L on the 1-bar convention)
Worked example 3 — copper and concentrated nitric acid
Q. What mass of copper is needed to produce 1.12 L of nitrogen dioxide at S.T.P. with concentrated nitric acid? (Cu = 63.546)
Cu + 4HNO3 → Cu(NO3)2 + 2NO2 + 2H2O
Moles of NO2 = 1.12 ÷ 22.4 = 0.0500 mol
From the equation, 2 mol NO2 need 1 mol Cu, so moles of Cu = 0.0500 ÷ 2 =
0.0250 mol
Mass of Cu = 0.0250 × 63.546 = 1.59 g
Worked example 4 — identify the gas
Q. A colourless gas X turns moist red litmus paper blue and gives dense white fumes when a glass rod dipped in concentrated hydrochloric acid is held in it. Name X, give the equation for the white fumes, and state how X is dried.
X is ammonia, NH3 — it is the only common alkaline gas, so the litmus test alone identifies it.
NH3 + HCl → NH4Cl (the white fumes are solid ammonium chloride).
It is dried over quicklime; not over concentrated sulphuric acid or anhydrous calcium chloride, both of which react with it.
Mistakes that cost marks
- Saying dry HCl gas turns blue litmus red. It does not. The gas must dissolve in water and ionise first.
- Drying ammonia over concentrated sulphuric acid. Ammonia is basic; the acid would absorb it completely. Quicklime is the answer.
- Confusing "dehydrating" with "drying". Dehydration pulls hydrogen and oxygen out of a compound that had no water in it (sugar → carbon). Drying removes water that is already there.
- Writing that nitric acid gives hydrogen with metals. As an oxidising acid it normally gives NO or NO2 instead; only very dilute acid with Mg or Mn is quoted as an exception.
- Dissolving SO3 in water in the Contact process. It forms an uncondensable acid mist. Absorb in concentrated acid to make oleum, then dilute.
- Adding water to concentrated sulphuric acid. Always acid into water, slowly, with stirring.
- Not balancing the ammonia preparation. It is 2NH4Cl + Ca(OH)2, not NH4Cl + Ca(OH)2.
Where this chapter is examined
| Board / exam | How it appears |
|---|---|
| ICSE Class 10 | Named chapters. Expect labelled-diagram descriptions in words, equation writing, drying/collection reasoning and the identification tests |
| CBSE Class 10 | The same acids and their tests appear inside "Acids, Bases and Salts" rather than as separate chapters |
| Class 11–12 | Haber, Ostwald and Contact processes return with equilibrium and Le Chatelier reasoning attached |
| Competitive papers | Gas identification and the reagent-colour table are standard inorganic questions |
Qualitatively this is one of the most reliably scoring chapters in ICSE chemistry: the answers are fixed, so accurate recall of equations and conditions is rewarded directly. Write every equation balanced and state conditions above the arrow.
Balance and check every equation in this chapter. The Chemical Equation Balancer takes an unbalanced skeleton — for example Cu + HNO3 = Cu(NO3)2 + NO2 + H2O — and returns the balanced form, so you can verify each equation above before you commit it to memory.
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