Chapter in a nutshell: Sulphuric acid (H₂SO₄), the "King of Chemicals" or "Oil of Vitriol", is a strong dibasic (diprotic) acid manufactured industrially by the Contact Process. In its dilute form it shows all the typical reactions of an acid (with metals, oxides, hydroxides, carbonates, sulphites, sulphides), while in its concentrated form it shows three special characters — non-volatile (least volatile) acid, powerful dehydrating agent, and oxidising agent. It is identified by the white, acid-insoluble precipitate it gives with barium chloride or lead nitrate solution.
1. Introduction, Occurrence and Names
Sulphuric acid is the most important industrial chemical, earning the title "King of Chemicals" because of its vast use in the manufacture of fertilisers, paints, dyes, drugs, detergents, explosives and other acids. A nation's industrial strength was once judged by its sulphuric acid consumption.
- Molecular formula: H₂SO₄ | Molar mass: 98 g/mol (2×1 + 32 + 4×16)
- Common name: Oil of Vitriol
- It was historically prepared by the distillation of green vitriol, the hydrated ferrous sulphate crystal FeSO₄·7H₂O (this is why Glauber's acid was called "oil of vitriol"):
$$\mathrm{2FeSO_4 \cdot 7H_2O \xrightarrow{\;\Delta\;} Fe_2O_3 + SO_2\uparrow + 13H_2O + H_2SO_4}$$
Green vitriol = hydrated ferrous sulphate (FeSO₄·7H₂O). Blue vitriol / blue stone = hydrated copper sulphate (CuSO₄·5H₂O). White vitriol = hydrated zinc sulphate (ZnSO₄·7H₂O).
2. Manufacture — The Contact Process
Sulphuric acid is manufactured on a large scale by the Contact Process. The name comes from the catalytic step in which gases come in contact with a solid catalyst. There are four stages:
Stage 1 — Production of sulphur dioxide (the "Sulphur/Pyrite Burner")
SO₂ is obtained by burning sulphur in air, or by roasting iron pyrites (FeS₂):$$\mathrm{S + O_2 \longrightarrow SO_2}$$ $$\mathrm{4FeS_2 + 11O_2 \longrightarrow 2Fe_2O_3 + 8SO_2\uparrow}$$
Stage 2 — Purification of SO₂
The gases entering the catalytic chamber must be pure, especially free from arsenic impurities, because arsenic poisons (deactivates) the catalyst and reduces its efficiency. This is why great care is taken to purify the SO₂ and free it from dust and arsenious oxide.Stage 3 — Catalytic oxidation of SO₂ to SO₃ (Catalytic Chamber / Contact Tower)
The purified SO₂ is mixed with excess air and passed over the catalyst:$$\mathrm{2SO_2 + O_2 \;\underset{V_2O_5}{\xrightarrow{\;450^\circ C\;}}\; 2SO_3 \quad (\text{exothermic})}$$
| Condition | Value / detail |
|---|---|
| Catalyst | Vanadium pentoxide, V₂O₅ (or finely divided platinum, now rarely used) |
| Promoter | Potassium oxide (K₂O) — enhances catalyst activity |
| Temperature | ~450 °C |
| Nature of reaction | Exothermic — catalyst glows red hot, so only initial heating is needed |
Stage 4 — Absorption of SO₃ (Absorption Tower) and dilution
SO₃ is not absorbed directly in water. The reasons are crucial:- (i) The reaction of SO₃ with water is highly exothermic.
- (ii) A dense fog/mist of fine sulphuric acid particles forms which does not condense easily and escapes (it can also cause the disease "phosy jaw").
Instead, SO₃ is absorbed in concentrated sulphuric acid to form pyrosulphuric acid / oleum (H₂S₂O₇):
$$\mathrm{SO_3 + H_2SO_4(conc.) \longrightarrow H_2S_2O_7 \;\;(oleum)}$$
The oleum is then diluted with a calculated amount of water to obtain concentrated H₂SO₄ of the desired strength:
$$\mathrm{H_2S_2O_7 + H_2O \longrightarrow 2H_2SO_4}$$
Safety rule: Always add acid to water (slowly, with stirring), never water to acid. Adding water to concentrated acid causes a sudden, violent rise in temperature and the acid spurts out with the steam, causing dangerous burns.
3. Physical Properties
| Property | Description |
|---|---|
| State / appearance | Concentrated H₂SO₄ is a colourless, odourless, oily, viscous liquid |
| Density | Heavy — about 1.84 times that of water |
| Boiling point | Very high (~365 °C) → it is a non-volatile / least-volatile acid |
| Solubility | Miscible with water in all proportions; dissolution is highly exothermic |
| Nature | Hygroscopic (absorbs moisture from air → kept in airtight bottles) |
| Basicity | Dibasic / diprotic — gives two H⁺ per molecule |
| Conductivity | Pure anhydrous acid is a bad (poor) conductor of electricity |
$$\mathrm{H_2SO_4 \rightleftharpoons H^+ + HSO_4^-} \qquad \mathrm{HSO_4^- \rightleftharpoons H^+ + SO_4^{2-}}$$
The presence of H⁺ (hydronium) ions gives the dilute acid its acidic character. Pure (anhydrous) acid has strong covalent bonds and few ions, so it conducts poorly; dilution produces ions and conductivity.
Salts: Being dibasic, it forms two series of salts —
- Normal salt = sulphate (SO₄²⁻): e.g. Na₂SO₄ (formed with excess base)
- Acid salt = bisulphate / hydrogen sulphate (HSO₄⁻): e.g. NaHSO₄ (formed with limited base)
$$\mathrm{H_2SO_4 + NaOH \longrightarrow NaHSO_4 + H_2O \;\;(acid\;salt)}$$ $$\mathrm{NaHSO_4 + NaOH \longrightarrow Na_2SO_4 + H_2O \;\;(normal\;salt)}$$
The acid salt NaHSO₄ in water gives Na⁺, H⁺ and SO₄²⁻ ions (so its solution is acidic): $$\mathrm{NaHSO_4 \xrightarrow{H_2O} Na^+ + H^+ + SO_4^{2-}}$$
4. Properties of DILUTE Sulphuric Acid (Typical Acid Reactions)
Dilute H₂SO₄ shows all the characteristic reactions of an acid.
(a) With bases / alkalis (neutralisation → salt + water)
$$\mathrm{2KOH + H_2SO_4(dil.) \longrightarrow K_2SO_4 + 2H_2O}$$ $$\mathrm{2NaOH + H_2SO_4(dil.) \longrightarrow Na_2SO_4 + 2H_2O}$$ $$\mathrm{Cu(OH)_2 + H_2SO_4(dil.) \longrightarrow CuSO_4 + 2H_2O}$$ $$\mathrm{2NH_4OH + H_2SO_4(dil.) \longrightarrow (NH_4)_2SO_4 + 2H_2O}$$(b) With metallic oxides (basic oxide → salt + water)
$$\mathrm{CuO + H_2SO_4(dil.) \longrightarrow CuSO_4 + H_2O}$$ $$\mathrm{ZnO + H_2SO_4(dil.) \longrightarrow ZnSO_4 + H_2O}$$ $$\mathrm{MgO + H_2SO_4(dil.) \longrightarrow MgSO_4 + H_2O}$$ $$\mathrm{Al_2O_3 + 3H_2SO_4(dil.) \longrightarrow Al_2(SO_4)_3 + 3H_2O}$$Black-to-blue observation: When dilute acid is added to black copper oxide, the black colour disappears, leaving a blue solution (CuSO₄).
(c) With active metals (above H in activity series → liberate H₂ gas)
$$\mathrm{Zn + H_2SO_4(dil.) \longrightarrow ZnSO_4 + H_2\uparrow}$$ $$\mathrm{Mg + H_2SO_4(dil.) \longrightarrow MgSO_4 + H_2\uparrow}$$ $$\mathrm{Fe + H_2SO_4(dil.) \longrightarrow FeSO_4 + H_2\uparrow}$$(d) With metallic carbonates and bicarbonates (→ liberate CO₂)
$$\mathrm{K_2CO_3 + H_2SO_4(dil.) \longrightarrow K_2SO_4 + H_2O + CO_2\uparrow}$$ $$\mathrm{2KHCO_3 + H_2SO_4(dil.) \longrightarrow K_2SO_4 + 2H_2O + 2CO_2\uparrow}$$ $$\mathrm{Na_2CO_3 + H_2SO_4(dil.) \longrightarrow Na_2SO_4 + H_2O + CO_2\uparrow}$$ $$\mathrm{2NaHCO_3 + H_2SO_4(dil.) \longrightarrow Na_2SO_4 + 2H_2O + 2CO_2\uparrow}$$(e) With metallic sulphites and bisulphites (→ liberate SO₂)
$$\mathrm{K_2SO_3 + H_2SO_4(dil.) \longrightarrow K_2SO_4 + H_2O + SO_2\uparrow}$$ $$\mathrm{2KHSO_3 + H_2SO_4(dil.) \longrightarrow K_2SO_4 + 2H_2O + 2SO_2\uparrow}$$ $$\mathrm{Na_2SO_3 + H_2SO_4(dil.) \longrightarrow Na_2SO_4 + H_2O + SO_2\uparrow}$$ $$\mathrm{2NaHSO_3 + H_2SO_4(dil.) \longrightarrow Na_2SO_4 + 2H_2O + 2SO_2\uparrow}$$(f) With metallic sulphides (→ liberate H₂S, "rotten-egg" smell)
$$\mathrm{ZnS + H_2SO_4(dil.) \longrightarrow ZnSO_4 + H_2S\uparrow}$$ $$\mathrm{FeS + H_2SO_4(dil.) \longrightarrow FeSO_4 + H_2S\uparrow}$$ $$\mathrm{Na_2S + H_2SO_4(dil.) \longrightarrow Na_2SO_4 + H_2S\uparrow}$$(g) Effect on indicators
| Indicator | Colour change in dilute acid |
|---|---|
| Blue litmus | Turns red |
| Methyl orange | Turns red / pink |
| Phenolphthalein | Remains colourless |
5. Properties of CONCENTRATED Sulphuric Acid (Three Special Characters)
Concentrated H₂SO₄ shows three "special" behaviours not seen in the dilute acid.
(A) As a Non-volatile (Least Volatile) Acid
Because of its very high boiling point, hot concentrated H₂SO₄ can displace more volatile acids (HCl, HNO₃) from their salts (chlorides, nitrates) by double decomposition. The volatile acid escapes as vapour, driving the reaction forward.$$\mathrm{NaCl + H_2SO_4(conc.) \xrightarrow{<200^\circ C} NaHSO_4 + HCl\uparrow}$$ $$\mathrm{KNO_3 + H_2SO_4(conc.) \xrightarrow{<200^\circ C} KHSO_4 + HNO_3}$$
This property is used to prepare HCl, HNO₃ and acetic acid in the lab.
(B) As a Dehydrating Agent
A dehydrating agent removes the chemically combined elements of water (hydrogen and oxygen in the ratio 2 : 1) from a compound, causing a permanent chemical change. Concentrated H₂SO₄ has a great affinity for water.| Substance | Equation | Observation |
|---|---|---|
| Blue vitriol CuSO₄·5H₂O | $\mathrm{CuSO_4 \cdot 5H_2O \xrightarrow{conc.H_2SO_4} CuSO_4 + 5H_2O}$ | Blue crystals turn white powder |
| Sugar C₁₂H₂₂O₁₁ | $\mathrm{C_{12}H_{22}O_{11} \xrightarrow{conc.H_2SO_4} 12C + 11H_2O}$ | Black spongy charred mass, frothing, steam |
| Ethanol C₂H₅OH | $\mathrm{C_2H_5OH \xrightarrow[160{-}170^\circ C]{conc.H_2SO_4} C_2H_4 + H_2O}$ | Gives ethene |
| Formic acid HCOOH | $\mathrm{HCOOH \xrightarrow{conc.H_2SO_4} CO + H_2O}$ | Gives CO (one oxide of C) |
| Oxalic acid (COOH)₂ | $\mathrm{(COOH)_2 \xrightarrow{conc.H_2SO_4} CO + CO_2 + H_2O}$ | Gives CO + CO₂ (two oxides of C) |
(C) As an Oxidising Agent
Hot concentrated H₂SO₄ acts as a strong oxidising agent: it oxidises metals and non-metals, and is itself reduced to sulphur dioxide (SO₂) (the S goes from +6 to +4).$$\mathrm{Cu + 2H_2SO_4(conc.) \xrightarrow{\Delta} CuSO_4 + SO_2\uparrow + 2H_2O}$$ $$\mathrm{Zn + 2H_2SO_4(conc.) \xrightarrow{\Delta} ZnSO_4 + SO_2\uparrow + 2H_2O}$$ $$\mathrm{C + 2H_2SO_4(conc.) \longrightarrow CO_2 + 2SO_2\uparrow + 2H_2O}$$ $$\mathrm{S + 2H_2SO_4(conc.) \longrightarrow 3SO_2\uparrow + 2H_2O}$$ $$\mathrm{H_2S + H_2SO_4(conc.) \longrightarrow S\downarrow + SO_2 + 2H_2O}$$
Why conc. H₂SO₄ cannot dry H₂S or HBr: it oxidises them (H₂S → S, HBr → Br₂) instead of merely drying — so it is not a drying agent for these gases. Similarly it cannot dry ammonia (NH₃) because it reacts chemically with it to form ammonium sulphate.
(D) As a Drying Agent
Concentrated H₂SO₄ removes free moisture (uncombined water) from gases and other compounds without any chemical change. It is used to dry gases such as Cl₂, HCl, CO₂, SO₂, O₂ — but not H₂S, NH₃ (it reacts with these).6. Drying Agent vs Dehydrating Agent
| Drying agent | Dehydrating agent | |
|---|---|---|
| Removes | Free moisture from gases/compounds | Chemically combined H and O (ratio 2:1) |
| Type of change | Physical change only | Chemical change |
| Composition of substance | Stays unaltered | Totally altered |
| Example | Drying SO₂, CO₂, Cl₂ | Charring of sugar, blue vitriol → white |
7. Tests / Identification of Sulphuric Acid (Sulphate ion test)
A dilute mineral acid (or a sulphate solution) is identified by the white precipitate, insoluble in all mineral acids, that it gives with the following reagents:
(a) Barium chloride test
$$\mathrm{BaCl_2 + H_2SO_4(dil.) \longrightarrow BaSO_4\downarrow + 2HCl}$$ White precipitate of barium sulphate (BaSO₄), insoluble in dilute HCl/HNO₃.(b) Lead nitrate test
$$\mathrm{Pb(NO_3)_2 + H_2SO_4(dil.) \longrightarrow PbSO_4\downarrow + 2HNO_3}$$ White precipitate of lead sulphate (PbSO₄), insoluble in mineral acids.Distinguishing dilute H₂SO₄ from dilute HCl: Add BaCl₂ — H₂SO₄ gives a white ppt (BaSO₄); HCl gives no reaction (BaCl₂ + HCl → no ppt).
8. Uses of Sulphuric Acid
- Fertilisers — ammonium sulphate, superphosphate of lime (calcium dihydrogen phosphate + gypsum).
- Manufacture of other acids — HCl, HNO₃, acetic acid (non-volatile nature).
- Explosives (TNT, nitroglycerine), paints, dyes, drugs, disinfectants, detergents.
- Lead storage batteries / accumulators (electrolyte).
- Pickling of iron and steel — removes the oxide coating before plating/enamelling.
- Esterification — catalyst for carboxylic acid + alcohol → fruity ester.
Worked Examples
Example 1 — Mole/mass. What mass of zinc reacts completely with dilute H₂SO₄ to give 4.48 L of H₂ at STP? Zn + H₂SO₄ → ZnSO₄ + H₂. Moles H₂ = 4.48/22.4 = 0.2; ratio 1:1 → Zn = 0.2 mol → mass = 0.2 × 65 = 13 g.
Example 2 — Acid required. Mass of pure H₂SO₄ to neutralise 8 g NaOH? 2NaOH + H₂SO₄ → Na₂SO₄ + 2H₂O. Moles NaOH = 8/40 = 0.2; ratio 2:1 → H₂SO₄ = 0.1 mol → mass = 0.1 × 98 = 9.8 g.
Example 3 — Identify the role. In S + 2H₂SO₄ → 3SO₂ + 2H₂O, what role does H₂SO₄ play? S goes 0 → +4 (oxidised), while H₂SO₄ is reduced (S: +6 → +4); so H₂SO₄ acts as an oxidising agent.
Example 4 — Gas produced. Gas evolved when dilute H₂SO₄ is added to (i) Na₂CO₃, (ii) FeS, (iii) Na₂SO₃? (i) CO₂ (turns limewater milky); (ii) H₂S (rotten-egg smell, blackens lead acetate paper); (iii) SO₂ (turns acidified K₂Cr₂O₇ green).
Example 5 — CO₂ volume. Volume of CO₂ at STP from 10.6 g Na₂CO₃ + excess dilute H₂SO₄? Na₂CO₃ + H₂SO₄ → Na₂SO₄ + H₂O + CO₂. Moles = 10.6/106 = 0.1 → CO₂ = 0.1 mol → 2.24 L.
Example 6 — Dehydration. A blue crystal added drop-wise with conc. H₂SO₄. Identify and explain. It is blue vitriol, CuSO₄·5H₂O: CuSO₄·5H₂O → CuSO₄ + 5H₂O. Blue crystals lose water of crystallisation → white anhydrous CuSO₄; acid acts as a dehydrating agent.
Example 7 — Distinguishing gases. Why can't CO₂ and SO₂ be told apart by limewater? Both turn limewater milky. Distinguish using acidified K₂Cr₂O₇ — SO₂ turns it green (reducing action); CO₂ gives no change.
Key Terms — Quick Glossary
| Term | Meaning |
|---|---|
| King of chemicals | Sulphuric acid (vast industrial importance) |
| Oil of vitriol | Old common name for concentrated H₂SO₄ |
| Green vitriol | Hydrated ferrous sulphate, FeSO₄·7H₂O |
| Blue vitriol | Hydrated copper sulphate, CuSO₄·5H₂O |
| Contact process | Industrial manufacture of H₂SO₄ |
| Oleum / pyrosulphuric acid | H₂S₂O₇, formed by SO₃ + conc. H₂SO₄ |
| Catalyst (contact process) | Vanadium pentoxide, V₂O₅ (or Pt) |
| Promoter | K₂O — boosts catalyst activity |
| Dibasic acid | Acid giving two replaceable H⁺ ions |
| Normal salt | Sulphate, e.g. Na₂SO₄ |
| Acid salt | Bisulphate/hydrogen sulphate, e.g. NaHSO₄ |
| Dehydrating agent | Removes combined H and O (2:1) chemically |
Common Mistakes to Avoid
- Confusing drying with dehydrating. Drying removes free moisture (physical change); dehydrating removes chemically combined water (chemical change). Conc. H₂SO₄ does both — but the cause of charring/CuSO₄ whitening is dehydration, not drying.
- Writing SO₃ + H₂O → H₂SO₄ as the industrial step. SO₃ is never absorbed directly in water; it is absorbed in conc. H₂SO₄ to form oleum, which is then diluted.
- Forgetting conc. H₂SO₄ oxidises H₂S, HBr, NH₃ — so it cannot be used to dry these gases.
- Using dilute acid with metals to "show oxidising power". Dilute acid liberates H₂ (typical acid reaction); only the concentrated, hot acid acts as an oxidising agent (liberating SO₂, not H₂).
- Unbalanced equations. Watch coefficients: Cu + 2H₂SO₄(conc.); S + 2H₂SO₄ → 3SO₂ + 2H₂O; C + 2H₂SO₄ → CO₂ + 2SO₂ + 2H₂O.
- Saying BaSO₄ dissolves in acid. The white precipitate (BaSO₄/PbSO₄) is insoluble in all mineral acids — that is exactly why it is the confirmatory test.
- Adding water to acid. Always add acid to water; never the reverse (spurting/burns).
Likely Exam Questions (with crisp answers)
- Why is H₂SO₄ called the "King of Chemicals"? Because of its extensive use in manufacturing fertilisers, acids, explosives, dyes, drugs and many industrial processes.
- Name and write the formula of green vitriol. Hydrated ferrous sulphate, FeSO₄·7H₂O.
- Name the process and catalyst for manufacturing H₂SO₄. Contact process; catalyst = vanadium pentoxide (V₂O₅) [or platinum].
- What is the promoter in the contact process? Potassium oxide (K₂O).
- Why must SO₂ be purified before the catalytic chamber? Arsenic impurities poison the catalyst and lower its efficiency.
- Why is SO₃ not dissolved directly in water? The reaction is highly exothermic and forms a dense, non-condensing acid mist; instead SO₃ is absorbed in conc. H₂SO₄ to form oleum.
- What is oleum? Give its formula. Pyrosulphuric acid, H₂S₂O₇ (SO₃ + conc. H₂SO₄).
- Why is H₂SO₄ a dibasic acid? It ionises to give two replaceable H⁺ ions and forms two salts — bisulphate (NaHSO₄) and sulphate (Na₂SO₄).
- Why is conc. H₂SO₄ kept in airtight bottles? It is hygroscopic — it absorbs moisture from the air and gets diluted.
- Why is conc. H₂SO₄ not a drying agent for H₂S? It oxidises H₂S to sulphur: H₂S + H₂SO₄ → S↓ + SO₂ + 2H₂O.
- Which property allows H₂SO₄ to prepare HCl and HNO₃? Its non-volatile (least volatile) nature.
- What happens when conc. H₂SO₄ is added to sugar? A black spongy/charred mass of carbon forms (dehydrating agent): C₁₂H₂₂O₁₁ → 12C + 11H₂O.
- What happens to blue vitriol with conc. H₂SO₄? Blue crystals turn to white anhydrous CuSO₄ (dehydration).
- Give the test for sulphuric acid / sulphate ions. White precipitate (BaSO₄ or PbSO₄), insoluble in mineral acids, with BaCl₂ or Pb(NO₃)₂ solution.
- Which organic acid on dehydration gives both CO and CO₂? Oxalic acid: (COOH)₂ → CO + CO₂ + H₂O.
- Name the gas evolved when dilute H₂SO₄ reacts with a sulphite. Sulphur dioxide (SO₂).
- Why does dilute H₂SO₄ conduct electricity but pure acid does not? Dilution ionises the acid into H⁺ and SO₄²⁻ ions; pure anhydrous acid has few ions (mostly covalent).
- Write the equation for conc. H₂SO₄ acting as an oxidising agent on copper. Cu + 2H₂SO₄(conc.) → CuSO₄ + SO₂↑ + 2H₂O.
- Why are large quantities of H₂SO₄ used in the iron and steel industry? For pickling — to remove the oxide coating from the metal surface before plating/enamelling.
- What is the basicity of sulphuric acid? Two (dibasic / diprotic).
Chemical Equations & Formulas (quick reference)
Key formulas: H₂SO₄ (sulphuric acid, M = 98) · SO₂ · SO₃ · H₂S₂O₇ (oleum) · FeSO₄·7H₂O (green vitriol) · CuSO₄·5H₂O (blue vitriol) · ZnSO₄·7H₂O (white vitriol) · V₂O₅ (catalyst) · K₂O (promoter) · BaSO₄ / PbSO₄ (white test precipitates) · Na₂SO₄ (normal salt) · NaHSO₄ (acid salt).
Preparation from green vitriol
- 2FeSO₄·7H₂O → Fe₂O₃ + SO₂↑ + 13H₂O + H₂SO₄
Contact process
- S + O₂ → SO₂
- 4FeS₂ + 11O₂ → 2Fe₂O₃ + 8SO₂↑
- 2SO₂ + O₂ ⇌(V₂O₅, 450°C) 2SO₃
- SO₃ + H₂SO₄(conc.) → H₂S₂O₇ (oleum)
- H₂S₂O₇ + H₂O → 2H₂SO₄
Dilute acid — bases / oxides
- 2NaOH + H₂SO₄ → Na₂SO₄ + 2H₂O
- 2KOH + H₂SO₄ → K₂SO₄ + 2H₂O
- Cu(OH)₂ + H₂SO₄ → CuSO₄ + 2H₂O
- 2NH₄OH + H₂SO₄ → (NH₄)₂SO₄ + 2H₂O
- CuO + H₂SO₄ → CuSO₄ + H₂O
- ZnO + H₂SO₄ → ZnSO₄ + H₂O
- MgO + H₂SO₄ → MgSO₄ + H₂O
- Al₂O₃ + 3H₂SO₄ → Al₂(SO₄)₃ + 3H₂O
Dilute acid — metals (H₂)
- Zn + H₂SO₄ → ZnSO₄ + H₂↑
- Mg + H₂SO₄ → MgSO₄ + H₂↑
- Fe + H₂SO₄ → FeSO₄ + H₂↑
Dilute acid — carbonates / bicarbonates (CO₂)
- Na₂CO₃ + H₂SO₄ → Na₂SO₄ + H₂O + CO₂↑
- 2NaHCO₃ + H₂SO₄ → Na₂SO₄ + 2H₂O + 2CO₂↑
- K₂CO₃ + H₂SO₄ → K₂SO₄ + H₂O + CO₂↑
- 2KHCO₃ + H₂SO₄ → K₂SO₄ + 2H₂O + 2CO₂↑
Dilute acid — sulphites / bisulphites (SO₂)
- Na₂SO₃ + H₂SO₄ → Na₂SO₄ + H₂O + SO₂↑
- 2NaHSO₃ + H₂SO₄ → Na₂SO₄ + 2H₂O + 2SO₂↑
- K₂SO₃ + H₂SO₄ → K₂SO₄ + H₂O + SO₂↑
Dilute acid — sulphides (H₂S)
- ZnS + H₂SO₄ → ZnSO₄ + H₂S↑
- FeS + H₂SO₄ → FeSO₄ + H₂S↑
- Na₂S + H₂SO₄ → Na₂SO₄ + H₂S↑
Conc. acid — non-volatile
- NaCl + H₂SO₄(conc.) →(<200°C) NaHSO₄ + HCl↑
- KNO₃ + H₂SO₄(conc.) →(<200°C) KHSO₄ + HNO₃
Conc. acid — dehydrating
- CuSO₄·5H₂O → CuSO₄ + 5H₂O
- C₁₂H₂₂O₁₁ → 12C + 11H₂O
- C₂H₅OH →(160–170°C) C₂H₄ + H₂O
- HCOOH → CO + H₂O
- (COOH)₂ → CO + CO₂ + H₂O
Conc. acid — oxidising (reduced to SO₂)
- Cu + 2H₂SO₄(conc.) → CuSO₄ + SO₂↑ + 2H₂O
- Zn + 2H₂SO₄(conc.) → ZnSO₄ + SO₂↑ + 2H₂O
- C + 2H₂SO₄(conc.) → CO₂ + 2SO₂↑ + 2H₂O
- S + 2H₂SO₄(conc.) → 3SO₂↑ + 2H₂O
- H₂S + H₂SO₄(conc.) → S↓ + SO₂ + 2H₂O
Salt formation (dibasic)
- H₂SO₄ + NaOH → NaHSO₄ + H₂O (acid salt)
- NaHSO₄ + NaOH → Na₂SO₄ + H₂O (normal salt)
Tests
- BaCl₂ + H₂SO₄ → BaSO₄↓ (white) + 2HCl
- Pb(NO₃)₂ + H₂SO₄ → PbSO₄↓ (white) + 2HNO₃