Hydrogen Chloride (HCl)

Chapter in a nutshell: Hydrogen chloride (HCl) is a colourless, pungent-smelling, polar covalent gas that is about 1.28 times heavier than air and extremely soluble in water. In the lab it is made by warming sodium chloride (common salt) with concentrated sulphuric acid below 200 °C; its aqueous solution, hydrochloric acid, is a strong monobasic acid that reacts with metals, oxides, hydroxides, carbonates, sulphides, sulphites and thiosulphates, and can be oxidised to chlorine by MnO₂, PbO₂ and red lead.

1. The Hydrogen Chloride Molecule

• Glauber first prepared the acid (1648); Davy named it hydrochloric acid (1810); Lavoisier called it muriatic acid. Free HCl occurs in volcanic emissions; hydrochloric acid (0.2–0.4 %) is present in the gastric juice of mammals, aiding digestion.

2. General Methods of Preparation

(a) By synthesis (direct combination): Moist hydrogen burns in chlorine in the presence of diffused sunlight:

$$\mathrm{H_2(g) + Cl_2(g) \xrightarrow{diffused\ sunlight} 2HCl(g)}$$

Note: The reaction is explosive in direct/bright sunlight and negligible in the dark. With a catalyst such as activated charcoal, it proceeds even in the dark. A burning jet of hydrogen also continues to burn in chlorine, forming HCl.

(b) By heating a metallic chloride with conc. H₂SO₄:

3. Laboratory Preparation of Hydrogen Chloride Gas

Reactants: sodium chloride (common salt) + concentrated sulphuric acid.

Sodium chloride is preferred over other metallic chlorides because it is cheap and easily available. Conc. nitric acid is NOT used in place of H₂SO₄ because it is volatile and would escape along with the HCl gas.

Apparatus & procedure (described): Common salt is placed in a round-bottomed flask; conc. sulphuric acid is poured in through a thistle funnel dipping into the acid. The flask is heated gently. The reaction is slow in the cold; on warming, HCl gas is steadily evolved. The gas passes through a wash bottle of conc. H₂SO₄ to dry it, then into a gas jar by upward displacement of air.

Reactions (note the temperature control):

$$\mathrm{NaCl + H_2SO_4(conc.) \xrightarrow{below\ 200^\circ C} NaHSO_4 + HCl\uparrow}$$

$$\mathrm{NaHSO_4 + NaCl \xrightarrow{above\ 200^\circ C} Na_2SO_4 + HCl\uparrow}$$

Overall above 200 °C:

$$\mathrm{2NaCl + H_2SO_4(conc.) \xrightarrow{above\ 200^\circ C} Na_2SO_4 + 2HCl\uparrow}$$

The formation of sodium hydrogen sulphate (NaHSO₄) at low temperature shows the acid nature of NaHSO₄. Although reversible, the reaction goes to completion because HCl continuously escapes as a gas.

Why the temperature is kept below ~200 °C:

• Phosphorus pentoxide (P₂O₅): $\mathrm{2P_2O_5 + 3HCl \rightarrow POCl_3 + 3HPO_3}$
• Quick lime (CaO): $\mathrm{CaO + 2HCl \rightarrow CaCl_2 + H_2O}$

Both react chemically with HCl, so they cannot be used as drying agents.

Collection: HCl is collected by upward displacement of air (downward delivery) because it is heavier than air (≈1.28×). It is NOT collected over water because it is extremely soluble in water.

Identification of the full jar: Bring a glass rod dipped in ammonium hydroxide (ammonia) near the mouth of the jar. Dense white fumes of ammonium chloride confirm HCl:

$$\mathrm{NH_3 + HCl \rightarrow NH_4Cl}\quad(\text{or } \mathrm{HCl + NH_4OH \rightarrow NH_4Cl + H_2O})$$

4. Physical Properties

Density experiment: HCl poured into a jar holding a burning candle extinguishes the flame — the heavier gas displaces air and occupies the lower part of the jar, proving HCl is heavier than air.

Fountain experiment (proves extreme solubility): A dry flask filled with dry HCl has a two-holed stopper carrying a long jet tube (dipping into blue litmus solution) and a dropper of water. Pressing the dropper lets in a little water; HCl dissolves rapidly, sharply lowering the inside pressure. The higher outside pressure pushes the litmus up the jet tube as a red fountain — red because the solution is acidic, proving HCl is highly soluble in water.

5. Chemical Properties of Hydrogen Chloride Gas

1. Combustibility: Neither combustible nor a supporter of combustion — it extinguishes a burning splint.

2. Thermal dissociation: Above 500 °C it dissociates:

$$\mathrm{2HCl \xrightarrow{>500^\circ C} H_2 + Cl_2}$$

3. Action on metals: Metals above hydrogen in the activity series displace hydrogen when heated with HCl gas:

$$\mathrm{Mg + 2HCl \xrightarrow{\Delta} MgCl_2 + H_2}\qquad \mathrm{Zn + 2HCl \xrightarrow{\Delta} ZnCl_2 + H_2}$$ $$\mathrm{Fe + 2HCl \xrightarrow{\Delta} FeCl_2 + H_2}\qquad \mathrm{Ca + 2HCl \xrightarrow{\Delta} CaCl_2 + H_2}$$ $$\mathrm{2Na + 2HCl \rightarrow 2NaCl + H_2}$$

4. Reaction with ammonia: Two gases combine to a solid, giving dense white fumes of ammonium chloride:

$$\mathrm{NH_3(g) + HCl(g) \rightarrow NH_4Cl(s)}$$

5. Dissolution in water: Being polar covalent, HCl ionises in water to give a strong monobasic acid (hydrochloric acid):

$$\mathrm{HCl(g) + H_2O \rightleftharpoons H_3O^+ + Cl^-}$$

6. Effect on litmus (dry vs moist): Dry HCl gas and liquefied HCl have no effect on dry litmus (non-acidic, no free ions). Only in water does it turn blue litmus red, showing acids show their properties only in the presence of water.

7. Non-conduction: Liquefied HCl and HCl dissolved in toluene do not conduct electricity and do not turn litmus red — there is no H₃O⁺ — proving HCl is a covalent compound.

6. Hydrochloric Acid — Preparation & the Funnel Arrangement

The aqueous solution of HCl gas is hydrochloric acid. It is prepared by dissolving HCl gas in water until saturated; the concentrated acid contains about 36 % HCl by mass.

Because HCl is so soluble, passing the gas directly into water makes water rush up the delivery tube — back suction — which could crack the hot generating flask. To prevent this, an inverted funnel is fixed so its rim just touches the water surface.

Mechanism — how back suction is avoided: As gas dissolves, the funnel sucks water up inside it; the outside level falls, creating an air gap; pressures equalise and the water drops back. The funnel (a) prevents back suction into the flask and (b) gives a large surface area for absorption. An empty (anti-suction) flask between the generating flask and trough collects any sucked-back water.

Constant boiling mixture (azeotrope): On distillation the acid concentrates to 22.2 % HCl by mass, boiling at a constant 110 °C with no further change in composition — an azeotrope. Hence dilute HCl cannot be concentrated to pure HCl by distillation.

7. Chemical Properties of Hydrochloric Acid

Dilute HCl is a typical strong acid and shows all acidic reactions.

(a) Action on indicators:

$$\mathrm{ZnO + 2HCl \rightarrow ZnCl_2 + H_2O}\qquad \mathrm{CuO + 2HCl \rightarrow CuCl_2 + H_2O}$$ $$\mathrm{MgO + 2HCl \rightarrow MgCl_2 + H_2O}\qquad \mathrm{Fe_2O_3 + 6HCl \rightarrow 2FeCl_3 + 3H_2O}$$

(c) With hydroxides (neutralisation) → salt + water:

$$\mathrm{NaOH + HCl \rightarrow NaCl + H_2O}\qquad \mathrm{KOH + HCl \rightarrow KCl + H_2O}$$ $$\mathrm{Ca(OH)_2 + 2HCl \rightarrow CaCl_2 + 2H_2O}$$

(d) With active metals → metal chloride + hydrogen (colourless, odourless gas burning with a pop):

$$\mathrm{Mg + 2HCl \rightarrow MgCl_2 + H_2\uparrow}\qquad \mathrm{Zn + 2HCl \rightarrow ZnCl_2 + H_2\uparrow}$$ $$\mathrm{Fe + 2HCl \rightarrow FeCl_2 + H_2\uparrow}\qquad \mathrm{2Al + 6HCl \rightarrow 2AlCl_3 + 3H_2\uparrow}$$

(e) With carbonates and bicarbonates → CO₂ (turns lime water milky):

$$\mathrm{Na_2CO_3 + 2HCl \rightarrow 2NaCl + H_2O + CO_2\uparrow}$$ $$\mathrm{NaHCO_3 + HCl \rightarrow NaCl + H_2O + CO_2\uparrow}$$ $$\mathrm{CaCO_3 + 2HCl \rightarrow CaCl_2 + H_2O + CO_2\uparrow}$$

Confirmation of CO₂ in lime water: $\mathrm{Ca(OH)_2 + CO_2 \rightarrow CaCO_3\downarrow + H_2O}$ (milky).

(f) With sulphites and bisulphites → SO₂ (burning sulphur smell; turns acidified K₂Cr₂O₇ orange→green and KMnO₄ purple→colourless):

$$\mathrm{Na_2SO_3 + 2HCl \rightarrow 2NaCl + H_2O + SO_2\uparrow}$$ $$\mathrm{NaHSO_3 + HCl \rightarrow NaCl + H_2O + SO_2\uparrow}$$

(g) With sulphides → H₂S (rotten-egg smell; turns lead acetate paper black):

$$\mathrm{Na_2S + 2HCl \rightarrow 2NaCl + H_2S\uparrow}\qquad \mathrm{FeS + 2HCl \rightarrow FeCl_2 + H_2S\uparrow}$$ $$\mathrm{Pb(CH_3COO)_2 + H_2S \rightarrow PbS\downarrow + 2CH_3COOH}\ (\text{black})$$

(h) With thiosulphates → SO₂ + sulphur (yellow turbidity / yellow precipitate):

$$\mathrm{Na_2S_2O_3 + 2HCl \rightarrow 2NaCl + H_2O + SO_2\uparrow + S\downarrow}$$

This distinguishes thiosulphates from sulphites: sulphur (yellow ppt.) appears with thiosulphate but not with a sulphite.

(i) With nitrates: Dilute HCl does not normally react with nitrates, but lead nitrate and mercury(I) nitrate give white precipitates:

$$\mathrm{Pb(NO_3)_2 + 2HCl \rightarrow PbCl_2\downarrow + 2HNO_3}$$ $$\mathrm{Hg_2(NO_3)_2 + 2HCl \rightarrow Hg_2Cl_2\downarrow + 2HNO_3}$$

Lead chloride (PbCl₂) is a white precipitate insoluble in cold water but soluble in hot water.

(j) With silver nitrate → curdy white AgCl (soluble in excess ammonium hydroxide):

$$\mathrm{AgNO_3 + HCl \rightarrow AgCl\downarrow + HNO_3}$$ $$\mathrm{AgCl + 2NH_4OH \rightarrow [Ag(NH_3)_2]Cl + 2H_2O}\ (\text{diammine silver(I) chloride})$$

On exposure to light AgCl darkens: $\mathrm{2AgCl \rightarrow 2Ag + Cl_2}$.

8. Oxidation of Hydrochloric Acid to Chlorine

Concentrated HCl is oxidised to chlorine (a greenish-yellow, pungent gas) by strong oxidising agents — MnO₂, PbO₂, red lead (Pb₃O₄), K₂Cr₂O₇, KMnO₄ and bleaching powder:

$$\mathrm{MnO_2 + 4HCl(conc.) \xrightarrow{\Delta} MnCl_2 + 2H_2O + Cl_2\uparrow}$$ $$\mathrm{PbO_2 + 4HCl(conc.) \xrightarrow{\Delta} PbCl_2 + 2H_2O + Cl_2\uparrow}$$ $$\mathrm{Pb_3O_4 + 8HCl(conc.) \xrightarrow{\Delta} 3PbCl_2 + 4H_2O + Cl_2\uparrow}$$ $$\mathrm{K_2Cr_2O_7 + 14HCl(conc.) \xrightarrow{\Delta} 2KCl + 2CrCl_3 + 7H_2O + 3Cl_2\uparrow}$$ $$\mathrm{2KMnO_4 + 16HCl(conc.) \xrightarrow{\Delta} 2KCl + 2MnCl_2 + 8H_2O + 5Cl_2\uparrow}$$ $$\mathrm{CaOCl_2 + 2HCl(dil.) \rightarrow CaCl_2 + H_2O + Cl_2\uparrow}$$

The chlorine liberated turns moist starch-iodide paper blue-black:

$$\mathrm{2KI + Cl_2 \rightarrow 2KCl + I_2}\qquad \mathrm{I_2 + starch \rightarrow blue\text{-}black}$$

It also bleaches moist litmus: $\mathrm{Cl_2 + H_2O \rightarrow HCl + HClO}$, then $\mathrm{HClO \rightarrow HCl + [O]}$ bleaches the colour.

9. Aqua Regia

Aqua regia ("royal water") is a mixture of 3 parts conc. HCl + 1 part conc. HNO₃ by volume. The HNO₃ oxidises HCl, releasing nascent chlorine [Cl] and nitrosyl chloride (NOCl), which dissolve noble metals like gold and platinum:

$$\mathrm{HNO_3(conc.) + 3HCl(conc.) \rightarrow NOCl + 2[Cl] + 2H_2O}$$ $$\mathrm{Au + 3[Cl] \rightarrow AuCl_3}\ (\text{gold(III) chloride})$$ $$\mathrm{Pt + 4[Cl] \rightarrow PtCl_4}\ (\text{platinum(IV) chloride})$$

10. Uses & Tests

Uses of hydrochloric acid: lab reagent and aqua regia; manufacture of chlorine, ammonium chloride (dry cells), dyes, drugs, paints, AgCl; glucose from starch; pickling of steel before plating; purifying bone black; removing rust; calico printing and tanning; medicinally for low gastric acidity (aids protein digestion).

Tests for HCl gas / hydrochloric acid:

1. Turns moist blue litmus red.
2. Gives dense white fumes with ammonia: $\mathrm{NH_3 + HCl \rightarrow NH_4Cl}$.
3. Gives curdy white AgCl with AgNO₃, soluble in excess NH₄OH but insoluble in dilute HNO₃.
4. Gives white PbCl₂ with lead nitrate, soluble in hot water, insoluble in cold.
5. With conc. HCl + MnO₂ (heat) → greenish-yellow chlorine, turning starch-iodide paper blue-black.

Worked Examples

1. Mass of HCl from 117 g NaCl. $\mathrm{2NaCl + H_2SO_4 \rightarrow Na_2SO_4 + 2HCl}$. Moles NaCl = 117/58.5 = 2 mol → 2 mol HCl = 2 × 36.5 = 73 g HCl.

2. Volume of CO₂ at STP from 10.6 g Na₂CO₃. $\mathrm{Na_2CO_3 + 2HCl \rightarrow 2NaCl + H_2O + CO_2}$. Moles Na₂CO₃ = 10.6/106 = 0.1 → 0.1 mol CO₂ = 0.1 × 22.4 = 2.24 L at STP.

3. Mass of AgCl from excess AgNO₃ with 0.5 mol HCl. $\mathrm{AgNO_3 + HCl \rightarrow AgCl + HNO_3}$. 0.5 mol HCl → 0.5 mol AgCl = 0.5 × 143.5 = 71.75 g AgCl.

4. Volume of H₂ at STP when 6.5 g Zn reacts with excess dil. HCl. $\mathrm{Zn + 2HCl \rightarrow ZnCl_2 + H_2}$. Moles Zn = 6.5/65 = 0.1 → 0.1 mol H₂ = 2.24 L at STP.

5. Moles of HCl to dissolve 1 mol Pb₃O₄. From $\mathrm{Pb_3O_4 + 8HCl \rightarrow 3PbCl_2 + 4H_2O + Cl_2}$, 8 mol HCl per mol Pb₃O₄, releasing 1 mol Cl₂.

6. % HCl in concentrated acid. Concentrated acid is about 36 % HCl by mass; the azeotrope (110 °C) is 22.2 % HCl by mass, so a 1000 g bottle of conc. acid holds about 360 g HCl.

Key Terms — Quick Glossary

Common Mistakes to Avoid

• Writing HCl as ionic — it is polar covalent; it only forms ions in water.
• Saying dry HCl turns blue litmus red — only moist/aqueous HCl does.
• Using conc. HNO₃ instead of H₂SO₄ for lab preparation — HNO₃ is volatile and escapes.
• Trying to dry HCl with P₂O₅ or CaO — both react with HCl; use conc. H₂SO₄.
• Collecting HCl over water or by downward displacement — it must be collected by upward displacement of air (heavier, very soluble).
• Reversing the aqua regia ratio — it is 3 parts HCl : 1 part HNO₃, not the other way round.
• Confusing the sulphite test (SO₂) with thiosulphate (SO₂ + yellow S), or forgetting the below/above 200 °C condition (NaHSO₄ vs Na₂SO₄).

Likely Exam Questions (with crisp answers)

1. Name the reactants for the lab preparation of HCl. Sodium chloride (common salt) and concentrated sulphuric acid.
2. Why is conc. H₂SO₄ used and not conc. HNO₃? HNO₃ is volatile and would escape with the HCl gas; H₂SO₄ is non-volatile.
3. Write the equation below 200 °C. $\mathrm{NaCl + H_2SO_4 \xrightarrow{<200^\circ C} NaHSO_4 + HCl}$.
4. Write the equation above 200 °C. $\mathrm{2NaCl + H_2SO_4 \xrightarrow{>200^\circ C} Na_2SO_4 + 2HCl}$.
5. Why is the temperature kept below 200 °C? To avoid cracking glass, wasting fuel, and forming a hard crust of Na₂SO₄.
6. Which drying agent dries HCl and why not CaO/P₂O₅? Conc. H₂SO₄; CaO and P₂O₅ react chemically with HCl.
7. How is HCl collected and why? By upward displacement of air, because it is heavier than air and very soluble in water.
8. What does the fountain experiment prove? That HCl is extremely soluble in water (red fountain).
9. Why does the fountain turn red? The litmus solution becomes hydrochloric acid, which is acidic.
10. Why funnel arrangement when preparing hydrochloric acid? Prevents back suction of water into the hot flask and gives a large surface area for absorption.
11. What is aqua regia and which component is the oxidising agent? 3 parts conc. HCl + 1 part conc. HNO₃; HNO₃ oxidises HCl to chlorine.
12. Distinguish dil. HCl from dil. HNO₃ with one reagent. Add AgNO₃: HCl gives a curdy white precipitate (AgCl); HNO₃ does not.
13. Why does dry HCl not turn blue litmus red? It has no free ions; only in water does it ionise to give H₃O⁺.
14. Why is HCl a covalent compound although it forms ions in water? In liquefied form / in toluene it does not conduct or turn litmus red — no free ions.
15. Gas evolved when (i) Na₂CO₃ (ii) MnO₂ + conc. HCl (heat)? (i) CO₂ (odourless, turns lime water milky); (ii) Cl₂ (greenish-yellow, turns starch-iodide paper blue-black).
16. Name a white chloride soluble in hot water, insoluble in cold. Lead chloride (PbCl₂).
17. What is the constant boiling mixture of HCl? 22.2 % HCl by mass, boiling at 110 °C (azeotrope).

Chemical Equations & Formulas (quick reference)

• $\mathrm{H_2 + Cl_2 \xrightarrow{sunlight} 2HCl}$
• $\mathrm{NaCl + H_2SO_4 \xrightarrow{<200^\circ C} NaHSO_4 + HCl}$ ; $\mathrm{2NaCl + H_2SO_4 \xrightarrow{>200^\circ C} Na_2SO_4 + 2HCl}$
• $\mathrm{2P_2O_5 + 3HCl \rightarrow POCl_3 + 3HPO_3}$ ; $\mathrm{CaO + 2HCl \rightarrow CaCl_2 + H_2O}$
• $\mathrm{NH_3 + HCl \rightarrow NH_4Cl}$ ; $\mathrm{HCl + H_2O \rightleftharpoons H_3O^+ + Cl^-}$ ; $\mathrm{2HCl \xrightarrow{>500^\circ C} H_2 + Cl_2}$
• $\mathrm{Mg + 2HCl \rightarrow MgCl_2 + H_2}$ ; $\mathrm{Zn + 2HCl \rightarrow ZnCl_2 + H_2}$ ; $\mathrm{Fe + 2HCl \rightarrow FeCl_2 + H_2}$ ; $\mathrm{2Al + 6HCl \rightarrow 2AlCl_3 + 3H_2}$
• $\mathrm{ZnO + 2HCl \rightarrow ZnCl_2 + H_2O}$ ; $\mathrm{CuO + 2HCl \rightarrow CuCl_2 + H_2O}$ ; $\mathrm{Fe_2O_3 + 6HCl \rightarrow 2FeCl_3 + 3H_2O}$
• $\mathrm{NaOH + HCl \rightarrow NaCl + H_2O}$ ; $\mathrm{Ca(OH)_2 + 2HCl \rightarrow CaCl_2 + 2H_2O}$
• $\mathrm{Na_2CO_3 + 2HCl \rightarrow 2NaCl + H_2O + CO_2}$ ; $\mathrm{CaCO_3 + 2HCl \rightarrow CaCl_2 + H_2O + CO_2}$
• $\mathrm{Na_2SO_3 + 2HCl \rightarrow 2NaCl + H_2O + SO_2}$ ; $\mathrm{Na_2S + 2HCl \rightarrow 2NaCl + H_2S}$ ; $\mathrm{FeS + 2HCl \rightarrow FeCl_2 + H_2S}$
• $\mathrm{Na_2S_2O_3 + 2HCl \rightarrow 2NaCl + H_2O + SO_2 + S}$
• $\mathrm{Pb(NO_3)_2 + 2HCl \rightarrow PbCl_2 + 2HNO_3}$ ; $\mathrm{AgNO_3 + HCl \rightarrow AgCl + HNO_3}$ ; $\mathrm{AgCl + 2NH_4OH \rightarrow [Ag(NH_3)_2]Cl + 2H_2O}$
• $\mathrm{MnO_2 + 4HCl \xrightarrow{\Delta} MnCl_2 + 2H_2O + Cl_2}$ ; $\mathrm{PbO_2 + 4HCl \xrightarrow{\Delta} PbCl_2 + 2H_2O + Cl_2}$
• $\mathrm{Pb_3O_4 + 8HCl \xrightarrow{\Delta} 3PbCl_2 + 4H_2O + Cl_2}$
• $\mathrm{K_2Cr_2O_7 + 14HCl \xrightarrow{\Delta} 2KCl + 2CrCl_3 + 7H_2O + 3Cl_2}$ ; $\mathrm{2KMnO_4 + 16HCl \xrightarrow{\Delta} 2KCl + 2MnCl_2 + 8H_2O + 5Cl_2}$
• $\mathrm{2KI + Cl_2 \rightarrow 2KCl + I_2}$
• $\mathrm{HNO_3 + 3HCl \rightarrow NOCl + 2[Cl] + 2H_2O}$ ; $\mathrm{Au + 3[Cl] \rightarrow AuCl_3}$ ; $\mathrm{Pt + 4[Cl] \rightarrow PtCl_4}$