What you will learn

Describe the action of water on NaCl, MgCl₂, Al₂Cl₆, SiCl₄ and PCl₅. Link each product and observation to dissolution, hydration or hydrolysis, and explain why the solutions have different pH values. Exact pH depends on concentration and temperature; the ranges here are typical classroom descriptions, not fixed constants.

1. Three processes that must be distinguished

• Dissolution: a substance becomes dispersed in a solvent. An ionic lattice can separate into ions that already existed in the solid.
• Hydration: water molecules surround dissolved ions and may bond to them. Writing Mg²⁺(aq) means a hydrated ion, not a bare ion floating in water.
• Hydrolysis: reaction with water changes the chemical species. A hydrated metal ion can donate a proton from a coordinated water molecule; a covalent chloride can undergo bond replacement to form HCl and an oxygen-containing product.

Dissolution can be exothermic without producing an acidic solution. A clear solution also does not prove that no reaction occurred. Identify the particles and products rather than classifying every chloride by one rule.

2. Sodium chloride: dissolution, approximately neutral

White NaCl crystals dissolve to form a colourless solution. Water separates Na⁺ and Cl⁻ from the giant ionic lattice and hydrates the ions. Neither ion undergoes appreciable acid-producing hydrolysis under ordinary dilute conditions. The solution is approximately neutral: pH about 7 at 25 °C in pure water.

NaCl(s) → Na⁺(aq) + Cl⁻(aq)

This is a dissolution equation, not NaCl + H₂O → NaOH + HCl. Sodium chloride is not sodium metal: it does not fizz to produce hydrogen or form a strongly alkaline solution. Dissolved carbon dioxide or impurities can shift the measured pH slightly.

3. Magnesium chloride: hydrated ions and slight acidity

MgCl₂ dissolves readily to give a colourless solution. Dissolving anhydrous MgCl₂ commonly releases heat; the temperature change depends on whether the starting solid is anhydrous or hydrated. It does not normally give the vigorous HCl fumes associated with covalent chlorides.

MgCl₂(s) → Mg²⁺(aq) + 2Cl⁻(aq)

The Mg²⁺ ion is hydrated, commonly represented as [Mg(H₂O)₆]²⁺. Its positive charge attracts electron density from the oxygen atoms of coordinated water and weakens their O–H bonds. A small proportion lose a proton to surrounding water. This weak acid equilibrium produces a slightly acidic solution, often described as about pH 6 under classroom conditions, but not a universal exact value.

[Mg(H₂O)₆]²⁺(aq) + H₂O(l) ⇌ [Mg(H₂O)₅(OH)]⁺(aq) + H₃O⁺(aq)

For a simple comparison, say “dissolves, with only slight hydrolysis of the hydrated cation”. Do not write complete conversion to Mg(OH)₂ + 2HCl in cold excess water. Most magnesium remains as hydrated Mg²⁺; chloride is a spectator ion in this acid equilibrium.

4. Aluminium chloride: strongly polarising aluminium

Pearson names Al₂Cl₆, the molecular dimer representation of anhydrous aluminium chloride in suitable phases. In excess water the dimer is not retained. Aluminium becomes hydrated and chloride ions enter solution. Adding anhydrous aluminium chloride to water releases considerable heat and may give steamy acidic fumes near the initial contact. With sufficient water, the final solution is colourless and acidic.

Al₂Cl₆ + 12H₂O → 2[Al(H₂O)₆]³⁺ + 6Cl⁻
Overall hydration/species accounting; excess water. The physical phase of anhydrous aluminium chloride must be specified by the question.

The small Al³⁺ ion has greater charge density than Mg²⁺. It polarises the O–H bonds of coordinated water more strongly, so the hydrated aluminium ion is a stronger acid. A typical aluminium chloride solution may have pH about 2–3, depending on concentration. The first acid equilibrium is:

[Al(H₂O)₆]³⁺(aq) + H₂O(l) ⇌ [Al(H₂O)₅(OH)]²⁺(aq) + H₃O⁺(aq)

This equilibrium does not mean every Al³⁺ ion becomes solid Al(OH)₃. In the acidic final solution, hydrated aluminium species predominate and no hydroxide precipitate is required. Precipitation on adding alkali is a separate change. A simplistic complete equation Al₂Cl₆ + 6H₂O → 2Al(OH)₃ + 6HCl does not accurately describe the species in the usual acidic solution.

5. Silicon tetrachloride: vigorous hydrolysis

SiCl₄ is a colourless molecular liquid. It reacts vigorously with water, giving heat, white steamy fumes in moist air and a white silica-containing product. Its Si–Cl bonds are replaced as oxygen-containing products form. In excess water, the overall equation is conventionally written:

SiCl₄(l) + 2H₂O(l) → SiO₂(s) + 4HCl(aq)

The white material is often hydrated silica in practice; SiO₂ is the simplified overall product used here. HCl dissolves and ionises, making the aqueous phase strongly acidic, commonly around pH 1–2 for typical classroom mixtures. With limited water, hydrogen chloride can escape as gas; moist air produces a mist of acidic droplets, observed as steamy white fumes.

The pH belongs to the resulting aqueous phase, not to pure SiCl₄. Do not say it simply dissolves as Si⁴⁺ and Cl⁻; free aqueous Si⁴⁺ is not the appropriate product model.

6. Phosphorus(V) chloride: water amount controls the stage

PCl₅ is a white or pale solid that reacts vigorously with water. Heat and steamy acidic fumes may be observed. With a limited amount of water the initial overall product is phosphoryl chloride, POCl₃, plus hydrogen chloride:

PCl₅(s) + H₂O(l) → POCl₃(l) + 2HCl(g)
Limited water; HCl can escape. Where it dissolves, write HCl(aq).

With excess water, POCl₃ hydrolyses further, giving phosphoric acid and additional hydrogen chloride. The final aqueous mixture contains both H₃PO₄ and dissolved HCl, and is strongly acidic. Typical classroom descriptions use pH about 1–2, but dilution changes it.

POCl₃(l) + 3H₂O(l) → H₃PO₄(aq) + 3HCl(aq)

PCl₅(s) + 4H₂O(l) → H₃PO₄(aq) + 5HCl(aq)

The excess-water equation is the sum of the two stages. It is not PCl₅ + water → P(OH)₅, and hydrolysis does not produce hydrogen gas. Phosphorus remains at oxidation number +5 in the overall reaction; hydrolysis is not automatically a redox reaction.

7. Compare observations, products and pH

These approximate pH descriptions are not values prescribed for a fixed concentration by Pearson. A numerical pH needs concentration and suitable equilibrium information. Do not infer a strict SiCl₄ versus PCl₅ pH ranking from the formulas alone. Observations refer to controlled chemistry demonstrations, not instructions for handling reactive chlorides independently.

8. Four worked examples

Common misconceptions

• “Every chloride hydrolyses completely.” NaCl mainly dissolves; Mg²⁺ hydrolysis is slight; covalent chlorides can react vigorously.
• “Chloride ions make every solution acidic.” Hydrated cation hydrolysis or covalent chloride hydrolysis explains the acid production.
• “Al₂Cl₆ molecules remain unchanged in water.” Excess water produces hydrated aluminium species and chloride ions.
• “Al(OH)₃ must precipitate when aluminium chloride is added to water.” The acidic solution is better represented by hydrated aluminium ions and acid equilibria.
• “A chloride has one exact pH.” pH describes a solution and depends on its concentration and conditions.

Exam tips

• Separate the observation from the explanation: white fumes are an observation; HCl-containing acidic droplets explain them.
• For PCl₅, read whether water is limited or in excess before selecting products.
• Keep Al₂Cl₆ as Pearson’s specified starting representation but use hydrated ions for its aqueous products.
• State that Mg²⁺ hydrolysis is much weaker than Al³⁺ hydrolysis.
• Write a balanced equation and explain the species responsible for acidity, rather than only quoting a pH.

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