Chemistry Form 4 Notes : ACIDS, BASES AND SALTS

Chemistry Form Four Notes

Acids, Bases and Salts

A. Acids and Bases

At a school laboratory:

An acid may be defined as a substance that turns litmus red.
A base may be defined as a substance that turns litmus blue.

Litmus is a lichen found mainly in West Africa. It changes its colour depending on whether the solution it is in is basic/alkaline or acidic. It is thus able to identify whether:

An acid is a substance that dissolves in water to form H⁺/H₃O⁺ as the only positive ion/cation. This is called the Arrhenius definition of an acid. From this definition, an acid dissociates/ionizes in water releasing H⁺ ions as shown below:

HCl(aq) → H⁺(aq) + Cl^–(aq)

HNO₃(aq) → H⁺(aq) + NO₃^–(aq)

CH₃COOH(aq) → H⁺(aq) + CH₃COO^–(aq)

H₂SO₄(aq) → 2H⁺(aq) + SO₄^2-(aq)

H₂CO₃(aq) → 2H⁺(aq) + CO₃^2-(aq)

H₃PO₄(aq) → 3H⁺(aq) + PO₄^3-(aq)

A base is a substance which dissolves in water to form OH^– as the only negatively charged ion/anion. This is called the Arrhenius definition of a base. From this definition, a base dissociates/ionizes in water releasing OH^– ions as shown below:

KOH(aq) → K⁺(aq) + OH^–(aq)

NaOH(aq) → Na⁺(aq) + OH^–(aq)

NH₄OH(aq) → NH₄⁺(aq) + OH^–(aq)

Ca(OH)₂(aq) → Ca²⁺(aq) + 2OH^–(aq)

Mg(OH)₂(aq) → Mg²⁺(aq) + 2OH^–(aq)

An acid is a proton donor. A base is a proton acceptor. This is called the Bronsted-Lowry definition of acids and bases. From this definition, an acid donates H⁺ ions.

H⁺ has no electrons or neutrons; it contains only a proton.

Examples

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I. From the equation:

HCl(aq) + H₂O(l) ⇌ H₃O⁺(aq) + Cl^–(aq)

For the forward reaction (left to right), H₂O gains a proton to form H₃O⁺, so H₂O is a proton acceptor. It is a Bronsted-Lowry base.
For the backward reaction (right to left), H₃O⁺ donates a proton to form H₂O, so H₃O⁺ is a proton donor. It is a Bronsted-Lowry conjugate acid.
For the forward reaction, HCl donates a proton to form Cl^–, so HCl is a proton donor. It is a Bronsted-Lowry acid.
For the backward reaction, Cl^– gains a proton to form HCl, so Cl^– is a proton acceptor. It is a Bronsted-Lowry conjugate base.

Every base or acid in the Bronsted-Lowry definition thus must have a conjugate product or reactant.

II. From the equation:

HCl(aq) + NH₃(aq) ⇌ NH₄⁺(aq) + Cl^–(aq)

For the forward reaction, NH₃ gains a proton to form NH₄⁺, so NH₃ is a proton acceptor. It is a Bronsted-Lowry base.
For the backward reaction, NH₄⁺ donates a proton to form NH₃, so NH₄⁺ is a proton donor. It is a Bronsted-Lowry conjugate acid.
For the forward reaction, HCl donates a proton to form Cl^–, so HCl is a proton donor. It is a Bronsted-Lowry acid.
For the backward reaction, Cl^– gains a proton to form HCl, so Cl^– is a proton acceptor. It is a Bronsted-Lowry conjugate base.

4. Acids and bases show acidic and alkaline properties only in water but not in other solvents, for example:

(a) Hydrogen chloride gas dissolves in water to form hydrochloric acid. Hydrochloric acid dissociates/ionizes in water to free H⁺(aq)/H₃O⁺(aq) ions. The free H₃O⁺(aq)/H⁺(aq) ions are responsible for:

Turning blue litmus paper/solution red.
Showing pH values 1, 2, 3, 4, 5, 6.
Being good electrolytes/conductors of electricity/undergoing electrolysis.
Reacting with metals to produce/evolve hydrogen gas and a salt, as shown below:

Ionically:

For a monovalent metal: 2M(s) + 2H⁺(aq) → 2M⁺(aq) + H₂(g)
For a divalent metal: M(s) + 2H⁺(aq) → M²⁺(aq) + H₂(g)
For a trivalent metal: 2M(s) + 6H⁺(aq) → 2M³⁺(aq) + 3H₂(g)

Examples:

2Na(s) + 2H⁺(aq) → 2Na⁺(aq) + H₂(g)
Ca(s) + 2H⁺(aq) → Ca²⁺(aq) + H₂(g)
2Al(s) + 6H⁺(aq) → 2Al³⁺(aq) + 3H₂(g)

(v) React with metal carbonates and hydrogen carbonates to produce/evolve carbon(IV) oxide gas, water and a salt, as shown below:

Ionically:

For a monovalent metal: M₂CO₃(s) + 2H⁺(aq) → 2M⁺(aq) + H₂O(l) + CO₂(g)
MHCO₃(s) + H⁺(aq) → M⁺(aq) + H₂O(l) + CO₂(g)
For a divalent metal: MCO₃(s) + 2H⁺(aq) → M²⁺(aq) + H₂O(l) + CO₂(g)
M(HCO₃)₂(aq) + 2H⁺(aq) → M²⁺(aq) + 2H₂O(l) + 2CO₂(g)

Examples:

K₂CO₃(s) + 2H⁺(aq) → 2K⁺(aq) + H₂O(l) + CO₂(g)
NH₄HCO₃(s) + H⁺(aq) → NH₄⁺(aq) + H₂O(l) + CO₂(g)
ZnCO₃(s) + 2H⁺(aq) → Zn²⁺(aq) + H₂O(l) + CO₂(g)
Mg(HCO₃)₂(aq) + 2H⁺(aq) → Mg²⁺(aq) + 2H₂O(l) + 2CO₂(g)

(vi) Neutralize metal oxides/hydroxides to salt and water only, as shown below:

Ionically:

For a monovalent metal: M₂O(s) + 2H⁺(aq) → 2M⁺(aq) + H₂O(l)
MOH(aq) + H⁺(aq) → M⁺(aq) + H₂O(l)
For a divalent metal: MO(s) + 2H⁺(aq) → M²⁺(aq) + H₂O(l)
M(OH)₂(s) + 2H⁺(aq) → M²⁺(aq) + 2H₂O(l)
For a trivalent metal: M₂O₃(s) + 6H⁺(aq) → 2M³⁺(aq) + 3H₂O(l)
M(OH)₃(s) + 3H⁺(aq) → M³⁺(aq) + 3H₂O(l)

Examples:

K₂O(s) + 2H⁺(aq) → 2K⁺(aq) + H₂O(l)
NH₄OH(aq) + H⁺(aq) → NH₄⁺(aq) + H₂O(l)
ZnO(s) + 2H⁺(aq) → Zn²⁺(aq) + H₂O(l)
Pb(OH)₂(s) + 2H⁺(aq) → Pb²⁺(aq) + 2H₂O(l)

(b) Hydrogen chloride gas dissolves in methylbenzene/benzene but does not dissociate/ionize into free ions. It exists in molecular state showing none of the above properties.

(c) Ammonia gas dissolves in water to form aqueous ammonia which dissociates/ionizes to free NH₄⁺(aq) and OH^–(aq) ions. This dissociation/ionization makes aqueous ammonia:

Turn litmus paper/solution blue.
Have pH 8, 9, 10, 11.
Be a good electrical conductor.
React with acids to form ammonium salt and water only.

NH₄OH(aq) + HCl(aq) → NH₄Cl(aq) + H₂O(l)

(d) Ammonia gas dissolves in methylbenzene/benzene/kerosene but does not dissociate into free ions, therefore existing as molecules.

6. Solvents are either polar or non-polar.

A polar solvent is one which dissolves ionic compounds and other polar solvents.

Water is a polar solvent that dissolves ionic and polar substances by surrounding the free ions as shown below:

The oxygen atom is partially negative and the two hydrogen atoms are partially positive. They surround the free H⁺ and Cl^– ions.

A non-polar solvent is one which dissolves non-polar substances and covalent compounds.

If a polar ionic compound is dissolved in a non-polar solvent, it does not ionize/dissociate into free ions as shown below:

H-Cl (covalent bond)

7. Some acids and bases are strong while others are weak.

(a) A strong acid/base is one which is fully/wholly/completely dissociated/ionized into many free H⁺/OH^– ions, i.e.

I. Strong acids exist more as free H⁺ ions than molecules. For example:

HCl(aq) → H⁺(aq) + Cl^–(aq) (molecules) (cation) (anion)

HNO₃(aq) → H⁺(aq) + NO₃^–(aq) (molecules) (cation) (anion)

H₂SO₄(aq) → 2H⁺(aq) + SO₄^2-(aq) (molecules) (cation) (anion)

II. Strong bases/alkalis exist more as free OH^– ions than molecules. For example:

KOH(aq) → K⁺(aq) + OH^–(aq) (molecules) (cation) (anion)

NaOH(aq) → Na⁺(aq) + OH^–(aq) (molecules) (cation) (anion)

(b) A weak base/acid is one which is partially/partly dissociated/ionized in water into free OH^–(aq) and H⁺(aq) ions.

I. Weak acids exist more as molecules than as free H⁺ ions. For example:

CH₃COOH(aq) ⇌ H⁺(aq) + CH₃COO^–(aq) (molecules) (cation) (anion)

H₃PO₄(aq) ⇌ 3H⁺(aq) + PO₄^3-(aq) (molecules) (cation) (anion)

H₂CO₃(aq) ⇌ 2H⁺(aq) + CO₃^2-(aq) (molecules) (cation) (anion)

II. Weak bases/alkalis exist more as molecules than free OH^– ions. For example:

NH₄OH(aq) ⇌ NH₄⁺(aq) + OH^–(aq) (molecules) (cation) (anion)

Ca(OH)₂(aq) ⇌ Ca²⁺(aq) + 2OH^–(aq) (molecules) (cation) (anion)

Mg(OH)₂(aq) ⇌ Mg²⁺(aq) + 2OH^–(aq) (molecules) (cation) (anion)

8. The concentration of an acid/base/alkali is based on the number of moles of acid/base dissolved in a decimeter (litre) of the solution.

An acid/base/alkali with more acid/base in a decimeter (litre) of solution is said to be concentrated, while that with less is said to be dilute.

9. (a) (i) Strong acids have pH 1, 2, 3 while weak acids have higher pH 4, 5, 6.

(ii) A neutral solution has pH 7.

(iii) Strong alkalis/bases have pH 12, 13, 14 while weak bases/alkalis have pH 11, 10, 9, 8.

(b) pH is a measure of H⁺(aq) concentration in a solution.

The higher the H⁺(aq) ion concentration:

The higher the acidity.
The lower the pH.
The lower the concentration of OH^–(aq).
The lower the alkalinity.

At pH 7, a solution has equal concentration of H⁺(aq) and OH^–(aq).

Beyond pH 7, the concentration of OH^–(aq) increases as the H⁺(aq) ions decrease.

10. (a) When acids/bases dissolve in water, the ions present in the solution conduct electricity.

The more the dissociation, the higher the yield of ions and the greater the electrical conductivity of the solution.

A compound that conducts electricity in an electrolyte and thus shows high electrical conductivity is a strong electrolyte, while a compound showing low electrical conductivity is a weak electrolyte.

(b) Practically, a bright light on a bulb, a high voltage reading from a voltmeter, high ammeter reading from an ammeter, or a big deflection on a galvanometer is an indicator of a strong electrolyte (acid/base) and the opposite for weak electrolytes (acids/base).

11. Some compounds exhibit both properties of acids and bases/alkalis.

A substance that reacts with both acids and bases is said to be amphotellic.

The examples below show the amphotellic properties of:

(a) Zinc (II) oxide (ZnO) and Zinc hydroxide (Zn(OH)₂)

(i) When half a spatula full of Zinc(II) oxide is placed in a boiling tube containing 10 cm³ of either 2M nitric(V) acid or 2M sodium hydroxide solution, it dissolves in both the acid and the alkali/base to form a colourless solution, i.e.

(i) When reacting with nitric(V) acid, the oxide shows basic properties by reacting with an acid to form a simple salt and water only.

Basic oxide + Acid → salt + water

Examples:

Chemical equations:

ZnO(s) + 2HNO₃(aq) → Zn(NO₃)₂(aq) + H₂O(l)

ZnO(s) + 2HCl(aq) → ZnCl₂(aq) + H₂O(l)

ZnO(s) + H₂SO₄(aq) → ZnSO₄(aq) + H₂O(l)

Ionic equation:

ZnO(s) + 2H⁺(aq) → Zn²⁺(aq) + H₂O(l)

(ii) When reacting with sodium hydroxide, the oxide shows acidic properties by reacting with a base to form a complex salt.

Basic oxide + Base/alkali + Water → Complex salt

Examples:

1. When Zinc oxide is reacted with sodium hydroxide, the complex salt is sodium tetrahydroxozincate(II) complex salt.

ZnO(s) + 2NaOH(aq) + H₂O(l) → Na₂Zn(OH)₄(aq)

2. When Zinc oxide is reacted with potassium hydroxide, the complex salt is potassium tetrahydroxozincate(II) complex salt.

ZnO(s) + 2KOH(aq) + H₂O(l) → K₂Zn(OH)₄(aq)

Ionic equation:

ZnO(s) + 2OH^–(aq) + H₂O(l) → 2[Zn(OH)₄]^2-(aq)

(ii) When Zinc(II) hydroxide is placed in a boiling tube containing 10 cm³ of either 2M nitric(V) acid or 2M sodium hydroxide solution, it dissolves in both the acid and the alkali/base to form a colourless solution, i.e.

(i) When reacting with nitric(V) acid, the hydroxide shows basic properties. It reacts with an acid to form a simple salt and water only.

Basic hydroxide + Acid → salt + water

Examples:

Chemical equations:

Zn(OH)₂(s) + 2HNO₃(aq) → Zn(NO₃)₂(aq) + 2H₂O(l)

Zn(OH)₂(s) + 2HCl(aq) → ZnCl₂(aq) + 2H₂O(l)

Zn(OH)₂(s) + H₂SO₄(aq) → ZnSO₄(aq) + 2H₂O(l)

Ionic equation:

Zn(OH)₂(s) + 2H⁺(aq) → Zn²⁺(aq) + 2H₂O(l)

(ii) When reacting with sodium hydroxide, the hydroxide shows acidic properties by reacting with a base to form a complex salt.

Basic hydroxide + Base/alkali → Complex salt

Examples:

Chemical equations:

1. When Zinc hydroxide is reacted with sodium hydroxide, the complex salt is sodium tetrahydroxozincate(II) complex salt.

Zn(OH)₂(s) + 2NaOH(aq) → Na₂Zn(OH)₄(aq)

2. When Zinc hydroxide is reacted with potassium hydroxide, the complex salt is potassium tetrahydroxozincate(II) complex salt.

Zn(OH)₂(s) + 2KOH(aq) → K₂Zn(OH)₄(aq)

Ionic equation:

Zn(OH)₂(s) + 2OH^–(aq) → 2[Zn(OH)₄]^2-(aq)

(b) Lead (II) oxide (PbO) and Lead (II) hydroxide (Pb(OH)₂)

(i) When half a spatula full of Lead(II) oxide is placed in a boiling tube containing 10 cm³ of either 2M nitric(V) acid or 2M sodium hydroxide solution, it dissolves in both the acid and the alkali/base to form a colourless solution, i.e.

(i) When reacting with nitric(V) acid, the oxide shows basic properties by reacting with an acid to form a simple salt and water only. All other Lead salts are insoluble.

Chemical equation:

PbO(s) + 2HNO₃(aq) → Pb(NO₃)₂(aq) + H₂O(l)

Ionic equation:

PbO(s) + 2H⁺(aq) → Pb²⁺(aq) + H₂O(l)

(ii) When reacting with sodium hydroxide, the oxide shows acidic properties by reacting with a base to form a complex salt.

Chemical equation:

1. When Lead(II) oxide is reacted with sodium hydroxide, the complex salt is sodium tetrahydroxoplumbate(II) complex salt.

PbO(s) + 2NaOH(aq) + H₂O(l) → Na₂Pb(OH)₄(aq)

2. When Lead(II) oxide is reacted with potassium hydroxide, the complex salt is potassium tetrahydroxoplumbate(II) complex salt.

PbO(s) + 2KOH(aq) + H₂O(l) → K₂Pb(OH)₄(aq)

Ionic equation:

PbO(s) + 2OH^–(aq) + H₂O(l) → 2[Pb(OH)₄]^2-(aq)

(ii) When Lead(II) hydroxide is placed in a boiling tube containing 10 cm³ of either 2M nitric(V) acid or 2M sodium hydroxide solution, it dissolves in both the acid and the alkali/base to form a colourless solution, i.e.

(i) When reacting with nitric(V) acid, the hydroxide shows basic properties. It reacts with the acid to form a simple salt and water only.

Chemical equation:

Pb(OH)₂(s) + 2HNO₃(aq) → Pb(NO₃)₂(aq) + 2H₂O(l)

Ionic equation:

Pb(OH)₂(s) + 2H⁺(aq) → Pb²⁺(aq) + 2H₂O(l)

(ii) When reacting with sodium hydroxide, the hydroxide shows acidic properties. It reacts with a base to form a complex salt.

Chemical equation:

1. When Lead(II) hydroxide is reacted with sodium hydroxide, the complex salt is sodium tetrahydroxoplumbate(II) complex salt.

Pb(OH)₂(s) + 2NaOH(aq) → Na₂Pb(OH)₄(aq)

2. When Lead(II) hydroxide is reacted with potassium hydroxide, the complex salt is potassium tetrahydroxoplumbate(II) complex salt.

Pb(OH)₂(s) + 2KOH(aq) → K₂Pb(OH)₄(aq)

Ionic equation:

Pb(OH)₂(s) + 2OH^–(aq) → 2[Pb(OH)₄]^2-(aq)

(i) When half a spatula full of Aluminium(III) oxide is placed in a boiling tube containing 10 cm³ of either 2M nitric(V) acid or 2M sodium hydroxide solution, it dissolves in both the acid and the alkali/base to form a colourless solution, i.e.

(i) When reacting with nitric(V) acid, the oxide shows basic properties by reacting with an acid to form a simple salt and water only.

Chemical equations:

Al₂O₃(s) + 6HNO₃(aq) → 2Al(NO₃)₃(aq) + 3H₂O(l)

Al₂O₃(s) + 6HCl(aq) → 2AlCl₃(aq) + 3H₂O(l)

Al₂O₃(s) + 3H₂SO₄(aq) → Al₂(SO₄)₃(aq) + 3H₂O(l)

Ionic equation:

Al₂O₃(s) + 6H⁺(aq) → 2Al³⁺(aq) + 3H₂O(l)

(ii) When reacting with sodium hydroxide, the oxide shows acidic properties by reacting with a base to form a complex salt.

Chemical equations:

1. When Aluminium(III) oxide is reacted with sodium hydroxide, the complex salt is sodium tetrahydroxoaluminate(III) complex salt.

Al₂O₃(s) + 2NaOH(aq) + 3H₂O(l) → 2NaAl(OH)₄(aq)

2. When Aluminium(III) oxide is reacted with potassium hydroxide, the complex salt is potassium tetrahydroxoaluminate(III) complex salt.

Al₂O₃(s) + 2KOH(aq) + 3H₂O(l) → 2KAl(OH)₄(aq)

Ionic equation:

Al₂O₃(s) + 2OH^–(aq) + 3H₂O(l) → 2[Al(OH)₄]^–(aq)

(ii) When Aluminium(III) hydroxide is placed in a boiling tube containing 10 cm³ of either 2M nitric(V) acid or 2M sodium hydroxide solution, it dissolves in both the acid and the alkali/base to form a colourless solution, i.e.

(i) When reacting with nitric(V) acid, the hydroxide shows basic properties. It reacts with the acid to form a simple salt and water only.

Chemical equations:

Al(OH)₃(s) + 3HNO₃(aq) → Al(NO₃)₃(aq) + 3H₂O(l)

Al(OH)₃(s) + 3HCl(aq) → AlCl₃(aq) + 3H₂O(l)

2Al(OH)₃(s) + 3H₂SO₄(aq) → Al₂(SO₄)₃(aq) + 3H₂O(l)

Ionic equation:

Al(OH)₃(s) + 3H⁺(aq) → Al³⁺(aq) + 3H₂O(l)

(ii) When reacting with sodium hydroxide, the hydroxide shows acidic properties. It reacts with a base to form a complex salt.

Chemical equations:

1. When aluminium(III) hydroxide is reacted with sodium hydroxide, the complex salt is sodium tetrahydroxoaluminate(III) complex salt.

Al(OH)₃(s) + NaOH(aq) → NaAl(OH)₄(aq)

2. When aluminium(III) hydroxide is reacted with potassium hydroxide, the complex salt is potassium tetrahydroxoaluminate(III) complex salt.

Al(OH)₃(s) + KOH(aq) → KAl(OH)₄(aq)

Ionic equation:

Al(OH)₃(s) + OH^–(aq) → [Al(OH)₄]^–(aq)

Summary of amphotellic oxides/hydroxides

Oxide	Hydroxide	Formula of simple salt from nitric (V) acid	Formula of complex salt from sodium hydroxide
ZnO	Zn(OH)₂	Zn(NO₃)₂	Na₂Zn(OH)₄ [Zn(OH)₄]^2-(aq) Sodium tetrahydroxozincate(II)
PbO	Pb(OH)₂	Pb(NO₃)₂	Na₂Pb(OH)₄ [Pb(OH)₄]^2-(aq) Sodium tetrahydroxoplumbate(II)
Al₂O₃	Al(OH)₃	Al(NO₃)₃	NaAl(OH)₄ [Al(OH)₄]^–(aq) Sodium tetrahydroxoaluminate(III)

12. (a) A salt is an ionic compound formed when the cation from a base combines with the anion derived from an acid.

A salt is therefore formed when the hydrogen ions in an acid are replaced wholly or partially, directly or indirectly by a metal or ammonium radical.

(b) The number of ionizable/replaceable hydrogen in an acid is called the basicity of an acid.

Some acids are therefore:

(i) Monobasic acids generally denoted HX e.g. HCl, HNO₃, HCOOH, CH₃COOH.
(ii) Dibasic acids generally denoted H₂X e.g. H₂SO₄, H₂SO₃, H₂CO₃, HOOCOOH.
(iii) Tribasic acids generally denoted H₃X e.g. H₃PO₄.

(i) A normal salt is formed when all the ionizable/replaceable hydrogen in an acid is replaced by a metal or metallic/ammonium radical.
(ii) An acid salt is formed when part of the ionizable/replaceable hydrogen in an acid is replaced by a metal or metallic/ammonium radical.

Table showing normal and acid salts derived from common acids

Acid name	Chemical formula	Basicity	Normal salt	Acid salt
Hydrochloric acid	HCl	Monobasic	Chloride (Cl^–)	None
Nitric(V) acid	HNO₃	Monobasic	Nitrate(V) (NO₃^–)	None
Nitric(III) acid	HNO₂	Monobasic	Nitrate(III) (NO₂^–)	None
Sulphuric(VI) acid	H₂SO₄	Dibasic	Sulphate(VI) (SO₄^2-)	Hydrogen sulphate(VI) (HSO₄^–)
Sulphuric(IV) acid	H₂SO₃	Dibasic	Sulphate(IV) (SO₃^2-)	Hydrogen sulphate(IV) (HSO₃^–)
Carbonic(IV) acid	H₂CO₃	Dibasic	Carbonate(IV) (CO₃^2-)	Hydrogen carbonate(IV) (HCO₃^–)
Phosphoric(V) acid	H₃PO₄	Tribasic	Phosphate(V) (PO₄^3-)	Dihydrogen phosphate(V) (H₂PO₄^2-), Hydrogen diphosphate(V) (HP₂O₄^2-)

The table below shows some examples of salts.

Base/alkali	Cation	Acid	Anion	Salt	Chemical name of salts
NaOH	Na⁺	HCl	Cl^–	NaCl	Sodium(I) chloride
Mg(OH)₂	Mg²⁺	H₂SO₄	SO₄^2-	MgSO₄, Mg(HSO₄)₂	Magnesium sulphate(VI), Magnesium hydrogen sulphate(VI)
Pb(OH)₂	Pb²⁺	HNO₃	NO₃^–	Pb(NO₃)₂	Lead(II) nitrate(V)
Ba(OH)₂	Ba²⁺	HNO₃	NO₃^–	Ba(NO₃)₂	Barium(II) nitrate(V)
Ca(OH)₂	Ca²⁺	H₂SO₄	SO₄^2-	CaSO₄	Calcium sulphate(VI)
NH₄OH	NH₄⁺	H₃PO₄	PO₄^3-	(NH₄)₃PO₄, (NH₄)₂HPO₄, NH₄H₂PO₄	Ammonium phosphate(V), Diammonium phosphate(V), Ammonium diphosphate(V)
KOH	K⁺	H₃PO₄	PO₄^3-	K₃PO₄	Potassium phosphate(V)
Al(OH)₃	Al³⁺	H₂SO₄	SO₄^2-	Al₂(SO₄)₃	Aluminium(III) sulphate(VI)
Fe(OH)₂	Fe²⁺	H₂SO₄	SO₄^2-	FeSO₄	Iron(II) sulphate(VI)
Fe(OH)₃	Fe³⁺	H₂SO₄	SO₄^2-	Fe₂(SO₄)₃	Iron(III) sulphate(VI)

(d) Some salts undergo hygroscopy, deliquescence, and efflorescence.

(i) Hygroscopic salts/compounds are those that absorb water from the atmosphere but do not form a solution.

Some salts which are hygroscopic include anhydrous copper(II) sulphate(VI), anhydrous cobalt(II) chloride, potassium nitrate(V), and common table salt.

(ii) Deliquescent salts/compounds are those that absorb water from the atmosphere and form a solution.

Some salts which are deliquescent include sodium nitrate(V), calcium chloride, sodium hydroxide, iron(II) chloride, and magnesium chloride.

(iii) Efflorescent salts/compounds are those that lose their water of crystallization to the atmosphere.

Some salts which effloresce include sodium carbonate decahydrate, iron(II) sulphate(VI) heptahydrate, and sodium sulphate(VI) decahydrate.

(e) Some salts contain water of crystallization. They are hydrated. Others do not contain water of crystallization. They are anhydrous.

Table showing some hydrated salts.

Name of hydrated salt	Chemical formula
Copper(II) sulphate(VI) pentahydrate	CuSO₄·5H₂O
Aluminium(III) sulphate(VI) hexahydrate	Al₂(SO₄)₃·6H₂O
Zinc(II) sulphate(VI) heptahydrate	ZnSO₄·7H₂O
Iron(II) sulphate(VI) heptahydrate	FeSO₄·7H₂O
Calcium(II) sulphate(VI) heptahydrate	CaSO₄·7H₂O
Magnesium(II) sulphate(VI) heptahydrate	MgSO₄·7H₂O
Sodium sulphate(VI) decahydrate	Na₂SO₄·10H₂O
Sodium carbonate(IV) decahydrate	Na₂CO₃·10H₂O
Potassium carbonate(IV) decahydrate	K₂CO₃·10H₂O
Potassium sulphate(VI) decahydrate	K₂SO₄·10H₂O

(f) Some salts exist as simple salts while some as complex salts. Below are some complex salts.

Table of some complex salts

Name of complex salt	Chemical formula	Colour of the complex salt
Tetraamminecopper(II) sulphate(VI)	Cu(NH₃)₄SO₄·H₂O	Royal/deep blue solution
Tetraamminezinc(II) nitrate(V)	Zn(NH₃)₄(NO₃)₂	Colourless solution
Tetraamminecopper(II) nitrate(V)	Cu(NH₃)₄(NO₃)₂	Royal/deep blue solution
Tetraamminezinc(II) sulphate(VI)	Zn(NH₃)₄SO₄	Colourless solution

(g) Some salts exist as two salts in one. They are called double salts.

Table of some double salts

Name of double salts	Chemical formula
Trona (sodium sesquicarbonate)	Na₂CO₃·NaHCO₃·2H₂O
Ammonium iron(II) sulphate(VI)	FeSO₄(NH₄)₂SO₄·2H₂O
Ammonium aluminium(III) sulphate(VI)	Al₂(SO₄)₃(NH₄)₂SO₄·H₂O

(h) Some salts dissolve in water to form a solution. They are said to be soluble. Others do not dissolve in water. They form a suspension/precipitate in water.

Table of solubility of salts

Soluble salts	Insoluble salts
All nitrate(V) salts
All sulphate(VI)/SO₄^2- salts except	Barium(II) sulphate(VI)/BaSO₄, Calcium(II) sulphate(VI)/CaSO₄, Lead(II) sulphate(VI)/PbSO₄
All sulphate(IV)/SO₃^2- salts except	Barium(II) sulphate(IV)/BaSO₃, Calcium(II) sulphate(IV)/CaSO₃, Lead(II) sulphate(IV)/PbSO₃
All chlorides/Cl^– except	Silver chloride/AgCl, Lead(II) chloride/PbCl₂ (dissolves in hot water)
All phosphate(V)/PO₄^3-
All sodium, potassium and ammonium salts
All hydrogen carbonates/HCO₃^–
All hydrogen sulphate(VI)/HSO₄^–
Sodium carbonate/Na₂CO₃, potassium carbonate/K₂CO₃, ammonium carbonate (NH₄)₂CO₃	except All carbonates
All alkalis (KOH, NaOH, NH₄OH)	except All bases

13. Salts can be prepared in a school laboratory by a method that uses its solubility in water.

Soluble salts may be prepared by using any of the following methods:

(i) Direct displacement/reaction of a metal with an acid.

By reacting a metal higher in the reactivity series than hydrogen with a dilute acid, a salt is formed and hydrogen gas is evolved.

Excess of the metal must be used to ensure all the acid has reacted.

When effervescence/bubbling/fizzing has stopped, excess metal is filtered off.

The filtrate is heated to concentrate then allowed to crystallize.

Washing with distilled water then drying between filter papers produces a sample crystal of the salt, i.e.

M(s) + H₂X → MX(aq) + H₂(g)

Examples:

Mg(s) + H₂SO₄(aq) → MgSO₄(aq) + H₂(g)
Zn(s) + H₂SO₄(aq) → ZnSO₄(aq) + H₂(g)
Pb(s) + 2HNO₃(aq) → Pb(NO₃)₂(aq) + H₂(g)
Ca(s) + 2HNO₃(aq) → Ca(NO₃)₂(aq) + H₂(g)
Mg(s) + 2HNO₃(aq) → Mg(NO₃)₂(aq) + H₂(g)
Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g)
Zn(s) + 2HCl(aq) → ZnCl₂(aq) + H₂(g)

(ii) Reaction of an insoluble base with an acid

By adding an insoluble base (oxide/hydroxide) to a dilute acid until no more dissolves in the acid, a salt and water are formed. Excess of the base is filtered off. The filtrate is heated to concentrate, allowed to crystallize then washed with distilled water before drying between filter papers, e.g.

PbO(s) + 2HNO₃(aq) → Pb(NO₃)₂(aq) + H₂O(l)
Pb(OH)₂(s) + 2HNO₃(aq) → Pb(NO₃)₂(aq) + 2H₂O(l)
CaO(s) + 2HNO₃(aq) → Ca(NO₃)₂(aq) + H₂O(l)
MgO(s) + 2HNO₃(aq) → Mg(NO₃)₂(aq) + H₂O(l)
MgO(s) + 2HCl(aq) → MgCl₂(aq) + H₂O(l)
ZnO(s) + 2HCl(aq) → ZnCl₂(aq) + H₂O(l)
Zn(OH)₂(s) + 2HNO₃(aq) → Zn(NO₃)₂(aq) + 2H₂O(l)
CuO(s) + 2HCl(aq) → CuCl₂(aq) + H₂O(l)
CuO(s) + H₂SO₄(aq) → CuSO₄(aq) + H₂O(l)
Ag₂O(s) + 2HNO₃(aq) → 2AgNO₃(aq) + H₂O(l)
Na₂O(s) + 2HNO₃(aq) → 2NaNO₃(aq) + H₂O(l)

(iii) Reaction of insoluble/soluble carbonate/hydrogen carbonate with an acid.

By adding an excess of a soluble/insoluble carbonate or hydrogen carbonate to a dilute acid, effervescence/fizzing/bubbling out of carbon(IV) oxide gas shows the reaction is taking place. When effervescence/fizzing/bubbling out of the gas is over, excess of the insoluble carbonate is filtered off. The filtrate is heated to concentrate, allowed to crystallize then washed with distilled water before drying between filter papers, e.g.

PbCO₃(s) + 2HNO₃(aq) → Pb(NO₃)₂(aq) + H₂O(l) + CO₂(g)
ZnCO₃(s) + 2HNO₃(aq) → Zn(NO₃)₂(aq) + H₂O(l) + CO₂(g)
CaCO₃(s) + 2HNO₃(aq) → Ca(NO₃)₂(aq) + H₂O(l) + CO₂(g)
MgCO₃(s) + H₂SO₄(aq) → MgSO₄(aq) + H₂O(l) + CO₂(g)
CuCO₃(s) + H₂SO₄(aq) → CuSO₄(aq) + H₂O(l) + CO₂(g)
Ag₂CO₃(s) + 2HNO₃(aq) → 2AgNO₃(aq) + H₂O(l) + CO₂(g)
Na₂CO₃(s) + 2HNO₃(aq) → 2NaNO₃(aq) + H₂O(l) + CO₂(g)
K₂CO₃(s) + 2HCl(aq) → 2KCl(aq) + H₂O(l) + CO₂(g)
NaHCO₃(s) + HNO₃(aq) → NaNO₃(aq) + H₂O(l) + CO₂(g)
KHCO₃(s) + HCl(aq) → KCl(aq) + H₂O(l) + CO₂(g)

(iv) Neutralization/reaction of soluble base/alkali with dilute acid

By adding an acid to a burette into a known volume of an alkali with 2-3 drops of an indicator, the colour of the indicator changes when the acid has completely reacted with the alkali at the end point. The procedure is then repeated without the indicator. The solution mixture is then heated to concentrate, allowed to crystallize, washed with distilled water before drying with filter papers, e.g.

NaOH(aq) + HNO₃(aq) → NaNO₃(aq) + H₂O(l)
KOH(aq) + HNO₃(aq) → KNO₃(aq) + H₂O(l)
KOH(aq) + HCl(aq) → KCl(aq) + H₂O(l)
2KOH(aq) + H₂SO₄(aq) → K₂SO₄(aq) + 2H₂O(l)
2NH₄OH(aq) + H₂SO₄(aq) → (NH₄)₂SO₄(aq) + 2H₂O(l)
NH₄OH(aq) + HNO₃(aq) → NH₄NO₃(aq) + H₂O(l)

(v) Direct synthesis/combination.

When a metal burns in a gas jar containing a non-metal, the two directly combine to form a salt, e.g.

2Na(s) + Cl₂(g) → 2NaCl(s)
2K(s) + Cl₂(g) → 2KCl(s)
Mg(s) + Cl₂(g) → MgCl₂(s)
Ca(s) + Cl₂(g) → CaCl₂(s)

Some salts once formed undergo sublimation and hydrolysis. Care should be taken to avoid water/moisture entering the reaction flask during their preparation. Such salts include aluminium(III) chloride (AlCl₃) and iron(III) chloride (FeCl₃).

1. Heated aluminium foil reacts with chlorine to form aluminium(III) chloride that sublimes away from the source of heating then deposits as solid again:

2Al(s) + 3Cl₂(g) → 2AlCl₃(s/g)

Once formed, aluminium(III) chloride hydrolyses/reacts with water vapour/moisture present to form aluminium hydroxide solution and highly acidic fumes of hydrogen chloride gas.

AlCl₃(s) + 3H₂O(g) → Al(OH)₃(aq) + 3HCl(g)

2. Heated iron filings react with chlorine to form iron(III) chloride that sublimes away from the source of heating then deposits as solid again:

2Fe(s) + 3Cl₂(g) → 2FeCl₃(s/g)

Once formed, iron(III) chloride hydrolyses/reacts with water vapour/moisture present to form iron hydroxide solution and highly acidic fumes of hydrogen chloride gas.

FeCl₃(s) + 3H₂O(g) → Fe(OH)₃(aq) + 3HCl(g)

(b) Insoluble salts can be prepared by reacting two suitable soluble salts to form one soluble and one insoluble. This is called double decomposition or precipitation. The mixture is filtered and the residue is washed with distilled water then dried.

CuSO₄(aq) + Na₂CO₃(aq) → CuCO₃(s) + Na₂SO₄(aq)

BaCl₂(aq) + K₂SO₄(aq) → BaSO₄(s) + 2KCl(aq)

Pb(NO₃)₂(aq) + K₂SO₄(aq) → PbSO₄(s) + 2KNO₃(aq)

2AgNO₃(aq) + MgCl₂(aq) → 2AgCl(s) + Mg(NO₃)₂(aq)

Pb(NO₃)₂(aq) + (NH₄)₂SO₄(aq) → PbSO₄(s) + 2NH₄NO₃(aq)

BaCl₂(aq) + K₂SO₃(aq) → BaSO₃(s) + 2KCl(aq)

14. Salts may lose their water of crystallization, decompose, melt or sublime on heating on a Bunsen burner flame.

The following shows the behavior of some salts on heating gently or strongly in a laboratory school burner:

(a) Effect of heat on chlorides

All chlorides have very high melting and boiling points and therefore are not affected by laboratory heating except ammonium chloride. Ammonium chloride sublimes on gentle heating. It dissociates into the constituent ammonia and hydrogen chloride gases on strong heating.

NH₄Cl(s) ⇌ NH₄Cl(g) ⇌ NH₃(g) + HCl(g)

(sublimation) (dissociation)

(b) Effect of heat on nitrate(V)

(i) Potassium nitrate(V) (KNO₃) and sodium nitrate(V) (NaNO₃) decompose on heating to form potassium nitrate(III) (KNO₂) and sodium nitrate(III) (NaNO₂) and produce oxygen gas in each case.

2KNO₃(s) → 2KNO₂(s) + O₂(g)

2NaNO₃(s) → 2NaNO₂(s) + O₂(g)

(ii) Heavy metal nitrates(V) salts decompose on heating to form the oxide and a mixture of brown acidic nitrogen(IV) oxide and oxygen gases, e.g.

2Ca(NO₃)₂(s) → 2CaO(s) + 4NO₂(g) + O₂(g)

2Mg(NO₃)₂(s) → 2MgO(s) + 4NO₂(g) + O₂(g)

2Zn(NO₃)₂(s) → 2ZnO(s) + 4NO₂(g) + O₂(g)

2Pb(NO₃)₂(s) → 2PbO(s) + 4NO₂(g) + O₂(g)

2Cu(NO₃)₂(s) → 2CuO(s) + 4NO₂(g) + O₂(g)

2Fe(NO₃)₂(s) → 2FeO(s) + 4NO₂(g) + O₂(g)

(iii) Silver(I) nitrate(V) and mercury(II) nitrate(V) are lowest in the reactivity series. They decompose on heating to form the metal (silver and mercury) and the nitrogen(IV) oxide and oxygen gas, i.e.

2AgNO₃(s) → 2Ag(s) + 2NO₂(g) + O₂(g)

2Hg(NO₃)₂(s) → 2Hg(s) + 4NO₂(g) + O₂(g)

(iv) Ammonium nitrate(V) and ammonium nitrate(III) decompose on heating to nitrogen(I) oxide (relights/rekindles glowing splint) and nitrogen gas respectively. Water is also formed, i.e.

NH₄NO₃(s) → N₂O(g) + H₂O(l)

NH₄NO₂(s) → N₂(g) + H₂O(l)

Only iron(II) sulphate(VI), iron(III) sulphate(VI) and copper(II) sulphate(VI) decompose on heating. They form the oxide, and produce highly acidic fumes of acidic sulphur(IV) oxide gas.

2FeSO₄(s) → Fe₂O₃(s) + SO₃(g) + SO₂(g)

Fe₂(SO₄)₃(s) → Fe₂O₃(s) + SO₃(g)

CuSO₄(s) → CuO(s) + SO₃(g)

(d) Effect of heat on carbonates(IV) and hydrogen carbonate(IV)

(i) Sodium carbonate(IV) and potassium carbonate(IV) do not decompose on heating.

(ii) Heavy metal carbonate(IV) salts decompose on heating to form the oxide and produce carbon(IV) oxide gas. Carbon(IV) oxide gas forms a white precipitate when bubbled in lime water. The white precipitate dissolves if the gas is in excess. e.g.

CuCO₃(s) → CuO(s) + CO₂(g)

CaCO₃(s) → CaO(s) + CO₂(g)

PbCO₃(s) → PbO(s) + CO₂(g)

FeCO₃(s) → FeO(s) + CO₂(g)

ZnCO₃(s) → ZnO(s) + CO₂(g)

(iii) Sodium hydrogen carbonate(IV) and potassium hydrogen carbonate(IV) decompose on heating to give the corresponding carbonate(IV) and form water and carbon(IV) oxide gas, i.e.

2NaHCO₃(s) → Na₂CO₃(s) + CO₂(g) + H₂O(l)

2KHCO₃(s) → K₂CO₃(s) + CO₂(g) + H₂O(l)

(iii) Calcium hydrogen carbonate(IV) and magnesium hydrogen carbonate(IV) decompose on heating to give the corresponding carbonate(IV) and form water and carbon(IV) oxide gas, i.e.

Ca(HCO₃)₂(aq) → CaCO₃(s) + CO₂(g) + H₂O(l)

Mg(HCO₃)₂(aq) → MgCO₃(s) + CO₂(g) + H₂O(l)

15. Salts contain cations (positively charged ions) and anions (negatively charged ions). When dissolved in polar solvents/water, the cation and anion in a salt are determined usually by precipitation of the salt using a precipitating reagent.

The colour of the precipitate is a basis of qualitative analysis of a compound.

16. Qualitative analysis is the process of identifying an unknown compound/salt by identifying the unique qualities of the salt/compound.

It involves some of the following processes.

(a) Reaction of cation with sodium/potassium hydroxide solution.

Both sodium and potassium hydroxide solutions are precipitating reagents.

The alkalis produce unique colours of precipitates/suspensions when a few or three drops are added and then excess alkali is added to the unknown salt/compound solution.

NB: Potassium hydroxide is not commonly used because it is more expensive than sodium hydroxide.

The table below shows the observations, inferences/deductions and explanations from the following test tube experiments:

Procedure

Put about 2 cm³ of MgCl₂, CaCl₂, AlCl₃, NaCl, KCl, FeSO₄, Fe₂(SO₄)₃, CuSO₄, ZnSO₄, NH₄NO₃, Pb(NO₃)₂, Ba(NO₃)₂ each into separate test tubes. Add three drops of 2M sodium hydroxide solution then excess (2/3 the length of a standard test tube).

Observation	Inference	Explanation
No white precipitate	Na⁺ and K⁺	Both Na⁺ and K⁺ ions react with OH^– from 2M sodium hydroxide solution to form soluble colourless solutions.
No white precipitate then pungent smell of ammonia/urine	NH₄⁺ ions	NH₄⁺ ions react with 2M sodium hydroxide solution to produce pungent smelling ammonia gas.
White precipitate insoluble in excess	Ba²⁺, Ca²⁺, Mg²⁺ ions	Ba²⁺, Ca²⁺ and Mg²⁺ ions react with OH^– from 2M sodium hydroxide solution to form insoluble white precipitate of their hydroxides.
White precipitate soluble in excess	Zn²⁺, Pb²⁺, Al³⁺ ions	Pb²⁺, Zn²⁺ and Al³⁺ ions react with OH^– from 2M sodium hydroxide solution to form insoluble white precipitate of their hydroxides. The hydroxides formed react with more OH^– ions to form complex salts/ions.
Blue precipitate insoluble in excess	Cu²⁺	Cu²⁺ ions react with OH^– from 2M sodium hydroxide solution to form insoluble blue precipitate of copper(II) hydroxide.
Green precipitate insoluble in excess. On adding 3 cm³ of hydrogen peroxide, brown/yellow solution formed	Fe²⁺	Fe²⁺ ions react with OH^– from 2M sodium hydroxide solution to form insoluble green precipitate of iron(II) hydroxide. Hydrogen peroxide is an oxidizing agent that oxidizes green Fe²⁺ to brown Fe³⁺.
Brown precipitate insoluble in excess	Fe³⁺	Fe³⁺ ions react with OH^– from 2M sodium hydroxide solution to form insoluble brown precipitate of iron(III) hydroxide.

(b) Reaction of cation with aqueous ammonia

Aqueous ammonia is a precipitating reagent that can be used to identify the cations present in a salt.

Like NaOH/KOH, the OH^– ion in NH₄OH reacts with the cation to form a characteristic hydroxide.

Below are the observations, inferences and explanations of the reactions of aqueous ammonia with salts from the following test tube reactions.

Procedure

Put about 2 cm³ of MgCl₂, CaCl₂, AlCl₃, NaCl, KCl, FeSO₄, Fe₂(SO₄)₃, CuSO₄, ZnSO₄, NH₄NO₃, Pb(NO₃)₂, Ba(NO₃)₂ each into separate test tubes. Add three drops of 2M aqueous ammonia then excess (2/3 the length of a standard test tube).

Observation	Inference	Explanation
No white precipitate	Na⁺ and K⁺	NH₄⁺, Na⁺ and K⁺ ions react with OH^– from 2M aqueous ammonia to form soluble colourless solutions.
White precipitate insoluble in excess	Ba²⁺, Ca²⁺, Mg²⁺, Pb²⁺, Al³⁺ ions	Ba²⁺, Ca²⁺, Mg²⁺, Pb²⁺ and Al³⁺ ions react with OH^– from 2M aqueous ammonia to form insoluble white precipitate of their hydroxides.
White precipitate soluble in excess	Zn²⁺ ions	Zn²⁺ ions react with OH^– from 2M aqueous ammonia to form insoluble white precipitate of zinc hydroxide. The zinc hydroxide formed reacts with NH₃(aq) to form complex salts/ions.
Blue precipitate that dissolves in excess ammonia solution to form a deep/royal blue solution	Cu²⁺	Cu²⁺ ions react with OH^– from 2M aqueous ammonia to form blue precipitate of copper(II) hydroxide. The copper(II) hydroxide formed reacts with NH₃(aq) to form complex salts/ions.
Green precipitate insoluble in excess. On adding 3 cm³ of hydrogen peroxide, brown/yellow solution formed	Fe²⁺	Fe²⁺ ions react with OH^– from 2M aqueous ammonia to form insoluble green precipitate of iron(II) hydroxide. Hydrogen peroxide is an oxidizing agent that oxidizes green Fe²⁺ to brown Fe³⁺.
Brown precipitate insoluble in excess	Fe³⁺	Fe³⁺ ions react with OH^– from 2M aqueous ammonia to form insoluble brown precipitate of iron(III) hydroxide.

Note

Only Zn²⁺ ions/salts form a white precipitate that dissolves in excess of both 2M sodium hydroxide and 2M aqueous ammonia.
Pb²⁺ and Al³⁺ ions/salts form a white precipitate that dissolves in excess of 2M sodium hydroxide but not in 2M aqueous ammonia.
Cu²⁺ ions/salts form a blue precipitate that dissolves to form a deep/royal blue solution in excess of 2M aqueous ammonia but only blue insoluble precipitate in 2M sodium hydroxide.

All chlorides are soluble in water except Silver chloride and Lead (II) chloride (that dissolve in hot water). When a soluble chloride like NaCl, KCl, NH₄Cl is added to about 2 cm³ of a salt containing Ag⁺ or Pb²⁺ ions, a white precipitate of AgCl or PbCl₂ is formed. The following test tube reactions illustrate the above.

Experiment:

Put about 2 cm³ of silver nitrate(V) and lead(II) nitrate(V) solution into separate test tubes. Add five drops of NaCl/KCl/NH₄Cl/HCl. Heat to boil.

Observation	Inference	Explanation
White precipitate does not dissolve on heating	Ag⁺ ions	Ag⁺ ions react with Cl^– ions from a soluble chloride salt to form a white precipitate of AgCl.
White precipitate dissolves on heating	Pb²⁺ ions	Pb²⁺ ions react with Cl^– ions from a soluble chloride salt to form a white precipitate of PbCl₂. PbCl₂ dissolves on heating.

Note

Both Pb²⁺ and Al³⁺ ions form an insoluble white precipitate in excess aqueous ammonia. A white precipitate on adding Cl^– ions/salts shows Pb²⁺.

No white precipitate on adding Cl^– ions/salts shows Al³⁺.

Adding a chloride/Cl^– ions/salts can thus be used to separate the identity of Al³⁺ and Pb²⁺.

(d) Reaction of cation with sulphate(VI)/SO₄^2- and sulphate(IV)/SO₃^2- ions

All sulphate(VI) and sulphate(IV)/SO₃^2- ions/salts are soluble/dissolve in water except calcium sulphate(VI)/CaSO₄, calcium sulphate(IV)/CaSO₃, barium sulphate(VI)/BaSO₄, barium sulphate(IV)/BaSO₃, lead(II) sulphate(VI)/PbSO₄ and lead(II) sulphate(IV)/PbSO₃. When a soluble sulphate(VI)/SO₄^2- salt like Na₂SO₄, H₂SO₄, (NH₄)₂SO₄ or Na₂SO₃ is added to a salt containing Ca²⁺, Pb²⁺, Ba²⁺ ions, a white precipitate is formed.

The following test tube experiments illustrate the above.

Procedure

Place about 2 cm³ of Ca(NO₃)₂, Ba(NO₃)₂, BaCl₂ and Pb(NO₃)₂ in separate boiling tubes. Add six drops of sulphuric(VI) acid/sodium sulphate(VI)/ammonium sulphate(VI) solution. Repeat with six drops of sodium sulphate(IV).

Observation	Inference	Explanation
White precipitate	Ca²⁺, Ba²⁺, Pb²⁺ ions	CaSO₃ and CaSO₄ do not form a thick precipitate as they are sparingly soluble.

Ca²⁺(aq) + SO₃^2-(aq) → CaSO₃(s)

Ca²⁺(aq) + SO₄^2-(aq) → CaSO₄(s)

Ba²⁺(aq) + SO₃^2-(aq) → BaSO₃(s)

Ba²⁺(aq) + SO₄^2-(aq) → BaSO₄(s)

Pb²⁺(aq) + SO₃^2-(aq) → PbSO₃(s)

Pb²⁺(aq) + SO₄^2-(aq) → PbSO₄(s)

(e) Reaction of cation with carbonate(IV)/CO₃^2- ions

All carbonate salts are insoluble except sodium/potassium carbonate(IV) and ammonium carbonate(IV).

They dissociate/ionize to release CO₃^2- ions. CO₃^2- ions produce a white precipitate when the soluble carbonate salts are added to any metallic cation.

Procedure

Place about 2 cm³ of Ca(NO₃)₂, Ba(NO₃)₂, MgCl₂, Pb(NO₃)₂ and ZnSO₄ in separate boiling tubes.

Add six drops of potassium/sodium carbonate(IV)/ammonium carbonate(IV) solution.

Observation	Inference	Explanation
Green precipitate	Cu²⁺, Fe²⁺ ions	Copper(II) carbonate(IV) and Iron(II) carbonate(IV) are precipitated as insoluble green precipitates.
White precipitate	CO₃^2-	White precipitate of the carbonate(IV) salt is precipitated.

Note

Iron(III) carbonate(IV) does not exist.
Copper(II) carbonate(IV) exists only as the basic CuCO₃·Cu(OH)₂.
Both BaCO₃ and BaSO₃ are insoluble white precipitates. If hydrochloric acid is added to the white precipitate:

BaCO₃ produces CO₂ gas. When bubbled/directed into lime water solution, a white precipitate is formed.
BaSO₃ produces SO₂ gas. When bubbled/directed into orange acidified potassium dichromate(VI) solution, it turns green/decolorizes acidified potassium manganate(VII).

(f) Reaction of cation with sulphide/S^2- ions

All sulphides are insoluble black solids/precipitates except sodium sulphide (Na₂S) and potassium sulphide (K₂S). When a few drops of the soluble sulphide is added to a metal cation/salt, a black precipitate is formed.

Procedure

Place about 2 cm³ of Cu(NO₃)₂, FeSO₄, MgCl₂, Pb(NO₃)₂ and ZnSO₄ in separate boiling tubes.

Add six drops of potassium/sodium sulphide solution.

Observation	Inference	Explanation
Black precipitate	S^2- ions	CuS, FeS, MgS, PbS, ZnS are black insoluble precipitates.

Sample qualitative analysis guide

You are provided with solid Y (aluminium(III) sulphate(VI) hexahydrate). Carry out the following tests and record your observations and inferences in the space provided.

1(a) Appearance

Observations and inference (1 mark)

White crystalline solid. Coloured ions Cu²⁺, Fe²⁺, Fe³⁺ absent.

(b) Place about half a spatula full of the solid into a clean dry boiling tube. Heat gently then strongly.

Observations and inference (1 mark)

Colourless droplets formed on the cooler part of the test tube containing water of crystallization. Solid remains a white residue.

(c) Place all the remaining portion of the solid in a test tube. Add about 10 cm³ of distilled water. Shake thoroughly. Divide the mixture into five portions.