A-level Chemistry/AQA/Module 5/Periodicity

The Reaction Of Period 3 Elements With Water

Na, Mg, Al and Si are more electropositive than H and can reduce the water to hydrogen gas:

Na reacts vigorously with water to give the hydroxide and hydrogen:

2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g)

The resulting solution is strongly alkaline, and will have a pH of 14.

Mg reacts with cold water very slowly, but can react quickly with steam to give the oxide and hydrogen:

Mg(s) + 2H2O(l) → Mg(OH)2(aq) + H₂(g)

Mg(s) + H2O(g) → MgO(s) + H2(g)

The resulting solution is weakly alkaline, since the oxide is slightly basic (pH = 9).

Al and Si also react with steam under certain conditions.

P, S and Cl₂ do not reduce water to hydrogen gas. Phosphorus and sulphur do not react with water but chlorine will disproportionate to give an acidic solution:

Cl₂(g) + H₂O(l) → HClO(aq) + HCl(aq)

The resulting solution contains HCl(aq) and is thus acidic (pH = 2).

The reactivity of the elements of period 3 towards water thus decreases from Na to Si, and then increases from P to Cl. The resulting solutions become increasingly acidic.

The Oxides Of Period 3 Elements

Formation of oxides

All the elements in Period 3 except chlorine and argon combine directly with oxygen to form oxides.

4Na(s) + O₂(g) → 2Na₂O(s)

Na₂O is an ionic oxide. This reaction is very vigorous as sodium burns with a yellow flame

2Na(s) + O2(g) → Na2O2(s)

Na can also react with O2 to form Na2O2(sodium peroxide).

2Mg(s) + O₂(g) → 2MgO(s)

MgO is also an ionic oxide. This reaction is vigorous with a brilliant white flame forming a white ash of magnesium oxide.

4Al(s) + 3O₂(g) → 2Al₂O₃(s)

Al₂O₃ is mostly ionic, but there is significant covalent character. This reaction is initially vigorous.

Si(s) + O₂(g) → SiO₂(s)

SiO₂ is a giant covalent oxide. This reaction is slow.

P₄(s) + 5O₂(g) → P₄O₁₀(s)

P₄O₁₀ is a molecular covalent oxide. The oxidation number of P in this oxide is +5. This is a vigorous reaction forming masses of white fumes of phosphorus(V) oxide.

S(s) + O₂(g) → SO₂(g)

SO₂ is a molecular covalent oxide. Sulphur melts easily and burns with a blue flame forming sulphur(IV) oxide (sulphur dioxide), a colourless gas with a choking odour.

Another oxide, SO₃ is formed in a reversible process when SO₂ and O₂ are heated with a V₂O₅ catalyst (the Contact Process).

Physical properties of oxides

The physical properties of these oxides depend on the type of bonding.

Na₂O, Al₂O₃ and MgO are ionic oxides and hence have a high melting point. MgO and Al₂O₃ have a higher melting point than Na₂O since the charges are higher and the atomic radii smaller, resulting in a stronger electrostatic attraction (forces) between the ions.

SiO₂ has a giant covalent structure and hence a high melting point. There are strong covalent bonds between all the atoms and thus lots of energy is required to break them.

P₄O₁₀ and SO₃ are molecular covalent and so only intermolecular forces (Van Der Waals) exist between the molecules. The melting points are thus much lower. P₄O₁₀ is a much bigger molecule than SO₃ and so has a much higher melting point, as the van der Waal’s forces are stronger.

Element	Formulae of oxide	Structure of oxide	Melting point of oxide /°C
Na	Na2O	Ionic	1275
Mg	MgO	Ionic	2852
Al	Al2O3	Mostly Ionic	2072
Si	SiO2	Giant Covalent	1703
P	P4O10	Molecular Covalent	300
S	SO3	Molecular Covalent	-10

Acid-base character of oxides

Ionic oxides contain the O²⁻ ion. This is a strongly basic ion which reacts with water to produce hydroxide ions:

O²⁻(aq) + H₂O(l) → 2OH_-(aq)

Thus all ionic oxides are BASIC.

Covalent oxides do not contain ions, but have a strongly positive dipole on the atom which is not oxygen. This attracts the lone pair on water molecules, releasing H ions:

MO(s) + H₂O(l) → MO(OH)_-(aq) + H⁺(aq)

Thus all covalent oxides are ACIDIC.

Intermediate oxides can react in either of the above ways, depending on the conditions. They can thus behave as either acids or bases and are thus AMPHOTERIC.

Na₂O is a basic oxide. It dissolves in water to give an alkaline solution (pH = 14). It also reacts with acids:

Na₂O(s) + H₂O(l) → 2NaOH(aq)

Na₂O(s) + 2H⁺(aq) → 2Na⁺(aq) + H₂O(l)

MgO is a basic oxide. It is only slightly soluble in water and so the solution is only slightly alkaline (pH = 9). It reacts readily with acids:

MgO(s) + H₂O(l) == Mg(OH)₂(s) == Mg(OH)₂(aq)

MgO(s) + 2H⁺(aq) → Mg²⁺(aq) + H₂O(l)

Al₂O₃ is an amphoteric oxide. It is insoluble in water (pH = 7) but dissolves in both acids and alkalis:

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

Al₂O₃ + 3H₂O(l) + 6OH_-(aq) → 2Al(OH)3−6(aq)

Al₂O₃ + 3H₂O(l) + 2OH_-(aq) → 2Al(OH)₄-(aq)

SiO₂ is an acidic oxide. It is insoluble in water (pH = 7) but dissolves in hot concentrated alkalis:

SiO₂(s) + 2OH_-(aq) → SiO2−3(aq) + H₂O(l)

P₄O₁₀ is an acidic oxide. It dissolves in water to give acidic solutions and is also soluble in alkalis:

P₄O₁₀(s) + 6H₂O(l) → 4H₃PO₄(aq)   pH = 3

P₄O₁₀(s) + 12OH_-(aq) → 4PO3−4(aq) + 6H₂O(l)

SO₂ and SO₃ are acidic oxides. They dissolve in water to give acidic solutions, and also react with alkalis:

SO₂(g) + H₂O(l) $\rightleftharpoons$ H₂SO₃(aq)   pH = 2

SO₃(g) + H₂O(l) → H₂SO₄(aq)   pH = 1

SO₂(g) + 2OH_-(aq) → SO2−3(aq) + H₂O(l)

SO₃(g) + 2OH_-(aq) → SO2−4(aq) + H₂O(l)

SO₂ is a waste gas in many industrial processes. It is harmful because it dissolves in rain water to give acid rain. It can be removed from waste gases because it dissolves in alkali and so it is passed through an alkaline solution in waste gas outlets to minimise the amount which escapes into the atmosphere.

The acid-base properties of the oxides of Period 3 can be summarised in the following table:

Element	Na	Mg	Al	Si	P	S
Formulae of oxide	Na2O	MgO	Al2O3	SiO2	P4O10	SO2	SO3
Acid-base character of oxide	Basic	Basic	Amphoteric	Acidic	Acidic	Acidic
pH of solution when dissolved in water	12 - 14	8 - 9	7 (insoluble)	7 (insoluble)	2 - 4	2 - 4	1 - 3

The oxides therefore become more acidic on moving from left to right in the periodic table.

The Chlorides Of Period 3 Elements (off syllabus since Jan 2010)

As of January 2010, Period 3 chlorides are no longer on the AQA Chemistry Unit 5 syllabus

Formation of chlorides

All the elements of Period 3 except argon combine directly with chlorine to give chlorides.

2Na(s) + Cl₂(g) → 2NaCl(s)

NaCl is an ionic chloride. This is a very vigorous reaction.

Mg(s) + Cl₂(g) → MgCl₂(s)

MgCl₂ is also an ionic chloride. Vigorous reaction when the elements are heated.

2Al(s) + 3Cl₂(g) → 2AlCl₃(s)

AlCl₃ is covalent. It forms a polymeric structure in the solid state, turning quickly on heating into a dimeric gas (Al₂Cl₆). It thus behaves as a simple molecular chloride. This reaction is vigorous when the elements are heated under anhydrous conditions.

Si(s) + 2Cl₂(g) → SiCl₄(s)

SiCl₄ is a molecular covalent chloride. Slow reaction.

P₄(s) + 10Cl₂(g) → 4PCl₅(s)

This is a slow reaction. P₂Cl₁₀ is actually ionic in the solid state - it exists as [PCl₄]⁺[PCl₆]_- in the solid state.

Physical properties of chlorides

NaCl and MgCl₂ are ionic chlorides. Since a large amount of energy is required to separate the ions, the melting point is high.

AlCl₃ and SiCl₄ are molecular covalent chlorides, and so only intermolecular forces exist between the molecules. The melting points are thus much lower than the ionic chlorides.

AlCl₃ actually exists in polymeric form in the solid state, which is converted to a dimeric form in the gas phase. This is because the aluminium atom is electron deficient – it has only 3 of its four valence orbitals occupied, so it has an empty orbital with which it can accept a lone pair of electrons from a Cl atom on an adjacent monomer. At high temperatures, it reverts to a simple molecular structure.

PCl₅ is ionic so its melting point is thus high. On heating, however, it reverts to a simple covalent structure and sublimes.

Element Na Mg Al Si P Formula of chloride NaCl MgCl2 AlCl3 SiCl4 PCl5 Structure of chloride ionic ionic polymer molecular covalent Ionic Melting point of chloride /°C 801 710 184 58 162

Reaction of chlorides with water

The way in which chlorides react with water depends on the type of bonding present in the chloride:

Ionic chlorides dissolve in water to give neutral solutions:

NaCl(s) → Na⁺(aq) + Cl_-(aq)   pH = 7

MgCl₂(s) → Mg²⁺(aq) + 2Cl_-(aq)   pH = 7

Aluminium chloride reacts with water to give hydrated aluminium ions and chloride ions. The hydrated aluminium ions undergo deprotonation to give an acidic solution:

AlCl₃(s) + 6H₂O(l) → Al(H₂O)3+6(aq) + 3Cl_-(aq)

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

The other covalent chlorides react readily with water at room temperature to form the oxide or hydroxide and HCl(g). The HCl is formed as white misty fumes, and the observance of these fumes is a good indication that the chloride is covalent.

SiCl₄(l) + 2H₂O(l) → SiO₂(s) + 4HCl(g)   pH = 1 - 2

PCl₅(s) + 4H₂O(l) → H₃PO₄(aq) + 5HCl(g)   pH = 1 - 2

Covalent chlorides thus react with water to give acidic solutions. The acidity is due to dissolved HCl.

The water molecules attack the covalent chlorides by donating lone pairs of electrons into empty low-lying orbitals on the electropositive atoms. In the case of AlCl₃, there is an available 3p orbital, and in SiCl₄ and PCl₅ there are available d-orbitals:

  3s                        3p    3d

AlCl₃ [Ne] ↓↑ ↓↑ ↓↑

SiCl₄ [Ne] ↓↑ ↓↑ ↓↑ ↓↑

PCl₅ [Ne] ↓↑ ↓↑ ↓↑ ↓↑ ↓↑

It is the availability of these low-lying empty orbitals which enables these chlorides to react readily with water.