A-level Chemistry/AQA/Module 5/Periodicity

From Wikibooks, open books for an open world
Jump to navigation Jump to search

The Reaction Of Period 3 Elements With Water[edit | edit source]

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[edit | edit source]

Formation of oxides[edit | edit source]

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[edit | edit source]

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[edit | edit source]

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) 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)[edit | edit source]

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

Formation of chlorides[edit | edit source]

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[edit | edit source]

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[edit | edit source]

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.