What you'll learn

• How p-block chemistry changes from non-metallic behaviour at the top of groups to more metallic behaviour lower down.
• Why some p-block compounds break the “octet rule”, including ${Al}_2{Cl}_6$ and ${NH}_3\cdot {BF}_3$.
• How Group 4 oxides and chlorides show trends in bonding, acid-base behaviour and oxidation-state stability.
• How to predict observations for lead(II), chlorine/alkali, and halide reactions with concentrated sulfuric acid.

1. The p-block: the big idea

The p-block contains elements whose highest-energy electrons are in a p sub-shell. In Eduqas “Group 3, 4, 5…” p-block language, this means the boron group, carbon group, nitrogen group, and so on.

Going down a p-block group:

• atoms get larger
• ionisation energies generally decrease
• metallic character increases
• lower oxidation states often become more stable

2. The inert pair effect

In Groups 3, 4 and 5, the outer configuration is based on ${ns}^2{np}^x$. Lower down these groups, the ${ns}^2$ pair is less likely to be used in bonding. This is called the inert pair effect.

So the more stable lower oxidation states are:

• Group 3: ${Tl}^{+}$ is more stable than ${Tl}^{3+}$
• Group 4: ${Pb}^{2+}$ is more stable than ${Pb}^{4+}$
• Group 5: ${Bi}^{3+}$ is more stable than ${Bi}^{5+}$

3. Amphoteric behaviour: reacting with acids and bases

An amphoteric substance reacts with both acids and bases. In this topic, aluminium and lead chemistry are the key examples.

Aluminium and aluminium(III)

Aluminium metal reacts with acids, producing hydrogen:

$$2Al(s)+6H^{+}(aq)\to 2{Al}^{3+}(aq)+3H_2(g)$$

It also reacts with hot concentrated alkali, forming aluminate ions:

$$2Al(s)+2{OH}^{-}(aq)+6H_{2}O(l)\to 2[{Al(OH)}_4{]}^{-}(aq)+3H_2(g)$$

For ${Al}^{3+}(aq)$, adding aqueous sodium hydroxide gives a white precipitate of aluminium hydroxide:

$${Al}^{3+}(aq)+3{OH}^{-}(aq)\to {Al(OH)}_3(s)$$

In excess hydroxide, the precipitate dissolves:

$${Al(OH)}_3(s)+{OH}^{-}(aq)\to [{Al(OH)}_4{]}^{-}(aq)$$

Lead and lead(II)

Lead(II) hydroxide behaves similarly. Adding hydroxide to ${Pb}^{2+}(aq)$ gives a white precipitate:

$${Pb}^{2+}(aq)+2{OH}^{-}(aq)\to {Pb(OH)}_2(s)$$

In excess hydroxide, it dissolves to form a plumbate(II) complex:

$${Pb(OH)}_2(s)+2{OH}^{-}(aq)\to [{Pb(OH)}_4{]}^{2-}(aq)$$

4. Fewer than eight and more than eight electrons

Some Group 3 compounds are electron-deficient, meaning the central atom has fewer than eight electrons around it. Examples include ${BF}_3$ and ${AlCl}_3$.

Some elements in Groups 5, 6 and 7, especially from Period 3 downwards, can form compounds with more than eight electrons around the central atom. Examples include ${PCl}_5$, ${SF}_6$ and ${IF}_7$.

5. Dative bonding: ${Al}_2{Cl}_6$ and ${NH}_3\cdot {BF}_3$

A dative covalent bond, also called a coordinate bond, is a covalent bond where both shared electrons come from the same atom.

In ${AlCl}_3$, aluminium has only six electrons around it. Two ${AlCl}_3$ units can join to form ${Al}_2{Cl}_6$, using bridging chlorine atoms that donate lone pairs towards electron-deficient aluminium atoms.

In ${NH}_3\cdot {BF}_3$, ammonia donates the lone pair on nitrogen to electron-deficient boron in ${BF}_3$:

$${NH}_3+{BF}_3\to H_{3}N\to {BF}_3$$

6. Boron nitride: same formula type, different structures

Boron nitride contains alternating boron and nitrogen atoms, but it can exist in different giant covalent structures.

Hexagonal boron nitride, h-BN, is like graphite:

• layers of hexagonal rings
• strong covalent bonds within each layer
• weak forces between layers
• slippery, so useful as a lubricant
• electrical insulator because electrons are not delocalised as in graphite

Cubic boron nitride, c-BN, is like diamond:

• three-dimensional tetrahedral giant covalent lattice
• very hard
• high melting point
• used in abrasives and cutting tools

7. Group 4: oxidation states, oxides and chlorides

Oxidation states II and IV

Group 4 elements can show +2 and +4 oxidation states. Down the group, +2 becomes more stable and +4 becomes less stable. This is especially important for lead.

Carbon and silicon strongly favour +4. Lead favours +2, so lead(IV) compounds are often oxidising agents because they are reduced to lead(II).

Carbon monoxide can act as a reducing agent because carbon in CO is oxidised from +2 to +4 in ${CO}_2$. For example:

$${PbO}_2(s)+CO(g)\to PbO(s)+{CO}_2(g)$$

Lead(IV) oxide is reduced to lead(II) oxide.

Carbon dioxide and lead(II) oxide

${CO}_2$ is a simple molecular covalent substance. It is a gas at room temperature, has a low boiling point, and does not conduct electricity. It is an acidic oxide:

$${CO}_2(g)+H_{2}O(l)\rightleftharpoons H_2{CO}_3(aq)$$

It reacts with alkali, for example:

$${CO}_2(g)+2NaOH(aq)\to {Na}_2{CO}_3(aq)+H_{2}O(l)$$

$PbO$ is a solid with much more ionic character and a high melting point. It is amphoteric:

$$PbO(s)+2H^{+}(aq)\to {Pb}^{2+}(aq)+H_{2}O(l)$$ $$PbO(s)+2{OH}^{-}(aq)+H_{2}O(l)\to [{Pb(OH)}_4{]}^{2-}(aq)$$

Group 4 chlorides and water

The chlorides show a bonding trend from covalent molecular chlorides to more ionic chlorides.

${CCl}_4$ is a simple covalent molecular liquid. It does not hydrolyse with water.

${SiCl}_4$ is also simple molecular, but it hydrolyses rapidly in water, giving steamy fumes of hydrogen chloride:

$${SiCl}_4(l)+2H_{2}O(l)\to {SiO}_2(s)+4HCl(aq)$$

${PbCl}_2$ is much more ionic. It is a white solid, sparingly soluble in cold water but more soluble in hot water, and it does not hydrolyse like ${SiCl}_4$.

8. Lead(II) ion tests

For ${Pb}^{2+}(aq)$:

• with aqueous sodium hydroxide: white ${Pb(OH)}_2$ precipitate, soluble in excess hydroxide
• with chloride ions: white ${PbCl}_2$ precipitate, soluble in hot water
• with iodide ions: bright yellow ${PbI}_2$ precipitate, forming golden crystals on cooling from hot solution

The key ionic equations are:

$${Pb}^{2+}(aq)+2{Cl}^{-}(aq)\to {PbCl}_2(s)$$ $${Pb}^{2+}(aq)+2I^{-}(aq)\to {PbI}_2(s)$$

9. Chlorine with sodium hydroxide: disproportionation

A disproportionation reaction is one where the same element is both oxidised and reduced.

With cold dilute sodium hydroxide:

$${Cl}_2(aq)+2{OH}^{-}(aq)\to {Cl}^{-}(aq)+{ClO}^{-}(aq)+H_{2}O(l)$$

Chlorine goes from 0 to -1 in ${Cl}^{-}$ and from 0 to +1 in ${ClO}^{-}$.

With warm concentrated sodium hydroxide, chlorate(V) forms:

$$3{Cl}_2(aq)+6{OH}^{-}(aq)\to 5{Cl}^{-}(aq)+{ClO}_3^{-}(aq)+3H_{2}O(l)$$

This can be understood as further disproportionation of chlorate(I):

$$3{ClO}^{-}(aq)\to 2{Cl}^{-}(aq)+{ClO}_3^{-}(aq)$$

Chlorine and chlorate(I) ions bleach and kill bacteria because they are oxidising agents. In water:

$${Cl}_2(aq)+H_{2}O(l)\rightleftharpoons HCl(aq)+HClO(aq)$$

Chloric(I) acid, $HClO$, and chlorate(I), ${ClO}^{-}$, oxidise coloured molecules and essential bacterial cell components.

10. Sodium halides with concentrated sulfuric acid

Concentrated sulfuric acid first acts as an acid:

$$NaX(s)+H_2{SO}_4(l)\to {NaHSO}_4(s)+HX(g)$$

where X is Cl, Br or I.

Then the behaviour depends on how strongly reducing the halide ion is. Reducing power increases down Group 7.

For $NaCl$, steamy white fumes of $HCl$ form. Chloride is not a strong enough reducing agent to reduce concentrated sulfuric acid.

For $NaBr$, $HBr$ forms, then reduces sulfuric acid to sulfur dioxide:

$$2HBr(g)+H_2{SO}_4(l)\to {Br}_2(g)+{SO}_2(g)+2H_{2}O(l)$$

You may see orange-brown bromine and smell choking ${SO}_2$.

For $NaI$, $HI$ is an even stronger reducing agent. It can reduce sulfuric acid to ${SO}_2$, sulfur, and hydrogen sulfide:

$$2HI(g)+H_2{SO}_4(l)\to I_2(s)+{SO}_2(g)+2H_{2}O(l)$$ $$6HI(g)+H_2{SO}_4(l)\to 3I_2(s)+S(s)+4H_{2}O(l)$$ $$8HI(g)+H_2{SO}_4(l)\to 4I_2(s)+H_{2}S(g)+4H_{2}O(l)$$

Observations include purple iodine vapour or black iodine solid, yellow sulfur, and a rotten-egg smell from $H_{2}S$.