Oxides of Period 3 Elements
This lesson covers:
- Reactions of period 3 elements with oxygen
- The formulae and oxidation numbers of period 3 oxides
- The effect of structure and bonding on the reactions of period 3 oxides with water
- The transition from alkaline to acidic oxides across period 3
- The reactions of period 3 oxides with acids and alkalis
Reactions of period 3 elements with oxygen
The period 3 elements, except noble gas argon, react with oxygen in exothermic redox reactions to form oxides.
Sodium - Burns brightly in air with a characteristic yellow flame to form white sodium oxide.
2Na(s) + 1⁄2O2(g) ➔ Na2O(s)
Magnesium - Burns in oxygen with a bright white flame, forming white magnesium oxide powder.
2Mg(s) + O2(g) ➔ 2MgO(s)
Aluminium - Its powder burns brightly when heated in oxygen, producing white aluminium oxide powder.
4Al(s) + 3O2(g) ➔ 2Al2O3(s)
Silicon - When heated strongly in oxygen, silicon reacts to form silicon dioxide.
Si(s) + O2(g) ➔ SiO2(s)
Phosphorus - White phosphorus ignites spontaneously in air, giving off white smoke of phosphorus pentoxide.
4P(s) + 5O2(g) ➔ P4O10(s)
Sulfur - Its powder burns in oxygen with a blue flame to produce sulfur dioxide gas.
S(s) + O2(g) ➔ SO2(g)
Oxidation numbers in period 3 oxides
The table below shows the common oxide compounds formed by period 3 elements, along with their oxidation numbers:
| Period 3 element | Sodium | Magnesium | Aluminium | Silicon | Phosphorus | Sulfur |
|---|---|---|---|---|---|---|
| Oxide formula | NaO | MgO | AlO | SiO | PO | SO / SO |
| Oxidation number | +1 | +2 | +3 | +4 | +5 | +4 / +6 |
The oxidation number of the period 3 element in the oxide increases as you move across the period. This happens because the period 3 elements use all the electrons in their outer shell when bonding to oxygen.
The period 3 elements exist in positive states in the oxide because oxygen has the highest electronegativity of any period 3 element.
Reactions of period 3 oxides with water
The oxides show different reactions with water due to differences in their structure and bonding.
Sodium and magnesium oxides
Na2O and MgO have giant ionic lattice structures. When dissolved in water, these lattices separate into hydrated ions:
- Na2O(s) ➔ 2Na+(aq) + O2−(aq)
- MgO(s) + H2O(l) ➔ Mg2+(aq) + O2−(aq)
The oxide ions (O2−) react with water molecules to form hydroxide ions (OH−):
- O2−(aq) + H2O(l) ➔ 2OH−(aq)
This makes the solution alkaline. For example:
- Na2O(s) + H2O(l) ➔ 2NaOH(aq) (pH 12-14)
- MgO(s) + H2O(l) ➔ Mg(OH)2(aq) (pH ≈10)
The difference in pH between the solutions formed by MgO and Na2O is due to the difference in the solubility of their hydroxides; NaOH is more soluble than Mg(OH)2, resulting in a higher concentration of OH- ions and thus a higher pH.
Aluminium oxide
Al2O3 has an ionic structure with some covalent character. It does not dissolve or react with water due to its giant structure.
Silicon dioxide
SiO2 has a giant covalent structure, so it does not dissolve or react with water.
Phosphorus pentoxide and sulfur oxides
P4O10, SO2 and SO3 have simple molecular structures.
When water is added, they react vigorously and dissolve to form acidic solutions containing hydrogen ions (H+) that are donated as the acidic oxides dissolve.
- Phosphorus pentoxide P4O10 dissolves in water to form phosphoric (V) acid: P4O10(s) + 6H2O(l) ➔ 4H3PO4(aq)
This reaction produces an acidic solution with a pH of 2.
- Sulfur dioxide SO2 dissolves in water to form sulfurous acid: SO2(g) + H2O(l) ➔ H2SO3(aq)
- Sulfur trioxide SO3 dissolves in water to form sulfuric acid: SO3(g) + H2O(l) ➔ H2SO4(aq)
The sulfur oxides reacting with water create acidic solutions with a pH of 0-3.
Trend from alkaline to acidic oxides
As you move left to right across period 3, the oxides transition from alkaline to acidic:
- Alkaline oxides - Sodium oxide, magnesium oxide
- Acidic oxides - Silicon dioxide, phosphorus pentoxide, sulfur dioxide, sulfur trioxide
So the alkaline oxides on the left will react with acids, while the acidic oxides on the right will react with bases.
Aluminium oxide reacts with both acids and alkalis so it is classed as an amphoteric oxide.
Reactions of alkaline oxides with acids
The alkaline metal oxides Na2O and MgO neutralise acids to form salts and water.
For example, sodium oxide reacts with sulfuric acid to produce sodium sulfate and water:
- Na2O(s) + H2SO4(aq) ➔ Na2SO4(aq) + H2O(l)
For example, magnesium oxide reacts with hydrochloric acid to produce magnesium chloride and water:
- MgO(s) + 2HCl(aq) ➔ MgCl2(aq) + H2O(l)
Reaction of aluminium oxide with acids and alkalis
The amphoteric aluminium oxide neutralises both acids and alkalis.
The reaction with acids produces a salt and water.
For example, Al2O3 reacts with hydrochloric acid to produce aluminium chloride and water:
- Al2O3(s) + 6HCl(aq) ➔ 2AlCl3(aq) + 3H2O(l)
The reaction with alkalis produces an aluminate salt.
For example, Al2O3 reacts with hot, concentrated sodium hydroxide to produce sodium aluminate:
- Al2O3(s) + 2NaOH(aq) + 3H2O(l) ➔ 2NaAl(OH)4(aq)
Reaction of acidic oxides with alkalis
The acidic oxides SiO2, P4O10, SO2, and SO3 neutralise alkalis to produce salt solutions.
Silicon dioxide reactions
Silicon dioxide reacts as a weak acid with strong bases to produce a silicate salt.
For example, SiO2 reacts with hot, concentrated sodium hydroxide to produce sodium silicate:
- SiO2(s) + 2NaOH(aq) ➔ Na2SiO3(aq) + H2O(l)
Phosphorus pentoxide reactions
Phosphorus pentoxide reacts with water to produce phosphoric(V) acid (H3PO4). It is actually the phosphoric acid that then reacts with alkalis, not the phosphorus pentoxide itself.
Phosphoric acid reacts with alkalis to produce phosphate salts and water.
For example, H3PO4 reacts with sodium hydroxide to produce sodium phosphate and water:
- H3PO4(aq) + 3NaOH(aq) ➔ Na3PO4(aq) + 3H2O(l)
The reaction occurs in three stages because H3PO4 contains three -OH groups with acidic hydrogens. The hydrogens are replaced by sodium in stages forming the intermediate products NaH2PO4 and Na2HPO4.
Reactions of sulfur oxides
Sulfur dioxide and sulfur trioxide neutralise alkalis to form salts and water.
The reaction with sulfur dioxide produces sulfite salts.
For example, SO2 reacts with sodium hydroxide to produce sodium sulfite and water:
- SO2(g) + 2NaOH(aq) ➔ Na2SO3(aq) + H2O(l)
The reaction with sulfur trioxide produces sulfate salts.
For example, SO3 reacts with sodium hydroxide to produce sodium sulfate and water:
- SO3(g) + 2NaOH(aq) ➔ Na2SO4(aq) + H2O(l)
Summary of properties of period 3 oxides
The table below summarises some properties, chemical bonding and structures of the period 3 oxides:
| Period 3 oxide | Na2O | MgO | Al2O3 | SiO2 | P4O10 | SO2 | SO3 |
|---|---|---|---|---|---|---|---|
| Chemical bonding | Ionic | Ionic | Ionic (with a degree of covalent character) | Covalent | Covalent | Covalent | Covalent |
| Structure | Giant ionic | Giant ionic | Giant ionic | Giant covalent | Simple molecular | Simple molecular | Simple molecular |
| Acid/base nature | Basic | Basic | Amphoteric | Acidic | Acidic | Acidic | Acidic |
| pH of solution | 12-14 | 10 | 7 | 7 | 1-2 | 2-3 | 0-1 |