Metallic & Non-Metallic Oxides (DP IB Chemistry): Revision Note

Metallic & Non-Metallic Oxides

Oxides across a period

• The acid-base character of the oxides shows chemical trends in the periodic table

• The broad trend is that oxides change from basic through amphoteric to acidic across a period

• Amphoteric means a substance can act as both an acid and a base

  • For example, aluminium oxide can react with both HCl and NaOH

Acidic and basic nature of Period 3 oxides

• Oxides on the left side of Period 3 are basic

  • e.g. sodium and magnesium oxides

• Aluminium oxide is amphoteric

  • This means that it can act react with acids and bases

• Oxides on the right side are acidic

  • e.g. silicon dioxide, phosphorus pentoxide, sulfur dioxide and sulfur trioxide

• This trend from basic to acidic oxides can be explained by examining the bonding and electronegativity of the elements involved

Structure, bonding and electronegativity of Period 3 oxides

• Sodium oxide (Na₂O) and magnesium oxide (MgO)

  • Giant ionic solids with high melting points

• Aluminium oxide (Al₂O₃)

  • Giant ionic solid with a very high melting point

  • Aluminium oxide has some covalent character due to the small size and high charge of the Al³⁺ ion

• Silicon dioxide (SiO₂)

  • Giant covalent solid with a very high melting point

• Phosphorus pentoxide (P₄O₁₀) and sulfur oxides (SO₂ and SO₃)

  • Simple molecular covalent compounds with low melting points

• The electronegativity values of the elements increase across the period:

  • Na (0.9) < Mg (1.2) < Al (1.5) < Si (1.8) < P (2.1) < S (2.5)

• Oxygen has an electronegativity of 3.4

• This means the difference in electronegativity is greatest between oxygen and the metals on the left:

  • For Na: ΔEN = 3.4 - 0.9 = 2.5

  • For Mg: ΔEN = 3.4 - 1.2 = 2.2

  • For Al: ΔEN = 3.4 - 1.5 = 1.9

• These large differences result in electron transfer from metal to oxygen and the formation of purely ionic oxides.

• Moving across the period, the electronegativity difference decreases, so bonding becomes more covalent.

• The largest difference in electronegativity is between oxygen (3.4) and the metals

  • Na: 3.4 - 0.9 = 2.5

  • Mg: 3.4 - 1.2 = 2.2

  • Al : 3.4 - 1.5 = 1.9

• These large differences result in electron transfer from metal to oxygen and the formation of purely ionic oxides

• Moving across the period, the electronegativity difference decreases

  • This causes the bonding to become more covalent

Oxides with water

Basic oxides (Group 1 and 2 metals):

• Metallic oxides tend to form hydroxides in water, resulting in alkaline solutions

• Sodium oxide reacts with water to form sodium hydroxide

  • Sodium hydroxide is a strongly alkaline solution (pH ≈ 14)

Na₂O (s) + H₂O (l) → 2NaOH (aq)

• Magnesium oxide reacts with water to form magnesium hydroxide

  • Magnesium hydroxide is a weakly alkaline solution (pH ≈ 10)

MgO (s) + H₂O (l) → Mg(OH)₂ (aq)

• Oxide ions (O^2-) react with water to form hydroxide ions:

O^2- (aq) + H₂O (l) → 2OH^- (aq)

Acidic oxides (non-metals):

• Non-metallic oxides tend to form oxoacids in water, resulting in acidic solutions

• The oxides of phosphorus and sulfur form purely covalent molecules, and react with water to form acids

• For example, phosphorus pentoxide forms phosphoric acid

  • Phosphoric acid is strongly acidic (pH ≈ 2)

P₄O₁₀ (s) + 6H₂O (l) → 4H₃PO₄ (aq)

• Other non-metallic oxides, like nitrogen dioxide and sulfur oxides, also form acids

• Nitrogen dioxide forms nitric acid and nitrous acid

  • Nitric acid and nitrous acid are strongly acidic (pH ≈ 1)

2NO₂ (aq) + H₂O (l) → HNO₃ (aq) + HNO₂ (aq)

• Sulfur dioxide forms sulfurous acid:

  • Sulfurous acid is strongly acidic (pH ≈ 1)

SO₂ (g) + H₂O (l) → H₂SO₃ (aq)

• Sulfur trioxide forms sulfuric acid

  • Sulfuric acid is strongly acidic (pH ≈ 1)

SO₃ (g) + H₂O (l) → H₂SO₄ (aq)

Amphoteric oxides:

• Aluminium oxide (Al₂O₃) does not react with water but can react with both acids and bases

Environmental links

Acid rain

• Acid rain is caused by non-metal oxides dissolving in atmospheric water:

SO₂ (g) + H₂O (l) → H₂SO₃ (aq)

2NO₂ (aq) + H₂O (l) → HNO₃ (aq) + HNO₂ (aq)

• These reactions produce strong acids, lowering the pH of rainwater and harming ecosystems

Ocean acidification

• Ocean acidification occurs when CO₂ dissolves in seawater to form carbonic acid:

CO₂ (g) + H₂O (l) ⇌ H₂CO₃ (aq)

• The weak carbonic acid dissociates:

H₂CO₃ (aq) ⇌ H⁺ (aq) + HCO₃⁻ (aq)

• The increase in H⁺ concentration lowers the pH of oceans, impacting marine life

Predicting oxide properties

• Trends in ionisation energy, electron affinity, and electronegativity help predict oxide bonding and acid–base character

• Metals (like Na, Mg) form ionic oxides, typically basic in character

• Aluminium forms an ionic oxide with covalent character, which has amphoteric properties

• Non-metals (like Si, P, S) form covalent oxides, typically acidic in character

• The oxides become:

  • More ionic as you go down the group as the electronegativity decreases

  • Less ionic as you go across a period as the electronegativity increases

• These trends can be used to deduce the bonding, structure, and acid–base behaviour unfamiliar oxides