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S1.1 - Introduction to the particulate nature of matter S1.2 - The nuclear atom S1.3 - Electron configurations S1.4 - Counting particles by mass - The mole S1.5 - Ideal gases S2.1 - The ionic model S2.2 - The covalent model S2.3 - The metallic model S2.4 - From models to materials S3.1 - The periodic table - Classification of elements S3.2 - Functional groups - Classification of organic compounds R1.1 - Measuring enthalpy changes R1.2 - Energy cycles in reactions R1.3 - Energy from fuels R1.4 - Entropy and spontaneity AHL R2.1 - How much? The amount of chemical change R2.2 - How fast? The rate of chemical change R2.3 - How far? The extent of chemical change R3.1 - Proton transfer reactions R3.2 - Electron transfer reactions R3.3 - Electron sharing reactions R3.4 - Electron-pair sharing reactions

S3.1 - The periodic table - Classification of elements

3.1.1 Periodic Table Structure 3.1.2 Periodic, Group and Electron Configuration 3.1.3 Periodicity of Elements 3.1.4 Group Trends 3.1.5 Metallic to Non-Metallic Oxide Behaviour 3.1.6 Oxidation States 3.1.7 Ionization Energy Trends Exceptions (AHL) 3.1.8 Transition Element Properties (AHL) 3.1.9 Transition Element Oxidation States (AHL) 3.1.10 Colour and Transition Element Compounds (AHL)

Metallic to Non-Metallic Oxide Behaviour

Specification Reference S3.1.5

Quick Notes

  • Metal oxides (from Groups 1 & 2) are basic and react with water to form alkaline solutions.
  • Non-metal oxides (like CO2 and SO2) are acidic and form acidic solutions in water.
  • Amphoteric oxides (e.g., Al2O3) can act as both acids and bases.
  • The trend across a period moves from:
    • Basic → Amphoteric → Acidic
  • Environmental impacts:
    • Acid rain can form from SO2 and NO2 dissolving in rainwater.
    • Ocean acidification from CO2 dissolving in seawater, lowering pH.

Full Notes

Oxide Behaviour Across the Periodic Table

Oxides show a continuum from basic to acidic, reflecting the metallic to non-metallic character of the element:

Type of Element Type of Oxide Behaviour Example
Group 1 & 2 metals Basic Forms alkali in water Na2O, CaO
Transition metals Variable Often amphoteric ZnO, Al2O3
Non-metals Acidic Forms acids in water CO2, SO3

Oxide Reactions with Water

Group 1 Metal Oxide

ExampleSodium Oxide + Water

Na2O (s) + H2O (l) → 2NaOH (aq)

Group 2 Metal Oxide

ExampleCalcium Oxide + Water

CaO (s) + H2O (l) → Ca(OH)2 (aq)

These form alkaline solutions (high pH)

Non-Metal Oxides

ExampleCarbon Dioxide + Water

CO2 (g) + H2O (l) → H2CO3 (aq) (carbonic acid)

ExampleSulfur Dioxide + Water

CO2 (g) + H2O (l) → H2CO3 (aq) (carbonic acid)

These form acidic solutions (low pH).

Amphoteric Oxides

Example:Aluminium oxide (Al2O3)

Does not react with water.
Reacts with both acids and bases, showing amphoteric behaviour:

Environmental Issues

Acid Rain

Caused by acidic oxides SO2 and NO2 dissolving in water to form acids which can fall as acid rain:

Harms aquatic life, erodes buildings, damages forests.

Ocean Acidification

CO2 dissolves in seawater forming carbonic acid:

CO2 + H2O ⇌ H2CO3 ⇌ H+ + HCO3−

Lower pH affects marine organisms like corals and shellfish that rely on carbonate ions.

Linked Course Questions

Structure 2.1, 2.2 – Linked Course Question

How do differences in bonding explain the differences in the properties of metal and non-metal oxides?

Metal oxides have ionic bonding, forming giant lattices with high melting points; they are usually basic, reacting with acids. Non-metal oxides have covalent bonding, often forming small molecules with lower melting points; they tend to be acidic, reacting with bases. Amphoteric oxides, like Al2O3, have both ionic and covalent character. Their mixed bonding explains why they can react with both acids and bases.

Summary