Oxides | Chemistry for GCSE/IGCSE - Class 10 PDF Download

Classifying Oxides

• Oxides are compounds formed by combining oxygen with another element
• Examples of oxides are MgO, ZnO, K₂O, CO₂, SO₂, and H₂O
• Oxides can be categorized based on their acid-base properties

Acid and Basic Oxides

• Acidic and basic oxides exhibit distinct pH values and properties
• Variations in pH arise from their connection to either a metal or a non-metal
• The nature of the element influences the acidity or basicity of the oxide
• Metals form basic oxides while non-metals form acidic oxides:

Acidic Oxides

• Acidic oxides are produced when a non-metal reacts with oxygen.
• They interact with bases to create salts and water.
• When dissolved in water, they generate an acidic solution with a low pH.
• Common examples include CO₂, SO₂, NO₂, and SiO₂.

Basic Oxides

• Basic oxides form when a metal combines with oxygen.
• They react with acids to produce salts and water.
• When dissolved in water, they result in a basic solution with a high pH.
• Common examples include CuO and CaO.

Amphoteric Oxides

Neutral Oxides

• Some oxides, such as N₂O, NO, and CO, do not react with acids or bases and are considered neutral.

Amphoteric Oxides

• Amphoteric oxides can act as both acidic and basic depending on the reactant. They form salt and water in both cases.
• Examples of amphoteric oxides are zinc oxide (ZnO) and aluminum oxide (Al₂O₃).
• The hydroxides of zinc and aluminum also exhibit amphoteric behavior.
• An example is aluminum oxide acting as a base in the reaction: Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O.
• When reacting with hydrochloric acid: Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O
• Example of aluminium oxide behaving as an acid: Al₂O₃ + 2NaOH → 2NaAlO₂ + H₂O
• This dual behavior is not simply explained by proton exchange. The Lewis acid-base theory offers insight into these scenarios.