Oxides and Hydrogen Peroxide

Acidic Oxides  

- Oxides of non‑metals.  

- Called acidic oxides because they dissolve in water to form acidic solutions.  

- Also known as acid anhydrides.  

- React with alkalis to form salt + water.  

- Example: SO₂ + H₂O → H₂SO₃  

Peroxides  

- Oxides with a higher proportion of oxygen compared to normal oxides.  

- Example: Calcium peroxide (CaO₂).  

Neutral Oxides  

- Neither acidic nor basic.  

- Show no reaction with litmus.  

- Examples: CO, N₂O.  

Amphoteric Oxides  

- Behave as both acidic and basic oxides.  

- Can react with acids and alkalis to form salt + water.  

- With alkalis, they form complex salts.  

- Examples: Al₂O₃, ZnO.  

Higher Oxides  

- Oxides with a higher proportion of oxygen, similar to peroxides.  

- Unlike peroxides, they do not yield hydrogen peroxide when reacting with acids.  

- Show tendency to release oxygen on heating.  

- Act as oxidizing agents.  

- Examples: PbO₂, MnO₂, Pb₃O₄, Fe₃O₄.  

Hydrogen Peroxide (H₂O₂)  

- Second hydride compound of hydrogen after water.  

- Discovered by Louis Jacques Thénard when dilute H₂SO₄ was reacted with barium peroxide.  

- Exists naturally in trace quantities in air, dew, and snow.  

Preparation of Hydrogen Peroxide  

Laboratory Preparation:  

- Prepared by the action of dilute acid on a metallic peroxide.  

- Commonly uses barium peroxide (BaO₂).  

- Reaction produces barium sulphate (BaSO₄), which is insoluble and can be filtered off.  

Industrial Preparation:  

- Produced by the oxidation of propan‑2‑ol (isopropanol) with oxygen under slight pressure.