Oxygen (Descriptive Chemistry)

Dioxygen (O₂) and ozone (O₃)

Decomposition of hydrogen peroxide with manganese dioxide catalyst: 2H₂O₂ → 2H₂O + O₂

Fractional distillation of liquid air.

Respiration + life support systems Welding and cutting metals Rocket fuel oxidizer Steel manufacturing

Acidic oxides Basic oxides Amphoteric oxides Neutral oxides

Non-metal oxides that react with bases to form salt + water. Examples: SO₂, CO₂, NO₂, P₄O₁o

Metal oxides that react with acids to form salt + water. Examples: Na₂O, CaO, MgO

Oxides that react with both acids and bases. Examples: Al₂O₃, ZnO, PbO

Oxides that do not react with either acids or bases. Examples: CO, NO, N₂O

A triatomic allotrope of oxygen (O₃). Pale blue gas with a strong smell

O₂ + UV light → 2O O + O₂ → O₃

Powerful oxidizing agent Toxic in high concentrations Unstable, decomposes to O₂

Water purification Bleaching agent Air sterilization

Absorbs harmful UV radiation, protecting life on Earth

An oxide containing the O₂²⁻ ion (where each oxygen has a −1 charge). There are bonds between atoms of oxygen in between the element and oxygen. (Ba2O2, H2O2)

Typically, by reacting alkali metals (like Na, K) with excess oxygen. Example: 2Na + O₂ → Na₂O₂

Contains the O₂⁻ ion, where oxygen has a −½ oxidation state

Formed when alkali metals react with excess oxygen under high pressure. Example: K + O₂ → KO₂

Oxides where the metal-to-oxygen ratio is unusually high – less oxygen than expected. These are often non-stoichiometric compounds.

C₃O₂ (carbon suboxide), Cu₄O₃ (copper suboxide)

Compounds that contain more than one oxidation state of the same element bonded with oxygen

Fe₃O₄ (contains Fe²⁺ and Fe³⁺), Pb₃O₄

Oxides that contain oxygen in the typical −2 oxidation state. Example: H₂O, CO₂, MgO

Normal oxide: O²⁻ → oxidation state = −2 Peroxide: O₂²⁻ → each O = −1 Superoxide: O₂⁻ → each O = −½ Mixed/Suboxides: weird ratios or mixed oxidation states