18.9 Occurrence, Preparation, and Compounds of Oxygen

Properties and Preparation of Oxygen

Oxygen is a colorless, odorless diatomic gas ($O_2$) that makes up about 21% of Earth's atmosphere by volume. It supports combustion and is required for cellular respiration, making it one of the most important elements for life on Earth. Its slight solubility in water is what allows aquatic organisms to survive by dissolving enough $O_2$ for fish and other creatures to "breathe."

Preparation Methods

There are several common ways to produce oxygen gas in the lab and in industry:

1. Fractional distillation of liquefied air — Air is cooled until it liquefies, then slowly warmed. Because each gas in the mixture has a different boiling point, oxygen (boiling point −183 °C) can be separated from nitrogen (boiling point −196 °C) and other components. This is the main industrial method.

2. Electrolysis of water — An electric current splits water into hydrogen and oxygen gas: $2H_2O \rightarrow 2H_2 + O_2$

3. Decomposition of hydrogen peroxide — $H_2O_2$ breaks down into water and oxygen, often sped up by a catalyst like manganese dioxide ($MnO_2$): $2H_2O_2 \rightarrow 2H_2O + O_2$

4. Heating metal chlorates — Potassium chlorate ($KClO_3$) releases oxygen gas when heated: $2KClO_3 \rightarrow 2KCl + 3O_2$

Common Oxygen Compounds

• Water ($H_2O$) — A polar molecule that is liquid at room temperature. Its polarity makes it an excellent solvent for ionic and many polar compounds, which is why it's often called a "universal solvent."
• Hydrogen peroxide ($H_2O_2$) — A pale blue liquid and strong oxidizing agent. It readily decomposes into water and oxygen gas, which is why bottles of hydrogen peroxide slowly lose potency over time.
• Carbon dioxide ($CO_2$) — A colorless gas produced during combustion and respiration. It acts as a greenhouse gas in the atmosphere and has practical uses in fire extinguishers and carbonated beverages.

Oxygen's solubility in liquids and its reactivity are both influenced by atmospheric pressure and temperature. At higher pressures, more $O_2$ dissolves in water; at higher temperatures, less dissolves.

Metal Oxides, Peroxides, and Hydroxides

Metals react with oxygen to form several categories of compounds, each with distinct structures and uses.

Metal oxides form when metals react directly with oxygen gas: $2M + O_2 \rightarrow 2MO$

Common examples include magnesium oxide ($MgO$), calcium oxide ($CaO$), and iron(III) oxide ($Fe_2O_3$). These compounds are used as refractory materials (they withstand very high temperatures), as catalysts, and as pigments.

Metal peroxides contain the peroxide ion ($O_2^{2-}$), where two oxygen atoms are bonded together. Sodium peroxide ($Na_2O_2$) and barium peroxide ($BaO_2$) are typical examples. They're used as oxygen sources in emergencies (like in submarines) and as bleaching agents.

Metal hydroxides form when metal oxides react with water: $MO + H_2O \rightarrow M(OH)_2$

Sodium hydroxide ($NaOH$), potassium hydroxide ($KOH$), and calcium hydroxide ($Ca(OH)_2$) are all strong bases. They're widely used as pH regulators and in the production of soap and detergents.

Oxyacids vs. Oxyanions of Nonmetals

Oxyacids are acids that contain oxygen, a nonmetal, and hydrogen, with the general formula $H_aX_bO_c$ (where X is a nonmetal). Their acidic behavior comes from the release of $H^+$ ions when they dissolve in water. Key examples:

• Sulfuric acid ($H_2SO_4$)
• Nitric acid ($HNO_3$)
• Phosphoric acid ($H_3PO_4$)

Oxyanions are the negatively charged polyatomic ions that remain after an oxyacid loses its $H^+$ ions (deprotonation). Each oxyacid has a corresponding oxyanion:

• $H_2SO_4 \rightarrow SO_4^{2-}$ (sulfate)
• $HNO_3 \rightarrow NO_3^{-}$ (nitrate)
• $H_3PO_4 \rightarrow PO_4^{3-}$ (phosphate)

Oxyacids react with bases in neutralization reactions to form a salt and water. For example: $H_2SO_4 + 2NaOH \rightarrow Na_2SO_4 + 2H_2O$

Oxyanions can also combine with metal cations to form precipitates (insoluble solids). A classic example is barium sulfate: $Ba^{2+} + SO_4^{2-} \rightarrow BaSO_4$

Oxygen Allotropes and Redox Reactions

Oxygen exists as two allotropes, which are different structural forms of the same element:

• Diatomic oxygen ($O_2$) — The stable, common form we breathe. It's essential for combustion and respiration.
• Ozone ($O_3$) — A triatomic form that is less stable and more reactive than $O_2$. In the upper atmosphere, ozone absorbs harmful UV radiation. At ground level, it's a pollutant and a strong oxidizing agent.

In redox (reduction-oxidation) reactions, electrons transfer between chemical species. Oxygen frequently acts as the oxidizing agent, meaning it gains electrons from other substances. Combustion and rusting are everyday examples of redox reactions where oxygen is the oxidizer.

Cryogenics plays a role in oxygen's industrial use. Oxygen can be liquefied at very low temperatures (−183 °C) and stored in insulated tanks for medical, welding, and aerospace applications. The fractional distillation method described above relies on cryogenic cooling to separate air into its components.