Oxygen

Oxygen is the most abundant element on Earth. It makes up 21% of the air by volume. It is found both as a free element and in combined forms in nature.

General Properties of Group VI Elements

  1. They are non-metals and are solids at room temperature, except for oxygen, which is a gas.
  2. They don’t react with water, but oxygen and sulfur can react directly with hydrogen to form water and hydrogen sulfide.
  3. They act as electron acceptors and are generally oxidizing in nature.

Electronic Configuration and Bonding in Oxygen

Oxygen has an atomic number of 8, so its electronic configuration is 1s2 2s2 2p4. This means it needs two more electrons to complete its outer shell (octet).

Oxygen has six valence electrons and can achieve a stable configuration by:

  1. Gaining two electrons from metals to form an oxide ion (O2−).
    Example: Ca2+ + O2− → CaO
  2. Sharing electrons with non-metals to form covalent bonds.
  3. Forming covalent bonds with other oxygen atoms.

Laboratory Preparation of Oxygen

  1. By heating potassium trioxochlorate(V) (KClO3) in the presence of manganese(IV) oxide (MnO2) as a catalyst:
    2KClO3(s) → 2KCl(s) + 3O2(g)
  2. By decomposing hydrogen peroxide (H2O2) with MnO2 as a catalyst:
    2H2O2(aq) → 2H2O(l) + O2(g)

Industrial Preparation of Oxygen

  1. By electrolysis of water.
  2. By fractional distillation of liquid air, which involves:
    1. Liquefaction of air: Air is passed through sodium hydroxide (NaOH) to remove carbon dioxide. Then, it is compressed, expanded, and cooled until it turns into liquid air at -200°C.
    2. Fractional distillation: The liquid air is distilled. Nitrogen, which has a lower boiling point (-196°C), evaporates first. Oxygen remains and later boils off at -183°C.

Physical Properties of Oxygen

  1. It is a colorless, odorless, and tasteless diatomic gas (O2).
  2. It is slightly soluble in water.
  3. It is neutral to litmus paper.
  4. It becomes a pale blue liquid at -183°C.
  5. It is denser than air.

Chemical Properties

  1. Reaction with metals: Oxygen reacts with metals to form basic oxides.
    • 2Ca + O2 → 2CaO
    • 4K + O2 → 2K2O

    The oxides of very electropositive metals like K, Na, and Ca dissolve in water to form alkalis:

    • 2K2O + 2H2O → 4KOH
  2. Reaction with Non-Metals: Non-metals burn in oxygen to form acidic oxides (acid anhydrides), which dissolve in water to form acids:
    • S(s) + O2(g) → SO2(g)
    • SO2(g) + H2O(l) → H2SO3(aq)
    • P4(s) + 3O2(g) → P4O6(g)
    • P4O6(g) + 6H2O(l) → 4H3PO3(aq)
  3. Combustion of Hydrocarbons: Most hydrocarbons and compounds containing carbon, hydrogen, and oxygen burn in oxygen to produce carbon dioxide and water:
    • C2H5OH(l) + 3O2(g) → 2CO2(g) + 3H2O(l)

Uses of Oxygen

  1. Used in oxy-ethylene flames for welding and cutting metals.
  2. Essential for respiration in living organisms.
  3. Used in the steel industry to remove impurities such as carbon, sulfur, and phosphorus from pig iron.
  4. Used as a propellant in space rockets along with fuel.
  5. Used in manufacturing acids like tetraoxosulphate(VI), trioxonitrate(V), and ethanoic acid.

Test for Oxygen

To test for oxygen gas:

Oxides

Oxides are binary compounds formed when oxygen combines with another element. They are classified into five main types:

1. Basic Oxides

Formed from metals. They react with acids to form salt and water.

2. Acidic Oxides

Formed from non-metals. They dissolve in water to form acids and react with alkalis to form salts and water.

3. Amphoteric Oxides

Oxides that react with both acids and bases to form salts and water. Examples include ZnO, Al2O3, and PbO.

4. Neutral Oxides

These oxides are neither acidic nor basic and do not change litmus color. Examples: CO, H2O, N2O.

5. Peroxides

These contain a higher proportion of oxygen than normal oxides. Examples: Na2O2, CaO2, BaO2. They react with dilute acids to produce hydrogen peroxide: