INTRODUCTION TO OXYGEN

Occurrence

Oxygen occurs free in air and makes up about 21% by volume of the atmosphere. It also occurs in combined states as oxides, sulphates, carbonates, nitrates etc, and widely distributed in rocks, clay and mineral compounds. Oxygen is also common in
carbohydrates, fats, proteins and water. The amount of oxygen in the atmosphere can be increased by processes such as photosynthesis and can be reduced by processes like combustion, respiration, rusting, germination etc.

Laboratory preparation of oxygen

Oxygen can be prepared in the laboratory from oxygen rich compounds such as hydrogen peroxide (H2O2), potassium chlorate (KClO3), potassium manganate (VII) (KMnO4), and sodium peroxide (Na2O2). On a large scale oxygen is prepared by electrolysis of sodium hydroxide (NaOH) and fractional distillation of liquid air.

Preparation of oxygen from hydrogen peroxide (H2O2) Apparatus and materials

Flat bottomed flask, dropping funnel, delivery tube, gas jar, beehive (gas jar stand), water trough, water, hydrogen peroxide, manganese (IV) oxide, cork.

Setup of apparatus

Procedure

The gas jar is filled with water and inverted over a gas jar stand.

Manganese (IV) oxide is added in the flask. The apparatus is then arranged as shown above.

Hydrogen peroxide is added to the manganese (IV) oxide through the dropping funnel.

Observation

Effervescence occurs as a colorless gas (oxygen gas) is given off. The colorless gas is collected over water because it is slightly soluble in water or when it is not required dry.

Equation of reaction

Word equation

Symbol Equation.

If the oxygen gas is required dry, it is passed through a wash bottle containing concentrated sulphuric acid as shown below or through a U tube containing calcium chloride and then collected using a syringe.

Purpose of Manganese (IV) oxide

Manganese (IV) oxide acts as a catalyst to speed up the rate of decomposition of hydrogen peroxide. Without it, the decomposition would take place but at a very slow rate. The rate of decomposition can also be increased by exposing the hydrogen peroxide to sunlight.

Using the same setup as above, hydrogen peroxide is added drop by drop from the funnel into potassium permanganate in the presence of dilute sulphuric acid. Oxygen is liberated until when all the potassium permanganate is decomposed and turns colorless.

Equation

A catalyst is a substance that speeds up the rate of a chemical reaction but remains unchanged chemically at the end of the reaction.

N.B. When sodium peroxide is used for the preparation of oxygen, the sodium peroxide is placed in the flat bottomed flask and water in the dropping funnel, then the water is run into the sodium peroxide and effervescence occurs as oxygen is given off. Equation

Chemical test for oxygen

Oxygen relights a glowing splint. When a glowing splint is lowered into a gas jar containing oxygen, the glowing splint relights.

Preparation of oxygen from potassium chlorate

A mixture of potassium chlorate and manganese (IV) oxide is heated in a boiling tube as shown below.

Results

The potassium chlorate decomposes giving off colorless oxygen gas and white residue of potassium chloride.

Equation

Industrial preparation of oxygen

Oxygen is manufactured on a large scale by fractional distillation of liquid air. Air is passed through sodium hydroxide to remove carbon dioxide and through silica gel to dry the air i.e remove water vapor.( carbon dioxide and water vapor are removed because at low temperatures, they solidify and block the pipes).

Equation for removal of carbon dioxide.

The dry air containing nitrogen, oxygen and noble gases is repeatedly compressed at very high pressure (about 200 atmospheres) and cooled until liquid air is obtained. The liquid air appears blue because of the presence of oxygen. Fractional distillation of the liquid air is carried out to obtain oxygen which boils at -183˚C.

N.B. During fractional distillation of liquid air, nitrogen with the lowest boiling (-196˚C) point evaporates first followed by argon (boiing point -186˚C) and then oxygen (boiling point -183˚C)

Liquefaction of air

The dry air containing nitrogen, oxygen and noble gases is compressed at high pressures of up to 200 atmospheres. The compression of air takes place with evolution of heat. As the air comes out of the compressor, it passes through tinny nozzles and a cooling system, in the process the compressed mixture is cooled. The mixture is compressed and cooled repeatedly as the temperatures get lower and lower until when liquefaction of air occurs.

Properties of oxygen

 Physical properties

• It‘s odorless, tasteless and
• It‘s slightly soluble (sparingly soluble) in water. This is why it is collected over water.
• It is approximately the same density as that of
• It is a neutral gas i.e. it has no effect on litmus

Chemical properties

 Oxygen supports burning (combustion). Most metals and non metals burn in oxygen forming basic and acidic oxides respectively.

Reaction of oxygen with metals

 Metals burn in oxygen forming basic oxides which when dissolved in water forms basic or alkaline solution i.e. solutions the turn red litmus paper blue and have effect on blue litmus paper.

Magnesium

 Magnesium burns in air(oxygen) with a brilliant white flame forming white ash (powder) of magnesium oxide.

The magnesium oxide dissolves slightly in water forming an alkaline solution of magnesium hydroxide.

Sodium

 Sodium burns in excess oxygen with a bright yellow flame to form a yellow powder of sodium peroxide.

The sodium peroxide dissolves in water forming an alkaline solution of sodium hydroxide with evolution of a colorless gas that relights a glowing splint (oxygen).

However in limited supply of oxygen, sodium burns with a bright yellow flame to form white solids of sodium oxide.

N.B. Sodium and potassium are kept under oil/paraffin since they are very reactive and react violently with both air and water.

The sodium oxide dissolves in water forming an alkaline solution of sodium hydroxide only.

Calcium

 Calcium burns in air with a bright red flame forming white solids of calcium oxide.

The calcium oxide slightly dissolves in cold water forming an alkaline solution of calcium hydroxide. The calcium hydroxide appears cloudy due to the presence of undissolved calcium oxide.

Potassium

 Potassium burns in oxygen with a lilac flame (purple flame) forming white ash of potassium oxide.

The potassium oxide dissolves in water to form an alkaline solution of potassium hydroxide.

Iron

Oxygen burn is oxygen with a shower of bright sparks leaving behind blue black solids of Tri iron tetra oxide.

The oxide of iron formed is insoluble and therefore has no effect on litmus paper.

Copper

 Copper metal burns with a blue flame turning red hot, on cooling forms black power of copper (II) oxide.

The copper (II) oxide is insoluble in water and therefore has no effect on litmus paper.

Lead

 Lead melts into shinny beads and finally forming yellow powder of lead (II) oxide.

Reaction of non metals with oxygen

 Non-metals burn in air to form acidic oxides (acid andydrides). These oxides dissolve in water to form acidic solutions that turn blue litmus paper red and have no effect on red litmus paper.

An acid anhydride is an oxide of a non metal that dissolves in water to form an acid. Examples include:

Carbon

 Carbon burns in oxygen with an orange flame and bright sparks forming a colorless gas that turns lime water milky (carbondioxide gas).

The carbondioxide gas dissolves in water forming a weakly acidic solution of carbonic acid. This solution turns blue litmus paper pink and not red.

Phosphorus

 Phosphorus burns in air with a bright yellow flame forming white clouds (fumes) of a oxides of phosphorus. The white fumes is a mixture of phosphorus(V)oxide and phosphorus (III)oxide.

Phosphorus(V)oxide and phosphorus (III)oxide dissolve in water forming phosphoric and phosphorus acids respectively.

N.B. Phoshorus is kept in water because if it is in contact with air, it smolders (burns without a flame slowly giving off white smoke)

Sulphur

 Sulphur burns in air with a bright blue flame forming cloudy fumes which have a choking smell. The white fume is a mixture of sulphur dioxide and sulphur trioxide.

The sulphur dioxide and sulphur trioxide dissolve in water forming sulphurous and sulphuric acids respectivey.

Uses of oxygen

• Oxygen is used by living things in the process of Respiration.
• Oxygen is essential for combustion or burning therefore providing heat source for various activities like cooking, burning wastes etc.
• It is used to aid breathing where the natural supply of oxygen is insufficient e.g. in high altitude flying or climbing, and also in hospitals for patients.
• Oxygen when mixed with ethyne produces a very hot flame (oxy acetylene flame) which is used for welding and cutting heavy metals.
• Oxygen is used in the manufacture of steel i.e. in the conversion of pig iron to steel.
• Liquid oxygen is used as fuel in space rockets.
• Liquid oxygen can be used as explosives in mines when mixed with charcoal and petrol.

END OF UNIT