Oxygen and Its Compounds

Chemistry SS 2 Week 6

Topic: Oxygen and Its Compound

Introduction – Oxygen

Oxygen is a chemical element with symbol O and atomic number 8. It is a member of the chalcogen group on the periodic table and is a highly reactive nonmetal and oxidizing agent that readily forms compounds (notably oxides) with most elements. Oxygen is the third most abundant element in the universe and makes up nearly 21% of the earth’s atmosphere. Oxygen accounts for nearly half of the mass of the earth’s crust, two thirds of the mass of the human body and nine tenths of the mass of water. Oxygen is a colourless, odourless, tasteless gas. It is denser than air and only slightly soluble in water. A poor conductor of heat and electricity, oxygen supports combustion but does not burn. Normal atmospheric oxygen is a diatomic gas (O₂). Ozone is a highly reactive triatomic (O₃) allotrope of oxygen. Oxygen is a highly reactive element and is capable of combining with most other elements. It is required by most living organisms and for most forms of combustion.

Oxygen is extremely active chemically, forming compounds with almost all of the elements except the inert gases. Oxygen unites directly with a number of other elements to form oxides. It is a constituent of many acids and of hydroxides, carbohydrates, proteins, fats and oils, alcohols, cellulose, and numerous other compounds such as the carbonates, chlorates, nitrates and nitrites, phosphates and phosphites, and sulphates and sulphites.

Discovery

Oxygen was first discovered by Scheede in 1772. But the discovery of Oxygen was credited to Priestley in 1774 two after by heating the oxide of mercury. Lavosier, another scientist was the first to describe the major properties of this gas. He called this gas – acid producer because he obtained acids by heating several non-metals in oxygen and dissolving the oxide in water. The credit for discovering oxygen is now shared by three chemists: an Englishman, a Swede, and a Frenchman. Joseph Priestley was the first to publish an account of oxygen, having made it in 1774 by focussing sunlight on to mercuric oxide (HgO), and collecting the gas which came off. He noted that a candle burned more brightly in it and that it made breathing easier. Unknown to Priestly, Carl Wilhelm Scheele had produced oxygen in June 1771. He had written an account of his discovery but it was not published until 1777. Antoine Lavoisier also claimed to have discovered oxygen, and he proposed that the new gas be called oxy-gène, meaning acid-forming, because he thought it was the basis of all acids.

Occurrence

Oxygen occurs in the free state as a gas, to the extent of 21 per cent by volume or 23 per cent by weight in the atmosphere.

Combined Oxygen also occurs

• In water,
• In vegetable and animal tissues,
• In nearly all rocks and
• In many minerals.
• Oxygen occurs to a larger extent in the earth’s crust than any other element.

Laboratory Preparation of Oxygen

Preparation of Oxygen from Potassium Trioxochlorate(V)

The most common laboratory method of preparation of oxygen gas is by heating a mixture of potassium trioxochlorate(V), KClO₃, and manganese(IV) oxide, MnO₂, in a test tube. The potassium trioxochlorate(V) is decomposed in the presence of the heat to form potassium chloride, KCl, and the oxygen gas, O₂.

2KClO₃(s) + heat → 2KCl(s) + 3O₂(g)

The oxygen formed may be collected as,

• Wet Oxygen – The oxygen is collected over water or
• Dry Oxygen – The oxygen oxygen is collected over mercury

Oxygen gas formed may be transferred to the bottom of a water trough and collected at the top of an inverted glass cylinder as wet oxygen.

The oxygen from the heated test tube may also be conveyed via a delivery tube to a U-tube containing anhydrous calcium chloride. The anhydrous calcium chloride is a drying agent that removes moisture from the oxygen gas so that dry oxygen is formed. The dry oxygen is finally delivered via another delivery tube into a glass saucer containing mercury and collected at the top of an inverted test tube on the mercury.

Another laboratory preparation of oxygen is the preparation from hydrogen peroxide in the absence of heat by,

• decomposition in the presence of manganese(IV) oxide, MnO₂, catalyst and
• oxidation in the presence of potassium tetraoxomanganate(VII), KMnO₄.

The decomposition of hydrogen peroxide using manganese dioxide as a catalyst also results in the production of oxygen gas.

Industrial Preparation of Oxygen

The industrial method of making oxygen involves two stages:

• liquefaction of air and
• fractional distillation of liquid air

Liquefaction of air 

Gaseous air is initially passed through caustic soda to remove carbon(IV) oxide and then delivered into a compressor operating at a pressure of 200 atm where it is cooled and allowed to escape rapidly through a minute aperture into a low pressure chamber to form liquid air. The liquid air is then delivered to a fractionating column to undergo fractional distillation.

Fractional distillation of liquid air

The liquid oxygen is distilled to evolve, first, nitrogen at a boiling point of -196°C leaving behind oxygen-rich liquid. The liquid is further heated to -183°C to form gaseous oxygen which is dried, compressed and stored under pressure (at 100 atm) in a steel cylinder.

Physical Properties of Oxygen

Oxygen is

• A colourless gas, without smell or taste,
• it is slightly heavier than air,
• it is sparingly soluble in water,
• it is difficult to liquefy, boiling point -183 ^oC, and the liquid is pale blue in colour and is appreciably magnetic. At still lower temperatures, light-blue solid oxygen is obtained, which has a melting point of -218.4 ^oC.

Chemical Properties of Oxygen

• Reaction of oxygen with air

Oxygen gas does not react with itself or nitrogen under normal conditions. However the effect of ultraviolet light upon oxygen gas is to form the blue gas ozone, O₃, the second allotrope of oxygen. Another way to make ozone is by passing a silent electric discharge through oxygen gas. This can result in a solution containing up to 10% ozone.

• Reaction of oxygen with water

Oxygen gas does not react with water. It does, however, dissolve to the extent of about x g kg^-1 at 20°C (297 K) and 1 atmosphere pressure.

• Reaction of oxygen with the halogens

Irradiation of a low pressure (10-20 mm Hg) mixture of oxygen, O₂, and fluorine, F₂, gases at low temperature (77 – 90 K) affords the gas dioxygen difluoride, O₂F₂.

O₂(g) + F₂(g) → F₂O₂(g)

• Reaction of oxygen with acids

Oxygen gas does not react with most acids under normal conditions.

• Reaction of oxygen with bases

Oxygen gas does not react with most bases under normal conditions.

Test For Oxygen

Oxygen supports combustion so a good method of testing for oxygen is to take a glowing splint and place it in a sample of gas, if it re-ignites the gas is oxygen.

This is a simple but effective test for oxygen. There are sometimes confusions as the splint can create a pop very slightly on re-ignition, which sometimes is mistaken for hydrogen. A hydrogen pop is much more violent, sometimes enough to completely extinguish the splint.

Uses

• Oxygen is essential for life and it takes part in processes of combustion, its biological functions in respiration make it important.
• Oxygen is sparingly soluble in water, but the small quantity of dissolved oxygen in is essential to the life of fish.
• Oxygen gas is used with hydrogen or coal gas in blowpipes and with acetylene in the oxy-acetylene torch for welding and cutting metals.
• Oxygen gas is also used in a number of industrial processes.
• Medicinally, oxygen gas is used in the treatment of pneumonia and gas poisoning, and it is used as an anesthetic when mixed with nitrous oxide, ether vapour, etc.
• Carbon Dioxide is often mixed with the oxygen as this stimulates breathing, and this mixture is also used in cases of poisoning and collapse for restoring respiration.
• Liquid oxygen mixed with powdered charcoal has been used as an explosive.

The main applications of oxygen in order of importance are: 1) melting, refining and manufacture of steel and other metals; 2) manufacture of chemicals by controlled oxidation; 3) rocket propulsion; 4) medical and biological life support; 5) mining, production and manufacture of stone and glass products.
An emergency supply of oxygen automatically becomes available for the passenger in an aircraft when the pressure drop suddently. This oxygen is stored not as an oxygen gas but as the chemical sodium chlorate. Oxygen is used in cellular respiration and released by photosynthesis, which uses the energy of sunlight to produce oxygen from water. It is too chemically reactive to remain a free element in air without being continuously replenished by the photosynthetic action of living organisms. Another form (allotrope) of oxygen, ozone (O
3), strongly absorbs UVB radiation and consequently the high-altitude ozone layer helps protect the biosphere from ultraviolet radiation, but is a pollutant near the surface where it is a by-product of smog. At even higher low earth orbit altitudes, sufficient atomic oxygen is present to cause erosion for spacecraft.

Binary Compound of Oxygen (Oxide)

An oxide is a compound of Oxygen and another element (e.g. through the process of Combustion) Metals form Metal Oxides and Non-Metals form Non-Metal Oxides. Oxides can be classified as either Acidic, Basic, Amphoteric or Neutral

Types of Oxides

Based on their acid-base characteristics oxides are classified as acidic or basic. An oxide that combines with water to give an acid is termed as an acidic oxide. The oxide that gives a base in water is known as a basic oxide.

• Acidic oxides

Acidic oxides are the oxides of non-metals. When combined with water, they produce acids, e.g.,

Examples:
SO₂, SO₃, CO₂, NO₂

Properties:
1. Do not react with acids.
2. React with bases and alkalis to form salt & water.
3. Dissolve in water to form acidic solutions.
4. Usually gases at room temp

• Basic oxides

Basic oxides are the oxides of metals. If soluble in water they react with water to produce hydroxides (alkalies) e.g.,

Examples:
Na₂O, CaO, MgO, FeO, CuO

Properties:
1. Do not react with bases.
2. React with acids to form salt & water.
3. Basic Oxides are usually insoluble in water. Those that dissolve in water forms alkaline solutions

• Amphoteric oxides

Amphoteric oxides are metallic oxides, which show both basic as well as acidic properties. When they react with an acid, they produce salt and water, showing basic properties. While reacting with alkalies they form salt and water showing acidic properties, e.g.,

Examples:
ZnO, Al₂O₃, PbO,

Properties:
1. React with both acids and bases to form salt & water

• Neutral oxides

These are the oxides, which show neither basic nor acidic properties, that is, they do not form salts when reacted with acids or bases, e.g., carbon monoxide (CO); nitrous oxide (N₂O); nitric oxide (NO), etc., are neutral oxides.

Examples:
CO, NO, H₂O

Properties:
1. Neutral pH

Higher Oxides

• Peroxides and dioxides

A peroxide is a metallic oxide which gives hydrogen peroxide by the action of dilute acids. They contain more oxygen than the corresponding basic oxide, e.g., sodium, calcium and barium peroxides.

• Compound oxides

Compound oxides are metallic oxides and they behave as if they are made up of two oxides, lower and higher oxides of the same metal, e.g.,

Red lead: Pb₃O₄ = PbO₂ + 2PbO

Ferro-ferric oxide: Fe₃O₄ = Fe₂O₃ + FeO

On treatment with an acid, compound oxides give a mixture of salts.

 Preparation of Oxides

Oxides can be generated via multiple reactions. Below are a few:

By direct heating of an element with oxygen

Many metals and non-metals burn rapidly when heated in oxygen or air, producing their oxides, e.g.,

By thermal decomposition of certain compounds like hydroxides, carbonates, and nitrates