Chemistry of Important Industrial Gases, Chemistry tutorial


The study of kinetics theory of gases and the physical studies of solids and liquids have illustrated three states of matter. Solids and liquids are the things which we see around us. They are comprised of molecules which are held altogether by forces that let them take up definite volumes or shapes.

The presence of gases can merely be inferred based on color, odor or by using our sense of feeling. This is due to the gas molecules which encompass much kinetic energy and move around in all directions at great speed; therefore they encompass no definite shape.

The atmosphere around us is filled by air that is comprised of gases, namely nitrogen, oxygen, carbon (IV) oxide and the noble gases. The sun and the stars are supposed to have gases example: ozone and hydrogen. Gases are necessary component of our existence example -oxygen for respiration in animals and carbon (IV) oxide for photosynthesis in plants; and are as well significant industrial chemicals example - Sulphur (IV) oxide for the manufacture of tetraoxosulphate (VI) acid; and carbon (II) oxide and hydrogen as reducing agents and fuels for a few industrial processes.


Hydrogen is the lightest of the elements. Hydrogen comprises of around 1% of the earth crust. Though, it is found free only in extremely small amounts (0.01%) in the atmosphere and in volcanic gases, however large amounts take place in the sun and the stars. Hydrogen is, though, broadly distributed in combination with other elements example: in water, natural gas and petroleum. It is as well a constituent of most of the other organic substances example: Carbohydrates, proteins and fats, which are necessary components of all living things.

Hydrogen gas is one of the most broadly used gases in industry for generating fuels (example: kerosene, petrol and so on), fertilizers, solvents, explosives and plastics. Hydrogen is the unusual element. It consists of a single valence electron such as the Group-I alkali metals, however hydrogen exists, as a diatomic gas having non-metallic properties. On the other, the hydrogen atom is one electron less than having a noble gas configuration just similar to the halogens in Group VII. In the periodic table, hydrogen is generally positioned in Group I for ease, though in reality can be considered to make a family of its own.

Laboratory preparation of hydrogen:

Hydrogen can be formed in the laboratory by any of such three methods, that is:

I) Action of zinc on dilute hydrochloric acid or tetraoxosulphate (VI) acid

Zn + 2 HCl → ZnCl2 + H2

Zn + H2SO4 → ZnSO4 + H2

II) Action of Zinc or Aluminium by Sodium hydroxide

Zn + 2NaOH → NaZnO2 + H2          

2Al + 2NaOH + 2H2O → 2NaAlO2 + 3H2

III) Action of iron on steam:

3Fe + 4H2O ↔ Fe3O4 + 4H2

Industrial Preparation of hydrogen:

Whenever needed on a commercial level, hydrogen can be made by a variety of processes, of which the production from hydrocarbon-based sources has become extremely significant. 

I) From water gas (Bosch Process):

A large quantity of hydrogen is made up from cheap raw materials, namely water and coke (that is, coal). Steam is passed over red-hot coke at around 1,200°C to give a mixture of carbon (II) oxide and hydrogen, termed as water gas. The water-gas is then mixed by more steam and passed over a catalyst, iron (III) oxide or chromium (III) oxide at 450°C. The carbon (II) oxide that is a toxic gas, in the water-gas is transformed to carbon (IV) oxide by further yield of hydrogen.

H2O + C → CO + H2 (Water gas)

CO + H2 + H2O → CO2 + 2H2

The carbon (IV) oxide is eliminated by passing the gaseous products via water and potassium hydroxide. Any unreacted carbon (II) oxide is absorbed in the ammoniacal solution of copper (I) ethanrate.

II) From hydrocarbon sources:

Such routes have slowly replaced: the Bosch process as it utilizes cheaper hydrocarbons example:  methane, propane from natural gas and petroleum. Methane (acquired from the natural gas) is mixed by steam and heated to around 900°C over a nickel catalyst. The mixture of hydrogen and carbon (II) oxide generated is termed as synthesis gas.

CH4 + H2O → CO + 3H2

Methane steam Synthesis gas

The 'synthesis gas' is mixed by more steam and passed over iron (III) oxide as catalyst at 450°C. This second phase is precisely similar as the water gas-steam reaction in the Bosch method.

The second industrial procedure for hydrogen production is to begin with petroleum. The gases obtained from the oil refining method, mostly hydrocarbons, are treated by steam in the presence of the nickel catalyst.

C3H8 + 6H2O → 3CO2 + 10H2

The carbon (IV) oxide is eliminated as in the Bosch method.

III) By electrolysis:

Hydrogen is obtained as the by-product in the electrolysis of brine for the production of sodium hydroxide and chlorine.

Uses of hydrogen:

i) In the manufacture of ammonia in the Haber method.

N2 + 3H2 ↔ 2NH3

The ammonia is transformed to trioxonitrate (V) acid that is then employed in the manufacture of explosives, fertilizers, plastics, drugs and dyes.

ii) Hydrogen is employed to harden (that is, hydrogenation) vegetable oils (example: palm oil, cotton seed oil, corn oil and soyabean oil) to provide solid fats. The solid fats are employed for the soap and candle industries and as the margarine for food.

iii) Catalytic hydrogenation of the coal generates petroleum such as oil mixture termed as synthetic petroleum that can be fractionally distilled to get petrol, lubricants and so on. This petrol is more costly to produce than the ordinary petrol and is utilized only in countries having plenty of coal however no petroleum.

iv) Hydrogen is the component of many gaseous fuels example: water gas and coal gas. Liquid hydrogen is employed as a rocket fuel.

v) Hydrogen discharges a lot of heat whenever it burns therefore it is utilized in oxy-hydrogen flames to generate high temperatures of 4,000 to 5,000°C. This flame is employed in the welding of aluminium or high-chromium or iron alloys.


Oxygen is fundamentally the life-sustaining gas. Oxygen gas comprises, around 21% by volume of the atmosphere and is reasonably soluble in water to support the marine life. This is by far the richest element on earth and around 60% of all atoms in the earth's crust are oxygen atoms.

Oxygen in the atmosphere is being used through the respiration process though the percentage remains fairly constant due to the procedure of photosynthesis, whereby green plants discharge oxygen back to the atmosphere.

The oxygen atom, having an atomic number of 8, consists of 6 electrons in its outer orbital and is in Group VI of the periodic table. Oxygen is an extremely active element and thus doesn't exist in the elemental form however form a variety of ionic (example: MgO) and covalent (example: O2 and H2O) compounds. Oxygen gas supports the combustion of numerous substances. Non-metals such as carbon, sulphur and phosphorus burn in oxygen to make acidic oxides.

C + O2 → CO2

S + O2 → SO2

4P + 5O2 → 2P2O5

Metals such as sodium, calcium and magnesium bums in oxygen to form the basic oxides

4Na + O2 → 2Na2O

2Ca + O2 → 2CaO

Mg + O2 → 2MgO

Laboratory preparation of oxygen:

Oxygen gas is formed in the laboratory via two methods namely:

I) Thermal decomposition of salt rich in oxygen example: KClO3

2KClO3 → 2KCl + 3O2

II) Decomposition of peroxides example: H2O2

2H2O2 → 2H2O + O2

Industrial preparation of oxygen:

The main source of oxygen is air, which is a mixture of gases. The proficient industrial method for obtaining oxygen from air is a two-stage method.

I) Liquefaction of air:

Air, in its natural form, is filtered to take out dust and then passed via sodium hydroxide solution to eliminate carbon (IV) oxide present in air. It is then compressed to a pressure of around 200 atm, cooled and allowed to escape rapidly via a fine jet to an expansion chamber. The rapid expansion further cools the air. The cooled air is utilized to cool the air, coming to the expansion chamber. Whenever the cold, compressed air coming to the expansion chamber is cold adequate, liquefaction takes place and liquid air is obtained.

II) Fractional distillation of liquid air:

The liquid air obtained is then passed to the fractionating column for distillation. Nitrogen that is more volatile than oxygen distills out first, at a temperature of - 196°C and is collected as liquid nitrogen. At a temperature of -183°C, the oxygen distills out and is collected as the liquid oxygen. Liquid oxygen is stored in the strong steel cylinders until whenever it is needed for use.

21_Liquefaction of air.jpg

Fig: Liquefaction of air (simplified flow diagram).

Uses of oxygen:

Naturally, oxygen is employed to sustain life, as oxygen gas for respiration and constituent of water and some biologically significant molecules example: carbohydrates proteins, fats and so on. Moreover, oxygen is helpful in numerous other ways.

i) Oxygen is utilized to help breathing by hospital patients with respiratory problems and those in surroundings having little or no air example: high-altitude, high-mountain climbers, pilots and deep-sea divers.

ii) As oxygen readily supports combustion, it is employed in oxy-hydrogen and oxy-acetylene flames to get high temperatures for welding aims.

iii) Liquid oxygen and fuels are utilized as propellants for the space rockets.

iv) Oxygen is utilized in the steel industry for the elimination of carbon, sulphur and phosphorus impurities from the pig iron.

v) Oxygen is utilized in the manufacture of significant chemicals like tetroxosulphate (VI) acid, trioxonitrate (V) acid and ethanoic acid.


Nitrogen exists mostly in air and is the richest element in air as it comprises 78% of the earth's atmosphere. Similar to oxygen, it is one of the most significant substances on earth, as it is a necessary element of each and every living thing. Particularly, it is a part of the protoplasm, the substance of living cell and of proteins, the building blocks of the living things.

Nitrogen, as the diatomic molecule, is, extremely stable and relatively inert. Its presence in air in this form is extremely significant as it dilutes the oxygen to the point where combustion, respiration and oxidation take place reasonably slow in nature. Or else ignition and fire might take place arbitrarily. The inertness of nitrogen is as well utilized in industry as some drugs and goods are packaged under nitrogen to prevent the oxidation.

Nitrogen gas is utilized as one of the gases to fill electric light-bulbs because it doesn't react with the tungsten filament, even if hot.

In the combined form, nitrogen takes place richly in the earth's crust as trioxonitrate (V) salts and in organic matter like protein, urea and vitamin B compounds. Plants get their proteins via absorbing soluble trioxonitrate (V) salts via their roots. As plants eradicate nitrogen from the soil to make proteins, the soil becomes infertile and the trioxonitrate (V) salts are substituted by using nitrogen fertilizers like ammonium trioxonitrate (V) and diammonium tetraoxosulphate (VI).

Laboratory preparation of nitrogen:

Nitrogen is made in the laboratory by quite a few methods.

I) From air, by eliminating the other constituents of air: Carbon (IV) oxide and oxygen are eliminated by passing air via sodium hydroxide and heated copper turnings correspondingly. The nitrogen obtained by this process includes around 1% by volume of rare gases as impurities and is denser than the pure nitrogen.

II) Decomposition of the ammonium dioxonitrate (III), NH4NO2: Thermal decomposition of this unstable salt solution made via mixing ammonium chloride and sodium dioxonitrate (III) in solution, provides nitrogen.

NH4Cl + NaNO2 → NH4NO2 + NaCl

                                ↓ Heat

                             N2 + 2H2O

III) From ammonia: Nitrogen is released whenever ammonia is oxidized though hot copper (II) oxide

2NH3 + 3CuO → 3Cu + 3H2O + N2

Industrial preparation of nitrogen:

Nitrogen is prepared through the fractional distillation of liquid air as illustrated above. The nitrogen distilled out at - 196°C; though the nitrogen so obtained includes around 1% argon and oxygen.

The oxygen is eliminated via passing the mixture over heated copper. Nitrogen is stored in steel cylinders and sold as liquid nitrogen or as the compressed gas.

Uses of nitrogen:

i) Nitrogen is utilized in the manufacture of ammonia and thence trioxonitrate (V) acid and trioxonitrate (V) fertilizers. Ammonia is prepared industrially through the Haber process - direct combination of hydrogen and nitrogen. Most of the ammonia is employed to make nitrogen fertilizer example: diammonium tetraoxosulphate (VI) and ammonium trioxonitrate (V).

ii) Liquid nitrogen is utilized as a cooling agent.

iii) Due to its inert nature, nitrogen is employed:

  • As a carrier gas in the gas chromatography.
  • In industrial methods comprising easily oxidizable chemicals example: metal transistors.
  • As preservative to prevent the rancidity (that is, due to the oxidation of fats) in the packaged foods.


Chlorine is the most significant member of the halogen family found in the Group 7 of the Periodic Table. Each and every halogen atom consists of seven valence electrons and is all very reactive non-metals. Because of their reactive nature, the halogens are hazardous and so are not found free in nature, however are generally found combined with the metals. Chlorine is a greenish yellow gas. This is generally found in nature in the combined form as chlorides. The richest of these is sodium chloride or common salt, which is mainly found both in the sea and as salt deposits. Sea water having high concentration of sodium chloride is known as brine.

Laboratory preparation of chlorine:

Chlorine is made in the laboratory by the oxidation of concentrated hydrochloric acid by a strong oxidizing agent like manganese (IV) oxide, potassium tetraoxomanganate (VII) or lead (IV) oxide.

4HCl + MnO2 → MnCl2 + 2H2O + Cl2

KMnO4 + 16 HCl → MnCl2 + KCl + 8H2O + 5Cl2

Industrially, chlorine is prepared by the electrolysis of brine or the chlorides of molten sodium, magnesium or calcium. Chlorine is broadly utilized as:

i) A bleaching agent for the linen, cotton and wood-pulp.

ii) A germicide in the sterilization of water for domestic and industrial purpose and in the treatment of the sewage.

iii) In the manufacture of organic solvents (example: trichloromethane, CHCl3; tetrachloromethane, CCl4).

  • Chloroethene (mainly use for making polychloroethene, PVC, plastic).
  • Bleaching agents (example - bleaching powder and sodium oxochlorate (I) utilize in dye works and (laundries);
  • Potassium trioxochlorate (V): Basically use for making matches and fire-works; and sodium trioxochlorate (V); a weed killer.
  • Hydrochloric acid.

Oxides of Sulphur:

There are mainly two oxides of sulphur, sulphur (IV) oxide and sulphur (VI) oxide; both of which are gases.

Sulphur (IV) oxide:

Sulphur (IV) oxide is frequently found in the volcanic gases coal that contains sulphur in the combined form, makes sulphur (IV) oxide whenever burned. The presence of huge amounts of sulphur (IV) oxide in the atmosphere is one of the main causes of acid rain.

Sulphur (IV) oxide is made industrially by burning sulphur or metallic sulphides in air or oxygen.

S + O2 → SO2

FeS2 + 11O2 → 2Fe2O3 + 8SO2

Sulphur (IV) oxide is a strong reducing agent, in the presence of water, due to the formation of the trioxosulphate (IV) ion, SO32- that readily donates electrons to the oxidizing agent example: tetraoxomanganate (VII) ion MnO4-, and heptaoxodichromate (VI) ion; Cr2O72-

SO2 + H2O ↔ H2SO3 ↔ 2H+ + SO32-

It decolorizes the acidified KMnO4 solution by reducing it to the manganese (II) tetraoxosulphate (VI).

2KMnO4 + 5SO2 + 2H2O → K2SO4 + 2MnSO4 + H2SO4

It modifies the color of acidified K2Cr2O7 solution from reducing orange Cr (VI) to green Cr (III).

K2Cr2O7 + 3SO2 + H2SO4 → K2SO4 + Cr2(SO4)3 + H2O

The solution of sulphur (IV) oxide in water bleaches both the natural and artificial dyes; due to its reducing power.

Dye + SO2 + H2O → (Dye-O) + H2SO4

In the presence of a stronger reducing agent example: hydrogen sulphide or carbon; sulphur (IV) act as the oxidizing agent

SO2 + 2H2S → 2H2O + 3S

SO2 + C → CO2 + S

Uses of sulphur (IV) oxide:

i) Its most significant use is in the manufacture of the sulphur (VI) oxide that is the precursor compound for the manufacture of tetraoxosulphate (VI) acid.

ii) This is used as a bleaching agent for wool, silk, sponges and wood pulp throughout paper making.

iii) This is employed as a germicide and a fumigant example: for destroying termites, moulds and bacteria.

iv) This is utilized as preservatives for food particularly dried fruits and grains.

Sulphur (VI) oxide:

Sulphur (VI) oxide is manufactured industrially through direct combination of sulphur (IV) oxide and oxygen. The conditions for good yield of sulphur (VI) oxide are as follows:

- Use of catalyst like vanadium (V) oxide or Platonized asbestos. 

- A temperature of around 400 to 450°C and

- A slight pressure.

2SO2 + O2 + V2O5 (450oC) → 2SO3

This is the part of contact process for the manufacture of the tetraoxosulphate (VI) acid.

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