Dioxygen

Dioxygen

It occurs in three isotopic forms such as in molecular state, it exists as O₂ and is called dioxygen or simply oxygen. All the dioxygen in the atmosphere is believed to be due to the photosynthesis taking place in green plants in the presence of sunlight.

H₂O + x CO_{2  →}  (CH₂O)_x + xO₂

                 Sunlight
                         Photosynthesis

Preparation of Dioxygen

(1) By the decomposition of oxygen rich compounds

 

Compounds containing large amounts of oxygen such as chlorates, nitrates, permanganates, etc. give dioxygen on strong heating. The thermal decomposition of KClO₃, requires a temperature of 670-720 K.It can be carried out at a lower temperature of about 420K by the use of MnO₂ as a catalyst.
Heat.

2KMnO_{4^{– →}} K₂MnO₄ + MnO₂ + O₂
Pot. permanganate

2KClO_{3 →} 2KCl + 3O₂
Pot. chlorate

2KNO₃ → 2KNO₂ + O₂
Pot. nitrate

 

(2) By heating dioxides, peroxides and higher oxides

 

Dioxygen can be prepared by heating oxides of certain metals low in electrochemical series

and higher oxides of certain metals.

 

2HgO → 2 Hg + O₂ (g)

Mercuric oxide

 

2Ag₂O (s) → 4 Ag (s) + O₂ (g)

Silver oxide

 

3MnO₂ (s) → Mn₃O₄(s) + O₂ (g)

Manganese oxide

 

2Pb₃O₄ → 6 PbO(s) + O₂ (g)

2PbO₂ (s) → 6PbO(s) + O₂ (g)

lead dioxide

 

2BaO₂ → 2BaO + O₂ (g) 

 

(3) From Hydrogen peroxide

 

Hydrogen peroxide is readily decomposed into water and dioxygen by catalysts such as finely divided metals and magnanese dioxide.

 

2H₂O₂ (aq) → 2H₂O(l) + O₂(g)

(4) Laboratory method of preparation of dioxygen

 
(a) Thermal decomposition of potassium chlorate

In the laboratory, dioxygen is prepared by heating a mixture of potassium chlorate (4 parts) and manganese dioxide (1 part) in a hard glass tube to about 420 K. The manganese dioxide acts as a catalyst.

KClO₃ → 2 KCl + 3 O₂ 

The gas is collected by the downward displacement of water.

In the absence of MnO₂ catalyst, the decomposition takes place at 670-720 K . Therefore, MnO₂ acts as a catalyst and also lowers the temperature for the decomposition of KClO₃.

(b) By the action of water on sodium peroxide

Dioxygen can be prepared by the action of water on sodium peroxide.

2Na₂O₂(s) + 2H₂O(l) → 4NaOH(aq) + O₂(g)

Dioxygen can also be prepared by the action of acidified potassium permanganate on sodium peroxide.

2KMnO₄ + 3H₂SO₄ → K₂SO₄ + 2MnSO₄ + 3H₂O + 5(0)

[ Na₂O₂ + H₂SO₄ + O → Na₂SO₄ + H₂O + O₂ ] × 5
2KMnO₄ + 8H₂SO₄ + 5 Na₂O₂ → K₂SO₄ + 5Na₂SO₄ + 2MnSO₄ + 8H₂O + 5O₂ 

 

(5) Pure dioxygen from barium hydroxide

Pure dioxygen can be prepared by the electrolysis of a solution of Ba(OH)₂ using nickel or platinum electrodes.

(6) Industrial preparation

(i) Isolation from air: Oxygen is prepared from air by first removing carbon dioxide and water vapour and then by the fractional distillation of air. During this process, dinitrogen with less boiling point (78 K) distils as vapour while dioxygen with higher boiling point (90 K) remains in the liquid state and can be separated.

 

(ii) From water: Dioxygen can also be obtained by the electrolysis of water containing a small amount of acid or alkali.

 

2H₂O⇔ 2H₂(g)+ O₂(g) 

 

Dioxygen is collected at the anode while dihydrogen is liberated at cathode.

 

 

Properties of Dioxygen

 

Physical properties of Dioxygen

 

(i) Dioxygen is a colourless, tasteless and odourless gas.

 

(ii) It is slightly soluble in water and its solubility is about 3.08 cm³ in 100 cm³ of water at 293 K. This solubility is just sufficient for the vital support of marine and aquatic life.

 

(iii) It can be liquefied to a pale liquid under pressure.

 

(iv) It liquefies at 90 K and freezes at 55 K.

 

(v) Its boiling point is 90.2 K.

 

(vi) It has three stable isotopes ¹⁶O, ¹⁷O, ¹⁸O.

 

(iv) Molecular oxygen is paramagnetic inspite of having even number of electrons.

 

Chemical Properties of Dioxygen

 

The dioxygen is quite stable in nature and its bond dissociation enthalpy is very high. Therefore, it is not very reactive as such

O2 → O + O

Bond dissociation enthalpy = 493.4 kJ mol^-1

 

Dioxygen reacts at higher temperatures. However, once the reaction starts, it proceeds of its own. This is because the chemical reactions of dioxygen are exothermic and the heat produced during the reaction is sufficient to sustain the reactions.

 

(1) Action with litmus

Like dihydrogen, it is also neutral and has no action on blue or red litmus.

 

(2) Supporter of combustion

Dioxygen is a supporter of combustion but itself is not combustible.

 

(3) Reaction with metals

Dioxygen directly reacts with almost all metals (except Au and Pt)

 

(i) Active metals like Na, Ca react at room temperature to form their respective oxides.

4Na + O₂ → 2Na₂O 

2Ca +O₂ → 2CaO

However, sodium also reacts with dioxygen at 575 K to form sodium peroxide.

2Na + O₂ → Na₂O₂

                  Sod. peroxide

 

(ii) Magnesium burns in dioxygen to form magnesium oxide.

2Mg + O₂ → 2MgO

 

(iii) Metals like Fe, Al, react only on heating.

4Al + 3O₂ → 2Al₂O₃

4Fe + 3O₂ → 2Fe₂O₃

 

(iv) Less active metals like gold and platinum (noble metals) do not combine with dioxygen.

 

(4) Action with non-metals

Dioxygen reacts with a number of non-metals form their respective oxides.

2H₂ + O₂ → 2 H₂O (1073, electric discharge)

N₂ + O₂ → 2 NO

S + O₂ → SO₂

2C + O₂(limited) → 2 CO

C + O₂ ( excess) → CO₂

P₄ + 5O₂ → P₄O₁₀

 

(5) Reaction with compounds

Dioxygen is an oxidising agent and it oxidises many compounds under specific conditions. 

 

(i) With hydrogen chloride

Dioxygen combines with vapours of hydrogen chloride in the presence of cupric chloride catalyst at 700K to evolve chlorine.

4HCl + O₂ →  2Cl₂ + 2H₂O (700k , CuCl₂)

 

(ii) With ammonia

Dioxygen oxidises ammonia to nitric oxide in the presence of platinum gauze catalyst at 1073 K.

4NH₃ + 5 O₂ → 4 NO + 6H₂O

 

(iii) With sulphur dioxide

Dioxygen combines with sulphur dioxide at 723 K in the presence of finely divided platinum or vanadium pentoxide (V₂O₅) to form sulphur trioxide.

 

2SO₂ + O₂ → 2 SO₃  (723 k , Pt or V₂O₅)

(iv) With carbon disulphide

Carbon disulphide burns in dioxygen to form carbon dioxide and sulphur dioxide.

 

CS₂ + 3 O₂ → CO₂ + 2SO₂

 

(v) With metal sulphides

Many metal sulphides such as ZnS, HgS, etc. react with dioxygen at high temperatures to form metal oxides and sulphur dioxide.

 

2ZnS + 3O₂ → 2ZnO + 2SO₂

(vi) With hydrocarbons

Saturated as well as unsaturated hydrocarbons burn in excess of air or oxygen to form carbon dioxide and water. These reactions are called combustion reactions and are highly exothermic in nature.

Therefore, the hydrocarbons are used as fuels.

 

CH₄ + 2O₂ → CO₂ + 2H₂O  ΔH =- 890 kJ mol^-1

Methane

 

C₂H₆ + 7/2 O₂ → 2CO₂ + 3H₂O  ΔH =- 1580 kJ mol^-1

 

2C₂H₂ + 5 O₂ →  4CO₂ + 2H₂O  ΔH =- 1304 kJ mol^-1

 

^{Uses of Dioxygen}

1) Dioxygen is used in the oxy-hydrogen or oxy-acetylene torches which are used for welding and cutting of metals.

 

(2) It is used in metallurgical processes to remove the impurities of metals and non-metals by oxidation. 

 

(3) It is also used in making steel and in metal fabrication where it functions as an aid to combustion.

 

(4) It is essential for life support systems in hospitals and in underwater during diving and also by mountaineers and pilots at high altitudes. Oxygen cylinders are widely used in hospitals, high altitude flying and in mountaineering.

 

(5) It is also used for artificial respiration in case of surgery and heart ailments.

 

(6) Liquid oxygen is used as a rocket fuel. For example, hydrazine in liquid oxygen provides the tremendous thrust in rockets.

 

(7) It is used in the manufacture of large number of oxygen containing organic compounds such as phenol, ethylene oxide, sulphur dioxide, sulphuric acid, etc.

 

(8) It is used as an oxidising agent and bleaching agent.

 

(9) A mixture of carbon dust and liquid oxygen is used as explosive for coal mining.

 

(10) Oxygen-18 isotope is used as a tracer in the study of reaction mechanisms.

 

(11 )It is used on a large scale for the production of TiO₂ and synthesis gas (CO + H₂).

 

Simple Oxides

 

The binary compounds of oxygen with other elements are called oxides.

Oxygen combines with metals and non-metals to form their respective binary oxides. Oxides can be simple (e.g: MgO, Al₂O₃ etc.) or mixed (Mn₂O₃, Fe₃O₄,  Pb₃O₄ etc.). 

On the basis of acid-base characteristics, the oxides may be classified into the following four types :

 

(1) Acidic Oxides

The oxides which combine with water to give acids are called acidic oxides. These are generally the oxides of non-metals such as carbon, sulphur, phosphorus, etc. 

For example:

SO₂ + H₂O → H₂SO₃

                     Sulphurous acid

SO₃ (g) + H₂O (l) → H₂SO₄ (aq) 

                                    Sulphuric acid

CO₂ + H₂O → H₂CO₃ 

                         Carbonic acid

P₄O₁₀ + 6 H₂O → 4H₃PO₄

                              Phosphoric acid

N₂O₅ (l) + H₂O (l) → 2 HNO₃ (aq)

                                  Nitric acid

Cl₂O₇ (g) + H₂O (l) → 2HClO₄(aq)

                                                             Perchloric acid

These acidic oxides neutralise hydroxides to form salt and water.

SO₂ + 2NaOH → Na₂SO₃ + H₂O

 

P₄O₁₀(s) + 12NaOH(aq) → 4 Na₃PO₄ (aq) + 6 H₂O(l) 

 

(2) Basic Oxides

The oxides which combine with water to give basic solution are called basic oxides. These are mostly the oxides of metals. 

For example:

Na₂O + H₂O →  2NaOH

MgO + H₂O → Mg(OH)₂

Fe₂O₃(s) + 3H₂O(l) → Fe₂(SO)₄ (aq) + 3H₂O (l)

The basic oxides react with acids to form salt and water.

Na₂O + 2HCl → 2NaCl + H₂O

 

Fe₂O₃(s) + 3H₂SO₄(aq) → 2 Fe(SO₄)₃ (aq) + 3 H₂O (l) 

 

(3) Amphoteric Oxides

 

The oxides which show acidic as well as basic character are called amphoteric oxides. 

These are the oxides formed by elements like aluminium, zinc, tin and lead, etc, which are present on the border line between metals and non-metals.

For example: aluminium oxide (Al2O3) reacts with both acids (hydrochloric acid) and alkalies (sodium hydroxide).

Al₂O₃(s) + 6HCl(aq) + 9H₂O → 2[Al(H₂O)₆]³⁺ (aq) + 6 Cl¯ (aq)

(Basic)

Al₂O₃(s) + 6NaOH(aq) + 3H₂O → 2Na₃[Al(OH)₆] (aq)

 

Some other examples of amphoteric oxides are SiO₂, ZnO etc.

 

(4) Neutral Oxides

 

The oxides which neither react with acids nor with bases are called neutral oxides. These are neutral to litmus solution. For example: N₂O,  CO, NO, etc.