- 1 Dioxygen
- 2 Preparation of Dioxygen
- 2.1 (1) By the decomposition of oxygen rich compounds
- 2.2 (2) By heating dioxides, peroxides and higher oxides
- 2.3 (3) From Hydrogen peroxide
- 2.4 (4) Laboratory method of preparation of dioxygen
- 2.5 (5) Pure dioxygen from barium hydroxide
- 2.6 (6) Industrial preparation
- 3 Properties of Dioxygen
- 4 Uses of Dioxygen
- 5 Simple Oxides
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 CO2 → (CH2O)x + xO2
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.
2KMnO4– → K2MnO4 + MnO2 + O2
2KClO3 → 2KCl + 3O2
2KNO3 → 2KNO2 + O2
(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)
2Ag2O (s) → 4 Ag (s) + O2 (g)
3MnO2 (s) → Mn3O4(s) + O2 (g)
2Pb3O4 → 6 PbO(s) + O2 (g)
2PbO2 (s) → 6PbO(s) + O2 (g)
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
(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
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).
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.
SO2 + H2O → H2SO3
SO3 (g) + H2O (l) → H2SO4 (aq)
CO2 + H2O → H2CO3
P4O10 + 6 H2O → 4H3PO4
N2O5 (l) + H2O (l) → 2 HNO3 (aq)
Cl2O7 (g) + H2O (l) → 2HClO4(aq)
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.
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)
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.