`=>` Dioxygen can be obtained in the laboratory by the following ways :

(i) By heating oxygen containing salts such as chlorates, nitrates and permanganates.

`color{red}(2KClO_3 underset(MnO_2) oversettext(Heat)→ 2KCl+3O_2)`

(ii) By the thermal decomposition of the oxides of metals low in the electrochemical series and higher oxides of some metals.

`color{red}(2Ag_2O(s) → 4Ag (s) +O_2 (s) ; \ \ \ \ \ \ 2Pb_3O_4 (s) → 6PbO (s) +O_2 (g))`

`color{red}(2HgO (s) → 2Hg (l) + O_2 (g) ; \ \ \ \ \ \ \ 2PbO_2 (s) → 2PbO (s) + O_2 (s))`

(iii) Hydrogen peroxide is readily decomposed into water and dioxygen by catalysts such as finely divided metals and manganese dioxide.

`color{red}(2H_2O_2 (aq) → 2H_2 O (l) +O_2 (g))`

(iv) On large scale it can be prepared from water or air. Electrolysis of water leads to the release of hydrogen at the cathode and oxygen at the anode.

`=>` Industrially, dioxygen is obtained from air by first removing carbon dioxide and water vapour and then, the remaining gases are liquefied and fractionally distilled to give dinitrogen and dioxygen.

`=>` Dioxygen is a colourless and odourless gas.

`=>` It's solubility in water is to the extent of `3.08 cm^3` in `100 cm^3` water at `293 K` which is just sufficient for the vital support of marine and aquatic life.

`=>` It liquefies at `90 K` and freezes at `55 K`.

`=>` It has three stable isotopes : `color{red}(text()^(16)O, text()^(17)O)` and `color{red}(text()^(18)O)`.

`=>` Molecular oxygen, `color{red}(O_2)` is unique in being paramagnetic inspite of having even number of electrons.

`=>` Dioxygen directly reacts with nearly all metals and non-metals except some metals ( e.g., `Au`, `Pt`) and some noble gases.

`=>` Its combination with other elements is often strongly exothermic which helps in sustaining the reaction. However, to initiate the reaction, some external heating is required as bond dissociation enthalpy of oxgyen-oxygen double bond is high (`493.4 kJ mol^(–1)`).

`=>` Some of the reactions of dioxygen with metals, non-metals and other compounds are given below.

`color{red}(2Ca+O_2 → 2CaO)`

`color{red}(4Al +3O_2 → 2Al_2O_3)`

`color{red}(P_4+5O_2 → P_4O_(10))`

`color{red}(C + O_2 → CO_2)`

`color{red}(2ZnS +3O_2 → 2ZnO+2SO_2)`

`color{red}(CH_4+2O_2 → CO_2+2H_2O)`

`=>` Some compounds are catalytically oxidised. For e.g.,

`color{red}(2SO_2+O_2 overset(V_2O_5)→ 2SO_3)`

`color{red}(4HCl + O_2 overset(CuCl_2)→ 2Cl_2+2H_2O)`

`=>` In addition to its importance in normal respiration and combustion processes, oxygen is used in oxyacetylene welding, in the manufacture of many metals, particularly steel.

`=>` Oxygen cylinders are widely used in hospitals, high altitude flying and in mountaineering.

`=>` The combustion of fuels, e.g., hydrazines in liquid oxygen, provides tremendous thrust in rockets.

When a slow dry stream of oxygen is passed through a silent electrical discharge, conversion of oxygen to ozone (10%) occurs. The product is known as ozonised oxygen.

`color{red}(3O_2 → 2O_3 DeltaH^(⊖) (298K) = +142 kJ mol^(-1))`

● Since the formation of ozone from oxygen is an endothermic process, it is necessary to use a silent electrical discharge in its preparation to prevent its decomposition.

● If concentrations of ozone greater than `10` per cent are required, a battery of ozonisers can be used, and pure ozone (b.p. `385 K`) can be condensed in a vessel surrounded by liquid oxygen.

`=>` Pure ozone is a pale blue gas, dark blue liquid and violet-black solid.

`=>` Ozone has a characteristic smell and in small concentrations it is harmless.

`=>` If the concentration rises above about `100` parts per million, breathing becomes uncomfortable resulting in headache and nausea.

`=>` Ozone is thermodynamically unstable with respect to oxygen since its decomposition into oxygen results in the liberation of heat (`color{red}(ΔH)` is negative) and an increase in entropy (`color{red}(ΔS)` is positive). These two effects reinforce each other, resulting in large negative Gibbs energy change (`color{red}(ΔG)`) for its conversion into oxygen. It is not really surprising, therefore, high concentrations of ozone can be dangerously explosive.

`=>` Due to the ease with which it liberates atoms of nascent oxygen (`color{red}(O_3 → O_2 + O)`), it acts as a powerful oxidising agent. For e.g., it oxidises lead sulphide to lead sulphate and iodide ions to iodine.

`color{red}(PbS (s) +4O_3 (g) →PbSO_4(s) + 4O_2 (g))`

`color{red}(2I^(-)(aq) + H_2O (l) +O_3 (g) → 20 H^(-) (aq) +I_2 (s) +O_2 (g))`

`=>` When ozone reacts with an excess of potassium iodide solution buffered with a borate buffer (`color{red}(pH)` `9.2`), iodine is liberated which can be titrated against a standard solution of sodium thiosulphate. This is a quantitative method for estimating `O_3` gas.

`=>` Experiments have shown that nitrogen oxides (particularly nitric oxide) combine very rapidly with ozone and there is, thus, the possibility that nitrogen oxides emitted from the exhaust systems of supersonic jet aeroplanes might be slowly depleting the concentration of the ozone layer in the upper atmosphere.

`color{red}(NO (g) +O_3 (g) → NO_2 (g) +O_2 (g))`

`=>` Another threat to this ozone layer is probably posed by the use of freons which are used in aerosol sprays and as refrigerants.

`=>` The two oxygen-oxygen bond lengths in the ozone molecule are identical (`128` pm) and the molecule is angular as expected with a bond angle of about `117^o`. It is a resonance hybrid of two main forms.

`=>` It is used as a germicide, disinfectant and for sterilising water.

`=>` It is also used for bleaching oils, ivory, flour, starch, etc.

`=>` It acts as an oxidising agent in the manufacture of potassium permanganate.