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Oxygen, Its Properties, Preparations, And Uses

 

Oxygen is the most abundant element on earth. It is found almost in everything in nature and also in free state.

Laboratory Preparation of Oxygen

Oxygen can be prepared in a number of ways in the laboratory. These include:

(a). Heating a mixture of potassium trioxochlorate(V), KClO3 and maganese(IV) oxide, MnO2.

The MnO2 acts as a catalyst. The reaction is actually the decomposition of KClO3.

2KClO3(s) → 2KCl(s) + 3O2(g)

(b). The decomposition of hydrogen peroxide, H2O2 using manganese(IV) oxide - this does not require heating.

Hydrogen peroxide is added drop wisely unto MnO2.

2H2O2(aq) → 2H2O(l) + O2(g)

Or by the drop wise addition of hydrogen peroxide on acidified KMnO4

5H2O2(aq) + 2KMnO4(aq) + 3H2SO4(aq)  → K2SO4(aq) + 2MnSO4(aq) + 8H2O(l) + 5O2(g)

(c). The reaction between water and sodium peroxide.

Hydrogen peroxide is formed, and it immediately decomposes by the catalytic effect of the OH- ions in solution.

Na2O2(s) + 2H2O → H2O2 + 2Na+ + 2OH-

Then, 2H2O2 → 2H2O + O2

(d). Application of heat on trioxonitrate(V) salts of metals.

Trioxonitrate(V) salts of metals, e.g, sodium trioxonitrate(V), NaNO3 give off a part of their oxygen upon being heated. NaNO3 loses one third of its oxygen.

2NaNO3(s) → 2NaNO2(s) + O2(g)

(e). Application of heat on certain oxides of the least active metals, e.g., mercury oxide, HgO and silver oxide, Ag2O. These oxides undergo complete dissociation when heated.

2HgO(s) → 2Hg(s) + O2(g)

2Ag2O(s) → 4Ag(s) + O2(g)

(f). Application of heat on oxides of certain metals with more than one oxidation state, e.g., lead(IV) oxide, PbO2 and mangan- ese(IV) oxide, MnO2.

These oxides give off only a part of their oxygen when they are heated. Such reactions usually require very high temperatures.

2PbO2(s) → 2PbO(s) + O2(g)

3MnO2(s) → Mn3O4(s) + O2(g)

(g). The electrolysis of water.

Oxygen is produced at the anode - see the electrolysis of water for details.

Test for Oxygen

Oxygen can be distinguished from all other gases except dinitrogen oxide, N2O by its rekindling of a glowing splint of wood.

It is however distinguished from N2O by the following observations:

(1). Oxygen does not have smell, while N2O has a sweet, sickly smell.

(2). Oxygen produces brown fumes of nitrogendioxide, NO2 with nitrogen monoxide,

2NO(g) + O2(g) → 2NO2(g) (brown fumes),

while N2O does not.

Properties of Oxygen

Physical Properties:

(1). It is colourless, has no odour and is neutral.
(2). It is slightly soluble in water.
(3). It has almost the same density as air.
(4). It freezes at 54K.
(5). It has a boiling point of 90K at 1 atmosphere pressure (i.e. 760 mm Hg).
(6). It is very active and reacts with many metals and non-metals to form basic and acidic oxides respectively.

Basic oxide - examples include MgO, Na2O and CaO;
Acidic oxide - examples are CO2, SO2 and P4O10 (acidic oxides of non-metal are also called acid anhydrides).

Summary of Reactivity of Oxygen with Metals

K, Na, Ca, Mg, Al, Zn, Fe, Pb, Cu - - show decreasing readiness to form oxides when heated in air, with Cu the least reactive with oxygen.

Hg, Ag, Au - - these metals show the least readiness to form oxides. Their oxides are easily decomposed to the
metal and oxygen.

K, Na, Ca, Mg, Al, Zn - - the oxides of these metals are not reduced to the metals by heating in a stream of hydrogen,
carbon or carbon(II) oxide.

Fe, Pb, Cu - - the oxides of these metals are reduced to the metals by heating in a stream of hydrogen, carbon or carbon(II) oxide.

Note: the more readily a metal combines with oxygen to form an oxide, the less readily it will be reduced to the metal by either heating in a stream of hydrogen or CO.

Therefore, the oxides of K to Zn are not reduced, while those of Fe and below are reduced.

Chemical Properties:

(1). Reaction with compounds - most hydrocarbons and compounds of carbon, hydrogen and oxygen burn in oxygen to form carbon (IV) oxide and water.

CH4(g) + 2O2(g) → CO2(g) + 2H2O(g)

C2H5OH(l) + 3O2(g) → 2CO2(g) + 3H2O(g)

4NH3(g) + 5O2(g) → 4NO(g) + 6H2O(l)

Uses of Oxygen

(1). Used as an aid for breathing where problem of breathing arises. Example in high altitude flying or climbing, and when a patient is under anesthetics.

(2). In the oxyacetylene (i.e. oxygen-ethyne) flame - used in welding and cutting steel plate - due to the very high temperature of the flame (about 2200oC).

(3). Used in the L-D process for making steel.

Commercial Production of Oxygen from Liquid Air

The process involves air firstly liquefied by compression, cooling, expansion and successive cooling. Then by fractional distillation of the liquid air - oxygen is separated.

Liquid air contains mainly oxygen and nitrogen. Nitrogen evolves first at 77 K while oxygen evolves later at 90 K, 760 mm Hg.

 

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