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Chemical Preparation of Oxygen






The majority of the methods for obtaining oxygen fall into this category, and may be classified according to whether the parent substance is a normal oxide, a higher oxide, or a more complex oxygenated compound. Several of the processes can be extended to serve as methods for the extraction of atmospheric oxygen.


Preparation of Oxygen from Normal Oxides

Water. - Perhaps the most important process by which water can be made to yield its oxygen in a free state consists in electrolysis in the presence of an alkaline substance such as potassium carbonate or an alkali hydroxide. The oxygen obtained in this manner, if due precautions are observed, presents a high degree of purity, and is hence particularly suitable for metal cutting and welding. Care must be taken to avoid contamination with hydrogen during the process, owing to the dangerously explosive nature of the mixture.

For laboratory purposes a glass apparatus after the principle of a Kipp may be conveniently used. The electrolyte consists of a 30 per cent, solution of sodium hydroxide, whilst sheet-nickel plates serve as electrodes. The inner electrode functions as anode and the supply of oxygen regulates itself automatically, the liquid in the inner space surrounding the anode being gradually expelled, as in Kipp's apparatus, as the pressure of the gas above increases, until the anode is left high and dry, when, of course, electrolysis ceases. Very pure oxygen may be obtained by the electrolysis of barium hydroxide solution. To a certain extent the electrolysis of water is used for the commercial preparation of oxygen. To this end, containing vessels of iron are used, the electrodes consisting of this metal or of nickel. The electrolyte consists of 15 per cent, caustic soda solution, and the liberated hydrogen and oxygen are collected in separate dome-shaped vessels under a pressure of some 60 mm. of water. A higher pressure cannot safely be employed owing to the danger of mixing. By means of a metallic partition between the electrodes, this danger is still further minimised. The containing vessels are packed in wooden boxes with sand, whereby the heat of the reaction is conserved, the temperature rising to about 70° C. Each vessel yields 110 litres of oxygen per hour of purity 97 per cent. Dilute solutions of acids, particularly sulphuric acid, may be employed instead of alkalies, but the latter are preferable.

Electrolytic oxygen may contain as much as 4 per cent, of hydrogen. This may be removed as water by passage over platinised asbestos.

Attempts have been made to cheapen the process by producing electrolytic oxygen without the simultaneous liberation of hydrogen by the adoption of depolarising electrolytes, or cathodes; also of cathodes which absorb hydrogen and may subsequently be employed as elements in gas cells.

At high temperatures steam dissociates into hydrogen and oxygen, and these gases admit of isolation by taking advantage of the greater velocity of diffusion of the hydrogen. This by no means constitutes a convenient method of preparing oxygen, but the process may be facilitated by the introduction of some substance capable of uniting with the hydrogen. Thus steam is readily decomposed by chlorine when the two are passed through a red-hot porcelain tube. The reaction is accelerated by the presence of fragments of porcelain in the tube to increase the heating surface.

2H2O + 2Cl2 = 4HCl + O2.

The hydrochloric acid may be absorbed by passage of the resulting gases through water or caustic soda solution.

Silver oxide, Ag2O, is readily decomposed by heat, evolving oxygen, and the characteristic change in colour from brown to silver-white renders the reaction particularly suitable for lecture demonstration. The equilibrium pressures of oxygen have been measured up to 800° C., and are found to conform with the following law:

log p = 6.2853 - 2859/T

where p is the pressure in atmospheres, and T the absolute temperature.

The decomposition of mercuric oxide by heat has already been mentioned as the method by which Priestley was led to the discovery of oxygen. The oxide, which is yellow or brick red in colour, first becomes black - a reversible, physical effect. Oxygen is then evolved and a sublimate of mercury collects on the cooler parts of the containing vessel. The reaction is reversible, thus:

2HgO ⇔ 2Hg + O2.

In the following table are given the dissociation pressures of mercuric oxide between 360° and 480° C.

Temperature, ° C.Pressure, mm.Temperature, ° C.Pressure, mm.
36090440642
3801414601017
4002314801581
420387. . .. . .


The rate of decomposition is accelerated by suitable catalysts such as finely-divided platinum, ferric oxide, and manganese dioxide. Aluminium and stannic oxides are apparently without effect.

The normal oxides of several other metals behave in an analogous manner to mercuric oxide. Thus palladous oxide, PdO, decomposes when heated, yielding metallic palladium and oxygen, the reaction being reversible:

2PdO ⇔ 2Pd + O2.

At 877° C. the dissociation pressure of the oxide is 760 mm.

In the case of auric oxide, Au2O3, the reaction is not reversible. When heated at 150° to 165° C., oxygen is evolved and aurous oxide, AuO, remains. At 250° C. this latter oxide is completely converted into metallic gold. Similarly, platinum dioxide, PtO2, upon ignition evolves oxygen, a residue of metallic platinum being obtained together with a solid solution of either the monoxide or sesqui-oxide in the dioxide.

When chlorine is passed over zinc oxide at a high temperature oxygen is evolved, zinc chloride remaining behind.

2ZnO + 2Cl2 = 2ZnCl2 + O2.

Oxides of the alkaline earth metals, namely CaO, SrO, and BaO, may be similarlv employed, as also litharge, PbO, and cadmium oxide, CdO.

Preparation of Oxygen from Higher Oxides

In addition to their normal ones, many metals yield oxides in which the percentages of oxygen are. greater than correspond to the valencies of the metals as manifested in their more common salts. Such compounds are conveniently termed higher oxides, and may usually be made to part with their excess of oxygen either by heating alone or with sulphuric acid.

Manganese dioxide, MnO2, when heated to moderate redness, evolves oxygen and leaves a residue of the sesqui-oxide Mn2O3.

4MnO2 = 2Mn2O3 + O2.

The reaction begins at 530° C. in air, and if the temperature is raised to 940° C. the sesqui-oxide in turn decomposes, yielding a further supply of oxygen and a residue of trimanganic tetroxide, Mn3O4.

6Mn2O3 = 4Mn3O4 + O2.

The foregoing reactions at one time offered one of the cheapest methods of preparing oxygen for commercial purposes. The source of the dioxide was the mineral pyrolusite, but the high temperature required to extract the oxygen led to the superseding of this method by other more convenient processes.

When heated with concentrated sulphuric acid, manganese dioxide evolves oxygen, leaving a residue of manganese sulphate. The reaction takes place in two stages, namely:

Lead dioxide, PbO2, when heated above 310° C., decomposes, yielding oxygen and lead monoxide.

2PbO2 = 2PbO + O2.

Similarly red lead, Pb3O4, when strongly heated evolves oxygen, a residue of lead monoxide remaining. This reaction is reversible.

2Pb3O4 ⇔ 6PbO + O2.

At 530° C. the red lead may be completely converted into monoxide in a vacuum, but in the presence of air a higher temperature is essential, as is evident from the following data.

Temperature ° C445500555636
Dissociation pressure of Pb3O4 in mm.560183763


Alkali peroxides are rapidly decomposed by water, oxygen being evolved. In the case of sodium peroxide the reaction proceeds according to the following equation:

2Na2O2 + 2H2O = 4NaOH + O2.

The reaction is conveniently carried out in a flask fitted with a drop funnel through which the water is slowly admitted. Some hydrogen peroxide is simultaneously produced. The evolution of oxygen is facilitated by the addition of a catalyst, such as a salt of nickel, cobalt, or copper. When pressed into small blocks or cubes, the mixture of sodium peroxide and catalyst may be placed in a Kipp or other gas- generating apparatus based on a similar principle, and a steady supply of oxygen obtained. The commercial commodity known as " oxylithe " has the following composition:

Sodium peroxide98.32 per cent
Oxide of iron1.00
Copper sulphate0.68 per cent


and is very suitable for this type of reaction.

The preparation of small quantities of oxygen for laboratory purposes may be conveniently effected by gently warming a mixture of fused sodium peroxide with some salt containing water of crystallisation. For this purpose crystals of sodium carbonate or sulphate are very suitable. The oxygen is evolved in a steady stream which is readily kept under control.

By the action of acids upon alkali or alkaline earth peroxides, hydrogen peroxide is liberated, which immediately undergoes partial or complete decomposition according to circumstances. Thus oxygen is readily obtained by the employment in a Kipp of lumps of the mixture obtained by adding 100 parts of sodium peroxide and 25 parts of magnesium oxide to 100 parts of molten potassium nitrate. The liquid reagent consists of dilute hydrochloric acid. The magnesia does not serve as a catalyst; on the contrary, it is added as an inert diluent to moderate the violence of the reaction.

Hydrogen peroxide readily yields up its oxygen either under the influence of heat or of a catalyst. As examples of the last named, colloidal solutions of the platinum metals may be mentioned.

In neutral solution hydrogen peroxide is decomposed catalytically by lead dioxide, but in acid solution the action is different and quantitative. Thus, in the presence of nitric acid,

PbO2 + H2O2 + 2HNO3 = Pb(NO3)2 + 2H2O + O2.

Manganese dioxide behaves similarly in acid solution, and, if charged in lump form into a Kipp and subjected in the usual manner to the action of commercial hydrogen peroxide acidified with sulphuric acid, a steady stream of oxygen is obtained.

Instead of using free hydrogen peroxide, barium peroxide may be used, lumps of a mixture of barium peroxide, gypsum, and manganese dioxide being introduced into the Kipp, the liquid reagent consisting of hydrochloric acid.

Hydrogen peroxide reacts with potassium permanganate in acid solution evolving oxygen. In the presence of dilute sulphuric acid the reaction proceeds along the lines indicated by the equation

2KMnO4 + 5H2O2 + 3H2SO4 = K2SO4 + 2MnSO4 + 8H2O + 5O2.

For laboratory purposes a steady evolution of oxygen may be obtained by allowing a solution of 25 grams of potassium permanganate in 500 c.c. of water and 50 c.c. of concentrated sulphuric acid to flow from a dropping funnel into a litre flask containing 500 c.c. of hydrogen peroxide solution (10 vol.). No heat is required.

Hydrogen peroxide reacts in an analogous manner with potassium bichromate, evolving oxygen. A convenient way of preparing the gas in small quantities consists in adding 150 grams of concentrated sulphuric acid to hydrogen peroxide solution (10 vol.) and allowing the mixture to come into contact with crystals of potassium bichromate in a Kipp's apparatus. The crystals should be large and the process carried out with care in the cold, as otherwise the reaction is liable to be very violent. In order to prevent small pieces of the bichromate from falling into the lower chamber of the Kipp, a layer of small pieces of pumice may be introduced into the middle chamber prior to the admission of the salt. The reaction proceeds according to the equation

K2Cr2O7 + 4H2SO4 + 3H2O2 = K2SO4 + Cr2(SO4)3 + 7H2O + 3O2.

In alkaline solution potassium ferricyanide and hydrogen peroxide also yield a steady stream of oxygen which can be immediately checked by the addition of an acid.

The reactions involved appear to be represented by the following equations:

2K3Fe(CN)6 + 2KHO = 2K4Fe(CN)6 + H2O + O (nascent)
O + H2O2 = H2O + O2.

Experiment shows that it is the amount of alkali present that controls the reaction - an observation in harmony with the above equation.

With bleaching powder, hydrogen peroxide in acidified solution readily yields oxygen gas,

Ca(OCl)Cl + H2O2 = CaCl2 + H2O + O2.

Barium peroxide is readily decomposed by heat, the reaction being reversible:

2BaO2 ⇔ 2BaO + O2.

By continuous removal of the oxygen, therefore, decomposition will continue at constant temperature until the whole of the solid phase has been converted into the monoxide.

If, on the other hand, the pressure of oxygen in contact with the solid phase is increased beyond the dissociation pressure at a given temperature, the barium peroxide is regenerated, the foregoing reaction now proceeding from right to left.

In the following table are given the dissociation pressures of barium peroxide in contact with moisture at various temperatures ranging from 618° to 868° C.

Temperature ° C618655697737794835853868
Oxygen pressure (cm. mercury)11.326.865.41413787189371166
Water-vapour (cm. mercury)7.313.726.34798159195231


Boussingault attempted to use the foregoing reactions for the preparation of oxygen on a commercial scale, but found that after several reheatings the barium oxide lost its power of absorbing oxygen. This difficulty was eventually overcome by the brothers Brin, who formed a company for the preparation of oxygen for industrial purposes.

The barium oxide was obtained in a hard and porous condition by ignition of the nitrate. Pieces about the size of a walnut were heated to 600° C. in vertical steel retorts into which air, purified from carbon dioxide and from most of its moisture, was conducted under a pressure of about 10 lb. per sq. inch. After seven minutes the pressure was reduced to 4 inches (10 cm.) of mercury, the temperature remaining constant, whereon the absorbed oxygen was evolved. This process was repeated four times per hour, and a gas of 95 per cent, purity obtained.

As late as 1907 three works were producing 30,000 cubic feet of oxygen per day by this process, which is now, however, obsolete in Great Britain, having been superseded by the liquid-air process, which not only yields a cheaper but a purer gas, namely, 97 per cent, oxygen.

In 1913 a process was patented by which the oxygen of the air could be obtained by alternate oxidation and reduction of oxides of nitrogen. Vapour of nitric acid is passed over heated sulphuric acid whereby oxygen is liberated and water and nitrosulphuric acid are produced. The latter is treated with water yielding sulphuric acid and a mixture of nitric oxide and nitrogen peroxide, the last named being reconverted into nitric acid by solution in water in presence of air. The reactions may be represented as follows:
  1. 4HNO3 - 4NO2 + 2H2O + O2.
  2. 4H2SO4 + 4NO2 = 4H(NO)SO4 + 2H2O + O2.
  3. H(NO)SO4 + H2O = 2H2SO4 + N2O3.
  4. N2O3 + O2 + H2O = 2HNO3.

Preparation of Oxygen from more Complex Compounds

Many oxygenated salts and other compounds yield oxygen when subjected to the influence of heat, either alone or in contact with other substances. They may even yield oxygen at ordinary temperatures in contact with suitable catalysers.

When metallic chlorates are gently heated, oxygen is evolved, a chloride being generally left behind. The salt which, for various reasons, has been studied most carefully in this connection is potassium chlorate. The decomposition of this salt is liable to be explosive if the heating is carried out suddenly. This is readily demonstrated by allowing very small drops of molten chlorate on the end of a glass rod to fall on to the bottom of a test-tube heated to redness. Sharp detonations result. When heated to 357° C. this salt undergoes no perceptible decomposition, but the powder cakes together and when examined under the microscope shows signs of incipient fusion. The salt becomes liquid at a slightly higher temperature, and at 370° to 380° C. there is a rapid evolution of oxygen. Several reactions now begin to take place:

(1) The formation of perchlorate. This is a case of autoxidation, one molecule of chlorate oxidising three other molecules of chlorate to perchlorate and being itself reduced to chloride. Thus:

KClO3 + 3KClO3 = KCl + 3KClO4.

This reaction is exothermic, evolving 61,300 calories.

The velocity of formation of potassium perchlorate has been measured at 395° C. and the reaction shown to be tetramolecular and to proceed in accordance with the above equation.

(2) In addition to the foregoing reaction, potassium chlorate undergoes decomposition into oxygen and potassium chloride. This is a monomolecular reaction and proceeds according to the equation

2KClO3 = 2KCl + 3O2.

(3) If the temperature is raised sufficiently the potassium perchlorate formed in the first reaction begins to decompose, in the main according to the equation

KClO4 = KCl + 2O2.

This reaction is inappreciable even at 411° C., but readily proceeds to completion at 445° C. A small quantity of potassium chlorate is simultaneously regenerated.

A trace of chlorine is usually found in the oxygen obtained by heating potassium chlorate in glass apparatus, a larger amount being obtained with Jena glass than with either soda or Bohemian combustion glass. When the chlorate is decomposed in platinum vessels, however, chlorine is either not evolved, or only in infinitesimal quantities whether at atmospheric or under reduced pressure.

Sodium chlorate decomposes in a precisely similar manner to the potassium salt.

The foregoing method of preparing oxygen possesses two disadvantages. Not only is the evolution of the gas inclined to be violent and difficult to control, but the temperature at which the reaction takes place is too high to be satisfactorily carried out in a glass vessel. These difficulties are overcome by mixing the chlorate with manganese dioxide prior to heating, a procedure first described by Doebereiner in 1832. Under these conditions the evolution of oxygen is steady and commences at about 240° C. instead of 370° C

It is important to remember that commercial manganese dioxide is occasionally contaminated with carbonaceous material such as coal dust. Such a mixture is very liable to explode when heated with potassium chlorate owing to the rapid combustion of the carbon in the oxygen. Manganese dioxide should, therefore, always be tested beforehand and rejected for the preparation of oxygen if it is found to contain any carbonaceous matter.

The manner in which the manganese dioxide assists the decomposition of the chlorate has been made the subject of considerable controversy. The oxide may be used over and over again without any measurable diminution of its activity. It has been suggested from time to time that its action is purely mechanical analogous to that of sand, etc., in promoting the boiling of water. The analogy, however, is misleading, for reduction of pressure does not materially facilitate the evolution of oxygen from potassium chlorate, although it greatly reduces the boiling- point of water. Again, were the action purely mechanical, all other finely-divided substances, irrespective of their chemical composition, might be expected to act in a similar manner just as they are known to do in the case of boiling water. This, however, is not the case, for although oxides of iron, cobalt, nickel, cerium, and copper facilitate the reaction, the oxides of zinc, magnesium, etc., appear incapable of doing so.

The most probable explanation is that alternately higher and lower oxides of manganese are formed - the higher oxide by the oxidising action of the heated chlorate, and the lower oxide by the decomposition of the higher, either alone or in contact with a further supply of chlorate. Mention has already been made of the fact that, when potassium chlorate is heated alone, some perchlorate is formed through self-oxidation simultaneously with the evolution of oxygen. This reaction does not occur in the presence of manganese dioxide, since this oxide effects the decomposition of the chlorate into chloride and oxygen at a temperature considerably below that at which autoxidation of the chlorate proceeds at an appreciable rate.

That several other minor or side reactions should take place, in addition to the main cycle indicated above, is only to be anticipated. Thus the fact that the oxygen invariably contains traces of chlorine suggested that a peculiar form of ozone was produced rather than chlorine; but this is negatived by the results of M'Leod. Small quantities of potassium permanganate are also undoubtedly formed in the solid mass, for when potassium chlorate is fused with a very small quantity of manganese dioxide a pink colour is observable on cooling. When this pink mass is fused over a flame, oxygen is evolved, but the colour persists until nearly all the chlorate is decomposed; it then becomes greenish and ultimately brownish. If the dioxide is present in considerable quantity any pink colour is masked by the blackness of the mixture.

M'Leod explains these changes as follows:
  1. The dioxide acts on the chlorate forming permanganate, chlorine, and oxygen.
    2KClO3 + 2MnO2 = 2KMnO4 + Cl2 + O2.
  2. The permanganate then undergoes decomposition by the heat yielding potassium manganate, manganese dioxide, and oxygen.
    2KMnO4 = K2MnO4 + MnO2 + O2.
    The reaction begins at about 200° C., and is rapid at 260° C.
  3. The manganate is decomposed by chlorine yielding potassium chloride, manganese dioxide, and oxygen.
    K2MnO4 + Cl2 = 2KCl + MnO2 + O2.
Oxygen may also be obtained by heating the chlorates of other metals, notably barium, calcium, strontium, lead, and silver, or by decomposition of metallic bromates and iodates.

Aqueous solutions of alkali hypochlorites readily yield oxygen at the boiling-point under the influence of catalysers. This is easily accomplished by passing a current of chlorine through a concentrated solution of caustic soda at the boiling-point, to which a small quantity of a cobalt salt has been added. The cycle of reactions involving the liberation of oxygen may be represented as follows:

NaOH + Cl2 = NaClO + NaCl + H2O
NaClO + CoO = NaCl + CoO2
2CoO2 = CoO + O2.

An aqueous solution of bleaching powder, to which a trace of a cobalt salt has been added to serve as catalyst, readily evolves oxygen when warmed to about 80° C. The procedure may be varied by using a thin cream of bleaching powder in water and warming this on a water- bath to 70° or 80° C. in the presence of a small quantity of a cobalt salt. The mixture froths excessively, but this tendency may be overcome by addition of a little paraffin oil.

The mechanism of the process consists in the immediate conversion of the cobalt salt into an oxide which undergoes alternate reduction and oxidation. What the composition of the higher oxide may be is uncertain; probably it is either the sesqui-oxide, Co2O3, or the dioxide, CoO2. Assuming it to be the latter, the reactions taking place may be represented as follows:

2Ca(OCl)Cl + 2CoO = 2CaCl2 + 2CoO2
2CoO2 = 2CoO + O2.

The velocity of reaction indicates it to be monomolecular. Salts of nickel, copper, or iron may be used instead of those of cobalt, but are less active. The theory that the catalyst effects the decomposition by its own alternate oxidation and reduction is supported by the result of passing chlorine into a 50 per cent, solution of sodium hydroxide containing dissolved copper hydroxide; the blue solution at first deposits a yellow copper peroxide, which rapidly decomposes, evolving oxygen and regenerating the original solution. The effect of adding two catalysts to bleaching powder is remarkable. If the bleaching powder is made into a cream with water, oxygen may be liberated at the ordinary temperature by addition of a ferrous or manganous salt and in the presence of a copper or nickel compound. The best result is obtained with a mixture of ferrous and copper sulphates.

Practically the same reaction takes place when a stream of chlorine gas is passed through boiling milk of lime containing a trace of cobalt oxide as catalyst. The oxygen is steadily evolved.

2Ca(OH)2 + 2Cl2 = 2CaCl2 + 2H2O + O2.

Oxygen is readily evolved at the ordinary temperature on adding water to a mixture of bleaching powder and an alkali or alkaline earth peroxide in the presence of a catalyst such as ferrous or copper sulphate. If the solid mixture is pressed into small lumps or cubes, it may be used in a Kipp's apparatus and thus afford a convenient method of preparing the gas for lecture or laboratory purposes.

Ca(OCl)Cl + Na2O2 + H2O = Ca(OH)2 + 2NaCl + O2.

Concentrated sulphuric acid, when strongly heated, decomposes into water and a mixture of sulphur dioxide and oxygen.

2H2SO4 = 2H2O + 2SO2 + O2.

To this end the acid is allowed to drop on to a red-hot surface and the resultant gases treated with suitable absorbents to remove the sulphur dioxide and steam.

Concentrated nitric acid readily decomposes, when heated, into water, nitrogen dioxide, and oxygen. The two former are readily converted again into nitric acid by the action of the atmospheric air.

Alkali nitrates, when heated above their melting-points, yield the corresponding nitrite and oxygen; but the gas is contaminated with nitrogen resulting from partial decomposition of the nitrite. In the case of potassium nitrate the reaction may be represented by the equation:

2KNO3 = 2KNO2 + O2.

Priestley had noticed as early as 1772 that, when a lighted candle is lowered into the gas obtained by heating potassium nitrate, the flame "increased," indicating more intense combustion.

The decomposition of alkali nitrates appears to be a reversible reaction. When heated in oxygen at a pressure of 175 atmospheres at a temperature gradually rising from 395° to 530° C. during nine hours, sodium nitrite is almost completely oxidised to nitrate. Thus:

2[NaNO2] + (O2) = 2[NaNO3] + 45,000 calories.

Calcium nitrite undergoes oxidation to nitrate in similar circumstances.

Potassium permanganate decomposes when gently heated. The pure, dry salt shows signs of decomposition at 200° C. The reaction is appreciable at 215° C. and is complete at 240° C. The oxygen pressure of the residue corresponds with that of pure manganese dioxide up to 485° C. The heat of dissociation of potassium permanganate is 60,000 calories.

The reaction proceeds approximately according to the equation

2KMnO4 = K2MnO4 + MnO2 + O2,

a residue of potassium manganate and manganese dioxide being obtained.

When a mixture of manganese dioxide and sodium hydroxide is heated to dull redness in a current of air, sodium manganate is formed:

4NaOH + 2MnO2 + O2 = 2Na2MnO4 + 2H2O.

The absorption of oxygen begins at 240° C., the rate of absorption increasing with the temperature, the optimum temperature being 600° C. The product, on treatment with steam at 450° C., evolves oxygen, sodium hydroxide and manganese dioxide being regenerated:

2Na2MnO4 + 2H2O - 4NaOH + 2MnO2 + O2.

The foregoing reactions were made the basis of a commercial method for the preparation of oxygen from the air, but, owing to the short life of the solid phase, the process has not proved particularly successful.

Teissier and Chaillaux suggest the employment of barytes and manganous oxide which are heated together to redness with the production of manganese dioxide and barium sulphide:

BaSO4 + 4MnO = BaS + 4MnO2.

The temperature is now raised to white heat, whereby the dioxide dissociates. Thus:

4MnO2 = 4MnO + 2O2.

Finally steam is injected under pressure, reconverting the barium sulphide into sulphate and liberating hydrogen:

BaS + 4H2O = BaSO4 + 4H2.

These reactions are interesting as constituting one of the few commercial processes in which hydrogen is simultaneously obtained in equivalent quantity to the oxygen.

The alkali bichromates, when gently heated with concentrated sulphuric acid, are converted into chromium salts with liberation of oxygen. Thus:

2K2Cr2O7 + 8H2SO4 = 2K2SO4 + 2Cr2(SO4)3 + 8H2O + 3O2.

The change in colour undergone by the mixture during the reaction is very marked, the deep red of the bichromate giving place to the deep green of chromic sulphate.

Upon exposure to moist air cuprous chloride absorbs oxygen, being converted into the basic oxide Cu2OCl2. This, on heating to 400° C., yields free oxygen and a residue of cuprous chloride, from which the basic salt can be obtained again as indicated above. The initial supply of basic cuprous chloride may be conveniently obtained by heating a moist mixture of cupric chloride, sand, and clay in a current of steam at 100° to 200° C.

Other processes that have been suggested involve the use of nitrosulphonic acid and haemoglobin.

Plumboxan, a mixture of the manganate and meta-plumbate of sodium, namely, Na2MnO4.Na2PbO3, readily evolves oxygen when heated in a current of steam at 430° to 450° C. The plumboxan is regenerated at the same temperature by replacing the steam with air, the issuing gas, during the initial stages of regeneration, consisting of a fairly pure nitrogen.

The oxygen obtained by this process is very pure if the precaution is taken to remove the last traces of nitrogen from the pores of the plumboxan after regeneration by connecting to a vacuous vessel before introducing the steam. The chemical reactions taking place are very complex, and but imperfectly understood.

Orthoplumbates of the alkaline earth metals yield oxygen when strongly heated. The calcium salt, Ca2PbO4, is readily obtained by heating calcium carbonate and lead oxide in the presence of air at about 600° C. When heated more strongly, the salt dissociates, yielding free oxygen, the dissociation pressures being as follow:

Temperature, ° C8809401020106011001110
Pressure in mm.471123505579401040


Although a higher temperature is required for the preparation of oxygen by this method than is the case with barium peroxide, the calcium plumbate is more rapidly regenerated in the presence of air when the temperature is lowered; furthermore, it is not necessary to remove the carbon dioxide from the air as in Brin's process.

The reactions entailed may be represented by the equation

4CaCO3 + 2PbO + O2(from air) ⇔ 2Ca2PbO4 + CO2
2Ca2PbO4 ⇔ 4CaO + 2PbO + O2.

The oxygen may be derived from calcium plumbate, however, in other ways than by heat alone. One method consists in heating to about 700° C. in carbon dioxide:

4CO2 + 2Ca2PbO4 ⇔ 4CaCO3 + 2PbO + O2.

The residue is then heated successively in steam and air whereby the plumbate is reformed.

Another method consists in exposing the calcium plumbate to moist furnace gases at a temperature of about 80° to 100° C. The carbon dioxide is readily absorbed, the solid phase being converted into a mixture of calcium carbonate and lead dioxide. On raising the temperature, oxygen is evolved, the process being facilitated by the introduction of steam. The calcium plumbate is then regenerated by heating in air.
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