The Principles of Chemistry, Volume I

by Dmitry Ivanovich Mendeleyev

Nitric oxide is a colourless gas which is only slightly soluble in water ( 1 / 20 of a volume at the ordinary temperature). Reactions of double decomposition in which nitric oxide readily takes part are not known—that is to say, it is an indifferent, not a saline, oxide. Like the other oxides of nitrogen, it is decomposed into its elements at a red heat (starting from 900°, at 1,200° 60 per cent. give N₂ and 2N₂O₃, but complete decomposition into N₂ and O₂ only takes place at the melting point of platinum, Emich 1892). The most characteristic property of nitric oxide is its capacity for directly and easily combining with oxygen (owing to the evolution of heat in the combination). With oxygen it forms nitrous anhydride and nitrogen peroxide, 2NO + O = N₂O₃, 2NO + O₂ = 2NO₂. If nitric oxide is mixed with oxygen and immediately shaken up with caustic potash, it is almost entirely converted into potassium nitrite; whilst after a certain time, when the formation of nitric peroxide has already commenced, a mixture of potassium nitrite and nitrate is obtained. If oxygen is passed into a bell jar filled with nitric oxide, brown fumes of nitrous anhydride and nitric peroxide are formed, even in the absence of moisture; these in the presence of water give, as we already know, nitric acid and nitric oxide, so that in the presence of an excess of water and oxygen the whole of the nitric oxide is easily and directly converted into nitric acid. This reaction of the re-formation of nitric acid from nitric oxide, air, and water, 2NO + H₂O + O₃ = 2HNO₃, is frequently made use of in practice. The experiment showing the conversion of nitric oxide into nitric acid is very striking and instructive. As the intermixture of the oxygen with the oxide of nitrogen proceeds, the nitric acid formed dissolves in water, and if an excess of oxygen has not been added the whole of the gas (nitric oxide), being converted into HNO₃, is absorbed, and the water entirely fills the bell jar previously containing the gas.[60] It is evident that nitric oxide[61] in combining with oxygen has a strong tendency to give the higher types of nitrogen compounds, which we see in nitric acid, HNO₃ or NO₂(OH), in nitric anhydride, N₂O₅ or (NO₂)₂O, and in ammonium chloride, NH₄Cl. If X stand for an atom of hydrogen, or its equivalents, chlorine, hydroxyl, &c., and if O, which is, according to the law of substitution, equivalent to H₂, be indicated by X₂, then the three compounds of nitrogen above named should be considered as compounds of the type or form NX₅. For example, in nitric acid X₅ = O₂ + (OH), where O₂ = X₄, and OH = X; whilst nitric oxide is a compound of the form NX₂. Hence this lower form, like lower forms in general, strives by combination to attain to the higher forms proper to the compounds of a given element. NX₂ passes consecutively into NX₃—namely, into N₂O₃ and NHO₂, NX₄ (for instance NO₂) and NX₅.

As the decomposition of nitric oxide begins at temperatures above 900°, many substances burn in it; thus, ignited phosphorus continues to burn in nitric oxide, but sulphur and charcoal are extinguished in it. This is due to the fact that the heat evolved in the combustion of these two substances is insufficient for the decomposition of the nitric oxide, whilst the heat developed by burning phosphorus suffices to produce this decomposition. That nitric oxide really supports combustion, owing to its being decomposed by the action of heat, is proved by the fact that strongly ignited charcoal continues to burn in the same nitric oxide[62] in which a feebly incandescent piece of charcoal is extinguished.

The compounds of nitrogen with oxygen which we have so far considered may all be prepared from nitric oxide, and may themselves be converted into it. Thus nitric oxide stands in intimate connection with them.[63] The passage of nitric oxide into the higher degrees of oxidation and the converse reaction is employed in practice as a means for transferring the oxygen of the air to substances capable of being oxidised. Starting with nitric oxide, it may easily be converted, with the aid of the oxygen of the atmosphere and water, into nitric acid, nitrous anhydride, and nitric peroxide, and by their means employed to oxidise other substances. In this oxidising action nitric oxide is again formed, and it may again be converted into nitric acid, and so on continuously, if only oxygen and water be present. Hence the fact, which at first appears to be a paradox, that by means of a small quantity of nitric oxide in the presence of oxygen and water it is possible to oxidise an indefinitely large quantity of substances which cannot be directly oxidised either by the action of the atmospheric oxygen or by the action of nitric oxide itself. The sulphurous anhydride, SO₂, which is obtained in the combustion of sulphur and in roasting many metallic sulphides in the air is an example of this kind. In practice this gas is obtained by burning sulphur or iron pyrites, the latter being thereby converted into oxide of iron and sulphurous anhydride. In contact with the oxygen of the atmosphere this gas does not pass into the higher degree of oxidation, sulphuric anhydride, SO₃, and if it does form sulphuric acid with water and the oxygen of the atmosphere, SO₂ + H₂O + O = H₂SO₄, it does so very slowly. With nitric acid (and especially with nitrous acid, but not with nitrogen peroxide) and water, sulphurous anhydride, on the contrary, very easily forms sulphuric acid, and especially so when slightly heated (about 40°), the nitric acid (or, better still, nitrous acid) being converted into nitric oxide—

3SO₂ + 2NHO₃ + 2H₂O = 2H₂SO₄ + 2NO.

The presence of water is absolutely indispensable here, otherwise sulphuric anhydride is formed, which combines with the oxides of nitrogen (nitrous anhydride), forming a crystalline substance containing oxides of nitrogen (chamber crystals, which will be described in Chapter XX.) Water destroys this compound, forming sulphuric acid and separating the oxides of nitrogen. The water must be taken in a greater quantity than that required for the formation of the hydrate H₂SO₄, because the latter absorbs oxides of nitrogen. With an excess of water, however, solution does not take place. If, in the above reaction, only water, sulphurous anhydride, and nitric or nitrous acid be taken in a definite quantity, then a definite quantity of sulphuric acid and nitric oxide will be formed, according to the preceding equation; but there the reaction ends and the excess of sulphurous anhydride, if there be any, will remain unchanged. But if we add air and water, then the nitric oxide will unite with the oxygen to form nitrogen peroxide, and the latter with water to form nitric and nitrous acids, which again give sulphuric acid from a fresh quantity of sulphurous anhydride. Nitric oxide is again formed, which is able to start the oxidation afresh if there be sufficient air. Thus it is possible with a definite quantity of nitric oxide to convert an indefinitely large quantity of sulphurous anhydride into sulphuric acid, water and oxygen only being required.[64] This may be easily demonstrated by an experiment on a small scale, if a certain quantity of nitric oxide be first introduced into a flask, and sulphurous anhydride, steam, and oxygen be then continually passed in. Thus the above-described reaction may be expressed in the following manner:—

nSO₂ + nO + (n + m)H₂O + NO = nH₂SO₄,mH₂O + NO

if we consider only the original substances and those finally formed. In this way a definite quantity of nitric oxide may serve for the conversion of an indefinite quantity of sulphurous anhydride, oxygen, and water into sulphuric acid. In reality, however, there is a limit to this, because air, and not pure oxygen, is employed for the oxidation, so that it is necessary to remove the nitrogen of the air and to introduce a fresh quantity of air. A certain quantity of nitric oxide will pass away with this nitrogen, and will in this way be lost.[65]

The preceding series of changes serve as the basis of the manufacture of sulphuric acid or so-called chamber acid. This acid is prepared on a very large scale in chemical works because it is the cheapest acid whose action can be applied in a great number of cases. It is therefore used in immense quantities.

The process is carried on in a series of chambers (or in one divided by partitions as in fig. [50], which shows the beginning and end of a chamber) constructed of sheet lead. These chambers are placed one against the other, and communicate by tubes or special orifices so placed that the inlet tubes are in the upper portion of the chamber, and the outlet in the lower and opposite end. The current of steam and gases necessary for the preparation of the sulphuric acid passes through these chambers and tubes. The acid as it is formed falls to the bottom of the chambers or runs down their walls, and flows from chamber to chamber (from the last towards the first), to permit of which the partitions do not reach to the bottom. The floor and walls of the chambers should therefore be made of a material on which the sulphuric acid will not act. Among the ordinary metals lead is the only one suitable.[65 bis]