THEORETICAL.

The proportion of chlorine liberated during the decomposition of chlorates by heat depends mainly on the nature of the base and the mode of heating. In order to explain this, two theories have been suggested. Schulze (loc. cit.) supposed the chlorate to decompose entirely into chloride and oxygen, the chlorine resulting from the action of "nascent oxygen' upon the chloride. W. Spring and Prost (Bull. Soc. Chim, 1889, [iii], 1, 340), on the contrary, suggested that the chlorate decomposes entirely into oxide and chloric anhydride, Cl205, the latter immediately breaking up into chlorine and oxygen, more or less of the chlorine then reacting with the oxide to form chloride with the liberation of more oxygen.

It will be noticed that these explanations are in direct opposition, but in neither of the papers does there appear to be evidence that the suggested second action actually takes place under the conditions obtaining in the decomposition, nor does either deal with possible alternatives, of which there would seem to be two, namely, (1) the simultaneous formation of both oxide and chloride as direct products and (2) the simultaneous action of chlorine and oxygen upon the residue first produced.

For the purpose of discussion, it is convenient to classify the different reactions which might give rise to the formation of oxide and chloride (evolution of chlorine and oxygen) during a decomposition. (a) Chlorate giving chloride and oxygen.

(b) Chlorate giving oxide, chlorine, and oxygen.

(c) Chlorate acting upon chloride with liberation of chlorine.

(d) Oxygen and chloride giving chlorine and oxide, apart from reverse action (e).

(e) Chlorine and oxide giving oxygen and chloride, apart from reverse action (d).

(f) Simultaneous action of oxygen and chlorine, as in (d) and (e) combined.

Decomposition of Barium Chlorate.-Averaging three of the series of experiments given in Table I, we obtain the following:

The action between a gas and a solid is usually increased by rise of temperature, and increases with the time of contact and the concentration of the gas; the two latter factors are included in the numbers given in the third column (duration multiplied by pressure in atmospheres). According to Spring and Prost, the chloride is produced by the action of chlorine upon the oxide first formed. Comparing the slow decompositions at 1-2 mm. and at 1 atm., we find that a slight fall of temperature combined with a reduction of the time-concentration factor from 165 to 0.5 has very slightly increased the amount of chloride from 99.907 to 99.934 per cent. of the possible amount, instead of very greatly decreasing it. Hence the chloride must be formed in another way, and Spring and Prost's theory does not hold for this chlorate.

Any reabsorption of chlorine which occurs is evidently not complete at atmospheric pressure, and would be much less so when the concentration of the gas is reduced by expansion (compare the experiments proving reabsorption by heated glass); such reabsorption therefore necessitates an increase of free chlorine on reduction of pressure, but none occurs at 1-2 mm., hence reaction (e) cannot occur to an appreciable extent. It may be noted that both oxide and chlorine are very greatly diluted with chloride and oxygen respectively. Comparing the effect of rapid decomposition at atmospheric pressure with that of reduction of pressure to 1-2 mm., we see that in either case the time concentration factor has been reduced to about 05, yet this change has been accompanied by a slight decrease of chlorine in the latter case, but by a sevenfold increase in the former. The increase with rapid decomposition must therefore be due to the great rise of temperature instead of to rapidity of removal of the gaseous products as supposed by Spring and Prost. The increase does not really seem a necessary consequence of this theory, as the rapid formation of oxide would partly compensate for the decrease in the time, and the great rise of temperature might even cause more complete absorption by accelerating the reaction between oxide and chlorine. It seems probable that the proportion of free chlorine is not affected by variations of pressure and that the slight decrease at 1-2 mm. is due to reduction of temperature.

Schulze's hypothesis cannot apply to barium chlorate, for it has been shown that no chlorine is expelled from barium chloride by barium chlorate or any of its decomposition products under the conditions actually obtaining during a decomposition. It has also been noticed that the first bubbles of gas contained about the average proportion of chlorine, although only traces of chloride had then been formed. The chloride experiments exclude reactions (c) and (d), and show that, in this case, (f) coincides with (e), which has also been excluded; as chloride and oxide are actually formed, it is concluded

that reactions (a) and (b) occur during the decomposition, the average velocity of (a) being about 1000 to 1500 times that of (b) when the decomposition proceeds slowly, but at a higher temperature, when the decomposition is rapid, the ratio is only about 140: 1. These average velocities represent the number of molecules of chloride to each molecule of oxide.

From the heats of formation, we obtain :

(a)

(b)

2Ba(ClO3)2 2 BaCl2 + 602 + 438K.

( 2Ba(ClO3)2 = 2 BaO+2Cl2 + 502 – 972K.

2 Ba(ClO3)2 = 2 BaO2 + 2Cl2 + 402 – 624K.

The result of rapid decomposition thus appears to be merely an example of an endothermic reaction (b) gaining upon an exothermic one (a) when the temperature is increased.

Decomposition of Potassium Chlorate. In this case, the evidence is of a somewhat negative character, but as less than 0.007 milligram of chlorine is present in the 400 litres of gas (measured at about 530° and 1.5 mm.) from 1 gram of the substance, it seems extremely improbable that any appreciable amount is evolved at first. The last stages of the reabsorption would be exceedingly slow, as the oxide would then have been all but completely transformed into chloride; in the final residue, less than 0.002 per cent. of the potassium can remain as oxide. The improbability is increased by the fact that no such reabsorption was detected in the decomposition of barium chlorate at atmospheric pressure when the proportions of oxide and of free chlorine were respectively at least 50 times and 25,000 times those just given for potassium chlorate.

It appears that the direct decomposition into chloride and oxygen is the only one which need be considered, this proceeding at a rate at least 50,000 times as great as any reaction yielding chlorine.

General Considerations.

The experiments in this paper, and some already performed with lead chlorate and calcium chlorate but not yet published, tend to indicate the general conclusion that when a chlorate is heated it undergoes two simultaneous and independent decompositions, (a) into chloride and oxygen, (b) into oxide, chlorine, and oxygen; it remains to be shown that this view will harmonise with the results of Schulze's and of Spring and Prost's experiments. In each of these papers, a point is made of the increase of chlorine with increase of weakness of the base; now as the affinity for oxygen approaches that for chlorine, there would be more tendency for oxygen to attack the chloride, less for chlorine to react with the oxide, and more tendency for the oxide to be directly produced; this point therefore accords

VOL. LXXVII,

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equally with each of the three theories. The increase of chlorine with rapid decomposition has already been dealt with under barium chlorate; the only remaining point brought forward by Spring and Prost is the suggestion that the proportion of free chlorine is increased by addition of an acidic oxide because it combines with the liberated basic oxide and so prevents reabsorption of the chlorine first liberated. This cannot, however, be a generally correct explanation, for many such substances cause this action with potassium chlorate when 50-200° below the temperature at which this salt undergoes appreciable decomposition when heated alone (Baudrimont, J. Pharm., 1871, [iv], 14, 81 and 161; Fowler and Grant, Trans., 1890, 57, 272); the "liberated oxide" would then be nonexistent. It does not seem remarkable that these substances should expel chlorine and oxygen (chloric anhydride) from chlorates, as many expel sulphuric anhydride from sulphates.

Schulze also shows that the amounts of free chlorine obtained by the decomposition of the chlorates of sodium, barium, and strontium agree with those resulting when equivalents of the chlorides are heated with potassium chlorate, and that comparable results would probably be obtained with other metals. Double decomposition, however, would be expected to take place with the formation of a chlorate more readily decomposed than that of potassium; these experiments would thus be decompositions of the respective chlorates rather than a study of the action of "nascent oxygen." In section III this has been shown to be the case with barium chloride, and there seems little doubt that the same will apply with other metals, as reaction takes place at a relatively low temperature.

It thus appears that the theory of two independent decompositions is in harmony, not only with the present investigation, but also with the results supposed to support the two older theories.

The author desires to express his thanks for facilities afforded him in the Davy-Faraday Research Laboratory.

XVI. The Interaction of Sulphuric Acid and
Potassium Ferrocyanide.

By RICHARD HALIBURTON ADIE, M.A., B.Sc., and KENDALL COLIN
BROWNING, B.A.

THE interaction of concentrated sulphuric acid and potassium ferrocyanide seems to have been first investigated by Döbereiner (Schweigger's Journ., 1820, 28, 107), who stated that pure carbon monoxide is formed (compare Berzelius, ibid., 1820, 30, 57).

Fownes (Phil. Mag., 1844, [iii], 24, 21), apparently in ignorance of Döbereiner's previous work, to which he does not allude, found that nearly pure carbon monoxide is formed, accompanied at first with traces of hydrocyanic acid and carbon dioxide. The residue consisted chiefly of ferrous, ammonium, and potassium hydrogen sulphates. Towards the end of the reaction, ferric sulphate and sulphur dioxide were formed, and crystals of anhydrous iron ammonium alum deposited.

Merk (Repert. Pharm., 1839, 68, 190), by rapidly distilling potassium ferrocyanide with sulphuric acid, obtained a distillate containing a little prussic acid, thiocyanic acid, and formic acid; he also obtained a sublimate of needle-shaped crystals of ammonium sulphite. Everitt (Phil. Mag., 1835, [iii], 6, 97) first showed that, with dilute sulphuric acid in slight excess, hydrocyanic acid is given off and a new salt, Everitt's salt K,Fe(CN)。, left.

Wittstein (Vierteljahr. Pharm., 1854, 4, 515), Aschoff (Arch. Pharm., 1860, [ii], 106, 257), and Williamson (Annalen, 1846, 57, 225, and Memoirs Chem. Soc., 1848, 3, 125) also investigated the reaction. Considering the differences between the conclusions of these investigators and the almost complete absence of quantitative results, the authors have investigated the decomposition of potassium ferrocyanide by sulphuric acid of concentrations varying from that of the approximately pure acid containing 98 per cent. H2SO4, to that of the acid represented by H2SO4,18H2O. The potassium ferrocyanide used was recrystallised until no impurity could be detected. It was dried at 104-105° until its weight was constant, and kept in a desiccator over phosphoric oxide. The sulphuric acid used was freshly distilled and was free from dissolved sulphur dioxide.

Estimation of Carbon Monoxide.

A weighed amount of ferrocyanide was introduced into a small flask or wide test-tube, which could be heated in an oil-bath, and the apparatus completely filled with dry carbon dioxide. The sulphuric acid was then run in by means of a tap-funnel and the vessel heated, a slow current of carbon dioxide being maintained. The carbon monoxide was collected over aqueous caustic potash (1:2). Temperatures were taken during the reaction by means of a thermometer immersed in the acid.

Estimation of Hydrocyanic Acid.

The anhydrous salt was placed in a flask and the apparatus filled with hydrogen, free from oxygen. The acid was run in and the mixture distilled into two sets of potash bulbs. In order to keep the concentration of the acid constant, a small double surface condenser was used, and at the end of each estimation its temperature was