Apr., 1921
T H E J O U R N A L OF I N D U S T R I A L A N D ENGINEERING CHEMISTRY
The Melting Point
303
of Ammonium Sulfate1
By James Kendall and Arthur W.Davidson CHEMISTRY DEPARTMENT, COLUMBIA UNIVERSITY, New YORK,N. Y.
The implicit trust which most chemists place in tables of constants is unfortunately not always justified. Even for t h e most common chemicals, t h e current data for such fundamental properties as melting point and boiling point are often quite indefinite. For example, t h e melting point of potassium iodide2 is variously given as from 614" t o 723'. For the melting paint of ammonium sulfate, a substance which is produced by the ton daily, t h e figures fluctuate still more widely. The lowest recorded value is 140°, the highest 423"; truly a remarkable variation in a simple physicochemical "constant !" This particular case is cited by a recent writerS as "a striking example of t h e neglect of physical chemistry in Germany;" his own efforts t o furnish a solution, however, only serve t o confuse the problem still further. T o enable us t o escape the possible reproach t h a t physical chemistry is still more flagrantly neglected in this country, a brief discussion of the fusion phenomena of ammonium sulfate and a n explanation of t h e discrepancies in the literature may be here presented. For a more detailed examination of the system: ammonium sulfate-sulfuric acid, reference should be made t o a recent article by Kendall and Landon.' PREVIOUS INVESTIGATIONS
The results of previous investigators may first be summarized. Marchands in 1837 obtained a melting point of 140', which has been handed down ever since in all the textbooks and tables a s the melting point of the neutral salt. After more than 80 years, however, i t has been discovered6 t h a t the work of Marchand, owing t o his rather misleading method of expression, has been misinterpreted, 140" referring not t o the n e u t r a l salt, (NH4)2SOa, but t o the acid salt, NHd.HS04. Hodgkinson and Bellairs' described the use of dried and c a r e f u l l y melted ammonium sulfate in 1895, but gave no' value for the melting point. The objection was immediately raised by Smiths t h a t neutral ammonium sulfate does not melt when heated, but decomposes with lass of ammonia, leaving finally the acid salt, which melts a t 146'. This was confirmed by Reikg in 1902 and by Langmuir'o in 1920. Bridgman1' has reported t h a t acid ammonium sulfate is "entirely melted" a t 150 ', but gives no minimum value. Kendall and Landon obtained 146.9 * 0.5' as the melting point of the acid salt, but did not succeed in melting t h e neutral salt in a sealed tube even a t the boiling point of sulfur. CasparI2 states t h a t the neutral Received February 2, 1921. Kaye and Laby, "Physical and Chemical Constants," 1911, 115. Janecke, 2. angew. Chem., 33 (1920). 278. ' J . A m . Chem Soc., 42 (1920), 2131. 6 Pogg Ann.. 42 (18371, 556. Caspar, B e y . , SS (lOZO), 821. 7 Proc. Chem. Soc., 152 (1895), 114, 8 J...Soc. Chem. I n d , 14 (1895), 629. Monalsh , 23 (1902), 1033. ' 0 J. Am. Chem. Soc., 42 (1920). 282. '1 Pmc. A m . A c a d . S c i , 62 (19161, 125. 12 Loc. c i l . I
Q
salt sinters in a n open tube a t about 310°, melts at 336" t o 339', and decomposes a t 355' with evolution of gas; in a closed tube i t sinters at about 360' and melts a t 417" t o 423". Jiinecke,' finally, in an ambitious attempt t o define t h e essential features of the complete phase-rule diagram for the system HzSOp NHs, claims t o have obtained 251' for the melting point of the acid salt and 367" for the simultaneous melting and decomposition points of the neutral salt under atmospheric pressure. The essential source of t h e divergent values obtained is the instability of t h e neutral salt. All investigators agree t h a t t h e acid salt NHd.HS04 is quite stable a t its melting point; JBnecke even gives i t a definite boiling point of 490', a figure which, in view of the dubious character of his remaining results, must be regarded with considerable' reserve. Kendall and Landon's carefully determined value for the melting point of the acid salt (146.9') is in very good agreement with the results of all previous observers; how Janecke could possibly obtain a figure more than 100' higher (251"), unless he misread his thermometer by loo", must remain a mystery. The lzeutral salt, however, loses ammonia when heated, decomposition being appreciable* even a t 200". When the neutral salt is heated in a n opeiz tube, therefore, the determination of a true melting point is impossible, since the composition of t h e solid phase is changing from minute to minute through loss of ammonia. If this ammonia is allowed t o escape freely and the experiment persisted in long enough, the melting point of t h e acid salt will finally be obtained. If, on the other hand, the apparatus is so arranged t h a t the ammonia evolved is permitted t o accumulate above the salt, decomposition will cease before the acid salt is reached. Thus, Smith found t h a t when dry NHI gas was bubbled through melted NHh.HS04 considerable absorption took place even a t temperatures as high as 420'. The ammonia so taken up was evolved again, however. on passing a current of air through the apparatus, It is obviously even a t temperatures as low as 200'. futile, consequently, t o speak of determining the melting point of neutral ammonium sulfate under atmospheric pressure. This statement holds even if a pure ammonia atmosphere is ensured, for while i t is true that the mixture of neutral salt and acid salt3 obtained on heating will possess a definite vapor tension with respect t o ammonia a t a l t y fixed temperature, and a t s o m e fixed temperature will melt, yet it could only be by a n extreme coincidence t h a t fusion should take place at t h a t very temperature for which the vapor tension just equals one atmosphere and (as will appear below) the coincidence does not occur in practice in this particular case. The values 336" and 357" 1 L O C . CZl.
* Smith, Loc 3
cit. It may be mentioned here t h a t the existence of salts intermediate in
composition between the neutral salt (NHn)BO4 and the acid salt NHd.HSO4 (see Kendall and Landon, LOC.cit.) necessitates the decomposition taking place in stages, and not directly.
T H E J O U R N A L OF I N D U S T R I A L A N D E N G I N E E R I N G C H E M I S T R Y
304
obtained by Caspar and Janecke, respectively, for the melting point of neutral ammonium sulfate heated in open tubes must, therefore, refer t o perfectly indeterminate mixtures, not in equilibrium with their vapor phase. For the same reason the elaborate phase-rule diagram presented by Janecke is hopelessly in error, as may be seen by comparing it with the results obtained by Kendall and Landon with the use of sealed tubes. The actual melting point of neutral ammonium sulfate can, indeed, be determined only by heating the salt in a sealed tube with practically no free air space, t o avoid appreciable loss of ammonia., I t is true t h a t a melting point so obtained refers t o a pressure in excess of atmospheric, but the temperature of the equilibrium solid-liquid changes in general so slightly,' except for tremendous pressure variations,2 t h a t this is of no practical significance. The only definite value reported for the melting point of neutral ammonium sulfate in a sealed tube is t h a t of Caspar, 417" t o 423". I n view of the fact, however, t h a t Kendall and Landon failed t o obtain a melting point for a specimen of the salt suspended in a sealed tube in the vapor of boiling sulfur (445'), i t would seem t h a t Caspar's determination is doubtful. The most probable explanation is t h a t considerable decomposition of the salt occurred before fusion, owing to the air space in the sealed tube being too large, thus inducing too low a value for the melting point. The experiments described below conclusively prove t h a t Caspar's result is in error. E X P E R I M E N T A L PART
A pure sample of the salt was obtained in the form of very fine crystals b y rapid cooling from a concentrated hot aqueous solution, which was well stirred during the -precipitation. The crystals were washed with alcohol and ether successively, and desiccated over 99 per cent sulfuric acid. Small glass bulbs of the type shown in the diagram were packed with the crystals and then sealed off a t the point A, leaving as small a n air space as possible. A sealed bulb was attached t o a nitrogen-filled thermometer (reading t o 560') and suspended in a Pyrex test tube containing powdered anhydrous zinc chloride. This tube was air-jacketed with larger tubes and finally with a beaker, the whole being surrounded b y sheet asbestos, with glass windows for observation. T h e temperature was raised very gradually by means of a number of Bunsen burners to about 550°, the crystals thus being brought t o their fusion point in a bath of molten zinc chloride. I T h e first tubes, made of thin glass, exploded before the salt showed any signs of melting. Later attempts were consequently conducted with bulbs made from thick-walled capillary tubing, with better success. TWO concordant experiments gave melting points of 520" *
A
1 For the mean case, an increase of pressure of more than 30 atm. is required t o produce a change in the melting point of lo. See Findlay, "The Phase Rule," 1918,71. 2 Bridgman, Proc. N Q ~Acad. . Sci., 1 (1915), 514.
Vol. 1 3 , KO.4
5", but in view of the smallness of the bath and t h e uncertainty in the exposed stem correction for the thermometer, this value was regarded as only approximately accurate. The salt showed signs of softening below 500'. The final experiments were carried out with a much larger bath (a one-liter Pyrex beaker, thoroughly insulated with asbestos and provided with observation windows, containing a mixture of fused nitrates stirred by means of a motor-driven brass stirrer) and a calibrated platinum resistance thermometer, The temperature of the bath was allowed t o rise exceedingly slowly (not more t h a n 0.2' per min.) in the neighborhood of the melting point. The salt began t o soften perceptibly a t 490" and finally melted This value may, therefore, be given a t 513' * 2'. as the definite melting point of neutral ammonium sulfate, under a n ammonia pressure of considerably more t h a n one atmosphere.2 High a s this figure may appear in comparison with the results of previous investigators, i t is of interest t o note t h a t i t is still far below t h a t recently predicted by L a n g m ~ i r . ~According t o the octet theory of valence, the melting point of ammonium sulfate should be only a little below t h a t of potassium sulfate (1072'); in reality i t is more than 500' lower. We have here, indeed, the first known example of a n inorganic sulfate with a melting point below t h a t of t h e corresponding chloride. Langmuir, by the use of the same method as was employed in this work, has lately determined the melting point of ammonium chloride as 5 5 0 ° , under a n estimated pressure of 66 atmospheres. The difference between this value and t h a t here obtained for ammonium sulfate is not very large,"but i t is significant, since all other sulfates melt a t temperatures considerably higher than the corresponding chlorides. SUMMARY
Janecke recently pointed out the fact t h a t the melting point of ammonium sulfate is not accurately known, and attempted t o remedy the deficiency. It is demonstrated in this article t h a t Janecke's value for the melting point of acid ammonium'sulfate (251') is more than 100' too high, the correct figure being 146.9' * 0.5', while his value for the melting point of neutral ammonium sulfate (359') is more than 150" too low, the correct figure being 5 1 3 O * 2'. The extreme discrepancies recorded in the literature are shown t o be due essentially t o the instability of t h e ' neutral salt when heated in a n open tube. I n the light of the results here obtained, it would seem t h a t Janecke's plea for moreinvestigations of the physical properties of the common chemicals in everyday use might profitably be amended t o a plea for fewer investigations, of somewhat greater accuracy. 1 For the use of this apparatus we wish t o express our thanks to Professor C. D. Carpenter. 2 Since some loss of ammonia must have occurred before the salt melted, the value here determined is, strictly speaking, only a minimum figure. In view of the small air space left in the sealed tube, however, we feel confident that any change in composition of the salt before melting, a n d consequently any error in the melting point recorded due t o this cause, eannot be appreciable. The only factor t h a t might introduce a n y significant error is the solvent action of the partially molten ammonium sulfate on the glass, which was quite noticeable, b u t unavoidable under the conditions of experiment. 3 J . A m Chem. Soc., 42 (1920). 282.
