NOTES
April, 1957 The precipitate was also analyzed for chloride by titration with AgN03using thiocyanate as the indicator. Anal. Calcd. for Tetren.5HCl: C1, 47.7. Found: C1, 47.2, 47 .l. The low chloride analysis is due to incomplete drying of the sample. Attempts to obtain a completely dried sample always resulted in partial decomposition of the salt. b. Tetren Pentahydr0nitrate.-About 150 g. of the distilled amine was dissolvedoin 200 ml. of 95% ethanol and the solution cooled to 10 . One hundred ml. of concentrated HN03 was added dropwise so that the temperature did not exceed 20'. The white precipitate was filtered by suction, recrystallized five times from ethanol and water, and washed with alcohol and ether after the last recrystallization. The precipitate was dried by suction. A weighed sample of the salt was titrated potentiometrically with standard sodium hydroxide solution; 0.1289 g. of the salt required 5.20 ml. of 0.2318 M NaOH, indicating that the salt formed is the pentahydronitrate. Paper chromatographic tests on all salt samples indicated the presence of only one component. 11. Preparation of Solution.-The standard aqueous amine solution was prepared by dissolving a weighed amount of the nitrate (chloride) salt. The solution was standardized by potentiometric titration with standard sodium hydroxide solution. This solution was then diluted to give a 0.001 M tetren hydronitrate solution. A carbonate-free sodium hydroxide solution was prepared in the usual manner. After dilution it was standardized by titration with standard acid solution using modified methyl red as the indicator. A 0.05000 M sodium hydroxide solution was then prepared from this stock solution. 111. Determination of Dissociation Constants.-The dissociation constants of the tetren were determined by titrating 50 ml. of 0.001 M solution of the nitrate and chloride salts, with 0.05 M NaOH, ming a Beckman model G pH meter and a Beckman all-purpose glass electrode. The titration was accomplished in a 4-neck flask, three necks symmetrically located around the center neck. The center neck was equipped with a stirring rod for mechanical stirring. The calomel and glass electrodes were placed in two of the other necks. The sodium hydroxide was introduced from a 10-ml. buret into the fourth neck. Nitrogen was bubbled through this neck during the titration. The titrations were made at 25 f 0.1', 35 f 0.1' and 45 f 0.1". The p H meter was standardized a t each temperature Kith Beckman buffer solutions.
Discussion It was found that pK1 could not be calculated from the titration data as these concentrations did not make ft large enough. A separate titration was therefore performed in which the solution was made 0.005 M with HN03 and 0.001 M with tetren 5HN03. This solution was then titrated with 0.05 M NaOH. A similar method was used for the chloride salt. pK1 could also be calculated only a t 25", as the accuracy of the pH meter did not permit any calculations a t 35 and 45" because of the correction for [H+]which must be applied to C,. Table I lists the value calculated for one of the five constants of tetren from the experiment data. e
TABLE I DETERMINATION OF pK6 DISSOCIATION CONSTANTSOF TETREN AT 25' log
ctatren
0.0009183 .0009166 ,0009149 .0000132 .0009116
Cs
G.
pH
n __ 1 -?I
Ad
pKs
5723 0.62 9.81 $0.21 0.02 9.98 .54 9.89 4933 .07 .02 9.94 4218 .46 9.99 .05 .02 9.92 3534 .39 10.08 - .20 .01 9.87 2871 .32 10.16 - .32 .01 9.83 9.92
+ -
505
Where Ctetren = total concentration of tetren in solution C. = [tetren H + ] + 2[tetren H 2 + ] + 3[tetren H?+J
+ 4[tetren H4"+1 + 5[tetren Hb6+]
Table I1 lists the values determined for the five constants a t the three temperatures. TABLE I1 DISSOCIATION CONSTANTS AT 35 AND 45' 25'
PIG PK2 PKa PK4 PKb
2.65 4.25 7.87 9.08 9.92
350
450
2.43 3.99 7.54 8.81 9.65
3.74 7.26 8.53 9.38
Further thermodynamic data calculated from them have no more accuracy than that to which the change in pK with temperature is shown; they are therefore not reported. The values so calculated, however, are in line with those expected from a comparison with the lower polyamines. Acknowledgment.-The financial help of the Office of Ordnance Research, U. S. Army, is gratefully acknowledged in this and continuing investigations.
THE DIFFERENTIAL THERMAL ANALYSIS O F PERCHLORATES BY MEYERMELVINMARKOWITZ Department of Chemical Engdneering New York University, Uniuersit?l Heights, New Y d ~ kCity, N . Y . Received November 10, 1068
The significance of differential thermal analysis (DTA) curves may often be increased when taken in conjunction with other studies. By supplementing observed DTA curves with information derived from chemical and X-ray analyses, kinetic mechanismslJ and phase diagrams3t4 have been determined. In a recent paper on DTA by Gordon and Campbel1,s the consistent presence of an exothermic break in the DTA curves of the perchlorates of silver, the alkali metals and the alkaline earths is quite puzzling. Except with such substances as the hydrazine and ammonium nitrates and perchlorates which undergo vigorous oxidation reactions, and various azides, it is to be anticipated that the thermal decompositions of most inorganic materials which yield, a t least in part, gaseous products are endothermic. From results reported in the literature concerning the pyrolysis of potassium perchlorate, it is believed that the endothermic nature of the decomposition of this salt can be adduced and that the general exothermic phenomena observed for the perchlorates6 can be traced t o the occurrence of reaction product crystallization. It has been noted that in the heating of potassium perchlorate, there is a t first the appearance of a (1) R . K. Osterheld and L. F. Audrieth, THISJOURNAL, 86, 38 (1952). (2) R. K. Osterheld and M. M. Markowitz, ibid., 60, 863 (1956). (3) R. K . Osterheld and R. P. Langguth, ibid., 69, 76 (1955). (4) E. P. Partridge, V. Hioks and G. W. Smith, J . A m . Chem. Soc., 68, 454 (1941). (5) S. Gordon and C. Campbell, Anal. Chem., 27, 1102 (1955).
506
NOTES
liquid, followed by complete solution, and finally the deposition of a solid until complete solidification occurs.6-6 Decomposition proceeds partially through a chlorate intermediate, frequently found, although only in low concentrations. The first appearance of liquid in a heated sample is a function of both the temperature and the length of the heating period.'^^ Chemical analyses of samples a t the onset of liquefaction strongly suggest that melting ensues because of eutectic formation between the potassium perchlorate and its decomposition products, potassium chloride and potassium chlorate. Accordingly, the DTA curve will show a pair of endothermic breaks5which may correspond to rapid fusion plus slow concomitant thermal decomposition of the chlorate and perchlorate, then followed by the more rapid endothermic decompositions of the latter materials. As the potassium chloride concentration in the liquid phase increases with prolonged heating, the salt will start to be deposited from the melt. This latter effect will then result in an exothermic break corresponding t o the evolution of the latent heat of crystallization of the metal chloride but moderated by the continuous endothermic decomposition of the perchlorate. That a high concentration of potassium chloride may *beformed during the relatively short interval from the first endothermic break (designated in Fig. 3, reference 5, as “fusion,” 588”) to the exothermic break, may be shown by assuming liquid phase decomposition of the potassium perchlorate while being maintained at an average temperature of 610” for four minutes. Using the rate equation,* IC = 1.31 X 1015e--(70500/RT), a t 610°, the calculated first-order rate constant is 4.71 X set.-', which gives after the heating period, 67.7 mole yo KC1 and 32.3 mole yoIZC104. The potassium chloride content is, of course, augmented during the period of the exotherm, so that at the end of the exotherm the residue is essentially pure solid potassium chloride. Experimental proof for the validity of the aforementioned behavior may be found in the instances of rubidium, cesium, silver and magnesium perchlorates for which DTA’s6had been carried out beyond the melting points of the respective chlorides (715, 646, 455 and 714’). Clearly indicated are endotherms, immediately following the exotherms, terminating a t about 715, 647, 464 and 720°, respectively. The absence of an endotherm indicating the melting of lithium chloride (m.p. 614’) in the DTA for lithium perchlorate trihydrate is inexplicable. Approximate thermodynamic calculations9 (Table I) pertaining to the major reactions assigned to the breaks observed during the DTA of potassium perchlorate appear to substantiate an endothermic decomposition (reaction e), Table I. Negli(6) A. Glasner and L. Weidenfeld, J . Am. Chem. Soc., 74, 2467 (1952). (7) L. L. Bircumshaw and T. R. Phillips, J . Chem. Soc., 703 (1953). (8) A. E . Harvey, Jr., RI. T. Edmison, E . D . Jones, R . A. Seybert and K . A. Catto, J. A m . Chcm. Soc., 7 6 , 3270 (1954). (9) Data taken from 3’. D. Rossini, D. D. Wagman, W. H. Evans, 8. Levine and I. Jaffe, “Seiected Values of Chemical Thermodynamic Properties,” National Bureau of Standards Circular 500,U. S. Government Printing Office, Washington, D. C., 1952.
Vol. 61
gible heat capacity and solution effects were assumed over the temperature range involved (25700’). The low heat of fusion ascribed to potassium perchlorate, 1.7-2.6 kcal./mole, is justified on the basis of the crystallographic transition occurring a t 3OOo1O which denotes the onset of rotation of the perchlorate groups in the solid state. Such rotation generally leads t o low entropies of fusion, and consequently low heats of fusion.ll-ls I n the computation of Table I, an entropy of fusion of 2-3 entropy units was assumed. TABLEI
THE REACTIONSEQUENCEOF POTASSIUM PERCHLORATE
k
DURINQ DTA Obsd. DTA Temp., OC.
AH,
Equation=
a 300 KC104(s, 11)KC104(s, I) b 588 KC104(s, I) + KC104(1) c
+
610 KClOd(1) --f KCI(1) 202(g) d 660 KCl(1) +KCl(s) 1 = liquid, s = solid, g = gas.
Process
kcal./ mole
Transition
3.29
Fusion
1.7-2.6
Decompn.
1.7-0.8
Crystn.
t
-6.1
The vigorous evolution of chlorous fumes found for the hydrated perchlorates of copper, zinc and mercury may be attributed t o the hydrolysis of these salts of weak bases as previously reported for the cases of aluminum, magnesium and iron(II1) hydrated perchlorates producing fumes of perchloric acid or chlorine o x i d e ~ . ~ ~ The decomposition behavior of the perchlorates of strong bases is in distinction t o that of the alkali metal nitrates,16s16where melting occurs a t a temperature considerably below that of rapid decomposition to yield products of low melting points. In conclusion, it may be said that the utility of DTA is enhanced when used with other analytical techniques, that there are no true melting points characterizing the metal perchlorates, l7 and that the general exothermic phenomena evidenced on the DTA curves of perchlorates5 may be related t o the deposition of high melting point solid decomposition products (ie., the metal chlorides) from the firstformed melts. Acknowledgment.-Acknowledgment is made of the helpful discussions with Professor John E. Ricci, Department of Chemistry, New York University. (10) C. Finbak and 0. Hassel, 2.physilc. Chem., BSS, 25 (1937). (11) R. R. Wenner, “Thermochemical Calculations,” McGrawHill Book Go.,New York, N . Y . , 1941,pp. 23-6. (12) C. P. Smyth, Chem. Reus., 19,329 (1936). (13) S. Glasstone, “Textbook of Physical Chemistry,” D . Van Nostrand Co., Inc., New York, N . Y., Second Edition, 1946, PP. 422-4, 462-3. (14) G.G.Marvin and L. B. Woolaver, I n d . Eng. Chem., Anal. Ed., 17, 474 (1945). (15) N. V. Sidgwick, “The Chemical Elements and Their Compounds,” Vol. I , Oxford University Press, Oxford, 1951, pp. 700-1. (16) M. Centnersewer, J. chin. phys., 27, 9 (1930). (17) A. E. Simchen, A. Glasner and B. Fraenkel, Bull. Research Council Israel, 2, 70 (1952).
*
