September 1949
%
f
INDUSTRIAL A N D E N G I N E E R I N G C H E M I S T R Y
(20) Frary, F. C., IND. ENO.CHEM.,38, 129 (1946). (21) Frenkel, Ya. I., J.Phys. Chem. (U.S.S.R.), 9 , 3 9 2 (1945). (22) Frey-Wyssling, A,, and Mtlhlethaler, K., Vierteljahrsschr. naturforsch. Ges. ZQrich, 89, 214 (1944). (23) Fricke, R., and Jucaites, P., 2.anorg. u. allgem. Chem., 191, 132 (1930). (24) Fricke, R., and Wullhorst, B., Ibid., 205, 131 (1932). (25) Glasstone, S., Laidler, K. J., and Eyring, H., “The Theory of Rate Processes,” 1st ed., pp. 10-25, New York, MoGrawHill Book Co., 1941. (26) Haber, F., Naturwissenschaften, 13, 1007 (1925). (27) Harrington, R. A , , and Nelson, H. R., Am. Inst. Mining Met. Engm. Tech. Pub., 1158 (1940). (28) Hedvall, J. A., Proc. Symposium on Chemistry of Cements, Stockholm, 1938, 42-57. (29) Hedvall, J. A., and Leffler, L., 2. anorg. u. allgem. Chem., 234, 235 (1937). (30) Hendricks, G. W., Huffman, H. C., Parker, R. L., Jr., and Stirton, R. I., paper presented before Division of Petroleum Chemistry, at the 109th Meeting of the AM. CHEM.SOC., Atlantic City, N. J. (31) Hillier, J., Canadian Chem. Process Inds., 28, 128 (1944). (32) Hofmann, U., and Hoper, W., Naturwissenschajten, 32, 225 (1944). (33) Hutig, G. F., and Wittgenstein, E. von, Z . anorg. u. allgem. Chem., 171, 323 (1928). (34) Ivanov, A . S., Trudy Tsentral. Nauch.-IssZedovatel Lab. Kammel Samotsvetov Tr. “Russkie Samotsvety,” 1938, No. 4 , 44-58. 37, 158 (35) Jellinek, M. H., and Fankuchen, I., IND.EN*. CWEM., (1945). (36) Kinsinger, W. G., Hillier, J., Picard, R. G., and Zieler, H. W., J. Applied Phys., 17, 989 (1946). (37) Kobozev, N. I., Acta Physicochim. U.R.S.S., 21, No. 2, 289 (1946). (38) La Lande, W. A,, Jr., McCarter, W. S. W., and Sanborn, J. B., IND. ENG.CHEM.,36, 105 (1944). (39) Laubengayer, A. W., and Weisz, R. S., J. Am. Chem. SOC., 65,247 (1943). (40) Lehl, H., J.Phys. Chem., 40, 41 (1936). (41) Lohse, H. W., “Catalytic Chemistry,” Brooklyn, N. Y., Chemical Publishing Co., 1945. (42) Marton, C., and Sass, S., J. Applied Phys., 14, 522 (1943); 15,575 (1944): 16,373 (1945). (43) Megaw, H. D., Z . Krist., 87, 185J1934). (44) Nahin, P. G . ,and Huffman, H. C., unpublished work. (45) Palache, C., Berman, H., and Frondel, C., “Dana’s System of Mineralogy,” 7th ed.. Vol. 1, p. 520, New York, John Wiley & Sons, 1944. (46) Ibid., p. 525.
202?
(47) Parravano, N., and Onorato, E., Atti accad. nazl. Lincei, 10, 475 (1929). (48) Prebus, A. E., in Jerome Alexander’s “Cglloid Chemistry,’’ Vol. V, pp. 152-235, New York, Reinhold Publishing Corp. 1944. (49) Rathbun, M. E., Eastwood, M. J., and Arnold, 0. M., J Applied Phys., 17,759 (1946). (50) Reichertz, P. P., and Yost, W. J., J. Chem. Phys., 14, 495 (1946). (51) Ridgway, R. R., Klein, A. A., O’Leary, W. J., Trans. EZectrochem. SOC., 70, 16 pp. (preprint) (1936). (52) Rogers’ “Industrial Chemistry,” edited by C. C. Furnas et al., 6th ed., Vol. 11, pp. 949-50, New York, D. Van Nostrand Co., 1942. (53) Schoon, T., and Beger, E., Z . phusik. Chem., A189, 171 (1941). (54) Schoon, T., and Klette, H., Naturwissenschajten,29, 652 (1941). (55) Schwab, G. M., “Catalysis from the Standpoint of Chemical Kinetics,” p. 281, New York, D. Van Nostrand Co., 1937. (56) Shekhter, A , , Roginskii, S., and Isaev, B., Acta Physicochim. U.R.S.S., 20, 117 (1945). (57) Smekal, A., Z . Elektrochem., 35, 567 (1929). (58) Steinheil, A., Ann. Physik, 19, 465 (1934). (59) Stranski, I. N., 2.physik. Chem., 136, 269-78 (1928). (60) Taylor, H. S., in “Twelfth Report of the Committee on Catalysis,” Nat. Research Council, 1st ed., pp. 29-41, Ne= York, John Wiley & Sons, 1940. (61) Toropov, N. A,, and Stukalova, M. M., Compt. rend. acad sci. U.R.S.S., 24,459 (1939). (62) Ibid., 27, 974 (1940). (63) Turkevich, J., J . Chem. Phys., 13, 235 (1945). (64) Verwey, E. J. W., 2.Krist., 91, 317 (1935). (65) Weiser, H. B,, “Inorganic Colloid Chemistry,” Vol. 11, Chap 111,New York, John Wiley & Sons, 1935. (66) Weiser, H. B., and Milligan, W. O., J. Phys. Chem., 36, 3010, (1932). (67) Wyckoff, R. W. G., Science, 104,21 (1946). (68) Yamaguti, S., Sci. Papers Inst. Phys. Chem. Research ( T o k y o ) . 36,463 (1939). (69) Zworykin, V. K., and Hillier, J., in Jerome Alexander’s “Colloid Chemistry,” Vol. V I , pp. 118-59, New York, Reinhold Pub. Corp., 1946. (70) Zworykin, V. K., Morton, G. A., Ramberg, E. G., Hillier, J.. and Vance, A. W., “Electron Optics and the Electron Microscope,” p. 241, New York, John Wiley & Sons, 1945. (71) Ibid., pp. 281-337. RECEIVED J a n u a r y 17, 1948. Presented before the Division of Petroleum
Chemistry at the 111th Atlantic City, N. J.
Meeting of
the AMERICAN CHDXICAL SOCIETY,
Heat Contents and Heat of Formation of Magnesium Nitride a-
HIGH TEMPERATURE MEASUREMENTS D. W. MITCHELL
*
University of California, Berkeley, Calif T h e sys tem, magnesium-nitrogen, contains but one known compound, magnesium nitride. This nitride has received a moderate amount of experimental attention and a number of publications concerning its preparation, crystal structure, and thermodynamic properties are available in the literature. However, disagreement on the stability and lack of heat content data above 415” C. indicated the desirability of extending the present lcnowledge of this substance. This investigation includes a new determination of the heat of formation from heat of solution measurements and heat content measurements up to 1000” c.
T
I
HE magnesium used in the preparation of the magnesium
nitride in this investigation was electrolytic metal produced by the United States Bureau of Mines. I t s analysis was magnesium, 99.9%; copper, traces; aluminum, 0.003%; and lead, traces (spectroscopic). After a number of attempts metal-free nitride was prepared b y maintaining an atmosphere of pure nitrogen over fine filings of the metal contained in an iron boat a t 650” t o 700” C. for 3 or 4 hours and then raising the temperature to 950” C. for 12 hours. Complete elimination of metallic magnesium was particularly desirable because its presence in even very small quantities would have necessitated large corrections in subsequent heat of solution measurements and be-
I N D U S T R I A L A N D E N G I N E E R I N G CHEMISTRY
2028
cause metallic magnesium react,s with platinum and with silica, the only two materials convenient for use as container6 in making the high temperature heat content measurements, Sitrogen analyses were made by distillation of ammonia from solutions of the nitride and titration with st,andard hydrochloric acid; magnesium was determined by phosphate precipitation. hlicroscopic examination of the batch of magnesium nitride used in the t,herniochemical work revealed complete absence of unreacted magnesium. The analysis of this bat,ch was nitrogen, 27.497,; magnesium, 72.147,. This nitrogen percentage corresponds to 99.10% magnesium nitride. The remaining 0.90% was taken as magnesium oxide. A mixture of 99.107, magnesium nitride and 0.9070 magnesium oxide has a theoret,ical magnesium content' of 72.17% which checks tjhe actual magnesium content well wit'hin the accuracy of the analysis. Because magnesium nitride reacts rapidly with at>mospheric moistlure according to the equation Mg,N2
+ 3H20
=
3hIg0
Vol. 41, No. 9
28,000
24,000
90,000 W
dI \
s
i
16,000
2
12,000
I
s 8,000
+ 2NH3
the exposed material was handled in a dry box. Samples for analysis were put into weighed glass vials and samples for use in the calorimeters into suitable containers. The presence of an appreciable amount of magnesium oxide in iiitride produced from metal filings made in air was shown experimentally when t,he temperature of a furnace in which about 5 grams of fine magnesium filings had just been converted to nitride accidentally got out of control and the temperaturc rose to an estimated 1350" to 1400" C . for several hours. When the boat, that had contained the filings was removed from the furnace, in it was found a cottonlike skeleton of magnesium oxide. This skeleton weighed only 35 mg. but had the same dimensions as the original heap of metal filings. This weight of oxide amounts to about 0.5% of the weight of nitride theoretically produced from 5 grams of magnesium. The nitride was made with very carefully purified and dried nitrogen a,nd the only possible source of t,he oxygen was from the surfaces of the metal filings themselves. An interesting observation made while learning to prepare pure material was that magnesium nitride is strongly fluorescent when illuminated with ultraviolet, light. The fluorescent light is bright orange. Heats of solution were determined in the hydrochloric acid calorimeter of the Pacific Experiment, Station of the U. S. Bureau of Mines. The calorimeter and its use are described by Southard ( 1 7 ) . Heat contents were determined in the same laboratory by dropping the magnesium nit,ride contained in a silica capsule from a furnace a t the temperature being investigated into a copper calorimet,er. This calorimeter is also described by Southard (18). Experimental t,hermal values are given in terms of the defined calorie (1 calorie = 4.1833 international joules), all weights are corrected to vacuo and at'omic weights used are the 1947 International Atomic Weights. HEAT CONTEXTS ABOVE 298.16' K.
A nearly spherical silica glass capsule approximately 2 cm. in diameter and with a wall thiclmess of 0.1 em. was used as a container for the measurements of heat content,s ab0.i.e 298.16 K. Even with solid tamping, it was possible to pack only 1.4565 grams of magnesium nitride in the capsule because of the low bulk density of the powdered substance-about 0.4 gram per cubic centimeter. The capsule weighed 3.8888 grams a t the beginning of the experiments. The use of such a small quantity of substance limited the accuracy with which heat contents could be determined. A large capsule would have required extensive modification of the apparatus. The measured heat contents above 298.16' K., corrected for magnesium oxide cont.ent, are given in Table I. The same data are given graphically in Figure 1. It is evident from the positions of the experimental points in Figure 1 that three different allotropic forms of magnesium nitride exist within the temperature
4,000
,
0 200
400
600
800
TEMPERATURE,
Figure I .
O
1000
1200
K.
H e a t C o n t e n t s a b o v e 298.16 OK. of lMagnesium Nitride
range of these measurements. The form stable below the first transition point is referred to &A the alpha form, the intermediate one as the beta, and the one stable above the second transition temperature as the gamma form. The alpha to beta transition occurs at 823 =t3 " K. and the beta to gamma transition a t 1061 * 5" K. Both transitions are very sharp and show no signs of pretransition effects. Also shown in Figure 1 is a curve whose equation was calculated from Shun-ichi Satoh's (16) specific heat equation for magnesium nitride. Satoh's data obtained for the temperature region indicated represent the only high temperature heat experiments on t,he substance described in the literature.
TABLEI.
EXPERIRIEKTAL HEAT CONTENTS ABOVE OF MAGNESIUM ITRI RIDE
SO.
T , K.
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18
....
HT
-
ffzg8.16
846.5 1183.3
19;diz 16,869 11,592 7,659 4,118 13,220 22,625 15,360 25,208
1272.5 698.4 743.6 764.9 770.4 773.3 784.7
27,540 10,682 11,999 12,439 12,669 12,592 12,790
973.4 899~9 728.6 589.3 462.4 776.7
1090.1
....
....
298.16" K.
-
NO.
T,' K.
19 20 21 22 23 24 26 25 27 28 29 30
825.2 836.7 974.6 716.9 816.3 1170.3 1075.5 1024.5 1018.9 1030.2 1037.7
14,394 14,697 18,814 11,118 13,896 24,788 22,136 20,319 20,152 20,500 20,653
1065.1 1051.7 1057.4 821.3 825.5
21,786 21,095 21,170 14,026 14,459
ai
32 33 34 35
....
€IT
Hz~8.18
....
An interesting feature of Figure 1 is the supercooling of the gamma form indicated by the dashed line. N o beta nitride was observed until after the nineteenth experiment. Once the beta form had appeared, supercooling was not again observed. Objection might be raised to the location of the point a t the low temperature extremity of the metastable line since it occurs below the temperature of the alpha to beta transition and it is impossible for the stable alpha form to transform spontaneously to the unstable gamma form. The explanation for the location of this point lies in the experimental technique. Frequently the capsule was put into the furnace when the temperature was different from the experimental temperature selected and then the capsule
INDUSTRIAL AND ENGINEERING CHEMISTRY
September 1949
and furnace were brought to the chosen temperature together. For the point in question, the capsule was put into the furnace at a high temperature, and then cooled to the experimental temperature, thus enabling the nitride t o change over t o the gamma form a t some temperature above the alpha to beta transition and then to b e supercooled to the temperature of the experiment. Heat content equations were derived from smoothed plots of the experimental values for the three forms of the nitride. These equations are:
+
++ +
For Mg3Nz,(a), H~-H,ss.la = -7125 22.81T 3.65 X l O - T z MgENz(P), H T - H B ~ .=~ ~- 10,020 29.60T Mg,Ng(y), H T - H s ~ .= ~ ~-9695 29.54T *
c
The first of these equations fits the data t o within 0.9% and the second and third t o within 0.3%, excluding the highest value as was done in plotting the curves. Heats of transition were obtained by substituting the transition temperature in the heat content equations:
AH = 220 cal. AH = 260 oal. Entropies of transition are obtained by dividing the heats by the temperatures of transition: 220 - = 0.27 cal./deg./mole Mg3N2(a) = illgtNp(@) AS = 823 260 = 0.24 cal./deg./mole Mg,N@) = MgaNt(y) AS = Mg8Nz(a) = Mg8Nz(p) M&N,(@) = Mg3Nn(y)
m1
Differentiating the heat content equations gives specific heat equations:
C d a ) = 22.81 C,(@) = 29.60 C,(y) = 29.54
+ 7.30 X 10-ST
Table I1 contains heat contents and entropies above 298.16' K. at even hundred degree temperatures and at transition temperatures. The heat contents were obtained from the heat content equations and the entropies from the equation:
TABLE 11. HEATCONTENTS AND ENTROPIES ABOVE 298.16' K. OF
T,
K.
350 400
600 600
700 800
823
- HZSS
MAGNESIUM NITRIDE T, HT - Haes.ia
-
ST &ga.lO Cal./Moie ' Cal./O K./Mdle
HT
16
1310 2580 5190 7880 10630 13460 14120 (a)
4.04 7.45 13.27 18.15 22.40 26.18 26.99 (a)
K.
823 900 1000
1061 1061 1100 1200
-
ST 82Qar16, Cal./Mole 'Cal./O K./Mole 14340 (j3) 27.26 ( 8 )
16620 19580 21390 j3) 21650 ( y ) 22800 26750
29.64 32.76
E:% {?) 35.74 38.38
HEAT OF FORMATION
Matignon (19) obtainedavalueof AH = - 119,700 cal. per mole for the heat of formation of magnesium nitride from measurements of the heat of solution in sulfuric acid at a temperature not specified but probably 18" C. Matignon's magnesium contained aluminum for which no correction was made. The most recent work is that of Neumann, Kroger, and Kunz (14) who determined the heat of formation by measuring the heat of solution of the nitride in a hydrochloric acid calorimeter. Their magnesium was nitrided in a boat contained in a tube furnace a t a temperature of from 800" to 850' C. All material not accounted for as magnesium nitride from Kjeldahl analysis of the calorimeter solution after the thermal measurements were made
2029
was taken to be magnesium metal for which the heat of solutioii values were corrected. The heat of formation obtained by Neumann, Kroger, and Kunz is AH = -115,180 cal. per mole a t a temperature presumed to be 20" C. Neumann, Kroger, and Haebler (18) found for the heat of formation AH = - 117,830 and AH = -113,980 cal. per mole from direct combination of the elements in a high temperature bomb calorimeter. Brunner [as reported by Neumann, Kroger, and Kunz ( I d ) ] obtained a value of A H = -177,000 cal. for the reaction Mg3Nds)
+ 6H20(1) = 3 Mg(0HMs) + 2NHLaq)
Bichowsky and Rossini ( 2 ) recalculated Matignon's heat of formation value t o be AH = -137 kg.-cal. and that of Neumann, Kroger, and Kunz to be AH = -112 kg.-cal., calculated the heat of formation of the nitride from Brunner's data to br AH = -108 kg.-cal. and chose the mean of Neumann, Kroger, and Haebler's measurements, AH = -116 kg.-cal., as the best value. Shomate and Huffman ( 1 6 ) determined the heat of solution of magnesium metal in 1 N hydrochloric acid very accurately using the same apparatus as used in the present work. I n order to make the best use of Shomate and Huffman's accurate data it was decided to use a weight of magnesium nitride corresponding t o their weight of magnesium metal. The average weight of magnesium used was 0.5725 gram and the results obtained were for reaction of 1 mole of magnesium in 77.2 moles of hydrochloric * 4190 moles of water. Thus the equivalent amount, or 0.7923 gram, of magnesium nitride was to be dissolved in approximately 1843 grams of 1 N hydrochloric acid. Seven successful measurements of the heat of solution were made (the first run was unsuccessful due to mechanical trouble), the results appearing in Table 111. It is clear from the reaction of magnesium nitride with hydrogen ion, MgSN2
+ 8Hf
= 3MgC'
+ 2:"
that the same final state can be achieved by dissolving magnesium in acid and then dissolving ammonia in the same solution as is obtained by dissolving the nitride in acid. Therefore, if the heats of solution of magnesium, magnesium nitride, and ammonia are known the heat of formation of magnesium nitride can be calculated, for the heat of formation of ammonia is accurately known. Liquid ammonia was dissolved in the hydrochloric acid calorimeter by breaking glass bulbs containing the liquid beneath the surface of the acid solution. The bulbs consisted of sections of glass tubing sealed off at one end and having a capillary stem on the other. Ammonia was introduced through the capillary and condensed in the bulb, and then the capillary was sealed off. The bulbs were weighed before and after filling and the weight of ammonia determined by difference. The fraction of the ammonia that was present as vapor was calculated from the vapor density of ammonia and the space occupied by the vapor. Experimental results were corrected for the heat of vaporization of ammonia. The weights of ammonia dissolved were all within 13.1% of 0.2673 gram, the weight of ammonia corresponding to 0.7923 gram of magnesium nitride. The acid used in the calorimeter was prepared by dissolving magnesium metal in it t o the extent of 1 mole of magnesium in 77.2 moles of hydrochloride acid 4190 moles of water. The quantity of acid used was the same as for the nitride. One ammonia bulb was broken in a beaker of 1N hydrochloric acid solution t o see if any ammonia escaped without reacting with the solution. There was not the slightest odor of ammonia above the solution after the bulb was broken. Table IV contains the heat of solution of liquid ammonia data. The heat of formation of ammonia from the elements a t 1 atmosphere pressure has been accurately determined by Becker and Roth ( 1 ) t o be AHzQ&ls= -11,010 * 70 cal. per mole. By
INDUSTRIAL AND ENGINEERING CHEMISTRY
2030
The unccrtairity here is the square root of t,he sum of the squares of the unccrtaint.ics of t,he individual values.
TABLE111. HEATOF SOLVTIOSOF MAGNESICXI NITRIDEIN HYDROCHLORIC ACID"
(1MgsNz t,o 231.6HC1.12570HzO) Heat Evolved, ~ t , lvt, MgO, 1IgaN1, Due t o MgO MggN? Gram Gram
-sa'.*
~ t , H~~~ Sample, Evolved, No. Gram Cal. 0.7838 2223.26 0.0071 0.8000 2266.11 0,0072 2199.90 0,0070 0.7756 2225.23 0,0071 0.7863 0.0071 0 , 7 8 8 5 2232.33 2256.07 0.0072 0.7967 0 , 7 8 5 7 2226.99 0,0071 " Lverage heat of solution, AHzss.1s
No. 2 3 4
lVt. SHa,
Liquid 0.2671 0.2959 0.2625 5 0.2403 ti 0.2993 .I6.erage heat
H m t of Solution of MgsNr Cal./gram cal./mole 2854 35 2850 29 2854 14 2847.72 2848,79 2849.51 2852 17
0.7767 6.30 2216.96 0,7928 6.39 2259.72 2193,68 0,7686 6.22 2218.93 0,7792 6.30 6.30 2226.03 0.7814 6.39 2249.68 0.7895 6.30 2220.69 0.7786 = -287,894 * 197 cal./inole.
H e a t of Heat of E\olved Solution for l O O ~ , Solution of NHg(l), of S I I s ( l ) , Liquid Vapor "3, Cal. Cal./Gram Cal./Mole 0.0044 257.14 256.91 942.58 16,054 942.74 16,057 0.0038 283.60 282.54 0.0043 282.34 251.12 941.23 16,031 0,0044 231.70 230.47 941.85 16.042 0.0044 941.03 16,OZR 285.79 287.02 of solution, AHm8.18 = - 16,043 * 12 cal./mole. Heat Evol\ed. Cal.
means of heat capacities the heat of formation a t 298.16' K. i. calculated to be AH29~.16 = - 11,036 * 70 cal. per mole. The heat of solution of magnesium, determined in the same volume of 1 normal acid and with a weight of metal corresponding to the weight of magnesium nitride dissolved, is given bv Shomate and Huffnian ( 1 6 ) as AH293.16 = -111,322 * 41 cal. per mole. Cragoe (3) gives the following data on saturated ammonia at 77" F. (298.16°KK.): Heat content of vapor = 630.2 B.t.u./lb. Latent heat of vaporization = 501.7 B.t.u./lb and for the superheated vapor a t 1 atmosphere presswe and 77" F.: Heat content of vapor
=
657.4 B.t.u./lb.
The above heat contents are referred to -40" F. (233.16' K.) the zero heat point. From these values, the heat content of the saturated liquid at 77" F. is 630.2 - 501.7 = 128.5 B.t.u. per pound and the difference in heat content between ammonia vapor A t 298 16" K. and 1 atmosphere pressure and liquid ammonia a t 298.16" K. is 657.1 - 128.5 = 628.9 B.t.u. per pound, or for t k e reaction: X € l l ( g , 1 atm.) = ",(I),
lH~98.16
=
-5005
*
8 cd./mole
The uncertainty is estimated from the statement that the measurements were made to 0.1% or better. The heat of formation of magnesium nitride may now be cald a t e d from accurate data from the literature and from experiiiiental data obtained in this investigation. This is done with the following thermochpmical equations:
+ +
BHCl(aq) = 3L1gClz(aq) 3H2(g, 1 a h . ) 3H2(g, 1 atm.) S 2 ( g ,1 a t m . ) = 2?iH,(g, 1 aim.) 2KH,(g, 1 atm.) = 2NH3(1, 298.16" K.) 2NH3(1, 298.16" K.) 2HCl(aq) = 2P;HaCl(aq) Mg;~N?(cy) 8HCl(aq) = 31LIgClI (aq) 2X€IdX(aq) 3Mg(s)
- 333,866 * 123
+
- 22,072
+
++
3hZg(s)
AH288.16 = -110,240
+
*
- 10,010
* 16
- 32,086
=t
24
- 287,894 * 197
+ X2(g, 1 atin.) = hlg,Nn(a) AH1
140
=t
AH = AH1 $. AH2
+ AH4 - AH5
275 ral./mole
Vol. 41, No. 9
STABILITY O F MAGNESIUM NITRIDE
Data on the stabilit,y of magnesium nit,ride a t elevated temperatures are useful to metallurgists concerned wit,h the produc,tion of mayncsium. Several attempts to measure the dissociation of magnesium nitride have been made. Lipski (11) reports that the nitride dissociated at 460" C. until equilibrium was established a t a nitrogen pressure of 12 em. of mercury, which value was approached from both low and high pressures. Zhukov ( 1 0 ) obtained quit,e different results--he observed no dissociation at 1250' C. Hagg (6) found no measurable dis3ociation at, 985" C. Ficht'er and Scliolly ( 4 )state that the nitride dissociates a t reduced nitrogen pre-' abures at 1500" C . Lafitte, Elchardus, and Grandadam (10) obtained a series of values for the dissociation pressure of magnesium nitride a t various temperatures approaching equilibrium from both low and high nitrogen pressures. These investigators give the di+ sociation pressure as 2.0 mni. at 670" C. and 67 mm. of mwcury at 1040' C, The experiment,al heat. content and heat of formatiori data can be combined wit,h thermal dat'a on magnesium and nitrogen and with an estimate of the entropy of magnesium nitride to oht,ain t,he free energy of formation of the compound as a funct>ion of temperature. The entropy of magnesium nitride may be estimated from Kelley's (6) modifiration of Latimer's rulr for est,imating entropies of oxides? ~
3 2
S2ss.l~ = - Ii a In A,,
+ 23 tz b r In 16 + ( a + b ) SJ -
where a, = number of metal atoms in the chemical formulit b = number of oxygen atoms in t8hechemical formula A , = atomic weight of metal C - 6 ~ where C, = true specific heat a t 298.16" K. p = 6bP SO = a constant, de ending upon the type of formula, taken as -3.1 for magnesium nit,ride which gives satisfactory results for most oxides and probably is satisfactory for many nitrides. Substituting 14 for 16 (atomic weights of nitrogen and oxygen) and the constants for magnesium nitride in this expression gives Sz88,16 = 22.4 cal./deg./mole. Combining the entropy and heat of formation a t 298.16" K. with the high tempemture hea,t content dat'a for magnesium nitzide and the thermodynamic propert,ies of magnesium and nit,rogcri ( 7 , 8,9 ) in t,he usual way the following expression is obtained for tmhefree energy of formation of y-magnesium nitride from the elemerit,s: AFO = -224,240
+ 189.672' + 5.00 X
10-'TZ
+ 8.137'
111
T
With this free energy equation and the relationship betrseer~ standard free energy and the equilibrium constant, aF" = -R7' In IC, the temperature a t which the total dissociation pressure is 1 atmosphere was calculated t o be 1790" K . or about, 1520" C. This result should be epted with caution since it involvcv an estimated entropy for magnesium nitride at, 298.16" K. and thc, specific heat of y-magnesium nitride has been extrapolated over 500" at, high temperature. The effect of a moderate error i n estimating the entropy at 298.16' K. on the temperature of dissociation is not great. For example, an error in the entropy of 2 cal./deg./mole vould change the temperature calculated above by less than 50". The error caused by extrapolating the specific heat, may be greater. The free energy equation is good enough, however, to completely invalidate the latest published work on the dissociation of magnesium nitride, that of T,afit,te, Elchardus, and GrandRdam.
INDUSTRIAL AND ENGINEERING CHEMISTRY
September 1949
LITERATURE CITED (1) Becker, G.,and Roth, W . A., 2. Elektrochem., 40, No. 12,836 (1934). (2) Bichowsky, F. R., and Rossini, F. D., "The Thermochemistry of the Chemical Substances," pp. 115, 341, New Y o r k . Reinhold Publishing Gorp., 1936. (3) Cragoe, C. S., Natl. BUT.8tandurds ( U . S.), C ~ T C42 . (April 16, 1923). (4) Fichter, F.,and Scholly, C., Helu. Chim. Acta, 3,298 (1920) (6) Hagg, Gunnar, 2. Krist., 74,96 (1930). (6) Kelley, K.K.,U . S . RUT.Mines, BUZZ.350,46(1932). (7) Ibid., 371,37 (1934). (8) Zbid., 383,65 (1935). (9) I b a . , 434,88-9 (1941).
2031
(10) Lafitte, M. P.,Elchardus, E., and Grandadam, P., Rm. ind. rninJrale, 375,861 (1936). (11) Lipski, J., 2.Elektrochem., 5,189 (1909). (12) Matignon, Camille, Compt. rend. 154,1351 (1912). (13) Neumann, B.,Krijger, C . , and Haebler, H., 2. anorg. U . allgem. Chem., 204,90 (1932). (14) Neumann, B., Krijger, C . , and Kunz, H., Ibid., 207, 133 (1932). (15) Satoh, Shun-ichi, Sci. Papers Inst. Phgs. Chem. Research (Tokyo), 34,No.1,3992 (1938). (16) Shomate, C. H.,and Huffman, E. H., J. Am. Chem. Soc., 65. 1625 (1943). 32,442(1940). (17) Southard, J. C . , IND.ENG.CHEM., (18) Southard, J. C.,J . Am. Chem. SOC.,63,3142 (1941). (19) Zhukov, I. I., Ann. inst. anal. phys. chim., 3, 14 (1926). RECEIVED October 18, 1948.
Solubility of Hydrogen Sulfide in Sulfur ROCCO FANELLI Texas Guy Sulphur Company, 75 East 45th Street, New Y o r k 17, N. Y . The solubility of hydrogen sulfide in sulfur has been determined in the temperature range 120" to 445' C. for PH# Psvoper= Potmos. The solubility is anomalous in that it increases with rising temperature, passes through a broad maximum, and then falls away to its lowest value at the sulfur boiling point. The data indicate that the hydrogen sulfide reacts with sulfur to form hydrogen persulfides. The modifying effect of hydrogen sulfide on the sulfur viscosity is discussed in relation to its solubility in and its reaction with sulfur.
+
H
YDROGEN sulfide has been shown to have a tremendous modifying effect on the viscosity of sulfur (W,6). Simply bubbling the gas through liquid sulfur ?t atmospheric pressure reduces the viscosity from its maximum of 932 poises at 186188" C. t o less than 2 poises. The modifying effect is believed to be due t o the dissolved gas and its reaction with the sulfur t o form hydrogen persulfides. The only data on this solubility in the literature date back to 1897 when Pelabon (8) published four different values for the sole temperature of 440 C. In view of the paucity of data on this solubility and its bearing on sulfur viscosity, solubility determinations were undertaken throughout the temperature range in which sulfur is normally liquid for the system
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ABSORPTION BULB
The weighed Pyrex absorption bulb (Figure l), with a capacity of 43 cc., was charged with 50 to 60 grams of pure liquid sulfur (1) and allowed t o cool t o room temperature. After reweighing, the open lubricated stopcocks were set in place and the ground joints heated with a mioroflame. Small pieces of sulfur placed on top of the joints quickly melted and flowed evenly over the ground surfaces when the
H P
0 /I II , I
1.1
Figure 1. Absorption Bulb
stopcocks were raised slightly and then puaaed home. Solidification of this thin film of sulfur made a n extremely satisfactory seal. When cool the complete assembly was weighed, and the barometer and room temperature were read. The sulfur in the bulb was then carefully melted over a low flame before it was suspended in a small electric furnace kept at 125' C. All temperatures were controlled by Fenwall thermoswitches, and were given by an iron-constantan cou le inserted in the well of the bulb. Cylinder hydrogen sulfide, g t e r e d through glass wool, was passed a t the rate of 60 t o 120 bubbles per minute through the sulfur, heated to the desired temperature for the duration of the experiment, which varied from 16 to 90 hours. The gas issuing from the bulb passed through a weighed tube packed with glass wool and finally against a 2-om. head through a water trap. The run was terminated, after the barometer was read, by stopping the flow of hydrogen sulfide and closing both sto cocks. The bulb was immediately removed from the furnace at &e operating temperature, and its outlet arm connected to the inlet tube of a Schiff gas buret. The buret inlet tube wad fitted with a 15-cc. expansion bulb with intermediate stopcock. Thus any gas liberated during cooling and solidification of the saturated sulfur was collected over mercury, and its volume reduced to standard conditions. The cooled bulb with stopcocks closed was finally placed in a desiccator and weighed 6 hours later. This weight minus the original weight gave the amount of hydrogen sulfide left in and over the sulfur. The tube packed with glass wool, kept a t room temperature during the run, was also weighed, and an? sulfur in the water trap was filtered on a Gooch filter, dried at 70 to 8 0 " C., and weighed. Owing t o the low vapor pressure of sulfur throughout most of its liquid range, very little sulfur was found in the tube packed with glass wool. It rarely amounted t o more than 3 mg. and was usually about half this value. The sulfur found in the water trap was of the same order of magnitude. The total weight of hydrogen sulfide given by the summation of the separate determinations was corrected for replacement of the air originally present in the bulb by the hydrogen sulfide at the temperature of the run and under its partial pressure. LIBERATION OF HYDROGEN SULFIDE
The volume of hydrogen sulfide collected during cooling and solidification of the saturated sulfur varied from 0 t o 90 cc. per 100 grams of sulfur ( a t normal temperature and pressure). Sulfur saturated at elevated temperatures sometimes evolved up to 90% of the absorbed hydrogen sulfide when the bulb stop cocks were kept closed during cooling and incipient solidification. During this period no hydrogen sulfide appeared to be liberated until just prior t o and during solidification, when the gas was so copiously evolved t h a t the liquid appeared t o boil. At the first sign of gas liberation, the stopcock on the bulb arm con-
