PHASEBEHAVIOR AND THERMAL PROPERTIES OF SH,F-HF
August, lMl Ethane. -
1291
tion 14 may be a significant rontrihiitor to methane production even at 200". A t all temperatures the material balance of Carbon Monoxide.-Reactions (3), (6) and (7). CH, 2(C2Hs) is about 5 5 7 , of the CO yield. If Acetone.-acetone arid acetyl propioriyl are added, about SOYo CHB CH&O +CHBCOCHB (16) of the CO yield is accounted in terms of methyl CH1 + IC"yCO'I> +CHjCOCH, + CH3CO (17) radical utilization. The fate of the remaining 20% of methyl radicals is not known. The mechanism Ketene and Acetaldehyde.described does not account for the unsaturated hy2CHiCO +CHZCO CHaCHO (18) drocarbons or for CO?. The latter compound does CH?COCOCHj --+ CHZCO CHaCO (19) not appear to be a secondary product as its formaAcetyl Propiony1.tion is approximately a linear function of total CH, + CH2COCOCHj --+ CHaCH?COCOCHX 120) dose and pressure. In experiments 87A and 87B. the effect of dose 1:xcepi for. the formation of H2 and the hot methyl rcacticin (12), all of these reactions have rate was compared and relatively more CH, was found a t the low dose rate than a t the higher h e n propoied iii photolysis studies. At 25", reaction 14 must account for the hulk dose rate. This is expected from the mechanism of the methaiie, n ith reaction 12 accounting for and the contribution of reactioli I S to methane less than 20% of it and reaction 13 another 20% formation a t 120". The effect of temperature and pressure on the or so if one awumes Xusloos and Steacie's value of 0.G X IO-'' for the function 2 to represent yields of acetone and ketene can be ascribed t o the limiting value of the methane formed by the increased importance of reaction 17 to the proabstraction by thermal radicals. These reactions duction of acetone and reaction 19 to the producthus account for the observation of a low activation tion of ketene at elevated temperatures. It is clear that the basic free radical mechanism energy near room temperature and for the relative which has been postulated to explain the photolysis independence of the methane/ethane ratio with pressure, doie rate, etc The ratio of CH4/C2He of biacetyl, can also be used to explain qualitatively was significantly higher in the 12-liter flask where the major products in the radiolysis of biacetyl. the radical could escape to a region of lorn radical Small contributions to methane a t room temperaconcentration where reaction 13 would become ture and most of the hydrogen production may be significant. Fieaction 14 has been shown to he of ascribed to "hot" radicals, though other mechaimportance in the room temperature photolysis nisms have not been completely eliminated. No mechanism has been proposed for the formation of biacetyl12 and acetone.6 At elevated temperatures, reaction 13 is expected of unsaturated hydrocarbons and carbon dioxide, to become the major methane source. The evidence which products are not reported in the photolysis indicates it t o do so. The activation energy for of biacetyl. Acknowledgments.-The authors wish to thank methane formation was an average of 4.7 kcal. between 25 arid 120", and 7.0 kcal. between 120 N r . William Everette for aiding in the electron and 200". This latter figure agrees with the irradiatioiis with the microwave linear accelerator, photolysiq resiilts of Blaret and Bell,', but is lower and to Dr. K. L. Hall and Mr. Norman Shields of than those of Ausloos and Steacie. This, together the California Research Corporation for use of with the high values of the function Z a t 120 and their Van de Graaff accelerator and for aid in the 200" compared to those of Ausloos and Steacie12 irradiations with it. The authors also wish to and Sheats and Xoye715 indicate that under the thank Dr. Peter Ausloos of the National Bureau of high iiitwiiity conrlitioiis of our irradiations, reac- Standards for helpful suggestions and criticisms. CHa
+ c& +C2Hs
(15)
+
+
+
+
PHASE BEHAVIOR ,4SD THERMAL PROPERTIES OF THE SYSTEM KHdF-HF' BY ROBERTD. EULER
EDGAR F. WESTRUM, JR.
Deportment qf Chemistry, University of Michigan, Ann Arbor, Michigan Received October 12, 1960
The system NHIF-HF \?-as studied by thermal analysis between the limits XHIHFP and HF. S o indication of the composition NHilH2F3previously reported by Ruff and Staubc was found. Low t'emperature heat capacity measurements on four conipositions approximating NH~HBF, confirmed and extended the thermal analysis and revealed the existence of a solid solution at this composition. Several thermal anomalies were found between 180°K. and the melting point. The decomposition pressure of YHdHZF, \vas also determined.
Introduction Ruff alld St:iube? investigated the KH,F-HF sYs(1) From the dirmertation of R. D. Euler submitted in partial fulfillment of the requirements o f the Doctor of Pllilosophy Degree at tlie University of hlichigan. (2) o. ~ u f and f I,. stauk z. anorg. allgem. Chem.. 219, 399 (1933).
tern over a composition range from 50 to 85 mole yo HF. Their data indicated the occurrence of the congruently melting compounds XH4HF2, SH4HzF3, NH4H3F4and KH,HZFg, as well as a' solid phase invariance occurring st - 3" extending over the entire composition range investigated.
1292
ROBERTD. EULER AND EDGAR F. WESTRUM, JR.
Analysis of the NH4F-HF system was undertaken to extend the solid-liquid equilibrium over the entire composition range, with the further intention of investigating the unusual solid-solid transition calorimetrically. A thermal analysis of the system over the composition range from NH4F to NHdHF2 has been made recently in this Laboratory.s In extending the phase behavior study to higher relative H F concentrations, important deviations from the data of Ruff and Staube were found in many features of the system. Melting point maxima occur a t compositions corresponding to NH4H3F4and NH4H6Fs,but a t lower temperatures than those previously reported. Furthermore, neither the compound NH4H2F3nor the indicated (-3") solid-solid transition were found. Since the transition of original interest was not observed, four samples near NH4H3F4in composition were prepared for equilibrium calorimetric study to verify the new findings. Experimental
.-
Preparation of Ammonium Monohydrogen Difluoride NH4HFz waa prepared by addition of 48% aqueous H F to analytical reagent (NH4)zCOa in a large silver beaker to obtain a mole ratio of H F to NHa of slightly less than two. The resulting solution then was boiled to expel COn and allowed to cool to room temperature with the formation of large acicular crystals of NHIHFz. After decanting the supernatant liquid, the crystals were allowed to air-dry for an hour and then were placed in a polyethylene beaker and evacuated for 40 hours. Spectrographic analysis of XKHFa prepared by this method shows no silica or water and only a trace of silver. The samples were analyzed by titration of the excess H F with standardized 0.1 N NaOH solution using brom thymol blue as indicator. Phenolphthalein also was used as an indicator and yielded more reproducible results if the solution being titrated was cooled to the ice point. As a check on the acidimetry, the fluoride was determined gravimetrically by the precipitation of PbClF and good agreement (&0.15% of theoretical) was obtained. Preparation of Ammonium Trihydrogen Tetrafluoride The pre aration of the four calorimetric samples approximating KH4HaFh was accomplished by addition of H F to the NH4HF2 prepared and analyzed as described above. The NH4HFz was weighed into a small, Monel reactor, with a Teflon-gasketed port and a needle valve for addition of a weighed increment of vacuum-distilled high purity anhydrous HF. The samples of NH4H~Fdthus prepared by weight were further analyzed by acidimetric titration of IIF with standardized 0.1 N NaOH to the phenolphthalein end-point and by Kjeldahl nitrogen determination. Water was not detected with Karl Fischer reagent although the practical lower limit of sensitivity of the test on this substance was 0.1% by weight. Thermal Analysis Method.-The thermal analysis of the NH4F-HF system was achieved with a ten-junction, copperconstantan thermel with reference junctions a t the ice point and a millivolt potentiometer. The thermel was calibrated a t the ice point, the carbon dioxide sublimation point, the oxygen boiling point, and compared between 273 and 373OK. with a thermometer certified by the National Bureau of Standards. Deviations from a standard thermocouple table mere then interpolated from these data. The 155-ml. sample container consisted of a copper cylinder 5 cm. dia. by 7.5 cm. length and 0.7 mm. wall. Two copper mds, 0.7 mm. thick, were force-fitted against shoulders and silver-alloy-brazed in place. An axial Monel filling tube 9 mm. dia. by 6 cm. length was provided with a Teflon-packed needle valve a t the upper (free) end to facilitate the addition or removal of increments of H F gas. A thermocouple well of 2 mm. dia. thin-wall Monel entered off-center through the cover and was so inclined that the end was centered 1.3 cm. from the bottom. The interior was
.-
(3) E. Benjamins, G. A. Burney and E. F. Westrum, Jr.. unpublished data.
Vol. 65
deoxidized by heating in hydrogen and the exterior was buffed to a high polish. Samples of NH4HF2 were introduced into the container and vacuum corrected using the crystallographic density. Changes in composition were effected by addition of 99.8% anhydrous HF, distilled under vacuum into the sample container cooled with liquid nitrogen. The composition change was calculated on the basis of the difference in weight of the sample container before and after addition of HF. The container was packed into two coaxial glass tubes with glass wool. An electrical heater was wound on the inner tube and the system partially immersed in a Dewar containing an appropriate cooling bath and centering rings of Styrafoam. No shaking was found necessary during the cooling curves and supercooling was infrequent. Cryostat and Calorimeter .-The Mark I cryostat used was similar to the one described by Westrum, Hatcher and Osborne.' The adiabatic determinations of heat capacity were made by measuring the temperature increment produced by a measured electrical energy input. Current and potential measurements were made on the electrical heater during the input of energy and on the capsule-type platinum resistance thermometer during drift periods. An autocalibrated White double potentiometer was used in conjunction with a galvanometer having a sensitivity of 0.01 pv./mm. as used. Energy exchange between the calorimeter and surroundings was virtually eliminated by maintaining an adiabatic shield a t the calorimeter temperature. Durations of the inputs were given by a timer operated by a thermostated, vacuum-jacketed, tuning fork. The International Temperature Scale region was established by measurements of the resistance of the platinum thermometer (laboratory designation A-3) a t the boiling point of oxygen, the steam point, and the boiling point of sulfur performed by the National Bureau of Standards and evaluation of the constants in the Callender-Van Dusen equation. Below the oxygen point the thermometer was calibrated by comparison with the Bureau's scale.s Because an electrolytic reaction between lead-tin solder fused to copper and the NH4HaF4takes place, a silver calorimeter with a Teflon closure was constructed with all seams silver-alloy-brazed. This calorimeter (laboratory designation W-8) was basically similar to one already described6 and had a measured internal volume of 88.21 ml. Twenty-four stamped, perforated bilver vanes 0.02 mm. thick were stacked so as to fit firmly against the heater well and the cylindrical calorimeter wall. A Monel screw cap closure with a Teflon gasket was provided for introduction of the sample. This closure could be operated from outside the vacuum system and was vacuum-tight near 300°K. To positively prevent leakage at low temperature the closure was further secured by 50-50% by weight Pb-Sn solder. Small corrections were made for slight differences in weight of the Teflon gaskets using data on the heat capacity of molded Teflon.? Apiezon-T grease was used in weighed amount to establish thermal contact between calorimeter, heater sleeve and thermometer. The heat capacity of the calorimeter (empty except for helium gas) was determined separately. I t contributed from 1 to 25% of the total heat capacity measured. Loading the sample consisted of admitting the liquid into the calorimeter and screwing the cap against the Teflon gasket. The calorimrter then was placed in a glais system, cooled to liquid nitrogen temperature, the cap partially opened, and the system evacuated. Opening and closing of the calorimeter inside the glass system were accomplished with a special socket wrench attached to a running standard taper joint. One atmosphere pressure of helium was admitted, the calorimetrr then warmed to room temperature, and the cap rcclosrti. Excess helium pressure could escape through the threads during the warming period. The calorimeter then was removed from the glass system and the cap soldered to the calorimeter. The correction for solder was minimized by re(4) E. F. Westrum, Jr., J. B. Hatcher and D. W. Osborne, J. Chsm. Phys., 2 1 , 419 (1953). (5) H. J. Hoge a n d F. G. Brickweddc, J. Research Natl. Bur. Standards, 22, 351 (1939). (6) D. W. Osborne and E. F. Westrum, Jr., J. Chem. Phw., 2 1 , 1884 (1953). ( 7 ) G. T. Furukawa, R. E. McCoskey and G. J. King, J . Research Natl. Bur. Standards, 49, 273 (1952).
PHASE BEHAVIOR AND THERMAL PROPERTIES OF NH4F-HF
August, 1961
producing closely the weight of solder used on the empty calorimeter. A small correction for differences in the amount of heliurn present was made using the value for the density of liquid NH4HaF, a t 300'K. of 1.30 g./ml. obtained in a rough pycnometric determination.
]Results and Discussion Thermal Aiialysis Measurements.-Five separate series of determinations were made, each employing a new sample of NHdHFz. The results are summarized in Table I and Fig. 1. Over the composition range 50 to 83.3 mole % HF the measuiemerits mere made between 403 and 223°K.; from 83.3 to 100 mole % H F from the melting point to 163°K. TABLEI THERMAL ANALYSISDATA Mole %
HF
51.27 52 39 53.79 54.98 56.13 62.48 63.25 66.04 60.27
N.P.
T. OK.
305 91 305 49 391 78 388.90 386 21 349 0.5 338.2;) 303.4;'
Eutectic T , OK.
Mole %
HF
Series I 78.34 79.54 80.47 81.54 81.76 82.26 82.69 82.90 87.91 267.70 88.28 267.27 89.20 00.14
T,OK.
M.P.
Eutectic T , OK.
285.10 279.07 273.60 264.92 262.82 259.67 '761.40 262.07 200.15 193.95 174.82 172.82
259.52 258.02 258.57 258.57 258.65 258.82 258.67 258.43 171.32 172.08 172.15 167.55
Series I1 395.79 379.21 346.10 302.63 2%. 81; 285.63 276 1:I 270.40 278.15 281 ,911 289.78 292.25 "1.311 290.55 287.70
267.50 266.25 266.00 270.15 266.00
74.18 75.42
293.83 295.40
262.90 264.62 258.02 226.60 175.35
75.02 74. on
FOR THE
NH,F-HF SYSTEM Ruff and Staubel
NH4F-NH4HF2eutectic .., NH4HF2melting point 397.8"K. N&HF2-NH4H,F4 eutectic 292 NH4H2Famelting point 301 N&H3F4 melting point 303 NH4HaF4-NH4HsFseutectic 290 NH,HsFs melting point 303 NH4H6F8-HF eutectic ... Solid-solid trsngition 270 (50-85 mole yo HF) ( -3') 0 Including data from other work in as shown on Fig. 1.
This work
382.2 f 0 4°K. 399.3 f 0.1" 266.710 9
... 206 4 f 0 3 258.8 f 0 . 3 265 2 f 0 . 5 172.6 f 0 . 5
...
this Laboratory3
400
3 50 y' W
5
300
F
a
E W
200 255.67
Series V 255.42 255.75
EQUILIBRIUM TEMPER.4TURES
I-
Series IV 82.97 83.47 84.23 86.47 91.73
OF
3 250
266.40 267,07 266.75 265.80
TABLEI1 COMPARISON
a
Series 111 51.27 56.94 62.48 65.96 66.36 67.46 68.09 69.08 70.20 70. 99 71.59 73.44 76.54 76.68 77.02 77.81
1293
295.00 296 ,2.5
Several differences are evident between the data of Ruff and i$taube2 (dotted lines in Fig. 1) and those of the present investigation. The previously reported compound NH4HZF3 was not observed. Moreover the eutectic temperatures and the melting points of NH4H3F4and NH4H6Feoccur a t considerably lower temperatures than those previously reported. The reported solid invariance did not appear. h comparison of the data of Ruff and Staube with those of this investigation is presented in Table 11. A single explanation for the discrepancies could not he found..' Several factors may have contributed to the difference in results. The supposed solid invariance found by Ruff and Staube may have been in part confusion with the eutectic temperature of 266.5"K. The existence of 270, 292
60
80
IO0
MOLE PERCENT HF, Fig. 1.-Partial phase diagram for the syptem "&€IF. The open circles indicate thermal analysis data of this rpsearch, solid circles thermal analysis of Benjxmins, et ( t i . , 3 open squares calorimetric data from Samples A, B, C and I ), and solid squares those from the preliminary samplc. Dotted lines represent the diagram of Ruff and Staube.2
and 290°K. halts may possibly have been due to a thermal conduction fluid used in the thermocouple well, or an impurity in the sample itself, The NH4HFz sample used by Ruff and Staube melted a t 397.8"K., whereas 399.3"K. was the orthobaric melting point obtained by Benjamins, et aL13and confirmed here. This evidence suggests that their NH4HF2 was impure. The description of the experimental procedure of Ruff and Staube indicates that the H F used to alter the composition was poured into the copper system through a copper funnel, apparently in the open air. This method certainly involves some contamination by water since hydrogen fluoride is highly hygroscopic. If appreciable water were present in the system,
ROBERTD. EULERAND EDGAR F. WESTRUM, JR.
1294
one mould expect a break due to the ?JH4HF2-H20 eutectic a t 258.4°K.8 There is, however, no evidence that Ruff and Staube studied the system to such low temperatures. The existence of a hemi-hydrate of KH4H2F3was reported by Hassel and Kringstad in 1932.9 The article, however, is very vague about experimental methods and the agreement between theoretical and analytically determined compositions is rather poor. The discrepancy hardest to explain is that of the melting behavior over the composition range 65-86 mole % HF. As a further experimental check, compositions corresponding to the compounds SHdH3F4 and NH4H& were prepared and allowed to stand in tightly sealed polyethylene bottles at 297-300°K. After more than six months, the material in both bottles remained in a completely liquid state. Heat Capacity Data.-A sample of NH4H3F4, upon which pr eliniinary heat capacity measurements were made ovcr the range 200 to 320"K., indicated an anomaly at 230°K. and a melting point of 29535°K. with considerable prenielting. It was suspected that since this sample had been transferred several times in the air it was quite impure. Consequently, another sample was prepared and analyzed. This sample was prepared by weight to be slightly richer in HF than the stoichiometric composition since it was considered that the preliminary sample might have been low in H F content. This sample, designated as Sample B, showed the same melting characteristics as found in the preliminary measurements but showed aiiomalous behavior belo\$ the melting point a t 192 and 207°K. as well as a much smaller effect near the aforementioned 230°K. In light of this unusually complex behavior, it was decided to prepare several samples, varying slightly from the stoichiometric composition, and to investigate their thermal properties. Samples A, C and D were prepared by appropriately altering the composition by weight of the previous sample with either SHdHF2or anhydrous HF. Table I11 lists the four compositions investigated, together with a summary of the observed temperatures of melting and of heat capacity anomalies. The calorimetric samples TTaried in mass from 103 to 116 g. TABLE 111 C.ILORIMETRIC DATAON SH,H,F4 1 I o l r 7" iy P , Anomalies,
QLMYARY
OI
HF
K.
OK.
A B C
76 08 75 05 74 62
I>
I 3
294 2 295 9 295 4 294 3
192, 207, 254 192, 207, 232 192, 231 192, 213, 262
Sample
-'
52
Experimental values of the heat capacities of Samples A, B, C and D are listed in Table IV as the "molal" values of heat capacity and energy but are arbitrarily calculated on the basis of 97.064 g., the mole weight of SH4H3F4,although none of the ( 8 ) V. S Yatlov a n d I:. 11. Polyakova, Z h r ObshLhei K h l m , 15, 724 (1948). (9) 0 Hassel and H Kringstad Z anorg allgem Chem , 208, 382 (1912)
J.01. 65
samplea had exactly stoichiometric composition. The results are expressed in terms of the defined thermochemical calorie equal to 4.1840 abs. j . and the ice point of 273.15"K. The data are listed in chronological sequence so the approximate temperature increments usually may be inferred from the adjacent mean temperatures. Curvature corrections were applied in regions where the heat capacity was normal. Values of the heat capacit,y are considered to have a probable error of 0.1%. TABLE IV HEAT CAPlCITIES O F NHdH:jF* SANPLESO N NOLEWEIGHT = 97 064 g. Sample A-76.08 mole yo HF: I T , OK.; Cp, cal.,'(deg.
ARBITRARY
mole) ,Series
11 1b6.95 41.90 1201.82 39.69 1205.00 43.22 1207.55 63.89 1210.19 39.93 1 216.23 41.62 1225.06 43.92 1233.86 48.11 1242.37 56.43 i249.43 75.22 1 254.36 87.01 1257.55 63.56 1263.32 62.37 1271.26 74.17 1277.82 94.801 284.95 167.51290.96 396.1296.52 184.6 1303.10 55.20 1308.51 55.29 1 Series11 1154.27 32.30 1163.02 34.36 1171.80 35.93 1180.38 37.71 1188.24 41.22 1192.61 45.66 194.06 43.83 1195.30 61.22 1196.40 57.60 1197.79 38.42 1249.30 73.29 1251.80 84.52 1253.13 88.60 254.03 92.05 1255.14 85.081256.40 69.70 1257.89 60.39 1259.53 58.42 AHE R u n KO.1;Series 11x1 197.81 37.65 i200.34 38.33 1202.83 39.28 1204.39 40.761205.60 44.89 1206.46 58.20/207.02 99.92 1207.56 55.94 1208.43 46.43 1210.69 40.68 I ~
~
Sample B-75.05 mole % ' HF: Series I 198.59 40.14 i206.02 53.67 I 213.76 40.32 1222.41 40.31 123C.98 44.66 1240.05 44.26 1249.36 46.15 1 258.30 49.14 1266.81 53.83 1274.61 63.06 p231.38 83.88 (286.51 130.9 I 289.75 221.0 1292.36 493.1294.38 1000. 1293.33 1422.j300.94 58.70 I 309.17 55.24 1315.01 55.32 I Series I1 1155.52 31.34 1164.30 33.13 1 173.23 35.15 1181.99 37.72 1190.50 41.47 1199.30 40.00 1207.46 53.20 1 215.45 40.66 1224.42 42.85 1233.36 44.62 1242.28 44.62 1251.10 46.66 1 259.70 49.69 I A H f R u n KO. 11307.89 55.27 1313.70 55.34 1 Series III 1 65.51 15.04 170.08 16.08 /75.Y8 17.13 182.68 18.45 189.90 19.80~97.32 21.09 105.12 22.44 1113.59 23.91 1122.38 25.42,128.95 26.31 132.75 27.20 1138.85 28.30 1147.62 29.82 1156.29 31.38,164.82 33.30 1173.47 35.31 1182.09 37.63 1187.57 39.84 1189.86 41.32 1192.52 42.81 1195.61 41.93 1199.07 36.89 1202.37 37.84 ,205.65 50.40,207.09 140.6'207.39 187.0 1208.69 39.95 1 AHf Run KO. 2 1 ~
~
~
Sample C-74.62 mole % HF: Series 11148.50 29.63 1157.66 31.36 1 166.50 33.17 l175.23 3%27 1183.81 37.91 1188.84 39.96 1190.53 41.11 1 192.16 42,731193.88 37.601195.73 35.941201.05 36.24,209.36 36.99, 218.11 38.19 1223.81 39.961226.65 44.30 1229.28 73.5 i231.21 85.6 233.79 41.69 1239.77 42.501248.56 44.44 1257.50 47.47 1266.24 52.92, 274.63 63.38 ;282.00 90.3 1287.22 154.9 1291.20 384.1293.89 1029. 1 299.06 1 7 1 . 0 ~ S e r i e s I I ~ 6 2 . 6 114.38169.94 16.03176.31 17.14 $2.25 18.30 188.91 19.55 196.90 20.93 1105.17 22.33 1113.31 23.64 122.00 25.14 1131.03 26.58 140.02 28.11 1149.12 29.77 i 138.14 31.46 1167.00 33.31 I175.69 33.44 184.30 38.13 j189.40 40.34 l190.65 41.14 1191.47 41.90;192.28 42.821193.09 41,951197.40 36.04 1206.18 36.671213:iO 37.73 1223.24 39.801227.27 44.52 1228.70 53.31 1229.79 82.0 1230.41 217.1231.12 84.31AHfRun S o . 11300.01 5.5.281304.87 55.11: ~
~
~
Sample D-73.52 mole % HF: Series I 1128.04 26.06 1135.97 27.37 1 144.41 28.81 1153.18 30.38 1162.13 32.12 1170.98 34.04 1179.44 36.25 1 187.68 39.35 1192.97 39.96 1195.48 35.55 1197.73 35.17 1201.21 35.53 1 204.85 35.76 1206.58 36.50 1208.29 36.981211.6C 68.87 1219.07 40.03 1 228.41 42.05 1236.87 44.54 1245.26 49.00 1252.13 56.00,257.17 68.38 260.35 76.74 1261.97 77.41 1263.61 71.371265.42 65.51 1267.37 62.03 i 272.22 66.29 1279.12 84.9 1286.40 164.5 1292.04 466.1294.92 535. ,298.77 55.28 1303.37 55.21 ISeries 111188.13 39.60 1193.20 39.68 1 194.63 35.20 I 195.85 35.24 1200.86 35.31 1191.76 41.63 1193.32 39.11 1206.50 36.62 1 209.00 37.68/211.22 46.34 1212.71 115.1 1213.62 107.1 1213.15 39.90: ~
The Anomalies and Fusion Transition.-Since the heat' capacit'ies of the samples contain PO many anomalies iii a small temperature range! it is difficult to assign a curve for the normal or background beat capacity. Continuous "lat'tice" heat capacity curves, determined by extrapolating the heat capacity from above and below the anomaly to the temperahre of t,ransition were drawn for all maxima except, those of fusion. Enthalpies were calculated by numerical quadrature of these c o i l t'inuous heat capacity us. temperature curves and subtracted from the actual enthalpy inpiits to de-
PHASE BEHAVIOR S N D THERMAL PROPERTIES OF XHd17-HF
August, 1!161
termine the enthalpy associated with each maximum. For heat of fusion calculations, a linear extrapolation of the solid heat capacity of Sample D to the melting point and a constant value of 55.25 cal./(deg. mole) for the liquid heat capacity were used. Enthalpies were calculated from these straight lines and subtracted from actual energy inputs to determine the enthalpy of fusion for all four samples. Since linear extrapolations of the solid heat capacity for Samples A, B, C and D were nearly the same, the extrapolation giving the lowest background enthalpy (that of Sample D) was used. A summary of the enthalpy increments associated with each anomdy is presented in Table V. Figure 2 presents heat capacity curves for the four samples I
I
I
I
equilibrium values and were completely reproducible for a given sample. This may be verified by noting the series of repeated determinations shown in Fig. 2. TABLE V E ~ T H A L PINCREMENTS Y FOR TRARTSITIONS Sample
-4 B C D
Mole %
HF
--Temp. of anomaly (In OK.)192 207 213 231 284
76 08 75 05 74.62 73 52
155 66 348 122 109 42 78 177 102 168
Sample
c
D
8C
_. I
W
6
6C
-z
s4
40
2 2c
u
C
I
-100
I
I
I
150
200
250
300
T, O K , Fig. 2.---Heat rapacity of four samples approximating SH4H3F4in composition. (Note that the lower curves are successively displiiced by increments of 20 cal./( deg. mole) as a convenience in presenting the anomalies.
with the curves for Samples B, C and D displaced from that of the next higher H F composition by 20 cal. '((leg. mole) to more conveniently display the anomalies. Hence, true heat capacities can be read directly from this graph for Sample B (75.05 % HF) only. Enthalpy increments beneath each anomaly and the observed melting temperatures are indicated in Fig. 2. In order to obtain more accurate enthalpy of fusion data some enthalpy-type runs were made over the entire fusion region for the various samples. As a test of the heat capacity-type runs through these regions of relatively slow equilibrium, the enthalpy increments of these runs are compared in Table VI with the integrated heat capacity undrr the curve over the same temperature region. The heat capacity points and the enthalpy increments of the anomalies were determined as essentially
262
Fusion
186
4129 4558 4721 4378
TABLE VI ENTHALPY TYPLI t r ~ s sTHROUGH THE FUSIOX RELIOA (In cal./inole; arbitrary mole weight = 9i.064 g.)
A B (No. 1) (No. 2) 100
1295
Ti,'I
