Jan. 20, 1957
THERMODYNAMIC PROPERTIES O F
culated in this manner from the extrapolated slopes a t zero concentration of poly-soap are designated as po in Table V. These values of POin Table V are of the same order of magnitude as the values reported by Phillips and Mysels18for sodium lauryl sulfate, based on a similar calculation. However our values of Pocalculated from the slopes-are much higher than the values of p l and p 2 calculated on the basis of eq. 1. This discrepancy cannot be accounted for by assuming that the molecules are aggregated, since this assumption would affect the value of po almost as much as the values of p1 and 02. Thus, in the experiment with poly-soap R in 0.00391 JII KCI, the assumption of agpregation to the extent of an aggregate weight of 2,000,000 would increase POto 128, while increasing $1 to only 29.4. Unfortunately the value of p has no direct significance in terms of measurements other than by light scattering. The internal discrepancy between the values of p calculated from the same data by two different methods must cast some doubt on the validity of the theory on which these calculations were based. iVe also have attempted to interpret the present data according to the theory of Doty and Steiner.lg This theory, however, predicts a dissymmetry coefficient 245 < 1, while in all our experiments we have found z45between 1.0 and 1.2. Conclusions From the theoretical standpoint, it appears that the data obtained in aqueous salt solutions cannot be adequately interpreted by either of the theories considered; while the theories proposed for soap in the absence of salt do not permit direct experiment a1 verification. From a practical standpoint, the present data indicate that molecular weights of soap micelles, estimated from light scattering measurements in aqueous solutions, may be too low by a factor from 2 to 5 . Insofar as the present data can be applied to ordinary soaps and other micelle-forming substances, (19) P Doty and R F Steiner, J Chein P h y s , 20, 8 5 (1952)
[ C O N T R I B U T I O S FROM THE
DEPARTMEST OF
L ~ h l M O N I U hFLUORIDE l
287
the data obtained in dilute salt solutions indicate that in the presence of an electrolyte in the range 0.014.1 N , micelle molecular weights within 30% of the correct value will be obtained. Extrapolation to the critical micelle concentration, as recommended by Debye20and Mysels16should give molecular weights of this order of accuracy, with soaps of fairly high critical concentration such as sodium laurate or sodium lauryl sulfate. However, with soaps of critical concentration much below 0.01 M , this procedure may give somewhat more erroneous results . The data obtained in salt solutions of higher concentration (0.2-0.4 M ) indicate that association of the poly-soap does take place. It is likely that association or enlargement of ordinary soap micelles would also take place a t these salt concentrations. However, the addition of salt to bring the total electrolyte concentration (including unmicellized soap) up to 0.01-0.1 41would seem a safe practice. Any effects arising from hydrolysis of the polysoaps do not prevent transfer of the conclusions drawn from the poly-soap systems to ordinary soap systems, since hydrolysis should be comparable in both systems. Furthermore, the extent of hydrolysis must be comparatively small, since poly-soap solutions of 0.34-17, concentration (0.008-0.025 AI) have a pH of 9.5. The present data should thus have some bearing on the behavior of detergents such as sodium lauryl sulfate, etc. A sulfate ester type of poly-soap has been prepared in this Laboratory, but we have not had the opportunity to study its light scattering behavior. Acknowledgment.-\Ye wish to thank the General Latex and Chemical Corporation, of Cambridge, Massachusetts, for sponsorship of this work and for permission to publish this paper. The advice and encouragement given by Professor Paul Doty and Dr. Helga Doty, of Harvard, and by Professor J. P. Mason, of Boston University, are gratefully acknowledged. ( 2 0 ) P Debye, J Phys Collozd Chem , 53, 18 (1949), P Debye and E W Anacker, tbzd., 65, 6.14 (1951)
CAMBRIDGE, MASS.
CHEMISTRY O F THE UNIVERSITY O F M I C H I G A N ]
The Thermodynamic Properties of Ammonium Fluoride: Heat Capacity from 6 to 309OK., the Entropy, Enthalpy and Free Energy Function' BY EDWINBENJAMINS AND EDGAR F. WESTRUM, JR. RECEIVED SEPTEMBER 6, 1956 Measurements of the heat capacity of NHaF have been made by adiabatic calorimetry over the range 6 to 309°K. 1x1 contrast with the reports by others of a hump in the heat capacity near 24OoKc., no thermal anomaly was found in the entire range of measurement. The molal heat capacity and the derived molal entropy and enthalpy at 298.16°K. are 15.60 i 0.02 cal./deg., 17.20 f 0.02 cal./deg., and 2655 f 3 cal., respectively.
Introduction Although recent thermodynamic, spectroscopic and diffractional studies have contributed much to (1) Adapted from the dissertation of Edwin Benjamins submitted t o t h e Graduate School of the University of Michigan in partial fulBllment of the requirements of the Ph.D. degree. This work was SUPported by t h e U. S. Atomic Energy Commission.
the interpretation of the mechanism of the pseudorotational transitions in ammonium chloride, bromide and iodide, the treatment of the fluoride has been less decisive. Simon, Sinison and Ruhemann? determined the low temperature heat capacities (2) F. Simon, C1. von Simson and M. Ruhemann, 2.p h y s i k . Chem.
lZ9,339 (1917).
of NHlF and the other ammonium halides and reported thermal anomalies of varying magnitude in all of them in the vicinity of 245'K. Since the anomaly claimed near 241'K. in ammonium fluoride was rather smaller than that in the other halides, they reached no conclusion as to its cause. Shulvas-Sorokina and Evdokimov3 investigated tlie temperature dependence of the dielectric constant of ammonium chloride and ammonium fluoside and thereby confirmed the reported transformation in the chloride and claimed that the fluoride also undergoes a transformation a t 246'K. Some authors of recent papers, monographs and texts concur in the view that ammonium fluoride would not be expected to undergo a transformation similar to that of the other halides because of its relatively strong hydrogen bonding and the stability of its wurtzite-type lattice4; however, confusion continues to exist. In the other ammonium halides, each N H 4 +ion is surrounded by an octet of anions, and the phase transitions occur as a consequence of the disordering of the equilibrium orientation of the ammonium ions. A different explanation must be sought to explain a phase change in NH4F, if genuine, since only four tetrahedrally disposed nearest neighbors occur in its hexagonal ~ t r u c t u r e . ~Plumb and Hornig have noted that the infrared spectrum is "consistent with a phase change and may indicate a certain amount of disordering in the room-temperature phase.'I5 Their X-ray diffractional studies, however, showed a structure a t room temperature which is similar to that deduced from their infrared spectrum a t - 1%". The present work was undertaken as part of ail extensive study of the thermodynamic properties of hydrogen-bonded substances to ascertain the reality of the reported transition in pure arnmoniuni fluoride. Experimental Preparation of Anhydrous Ammonium Fluoride ( NHaF).A stoichionietric excess of 48y0 aqueous reagent hydrofluoric acid was added t o solid reagent ammonium carbonate in a large silver beaker, and the solution was boiled vigorously t o expel carbon dioxide. Gaseous ammonia was bubbled into the solution until P H 8 was achieved. Gradual cooling yielded acicular crystals of N H a F which were separated from the mother liquor by filtration through a polyethylene Biichner-type funnel and washed with methanol saturated with ammonia. To avoid the dissociation of wet ammonium fluoride into ammonia and ammonium nionohydrogen difluoride,6 the crystals were dried in a polyethylene train under infrared radiation by streaming anhydrous ammonia over them for more than 24 hours. T h e dried salt was transferred t o stoppered polyethylene bottles in a dry-box and stored under an atmosphere of anhydrous ammonia in a desiccator charged with anhydrous calcium oxide; i t was highly evacuated t o rcinovc excess ammonia prior t o weighing. Spectrochemical analysis revealed the absence of silicon :tiid other nictallic contaminants; less than 1 p.p.m. of silver IWS detected. \'olumetric analysis of t h e product using the IVillnrd and Winter method' for fluoride and t h e distillation procedure of a Kjeldahl nitrogen determination gave the following results: fluoride (as HF), 51.46 =t 0.15'%; animonia, 48.78 i. 0.lO"i;l (theoretical, 51.29 arid 48.7170,
(3) R . D. Shulvas-Sorokina a n d V . G . Evdokimov, Acta Physicochirn U R S S , 11, 291 (1939). ( 4 ) \ V Zachariasen, Z . p h y s i k . Chem., 12'7, 218 (1927). ( 5 ) R.C . Plumb and D . F. Hornig, J . Chem. Phys., 83, 947 (195.5). ( G ) A . C . Shead and G . F. Smith, THISJ O U R N A L , 63, 483 (1931). (7) IT. 11. Willard and 0. B. Winter, I t i d . E n g . Chem., A n a l . Ed., 6 , 7 (1033).
respectively). Titratioii of tlie frce acid usiug a ,611 tiictcr indicated less than 0.05'% excess H F . T h e absence of any abnormality in the heat capacity measurements near the ice point is probably evidence t h a t the sample was anhytlrous. -4lthough direct combiuations of anhydrous amnionia with anhydrous IIF and with cryst;tllitie KH4HF2were :tlw :tttempted, neither yielded a crystalline product :is suitable for thermodynamic ineasuremeut as the ;tqueous inethotl tlescribed. Cryostat and Calorimeter.--The Mark 1 liquitl Iieliutii cryostat employed in these meitsurements was sitni1:tr ti, oiic described by \Vestrum, Hatcher a n d Osborne6 antl is detailed elsc\vhcre.~ T h e adiabatic manner of operation miployed in making heat capacity meitsurements, the cletermiuatioii of tlie temperature scale above 10°K. by the N i t tional Bureau of Standards,'o the provisional scale einployetl below 10°l(., and confirmation of the accuracy of tlie calorimeter with benzoic acid also hnve beeii clescribetl .g T h e gold-plated copper calorimeter (laboratory designation W-6) is 3.8 cm. in diameter and 7.7 cm. in length. I t is quite sirniltir in general design t o one already desrribctl,ll except that to facilitate loading i t n i l sealing tliermally U I I stable coiiipountls, the tiinnieter of the copper cover was rctlucetl t o 1.I) cni. and removed from the main portion of the calorimeter b>-a short tube of hlonel of 0.28 mrn. tliickiiess which proviiies :i therrnal dam during the sealing opcration with Cerroseal (SOYo tin and 50';; indium) solder. This "cupola" docs not interferc with the establishment of thermal equilibrium in the cryostat, yet avoids appreciable lieatiiig of calorimeter :tiid sample on sealing :mil opcning the calorimeter. The thickness of the copper shell is 0.37 mm. A copper cone provides thermal contact with the adiabatic shield for cooling the calorimeter. Eight radial vanes of 0.062 mm. copper foil aid the rapid achievement of thermal equilibrium. Pure helium gas, at a pressure of oiie atiiiosphere a t 25', is used t o provide thermal conductivity between thc calorimeter and the sample. A small holc drilled in thc end of a thin-wall Monel stud permits evacuittion, filling with lieliuni itnd sealing with solder. A n axial re-entrant well contains a 25-ohm capsule-t~-l)e platiuuni resistance thermometer with a cylindrical copper heater sleeve wound bifilarly with 160 ohms of No. 40 B . & S. gagc double Fibrcglas-irisulateci constantan wire cementetl with b;tked Forinv;tr in conforming double thread grooves. T h e glass he:iti through which the thertnometer leads arc sealed is entirely within :in enlarged portion of the well. Lead wires t o both thermometer and heater a r e brought t o temperature equilibrium with the outside of the caloriineter by a small copper spool around which the leads are ceinented with Formvar. T h e spool is bolted in place over the end of the well with 00-90 brass cap screws. A copper-constantan differential thermocouple junction co:ttcd with baked Forrnvar makes a snug fit in a small tube soltiercti t o the side of the calorimeter. Lubriscal stopcock grease is used to establish thermal contact between the calorimeter, thermometer, heater, spool and thermocouple. The weights of grease and solder a r e maintailled c o n s t m t for the me;tsureinents on the empty antl loaded calorimeter. T h e mass of t h e crnpty calorimeter iricluding thermometer, hcxtcr, spoon solder and grease is 87.0 g. Two itidepcnclent sets of determinatioris on the heat c a p x i t y of the empty cdorilneter were made. H e a t Capacity Measurements.---.\ 26.71 I-g. smiplc (oacuo) of anhydrous anirnonium fluoritle w:ts weighetl iiito calorimeter \$:-E in :t dry-hos. Corrections for buciyancy were based on it crystal1ogr;tpliic density of 1 .(KW ~ . / c c . ' ~ The initial coolin of tlit s:irnple at an average rate o f 0 . 5 " / min. revealed no indication of a t1ierm:il ;~nom;ily. T h c subsequent heat capacity nic:isurenients o f Scrics 1 d s i i failed to reveal any t r x c of the previously reported Iiunip. After very slow cooling a t a11 average rate of O.O:3"/initi. the redeterminations of Series I1 confirmed the absence of ciay anomalous behavior. ( 8 ) E. F. Westrum, Jr., J. B. Hatcher and D W. Osborne, J . Chem. P h y s . , 21, 419 (1953). (9) E. F. Westrum, Jr.. a n d A. P.Beale, Jr., t o be published. (10) H. J. H o g e a n d F. G. Brickwedde, J . Reseavch S o l i . Bur. Standards, 22, 351 (1939). (11) D. W. Osborne and E. F. Westrum, J r . , J , Chew%.P h r s . , 21, 1884 (19%). (121 P. Wulff a n d 11. K. Cameron, Z.p l i y s z i , Cheiii.. A b t . E , 10,317 (1830).
THERMODYN.\MIC PROPERTIES OF AMMONIUM FLUORIDE
Jan. 20, 195T
Results and Discussion The individual results of the heat capacity determinations are presented in chronological order in Table 1. The approximate temperature increment of each measurement can be inferred from the adjacent mean temperature values. The heat capacities are computed on the basis of a molecular weight of 37.04 g., the defined thermochemical calorie equal to 4.1840 absolute joules, and an ice point of 273.16'K. MOLAL IlEAT
TABLE I CAPACITY O P AMMONIUhl FLUORIDE C" cal. deg.-l I
T , OK.
CD
cnl. deg.-'
(Series 1) 111.19 7.539 120.93 8 158 8.071 129.55 137.81 9.157 145.96 9.623 154.13 i n . 06 162.37 10.48 i n . go 171.18 180.18 11.33 189.28 11.73 198.55 12.12 206.70 12.4.3 215.25 12.80 223.77 13.12 13 45 232.64 241.54 18.8') 14.11 250.28 259.00 14,37 267.75 14.67 276.65 14.95 15.23 285.79 295.05 15.51 304.36 15.79
CD
T,OK. cal. deg.-' (Series 2) 54 87
62 69 76 84 93 102 111 119 230 238 247 255 263 271 270
07 27 69 64 38 34 01
33 44 97 34 57 67 64 50
3.487 4,009 4.559 5.124 5.730 6.348 6.950 7.531 8.061 13.36 13.68 13.97 14.25 14.51 14.79 1.5.02
T , OK.
(Series 3) 6.01 0.051 7.25 ,028 9.41 ,079 11.41 ,138 13.10 .210 14.38 .283 15.59 ,363 16.89 ,460 18.30 .555 19.99 ,688 22.07 .863 24.58 1.078 27.45 1.323 30.59 1,589 33 85 1.853 2.119 37.22 40.83 2.383 44.79 2.691 3.027 49.24 51.21 3.407
Corrections have been applied t o the heat capacity values for curvature (;.e., for the finite temperature increments employed in the measurements) and for the slight differences between the amounts of helium gas in the loaded and empty calorimeter. The precision of the data above 40'K. is of the order of 0.03% The molal thermodynamic functions Soand H" HOowere computed by numerical quadrature of C, versus log T and T, respectively; however, all values were obtained from a single plot of C, versus log T. These functions, together with the free energy function and values of the molal heat capacity read from the smooth curve through the experimental points, are tabulated a t selected temperatures in Table 11. These heat capacity values are considered to have a probable error of 0.1% above 25OK. ; a t 10'K. the probable error may be 1%, and a t 5OK. it may be 5% as a consequence of the decreased sensitivity of the resistance thermometer and the provisional nature of the temperature scale below 10'K. Entropies and enthalpies a t 10'K. are calculated by extrapolation of the heat capacity to O O K . using a T-cubed dependence and assuming Sooto be zero. The expectation of no residual (zero-point) en-
289
TABLE I1 MOLALTHERMODYNAMIC FUNCTIONS OF AMMONIUM FLUORIDE
T
02.
CP, cal. deg.-l
10 15 20 25 30 35 40 45 50
0.094 .323 ,689 1.114 1.540 1.942 2.318 2 707 3,080 3.848 4.616 5.374
GO 70 80 90 100 110 120 130 140 150 160 170 180 190 200 210 220 230 240 250 280 270 280 290 300 273.16 298.16
6.108 6.802
7.465 8.100 8.700 9.282 9.837 10.36 10.85
11.33 11.76 12.18 12.59 12.98 13.35 13.71 14.06 14.40 14.73 15.03 I5,3F 15.65 14.84 15.60 &
H o - Ho',
SO,
cal. de&-'
cal.
0.0312 .lo53 .2465 ,4456 .6866 0.9545 1,2381 1.5335 1,8383 2.4679 3.1188 3 7847 4.4803 5.1404 5.8199 8 , 4968 7.1691 7.8353 8.4948 9.1464 9,7893 10.423 11.047 11.661 12.265 12.880 13,446 14.021 14.588 15.146 15.696, 16.238 16.771 17.297 15.888 0.02 17.201 i 0.02
0.234 1.189 3.688 8.189 14.829 23.548 34.184 48.745 61.22 95.88 138.20 188.15 245.59 310.19 381.54 459.38 543.4 633,3 728.9 829.9 930.0 104G.9 1162 3 1282.0 1405.9
-(PO
-
Ho')/T, cal. deg.-l 0.0078 ,0280 ,0621 .1180 ,1923 ,2817 ,3835 ,4947 ,6138 ,8398 1.1445 1.4328 1.7313 2.0383 2.3514 2.6686
2.9890 3.3116 3.6353 3.9594 4.2830 4,607 4.930 5,251 5.570 5.888 6,204 6.518 6,829 7.138 7.44.5 7.730
1533.8 1665.4 1800.7 1939.6 2082.0 2227.0 2376.6 2528,A 8.052 2683.7 8,331 7.542 2274.4 2654.9 i 3 8.297 & 0.01
tropy in NH4F is confirmed by the determination of the equilibrium dissociation pressures over the system NH4F-NH4HF2together with a determination of the phase diagram of the binary system between these compositions and other related data.I3 A comparison of the data obtained in this study with those of Simon, Simson and Ruhemann2 is presented in Fig. 1. No evidence of a hump is pres-
I 10 200
220
240
T,
260
275
OK.
Fig. 1.-The molal heat capacity of ammoniun: fluoride near 240'K. The triangles indicate the data of Simon, Simson and Ruhemann; the open and solid circles represent, respectively, series I and I1 of this work. (12) E Benjamins and E.
F Westrum, J r , t o be puhlkhed
GEORGE -1. CASTELLION ASD 11.. - ~ L B E R TNOYES,J R .
290
ent in our measurements, and the heat capacity data show the typical sigmoid curve anticipated in the absence of anomalous behavior. The heat capacity hump and the dielectric constant anomaly previously reported by others may be attributable to the presence of water or other impurities in their preparations. This hypothesis is consistent with the preparative methods described, but cannot be tested further because of the scant analytical data provided. Although thermal analysis14 of the NH4F-H20 system indicated simple eutectic behavior with a eutectic temperature of 24GOK. involving the phases NHIF and NH4F-Hi0, more (14) V S E'atloi and E M Polyakova J Criz C h i n U S S R , 16, 714 (1945)
[ C O N T R I B U T I O N FROM THE
recent studies by Zaromb and Brill'; have indicated that solid solutions exist in the water-rich regioii of the system. No corresponding exarninatioii of the ammonium fluoride-rich region of the phnsi. diagram has been published. Acknowledgment.-The authors acknowledge with gratitude the partial financial support of the Division of Research of the United States A1toniic Energy Commission, and cooperation in the provision of liquid helium by Dr. D. I\'. Osborne of thc Chemistry Division of .irgonne Satioiial Lxboratory. (1.5) S . Zaromb and R . Brill, .I. Cheiii. Phvs., 24, 89.5 (I!l.j(;)
.-INS ARBOR,MICHIGAS
DEPARTMENT O F CHEMISTRY,
Photochemical Studies. LII.
VCJl. 79
UNIVERSITY O F ROCHESTER]
The Nitrous Oxide-Ethane System1
BY GEORGE4. C A S T E L L I O N
AND
1%'.
A L B E R T X O Y E S , JR.
RECEIVED AUGUST27, 1956 The ratio of nitrogen formed photochemically at about 1900 A. in pure nitrous oxide t o t h a t formed under identical corn tiitions in a nitrous oxide-ethane mixture is 1.4. This indicates t h a t the primary photochemical process is probably N 2 0 = Sz 0 and that oxygen atoms react much more rapidly with ethane than with nitrous oxide. Several products are forrnctl including ethylene, butane, carbon monoxide, hydrogen, methane and probably ethanol a n d acetaldehyde. A complete elucidation of the mechanism is not possible, due to the fact t h a t oxygen atoms seem t o react more rapidly with one or more o f the products than they do with ethane itself. More ethylene is formed than one would expect from the amount of butane. I h e to secondary reactions which must involve the initially formed products i t has not been possible to obtain precise inforinntion about the reactions of oxygen atoms and of hydroxyl radicals.
+
Experimental Introduction An aluminum spark, which gives several intense lines beThe absorption spectrum of nitrous oxide probably cponsists of a continuum extending from about tween 1850 and 1990 A., was used as a source of radiation. The general design of the spark was similar t o t h a t developed 3000 A. to shorter wave lengths.? by W i g and Kistiakowsky.s Tile transformer has a capacThe photochemistry has been studied with radia- ity of 10.4 KVA4with a 17600 v . secondary when operated tion directly absorbed by nitrous oxide3 as well as on 115 v. ilctually it was operated on 220 v. primary cirthrough mercury ~ensitization.~ The reaction cuit with 0.032 mfd. capacity in parallel with the spark. .lir played directly on the spark for cooling purposes and to products (at least of the unsensitized reaction) are jets increase steadiness of operation. nitrogen, oxygen and nitric oxide. Oxygen atoms The method of focal isolation was used to obtain approsimonochromatic radiation. The change in focnl must be formed in the primary p r o c e ~ s . ~ ~ , ~ , ~ mately ,~ One of the great gaps in the interpretation of length with wave length of a quartz lens 3 cm. in diameter with a focal length of about 18 cm. was used to isolate radiaoxidation reactions of organic molecules lies in the tion of the desired wave length. T h e center of the lens was behavior of hydroxyl radicals5 The present work blocked off. A camera iris was placed between the focus was started in the hope that oxygen atoms from and the lens and t h e size of opening adjusted to allow maid\nitrous oxide would react with hydrocarbons, such light of wave length below 2000 b. t o pass. The focus was as ethane, to give hydroxyl radicals and that the located either by a piece of canary glass (Corning No. 9750) spectrobehavior of the latter could be investigated. This or by a piece of glass "silvered" with mercury. showed t h a t the 1850 k . line of aluniinun~did enter hope has not been realized because oxygen atoms gram the cell. also react rapidly with some of the reaction prodT h e portion of the glass vacuum system of the reaction ucts. Thus the over-all reaction is complex but cell and the nitrous oxide and ethane purification, storage and measuring system were kept as free from mercury as certain conclusions about it are possible. (I) This work c,f
was supported in p a r t by Contract with t h e Office N a v a l Research, United States S a r y , and with the Office of Air
I
