104
GERARD W. ELVERUM, JR.,AND DAVID M. MASON
MELTING POINT MEASUREMENTS OF THE SYSTEM
Vol. 60 HN03-N204-H201
BYGERARD W. ELVERUM, JR.,AND DAVIDM. MASON^ Jet Propulsion Laboratory, Pasadena, California Received Julu 6 , 1966
The solid-liquid equilibrium phase behavior of the HNOa-N204 system was determined from melting point measurements near atmos heric pressure over the range of composition 0 to 100 weight % N2O4. The existence of a solid com ound is indicated whicl! consists of 2 moles of "03 and 1 mole of NzO,. The solid-liquid hase behavior for the. HN88-N~0,-H~0 system was similarly measured over the range of composition of 0 to 25 weight % g204 and of 0 to 18 weight % HzO. The minimum melting points in the composition range studied for the HNOs-N2O4 system were -65" at the eutectic mixture of 25.6 weight % NzOc and for HNO~-N~O~-HZO system about -78" at the eutectic mixture of 18 weight % N?OIand 4 weight % HzO. The interpretation of these data in the light of chemicalequilibria which occur in nitric acid solutions is hscussed.
I. Introduction In connection with a program a t the Jet Propulsion Laboratory for the study of several of the physical and chemical properties of fuming nitric acid,8 the equilibrium melting points of the system HNOa-N204-H20 have been measured. Data on the solid-liquid and liquid-liquid phase relations of the binary system HNOsN204 were published by Pascal and Garnier4 in 1919. Based on their interpretation of cooling curve thermal data, these data indicated that a eutectic mixture containing 18 weight % N204 froze at a temperature of -73", and that a solid compound corresponding to the formula N204.NzO5.H2O should exist at 42.2 weight yo Nz04. Measurements of the liquid-solid phase equilibria in the system HPT.'33-H2O were made by Kuster and Kremann.5 These data have been used for the HN03-H20 boundary of the ternary diagram given in this paper. The freezing point of pure HN03has been found by Dunning and NuttB to be -41.62 f 0.05' based on the maximum freezing point obtained a t 50 mole % N206 for the binary system N205-H20. Forsythe and Giauque' report a value of -41.59". The value found in the present study is -41.7'. The work of Pascal and Garnier4indicates that a two-liquid phase region exists in the binary system HN03-N204 above 48 weight yo N204. This region was studied in detail by Klemenc and Spiesssa and by Corcoran, et a1.,8b and both sets of data are reproduced in the present paper to show the complete two-phase data for the HN03-Ne04 system. The freezing point of pure N204 has been re(1) Thia paper presents the reaults of one phase of research carried out a t the Jet Propulsion Laboratory, California Institute of Teohnology, under Contract No. DA-04-495-0rd 18, sponsored by the Department of the Army, Ordnance Corps. (2) Department of Chemistry and Chemical Engineering, Stanford University, Stanford, California. (3) The term fuming nitric acid refers to the ternary system " O r NtOrHzO in the range of composition rich in "01. Compositions in general are expressed on a formal baais unless otherwise stated. Le., in terms of the formula of a compound disregarding the molecular species that may result in its solution. (4) P. Pascal and. Gamier, Bull. doc. chim. France, 85, 309 (1919). (5) F. W. Kuster and R. Kremann, Z . anarg. allgem. Chem., 41, 1 (1904). (6) W. J. Dunning and C. W. Nutt. Trans. Furoday Soc., 47, 15 (1951). (7) W. R. Forsythe and W. F. Giauque, J . Am. Chem. SOC.,64, 48 (1942). (8) (a) A. Klemenc and Th. Spiess, Monotsh., 77, 216 (1947); (b) W. H. Corcoran, H. H . Reamer and B. H. Sage, Ind. Eno. Chem., 46, 2541 (1954).
ported by Whittaker, et al.,B 9s -11.23'. Giauque and Kemp'O give - 11.20' as the equilibrium melting point, and the value of -11.2' from the present investigation agrees with these values.
11. Description of Apparatus and Procedures The apparatus used in the determination of the melting points consisted of a vacuum jacketed glass tube. Temperatures were measured by means of a calibrated copperconstantan thermocouple, one end of which was located in the tip of a thin glass well entering the bottom of the glass tube. Equilibrium in the solutions was facilitated with agitation by means of a coiled-glass stirrer which moved up and down. Its displacement was such that the stirrer agitated the contents of the tube including the region adjacent t o the thermocouple well a t the bottom of its stroke. The stirrer was connected to a glass-enclosed steel rod, the upper end of which was attached t o a stainless-steel spring. The rod was actuated by a solenoid the current of which was controlled by a motor-driven rheostat. The apparatus was flushed with dry nitrogen, and solutions were added through a small opening in the top of the tube by means of a hypodermic syringe to which a long stainless-steel capillary tube was attached. Several series of measurements were made by successively adding liquid N204t o a weighed amount of HNOs in the melting point ap aratus by means of a calibrated syringe assembly. H28from an ice-bath was circulatedo through the syringe jacket to maintain the Ns04 near 0 . The density of NzOl at this temperature was determined at this Laboratory to be 1.489 ./cc. This syringe was similarly used to add water to H#O, solutions containing known amounts of NzO~. Another series of measurements was obtained by making up individual solutions by weight. Because of the extensive tendency-of HNOt solutions to supercool, the method of measunng temperature vs. time for the melting process instead of the freezing process was used t o establish the phase equilibria. In some measurements in which there was only a small change in solubility of the solid phase in the liquid phase over a relatively large temperature interval, the melting point was determined by observing under magnification the disap earance of the last few crystals and the reappearance of tge f i s t few crystals while the temperature of the outer bath was allowed to change very slowly.
111. Materials The " 0 8 was prepared by vacuum distillation at room temperature of a mixture of reagent-grade concentrated sulfuric acid and potassium nitrate. The colorless product was collected in a container at about -70" and stored at about -20". HzO used was distilled in the laboratory. Commercial NzO4was purified by bubbling oxygen through it and passing the vapors through a furnace at 300" t o oxidize any NO present and decompose any N,Os. The remainder of the purification train was .essentially that described in reference ll. The NzOcobtained froze to a colorless solid at -11.2'. (9) A. G. Whittaker, R. W. Sprague. A. Skolnik and G. B. L. Smith, J . A m . Chem. SOC.,74, 4794 (1952). (10) W. F. Giauque and J. D. Kemp, J . C h e m Phys., 6, 40 (1938). (11) A. KIemen, "Die Behandlung und Reindarabllungvon Gmm," Edwards Brothers Inc.. Ann Arbor, Michigan, 1943.
MELTINGPOINT MEASUREMENTS OF SYSTEMHN02-N204-H20
.Jan., 1950
105
IV. Results The complete phase diagram showing the solidliquid phase behavior a t 1 atmosphere pressure and liquid-liquid phase behavior near bubble point for the binary system HNO$-N204is presented in Fig. 1. The experimental data used t o construct the 1
LIQUID
so
_. --Ref
0
5
MOLE Yo N 2 0 4 .
Fig. 2.
MOLL % N20,.
Fig. 1.
solid-liquid phase portion of the diagram are given in Table I. This binary system exhibits two eutectic points, one at 25.6 weight % Nz04and -65', TABLE I MELTINQPOINTSOF HNOs-N2O4SYSTEM Nz04,mole %
M.P.. 'C.
N 1 0 4 , mole %
0 3.27 4.55 6.31 6.66 10.61 10.91 12.77 14.78 16.67 17.20 18.73 20.21 21.10 22.93 24.40 25.22 25.35 26.02 26.46 26.68
-41.7 -43.1 -43.8 -45.2 -45.4 -49.5 -49.9 -52.4 -55.6 -59.3 -60.2 -64.5 -59.8 -55.4 -51.9 -50.0 -49.5 -49.2 -48.8 -48.4 -47.7
27.46 29.21 29.83 31.55 32.24 33.19 33.53 33.93 35.48 35.72 36.33 38.89 41.21 42.10 50.67 94.18 96.56 97.53 98.45 99.19 100.00
M . p . , OC.
-47.1 -45.7 -45.5 -45.4 -39.0 -45.3 -37.1 -45.6 -37.6 -37.6 -34.0 -26.9 -18.6 -16.4 -12.6 -12.7 -12.5 -12.2 -12.0 -11.7 -11.2
and the other at 43 weigb.t yo N204and -45.7". I n Fig. 2, which gives in detail the phase diagram for the region between 93 and 100 mole yoNz04,it can be seen that an invariant point at 1 atmosphere exists at a temperature of - 12.7' and a concentration of 95.0mole yoN20a. At this point two liquid
phases are in equilibrium with solid N204. The addition of HN03 above 5.0 mole % a t the temperature -12.7' decreases the amount of solid N204until a t approximately 44.2 mole % N204 the concentration of the second liquid phase of the invariant system is reached (Fig. l), and further addition of "03 results in a single liquid system. The letters A and I3 are used in Figs. 1 and 2 to de3ignate the homogeneous liquids on the "03rich and N204-richsides, respectively, of the critical solution point near 62 mole % N204 and 61 '. The existence of a solid compound consisting of 2 moles of HN03 and 1 mole of N204as was indicated by the original data of Pascal and Garnier3 is definitely established (Fig. 1). A slight break i n the curve at 26.3 mole % N204and -48.7' seems to indicate the possibility of two crystalline forms for this 2:l compound which are designated by a and p. The possibility that this break is indicative of a compound having a 3 : 1 mole ratio of "03 to Nz04is precluded by the fact that -65' eutectic points were measured at concentrations &s high as 29.21 mole % N204 (Fig. 1). The two points shown by squares in Fig. 1 are included since a check of the data showed no detectable error in either their compositions or their melting points. These points suggest the possibility of the existence of a metastable condition for the compound N204. 2HN03. A portion of the HN03-N204-Hz0 system showing curves of constant melting point is presented in Fig. 3. A ternary eutectic point is indicated in the region around 18 weight % N204and 4 weight yo H2O and at a temperature near -78'. The original data from which Fig. 3 was constructed are given in Table 11.
V. Evaluation The initial portions of the phase diagrams for the systems HN03-KN03 and HN03-NH4N0a6and the
GERARD W. ELVERUM, JR.,
IO6
AND
DAVID M. MASON
48LI17-m7 -
VOl. 00
4
0
-45
2, ‘-50
w
a
3
ta L w -55 W
c
MOLE
Yo N204, ns ( m o l e fraction solute).
Fig. 3.
region of the phase diagram for HNOrN204 in which the solid phase is pure HN03 are shown in Fig. 4. Also shown are three curves based on the thermodynamic equation’2 dlua __ dT
- AH,,, -. - R2’2
(1)
where AHfuis the molar heat of fusion of the pure solvent (HN03) at the temperature T, and a is the activity of the solvent. For very low concentrations of solute, the heat of fusion may be taken TABLE I1 MELTING 1’OINTS N?O4, mole %
4.55 4.15 3.82 3.52 3.23 2.94 2.50 10.61 9.51 8.63 7.62 6.42
OF HNOI-NIOI-II?~SYSTEM H?O, mole To M.p., “C.
0 8.40 16.05 22.50 28.93 35.37 44.87 0 10.31 18.66 28.23 39.48
-43.8 -48.2 -57.6 -68.7 -54.0 -44.8 -30.2 -49.5 -60.4 -69.4 -50.5 -41.3
as a coiistaiit, but at higher concentrations one must take into account the change in the molar heat of fusion of the pure solvent due to the change in melting temperature and also to the partial molar heat of solution of the solvent in the solution at the concentration of interest; e.g. AHru = AHoru + AC, ( T - To)+ AH8 (2) where AHoruis the heat of fusion of pure solvent a t its melting point T O ;AC,, is the difference in molar (12) S. Glasstone, “Thermodynarnico for CLomiste,” D. Van Nostrand Co., New York, N. y., 1017.
Fig. 4.
isobaric heat capacity between the liquid and the solid solvent; and AHs is the partial molar heat of solution of the solvent in the solution a t the melting point T. Over temperature intervals of the order of 30°, ACp may be taken as constant. Values for AHs were calculated from data reported in reference 13 which gives the heats of solution of Nz04 in HN03 at 0”. Since the values of the heat capacities of the solutions13 were nearly equal t o the heat capacity of pure HNOI the heats of solution at the melting point were taken to be equal to those reported at, 0” for making the correction in equation 2. Values of AHBcould be compared with values of the melting point T obtained from the phase diagram of HN03-N204at the corresponding concentrations. This comparison showed that up t o about 15 mole % N2O4the heat of solution could be represented fairly well by AH0
N
- K ( T - To)’
(3)
where K = 1 cal./deg.2 mole. Therefore, if - T’ = T - TO,Equation 1 may be expressed as d In_ a = s r . + ACp (-T’)-_ dT’
Integrating, expanding, and letting y
+ ...
(4)
R(To - 2”)’
=
T’/TO
(5 )
(1s) D. M. Mason, K. Boornan and G. W.Elverurn. Jr., “Heats of Solution of Nitrogen Dioxide in the Nitria Acid-Nitrogen DIoxide System a t 0’” (Progress Report No. 20-216,Jet Propulsion Labor& tory, Pasadena, lQG4).
#Jan., 1956
MELTING POINTMEASUREMENTS OF SYSTEM HNOZ-NzO4-H2O
If one now makes the assumption that the activity of the HN03 species is equal to the actual mole fraction of the HN03 species existing in the solution (ie., the solution behaves in accordance with Raoult's law), then one may replace a by the mole fraction of the species "01 (ie., 1 X,), where X . is the effective mole fraction of solute due to total number of solute species resulting from solution. In Fig. 4, n, represents the formal mole fraction of solute as N204, KNO3 or NH4N03. A series of values of X , for these nitric acid systems may now be found from the equation
107
tion of HNO, of 4 mole %, an apparent value of CY to be expected at 1 mole % of added NO3- would be about 0.1. At higher concentrations of solute, the freezing point curve becomes parallel with the line a = 1, p = 1, k = 2, indicating that the selfionization has been suppressed, that 2 molecules of HN03 are associated with one of the ions of KNOa or NH4N03, and that this associatiori is nearly complete. Raman spectral studies have also shown that the following equilibrium association esistsI6 2"Os
+ N0a-f-
NOa-.(HNOs)z
(9)
and the spectral data indicate that this equilibrium lies quite far to the right, a fact which is consistent with p = 1 and k = 2 in Equation (7). where a may be defined as the fraction of total soThe initial phase behavior of the system NzO4lute molecules ionized, j as the number of ions HN03 can be seen to lie along the curve CY = 0, formed per solute molecule ionized, k as the num- p = 0, k = 0 out to about 0.025 mole fraction ber of solvent molecules associated with one of the solute. The phase diagram shows that beyond a solute ions, and p as the fraction of one of the few mole per cent. the curve rapidly bends down to ionic species actually associated with solvent niole- parallel the curve a = 1, p = 1, k = 2, supporting cules. For the solutes KW03, NH4N03 and N204, the known fact that N204 becomes ionized in soluthe value of j is 2, and equation 6 reduces to tion according to the equilibrium N~04
From equations 5 and 7 a series of curves may be obtained of T us. n,. for various values of a, B and k. Three of these curves are presented in Fig. 4. The value of AHoiu for nitric acid was taken to be 2503 cal./mole, and the values used for the heat capacities of solid and liquid nitric acid were 15.82 and 26.70 cal./mole, re~pectively.~ It can be seen that both KNOa and ",NO3 give the same curve for the freezing point depression. Inasmuch as both KN03 and NH4N03may be expected to be completely ionized at low concentrations in HN03, one would expect (assuming no association) that these data would lie along the curve for a = 1, p = 0, k = 0. This curve is not shown in Fig. 4, but at low concentrations it would lie very close to the CY = 1, p = 1, k = 2 :urve shown. However, it can be seen that the actual freezing point depression approaches tangen,ially the curve CY = 0, p = 0, IC = 0 for small contentrations of solute. The explanation for this )ehavior is to be found in the measurements of the taman spectra of "03, l4 which indicate that quid HNOI undergoes self-ionization to the extent f about 3 to 5 mole yoaccording to the expression 2HN03
NOz+
+ NO,- + HzO
(8)
reasurements of the specific conductance also (pport the existence of this ionization.16 It can seen that initially the addition of NO,- would ift this equilibrium to the HN03 side of equation IO that the net number of particles of the group 12+, NOS- and HzO would increase very slowly first. The initial freezing point depression uld thus essentially be due to the ions K+ or 14+and would therefore lie along the line a = 0, = 0, k = 0, as shown. Assuming a self-ioniza)
E. D. Hughes, C. K. Ingold and R. I. Reed, J . Chsm. Soe., 2400
). G . D. Robertson, D. M. Mason and B. H. Sage, "Electrolytic .uctance of the Ternary System of Nitric Acid-Nitrogen Dioxideat 32'F. and Atmospheric Pressure" (Progress Report No. 20let Propulsion Laboratory, Pasadena, 1Y51)4
+ Nos-
(10)
The existence of this equilibrium has been shown by Raman spectral studies of the NZO4-HNO3system" and is supported by electrical conductance measurements.16 The existence of the solid compound Nz04-2HN03also gives evidence that association takes place in solution according to equation 9. The fact that the Nz04curve initially follows the curve a = 0, p = 0, k = 0 may thus be explained by the suppression of the self-ionization of "03 by Nos- ions. It may be noted that approximately twice as many moles of Nz04 as KN03 or NH4NOaare required to suppress the selfionization of "03. This result would indicate that N204 was only about 50% ionized in solution. It can be seen, however, that the curve at higher concentrations does not follow parallel to the curve a = 0.5, p = 1, k = 2 (as it should for 50% ionization) but that it becomes parallel to the curve a = 1, /3 = 1, k = 2, indicating nearly 100% ionization. This situation suggests that the ionization of Nz04decreases as the solution becomes more dilute, a result which would not normally be expected. If however it is assumed that the ionization of Nz04 according to equation 10 is not complete compared with KN03 ,which may be assumed to be completely ionized, then it follows that the n'oa- ions from the self-ionization of nitric acidlo shift the equilibrium of equation 10 to the N204 side at very low concentrations of N204. An ionization of the order of 90% at a mole fraction of Nz04 of 0.1 would be reduced to an ionization of about 5070 at a mole fraction of Nz04of 0.01. Thus, although the N204 curve nearly follows the a = l, p = 1, k = 2 slope at high concentrations to the eutectic point, it requires a larger amount of Nz04 t>han of KNOa, to suppress the self-ionization of HN03. The initial slope of a = 0, p = 0, k = 0 also precludes the equilibria NzO,
j )
A r
NO+
2N02
J. Chedin and S. FBnBant, Compl. rend.,
(11)
998, 242 (1949). (17) J. D. 8. Coulden and D. J . Milleo, J . Chem. SOC..2620 (1860).
(16)
108
RUSSELLK. EDWARDS AND JAMES H. DOWNING
from lying more than a few per cent. to the right axccpt i n extremely dilute solutions. A dashed curve based on equations 5 and 7 is shown in Fig. 2. The value usedI0 for the heat of fusion of N204 was 3502 cal./mole (as N204). It can be seen that the depression of the melting point initially follows fairly close to of N204 by " 0 3
Vol. 60
the curve for CY = 0, p = 0, 2G = 0. The behavior of the experimental curve, however, shows a positive deviation from Raoult's law a t higher concentrations of HN03. The association of the polar "Os molecules at higher concentrations which finally result in the separation of a second liquid phase accounts for this behavior.
THE THERMODYNAMICS OF THE LIQUID SOLUTIONS IN THE TRIAD Cu-Ag-Au. I. THE Cu-Ag SYSTEM' BY RUSSELLK. EDWARDS AND JAMES H. DOWNING Contributionfrom the Department of Chemistry, Illinois Institute of Technology, Chicago, Ill. Rscdved July 7, 1966
The thermodynamics of the liquid Cu-Ag s stem have been investigated as part of a eneral study in the Cu-Ag-Au triad to consider the energetic relationships andTchemica1bonding among these elements. '$he study was conducted by the method of determination of the partial pressures over the liquid solutions as a function of composition and temperature. Partial pressures were measured by the molecular effusion technique and were related to the vapor pressures of the pure liquids similarly measured, and activities were calculated. The related thermodynamic properties for the mean temperature 1428'K. are reported, based on the temperature and composition dependencies of the activity data. The activities of both components demonstrate marked positive deviation from ideal solution behavior in the system, as might have been expected in view of the fact that a wide miscibility gap exists in the solid state for this system. Large positive values for partial and integral enthalpies of mixing were found. The partial and integral excess entropies of mixing are positive.
Introduction The investigation in the liquid Cu-Ag system was chosen to institute a general thermodynamic study in the triad, Cu-Ag-Au, to consider the energetic relationships and chemical bonding among these elements. The atomic radius for Cu is about 12% less than that of Ag, and the atomic radius for the latter is practically identical to that of Au. On the other hand, the cohesive energy of Cu is about 17% greater than that of Ag whereas the value for Au is about 27% greater than that of Ag. An investigation of the binary permutations among the three elements offers an opportunity of observing the effects of the major variables-cohesive energy and atomic radius-while other variables, in particular valence, can reasonably be expected to remain relatively constant. The study was conducted by the method of determination of the partial pressures over the liquid solutions as a function of composition and temperature. Partial pressures were measured by the molecular effusion technique and were related to the vapor pressures of the pure liquids similarly measured, to obtain self-consistent data from which activities were then calculated. It was imperative that activities be determined only from self-consistent data obtained by relative measurements under identical conditions since reported2 absolute vapor pressure data by various different investigators have varied by as much as 100%. The thermodynamic properties for the liquid solutions were calculated from the temperature and composition dependencies of the activity data. In(1) (a) Presented a t the 126th Meeting of the American Chemical Society. New York, September, 1954. (b) Based on part of a thesis by J. H. Downing, submitted to the Illinois Institute of Technology in partial fulfillment of the requirements for the Ph.D. degree, May, 1954. (c) This work waa supported by the U. 8. O5ce of Naval Research, U. 8. Navy, through Contract N7-onr-329, Task Order 11, and Contract NONR 1406, Task Order 11. (2) H. N. Hersh, J . Am. Chrm. Boc., 71, 1529 (1953).
asmuch as the vapor pressure of Cu in the temperature range studied (1300 to 1700OK.) is approximately a factor of ten lower than that of Ag, Cu was a minor but not negligible constituent in most of the effusates. Consequently the activities of Ag were subject t o more accurate direct determination than were the activities of Cu, and the best values of the activities of Cu were taken from GibbsDuhem integration of the Ag activity data. However, in addition, directly obtained Cu activity data were evaluated in order to check the validity of the Gibbs-Duhem integration.
Experimental Apparatus.-The vacuum apparatus used for the effusion measurements is shown in Fig. 1. The essential features of the apparatus are (a) an evacuated porcelain tube section heated by, (b) a furnace, shown schematically, havin Globar elements, (c) a graphite crucible wit,h a small elusion orifice, (d) a thermocouple, enclosed in a silica glass protection tube, for measuring the temperature in the region of the crucible and (e) a water-cooled cold h g e r which served to collect samples of the Cu-Ag gas mixture which effused from the crucible as a molecular beam. The crucible was detachable through a polished sliding fit from the gra hite support rod shown in the figure. The support rod itsel! was capable of being moved as desired while under vacuum by means of an external electromagnetic coil acting on an iron section a f i e d to the end of the rod. Vacuum pum ing from both ends of the apparatus was used to ensure a &@speed system. The mercury diffusion pumps used were isolated from the system by liquid nitrogen traps and so operated as to rohibit exposure of the system to any mercury vapor. eacuum pressures were measured by means of a Pirani gage mounted just outside the main vacuum chamber along one of the 20 mm. tubing sections and also by means of an ionization gage similarly mounted along the other 20 mm. tubing section. Dynamic vacuum pressures of less than mm. were maintained during the vapor pressure measurements. Furnace Tern erature Control.-The furnace temperature was controfled by means of a Micromax automatic potentiometer with recorder and controller, operating on the signal from a Pt, Pt-Rh (10%) thermocouple located in the vicinity of the furnace heater elements. In general this arrangement could hold the temperature of the heater element environment to f 5 O , and the t.emperature of the
