THERMODYNAMICS OF SODIUM CHLORIDE SOLUTIONS
Dec., 1934
nitrate with nitron.' Fluorine was determined by precipitation as lead chlorofluoride. In order that the percentage of both nitrogen and fluorine could be calculated, the vapor density of the gas was taken as 82.0. The results are shown in Table 11. The letter A represents the number of gram atoms of nitrogen in 22.4 liters of the gas at standard conditions, and B is the corresponding quantity of fluorine.
N >%
16.8 17.0
TABLE I1 ANALYSISOF THE GAS F, % A 23.15 0.984 23.13 ,996
B
0.999 .998
2637
The selection of a good name for this compound is not an easy task. Most of the chemists who have been consulted agree that a correct but not very descriptive name would be nitrogen trioxyfluoride. Two substances which are related to NOaF; namely, nitrosyl fluoride6 and nitryl fluoride),t have been known for some time. The author wishes to thank Professor W. C. Schumb and Dr. N. A. Milas for their interest and helpful advice.
Summary and Conclusions The reaction of fluorine with moderately dilute In addition to one atom of nitrogen and one of nitric acid produces a gaseous compound which fluorine, found by the analyses, each molecule of has the formula NOSF. When heated, the gas the gas must have contained three oxygen atoms. This conclusion may be reached by considering explodes, but it may be handled without difficulty the properties of the possible compounds of the at room temperatures. The substance has an irritating odor, is colorless and boils a t about elements which could have been present: nitrogen, -42'. It is moderately soluble in water, with fluorine, oxygen, hydrogen. A reaction such as it reacts slowly, liberating oxygen. With which that observed in the case of sodium hydroxide solutions of potassium iodide and of potassium could have resulted only from the action of N03F hydroxide, it reacts according to the equations or one of its hydrates. However, the vapor NOsF 31-+ IsFNO*(1) density of the gas, about 82.0, permits one to NOaF 20H-41 / 2 0 2 F' NOSH20 (2) decide that the substance was not hydrated and (5) G. Gore, J . Chem. Soc., 92,391 (1869). that its formula was NOaF, a compound of which (6) H.Moissan and P. Leheau, Comfit. rend., 140, 1576 and 1621 the molecular weight is 81.0. (1906).
+
+ + + +
CAMBRIDGE,' MASS
(4) Busch, Bsr., 38, 861 (1905).
[CONTRIBUTION
+
292 FROM TIiE DEPARTMENT OF CHEMISTRY, UNXVERSITY
+
RECEIWD OCTOBER 13, 1934
OF PITTSBURGH]
Integral Heats of Dilution, Relative Partial Molal Heat Contents and Heat Capacities of Dilute Aqueous Sodium Chloride Solutions1 BY E. A. GULBRANSEN AND A. L. ROBINSON I. Introduction The thermochemical properties of aqueous sodium chloride solutions have been studied extensively recently2 over rather wide temperature and concentration ranges but except a t 250a there are no reliable measurements of heat contents or heat capacities a t concentrations below 0.05 m. It is in this dilute region that measurements are particularly important for extrapolation to infinite dilution and for comparison with the predictions of current theories of strong electrolytes. In this paper are reported meas(1) From the thesis of E. A. Gulbransen presented to the Graduate Schoot of the University in partial fulfilment of the requirements for the degree of Doctor of Philusuphy, Augnst, 1934 (2) Randall and Bisson, Tms JOURNAL,42, 347 (1920) : Richards and Gueker, ;&id., 61,712 (1929); end other references. (3) Robinson, ibid., 64, 1311 (1932).
urements of intermediate heats of dilution of aqueous sodium chloride solutions at 25, 20, 15 and 10' from 0.8 m (0.4 m a t 15 and 10") to 0.00016 m (0.00032 m a t 10'). From these rerelative sults integral heats of dilution (AH,,,),* partial molal heat contents (El and and relative partial molal heat capacities (efi, have been calculated for the temperature and concentration ranges of the measurements.
11. Apparatus and Experimental Procedure A differential adiabatic calorimeter similar to the one
-
-
The integral heat of dilution is the heat absorbed when a solution containing one mole of solute is diluted t o m 0. AH, -niZi 1:(nl is the number of moles of water associated with one mole of solute). ( b ) Where possible the nomenclature of Lewis and Randall, "Thermodynamics," McGrgw-Hill Book Co., New York City, 1928, is used, (4)
-
E. A. GULBRANSEN AND -4. L. ROBINSON
2638
developed by Lange and his co-workers' was employed. The following details indicate the principal changes (mostly minor) incorporated into the present apparatus. The thermopile, built according t o the design of Lange and Streeck,'" consists of 1330 iron-constantan thermocouples and has a total resistance of 70 ohms. It is connected directly to a Leeds and Northrup narrow coil type HS galvanometer which has a sensitivity of 0.089 microvolt per mm. deflection a t a scale distance of one meter. With a scale distance of 6 meters the maximum sensitivity used in this work was equivalent to 0.0003 calorie per mm. deflection corresponding to a temperature change of about 2.5 X 10-7'. For most of the measurements a sensitivity of about one-fourth of this value was sufficient and was obtained by use of a shunt of proper resistance. To obtain a continuous reading of the galvanometer before, during and after a dilution a semi-automatic recording device6' driven by an electric clock motor was used. The calorimeter is a n unsilvered Dewar flask of about 2.5 liters capacity. Each half of the calorimeter contained 1000 cc. of water or solution, depending on whether a first or second dilution was being carried out. The pipets, holding the solution to be diluted, have volumes of 25.67 * 0.02 cc. The heating elements, used for compensating the heat effect of a dilution and for electrical calibration of the calorimeter, are of cupron wire enclosed in glass capillaries a few hundredths of a millimeter larger in internal diameter f h a n the wire. The resistances of the two heaters are matched and have remained constant to within 0.02% over a period of months. To switch .the heating
Vol. 56
calorimeter, are driven by a governed motor of the universal type through a chain drive. The heat conductivity constant of the calorimeter is 0.006 minute-'. The temperature difference between the calorimeter contents and the outside bath is measured by a 24-junction copper-advance thermopile. The adiabatic control is constant t o *O.OO0lo. Since the two halves of the calorimeter are very nearly symmetrical (to within 1 or 2%) disturbances due to a difference in temperature between the calorimeter and the outside bath are negligible. For most of the dilutions the temperature changes produced in the calorimeter were less than lo-*'; the method of measurement is practically "isothermaladiabatic" calorimetry. The toluene-mercury regulator for the outer bath is constructed of soft copper tubing coiled into a form convenient for the dimensions of the bath.' The tubing has a volume of 1470 cc. and a surface area of 5720 sq. cm. To hold the temperature of the bath constant (for runs near room temperature) the heat generated by stirring is balanced by a fine stream of cold water controlled by electromagnets activated by the regulator or (for runs above or below room temperature) by fine streams of cold or hot water. The regulator responds t o a temperature change of 5 X within three seconds. The experimental procedure and method of determining the value of the heat effect are the same as previously deMallinckrodt c. P. sodium chloride was recrystallized three times and dried for twenty-four hours at 220'.
111. Results
TABLE I HEATSOF DILUTIONOF AQUEOUSSODIUM CELOMDESOLUTIONS -25-' mi
0.816 .SI6 .404 .404 .2 .2
mt
0.02025 .0425 .01025 .02025 ,00512 .0101 .1 .00257 .I .00507 .05 .001285 .05 .002535 .025 ,000642 ,025 .00128 ,0125 .000322 .0125 .000634 .00625 .000161 ,00625 .000318
-000-
Runs
3 3 4 4 5 5 7 7 10
9 13 10
5 5 6 6
7 - 1 0 0 -
-'51-
Cal./mole NaCl
+39.6 *O. 7 $54.0*0.5 -30.6*0.8 -18.1*0.4 -57.1 " 0 . 5 -48.8 * 0 . 6 -65.1 * 0 . 3 -55.9 * O . 4 -56.2 * O . 8 -51.1 * 0 . 6 -46.4 * 1 . 4 -41.1 * l . 5 -34.6 * 1.0 -30.8 *O. 7 -27.6 *I. 6 -24.6 f 1 . 6
2 2 2 2 2 2 4 4 2 2 2 2 2 2 4 4
f 7 3 . 7 *O .6 f86. 8 * O . 4
-6.45*0.1 f 6 . 2 *0.1 -41.8*0.1 -32.7*0.1 -50.3*0.1 -43.5 * 0 . 7 -49.2*0.1 -43.9 '0.9 -45.9 * O . 5 -39.9 * O . 9 -32.2*0.1 -30.4*0.6 -25.3k0.6 -20.2 * O . 3
current into the heaters, and t o control accurately the time of heating by means of an electric clock, a special electrical circuit" is used by means of which the time of heating can be determined to 0.01 second. The mirror image stirrers, one in each half of the (6) (a) Lange and Monheim. 2. fihysik. Chem., 1498, 61 (1930); (b) Lange a n d Robinson, Chcm. Rm.,9, 89 (1931); (c) Lange and Streeck, 2.physik. Chcm., l P I A , 1 (1931).
2 2 2 2 2 2 3 3 2 2 4 4 3 2
-25.5 "0.0 -34.4 1 0 . 2 -20.7 * O . 4 -14.2 * O . 2 -40.3 * O . 3 -32.7*0.1 -39.1 *O. 5 -34.7 *O. 9 -38.0*0.2 -33.1 * O . 1 -32.5*0.4 -29.0 *O. 4 -23.7 *0.8 -20.2 *o. 2
2 2 2 2 3 2 3 2
2 2 2 2
+60.8 *O. 3 f68.3 1 0 . 2 - 1.9*0.1 f 4.910.1 -26.3 * O . 4 -22.110.3 -28.9 * O . 6 -24.6 * O . 7 -30.6 "0.0 -28.1*0.3 -24.7 * 0.1 -23.3 * O . 1
The first two columns of Table I " give the initial and final concentrations of a dilution, respectively. The remainder of the table indicates the average Of the n 1 d X r of duplicate the intermediate heat of dilution and the probable (7) The Ruud Manufacturing Company very kindly aided in the forming of this coil.
THERMODYNAMICS OF SODIUM CHLORIDESOLUTIONS
Dec., 1933
2639
error a t each temperature. It is estimated that mittedly the temperature range of these measuresystematic errors (errors in the measurement of ments is too small to allow a definite decision. concentration, time, potential, resistance, etc.) The slopes of the straight lines so obtained (az/ are less than one %. d T = E,,* - E&) have been plotted against ml/' Treatment of Results.-The data of Table I in Fig. 2, where for comparison also are shown the are plotted against the square root of the results of Randall and Rosinis and Harned and molality in Fig. 1. Below ml/' = 0.1 the experi- Nims'o and the limiting law slope according to mental points lie on a straight line within the limit of experimental error and the curves have been extrapolated linearly to infinite dilution to , -100 obtain integral heats of dilution by applying the method of least squares to all measurements below ml/'= 0.1. An additional uncertainty E of about one calorie is introduced into the values g of the integral heat of dilution by this extrapolation. There is some evidence of continuously changing slope to m = 0 but the heat effects below ma/'= 0.1 are too small to decide this question definitely. The relative partial molal heat content of 0 0.2 0.4 0.6 0.8 1.0 sodium chloride (z2)was evaluated from z2 = ??a'/*, - (m1/'/2)d/dm1'p ( w m ) * 8 The Fig. 1.-Heats of dilution of aqueous sodium chloride soluof the smoothed curves of Fig. 1were determined tions: 8,25', 0,loo,6,15O, 0 , loo. Measurements below from a large scale plot by use of the differentiated m1/2= 0.07 are shown in the larger scale plot where each form of the L~~~~~~~interpolation formula. curve is offset 5 calories from the preceding curve for conElvalues were obtained from the relation 1,= venience in plotting. ( - A H , - E,)/n, where 1c1 is the number of moles the theory of Debye and Hiickel. A straight of water associated with one mole of sodium line with a slope of 15.8 expresses the results chloride. In Table I1 are listed values of AH,, E,, = 15.8 ml/* cal. mol.-l deg.-'. Eland 12 for various concentrations in the range The values a t 25' given in Table I1 agree of these measurements. well with those previously reported by Robinson3
2 5
-
TABLE Ir INTEGRAL HEATSOF DILUTION AND RELATXVE PARTIAL MOLAL HEATCONTENTS OF AQUEOUS SODIUM CHLORIDE SOLUTIONS TI eal./mole, HzO -AH,, cal./mole, NaCI-L;, cal.~mole,.NaCl-
25" 0,0001 4.2 .0005 9 . 4 ,001 - 1 8 . 2 .005 - 2 9 . 6 .01 -39.3 .06 -71.3 .l -86.5 .2 -87.6 .404 -69.4 .816 -11.7 m
-
-
20'
- 3.7 - 8.2 -11.6 -26.0 -36.2 -63.5 -70.0 -68.3 -42.3 f23.5
15'
3.4 7.7 -10.8 -24.3 -33.3 -54.2 -58.3 -46.3 6.6
-
-
10'
2.3 4.3 7.5 -16.9 -23.9 -40.6 -38.6 -18.3 27.2
-
25' 6.3 14.8 20.1 42.2 55.8 93.7 99.6 85.0 30.0 -123.3
20 5.3 13.1 17 9 38.1 51.0 76.3 76.6 54.2 -19.4 -169.7
z2
15'
4.7 12.0 16.5 35.8 46.0 65.0 57.7 9.4 -89.9
To obtain (t& - E:*) the values of Table I1 have been plotted against the temperature for the concentrations listed in Table 11. The d a h do not cover a suEcient temperature range to permit the drawing of anything but straight lines. A slight but definite s curvature is apparent a t all concentrations and to be more than can be accounted for by experimental emor but ad(8) Rossini. Bur. Standards 1.fisedrch, 6, 799 (1931).
10'
3.0 8.3 11.6 25.6 34.8 43.2 30.2 20.0 -130.0
-
-
25' -0.0000038 ,000049 .00012 ,0011 .0030 .020 .024 .0094 .28 1.19
-
-
io0
----
-0.0000036 .000044 .00011 ,0011 .0027 .012 .051 .44 2.11
150 -0.0000024 .OOOOY9
-
-
.00010 .0010 .0029 .0098 00011 .13 .70
.
100
-0,0000013 ,000036 .000072 .00078 ,0020 .0023 .015 .14 .65
-
-
-
and are also in good agreement with the calorimetric results of Young and Vogelll for the concentration range in which the two series of measurements overlap (0.8 to 0.05 m)l2 but differ by (9) Randall and Ro=ini, THIS JOURNAL, 61,323 (1929). (10) Harned and Nims, ibid., 64,423 (1932). (11) Young and Vogel, ibid., 64,3030 (1932). (12) The values for A H , and are also in agreement with the values at 18' calculated by Rossini [Bur. Standards J . Hescarch, 6 , 791 (1031)Jfrom the calorimetric data of Richards and Rowe [THIS JOURNAL, IS, 770 (1921); Wiist and Lange, 2. physik. Chcm., 116, 161 (1925); Pratt, J . FranklinInsf.,l M , 663 (1918): d a2.l.
r*
E. A. GULBRANSEN AND A. L. ROBINSON
2640
Vol. 56
1Q to 40% from the results of Harned aad NimslO obtained from e. m. f. meaeurements. It is estiinated that the Ls valdes of Table I1 are ackurate to within 5%. It is more diftEcult to estimate the probable error in (e,$ - E:,), but it is probably not more than 15%.
Randall and Russinia have calculated limiting law values for the relative partial molal heat capacities of strong electrolytes by a double differentiation of the Debye-Hii~kel'~free energy BHpressiofi. For the dielectric constant of water they used the equation of AdamsZ0which is based on the measurements of Kockel.21 For uni-uniHarned Randall and and Nims Rossin i valent electrolytes they obtained C,, = 6 ml/' La Mer and Cowperthwaite2a have recently derived a limiting law for heat capacities DtsbYkHddkel in the same manner but have taken account of limiting the correction pointed out by Gatty16and Scatchlaw ard14 atld have used the dielectric constant data of Wyman. l7 For uni-univalent electrolytes they obtain e,, = 13.25m1/' E&. The agreement of the experimental value of the slope (15.8) with this theoretical value is largely fortuitous since a 0 0.2 0.4 0.6 term in d2D/dT2in the eltpression of La Mer and d 2 . Cowperthwaite has approximately this same magFig. 2.-Relative partial molal specific heat of sodium nitude (but opposite sign) and althaugh dD/dT chloride in dilute aqueous solutions. for water is probably known with an accuracy of IV. Discussion 5%,24 dzD/dT2 is uncertain to within 50% or The Debye-Hiickel limiting law value for the more. The temperature coefficient of (Ebz - C&) idtegrsl heat of dilution of strung electrolytes as for uni-univalent electrolytes is 0.16 cal. mole-1 derived by BjerrumlSand modified by 9catcha1-d~~oC.-2 between 10 and 25: according to the equaand Gatty16 to take account of the fact that the tion of La Mer and Cowperthwaite but this dilution is cartied aut a t constant pressure is value is even more uncertain than the theoretical value of the slope of (Efi2 - e$,) vs. ml/' beA B m = 2.~89/108(e3NJ/~/Lla/~)(~/10~~T)1/~(~,z~)L/n( 1+ T/D(dD/dT) T/3V(dV/dT))cal./molelB cause of the large uncertainty in d2D/dT2. Drawus. T is of The term la dV/dT reduces the value of the limit- ing straight lines for the plots of course equivalent to a decision to interpret (Cp, ing slope by 7.2% at 25'. At lower temperatures as having constant values within the experithe effect is much smaller since at 3.98' dV/dT mental error for the temperature range of these = 0. To evaluate the theoretical limitihg slope the dielectric constant measurements of Wyman17 measurements. The experimental slope of 15.8 were used and the term in dV/dT was calculated is lower than the value of 21 found by Randall from data givefi in the "International Critical and Rossinigfrom calorimetric measurements and Tables."I* The theoretical slopes calculated in the value of 24.8 obtained by Harned and Nims'O from e. m. f. data. The measurements of these this manner are -480, -436, -412 and -360 at 25, 20, 15 and 10'; the corresponding values investigators extended to 0.05 m. The zero tem- et,) agrees with of the: experimental slopes are -418, -370, -340 perature coefficient for previous investigations. and -239. Some significance may be attached To evaluate it is necessary to know E;*. to the fact that when the theoretical slope changes The integral heat of solution of sodium chloride rapidly the experimental slope also changes is the heat absorbed when one mole of the salt is rapidly. An inflection point for both the theodissolved in an infinite amount of water and is retical and experimental SIQP~S is indicated a t equal to (H, Extrapolating the heat 17-1 8'.
+
+
+
Et,)
z&').
(13) Bjerrum, Z ghysik. Chum., 119, 145 (€926). 118,2037 (1931) (14) Scatchard, THISJOURNAL, (15) Gatty, P h d . M a g , 11, 1 W 2 (1831). (16) u, is the number and z, is the valence of ions of the ith kind; the other symbols have their usual significance. (17) Wyman, Phys R e v , SS, 623 (1930) (18) "Iaternational Critical Tables," McGraw-Hi11 Book CU , New pork City
(19) Debye and Huckel, Physrk. Z.,S4,185 (1923). (20) Adams, TEIS JOURNAL, 48,621 (1926) (21) Kockel, A n n . P h y s z k , 77, 417 (1926). ( 2 2 ) The use of Wyman's dielectric constant data in the equation of Randall and Rossini gives a valde of 12 4 for the limiting slope. (23) La Mer and Cowperthwaite, TAIS IOURNAI,66, 1004 (1933)
(24) Lange and Robinson, rbrd., 61, 2811 (19301
THELATENTHEATOF FUSION OF DzO
Dec., 1934
2641
of solution data of Lipsett, Maass and Johnson26 Ramage% found -21 and Randall and RossiniO at 25 and 20’ to infinite dilution with the aid of give -23.3 calories per mole from the extrapolathe data presented here the integral heat of solu- tion of direct calorimetric measurements. E$, 15.8 rn’l’. tion is found to be 916 and 1087 calories per mole may be calculated from E*, = -21.9 at the temperatures 25 and 20°, respectively. We are greatly indebted to the National ReThe heat of solution data of Cohen and K 0 0 y ~ ~search Council for two Grants in Aid which made treated in the same manner give 916 and 1083 possible the construction of the calorimeter used calories.27 The temperature coefficient of the in this investigation. E f ) . Assumintegral heat of solution is (& Summary ing a linear variation of the integral heat of soluHeats of dilution of aqueous sodium chloride sotion over this small temperature range and sublutions have been measured at low concentrations stituting the value of e,, given in the “Internaa t intervals of 5’ from 25 to IOo and have been tional Critical Tables”18the value of -21.9 caloto infinite dilution to obtain integral extrapolated ries per mole is obtained for Randall and heats of dilution. The values of relative partial (25) Lipsett, Johnson and Maass, TRISJOURNAL, 49, 925 (1927); 49, 1940 (192’7). molal heat contents and partial molal heat capaci(26) Cohen and Kooy, 2. physik. Chem., is#, 273 (1928). ties have been calculated for the concentration (27) The heat of solution data of Wiiist and Langela when combined with our dilution measurements at 25’ give 926 calories per mole and temperature ranges of the measurements.
+
-
sodium chloride for the integral heat of solution. We have used the data of Lipsott, Maass and Johnson and Cohea and Kooy breausc they are available at two temperatures.
[CONTRIBUTION FROM THE
(28) Randall and Ramage, TXIS JOURNAL, 49, 93 (1927).
RECEIVED OCTOBER13, 1934
PITTSBURGH, PA.
DEPARTMENT OF CHEMISTRY,
COLUMBIA UNIVERSITY]
The Freezing Point of Mixtures of H,O and D,O. The Latent Heat of Fusion of D,O BYVICTORK. LAMERAND WELDON N. BAKER
In a preliminary communication, we pointed out that the solid phase which separates on freezing a mixture of the isotopic waters, HzO and DzO, is a solid solution and that we had been unable to effect a perceptible separation by slow crystallization of a mixture containing 40% DzO. Since the two components, A = HzO and B = DzO, differ only in the isotopic masses of the hydrogen atoms, H and D, it is reasonable to expect that the solid phase will behave as an ideal solid solution. The system should therefore furnish an unusually appropriate example for testing the equations of the perfect solid solution recently developed by S e k 2 In Table I we summarize our cryoscopic measurements obtained with a 15-cc. sample using the ordinary Beckmann technique. The experimental data of the two series of experiments, as shown in Fig. 1, fall upon a smooth curve which may be represented satisfactorily by the formula At = 39.42AS - 38.8( By correcting the formula of Lewis and LutenS (relating AS, the increase in specific gravity, to NB the mole fraction of D20)
from AS = 0.1056 for 100% DzO,to A S = 0.1079 a t 25°,4 we obtain the relation At = 4.213 NB 0.411 NB2.
(I) La Mer, Eichelberger and Urey, TEISJOURNAL, 66,248 (1934). (2) Seltz, ibid., 66,307 (199SY.
(5) Washburn, Smith and Frandsen, Bur. Stds. J . Res., 11, 453 (1933). (6) Lewis and Macdonald, T m s J O ~ N A 46, L , 3068 (1933).
Aa2.
(31 Lewisand Luten, fbid., 46,6061 (1933).
TABLE I OBSERVED FREEZING POINTS OF D20-Ha0 MIXTURES S At A1 At NB N
=
9.377AS
-
Obs. inAI = Atcreasein 39.42A.S- 4.213 NBP 38.8 (AS)* 0.411 N ~ s
-
Sp. gr.25/25 1.01 (AS)% f. p.. ‘2.
1.001376 0.0129 1.01644 .I539 1.02135 .1997 1.04411 .4117 1.04-456 .4158 .6915 1.06351 .8282 1.08918 1.10068 .9338 1.1079 1.OOO
0.0536 .6321 .824l 1.6701 1.679l 2.351 3.207 3.578
0.054 ,638 ,824 1.663 1.680 2.347 3.207 3.576 3.802
0.064 .639 ,825 1.665 1.681 2.348 3.207 3.576 3.802
The extrapolated value for the freezing point of pure DzOis 3.802’ (TB= 276.98) in good agreement with the preliminary measurements of Lewis and Macdonald (3.8u),5and the later measurements of Taylor and Selwood (3.82°).4 For the equilibrium state between a perfect (4) Taylor and Selwood, ibid., 66, 998 (1934).
