T H E PARTIAL PRESSURES OF WATER I N EQUILIBRIUM WITH AQUEOUS SOLUTIONS OF SULFURIC ACID' EDWARD M. COLLINS Thompson Chemical Laboratoru, Williams College, IVilliamstown, Mass. Received April 17, 1933
In connection with a method developed by Collins and Menzies for measuring aqueous tensions in salt hydrate systems, which is to be described in a subsequent paper, a knowledge was required of the partial pressures of water over aqueous solutions of sulfuric acid for various concentrations and a t various temperatures, Critical discussions of existing data have been given by Wilson (l), Greenewalt (2), and Hepburn (3), and tables have been compiled. The agreement between their values a t 25°C. indicates that this isotherm is fairly accurately located. At other temperatures, however, the tables of Wilson and Greenewalt do not agree to the extent expected if the claims of accuracy of each are considered, the disagreement in many cases being more than 4 per cent. We decided, therefore, to redetermine experimentally the vapor pressures over aqueous solutions of sulfuric acid. CHOICE OF METHOD
The dynamic boiling point method employed by Burt (4) presents many difficulties. Among the principal objections are superheating, change of concentration with boiling, and the existence of a pressure gradient. The gas-current saturation method used by Briggs (5), Sore1 (6), and others presents difficulties in temperature control and is accompanied by a change in concentration of the solution. The dew point method used by Hepburn (3) is practical for only moderate temperatures. The static methods used by Regnault (7), Bronsted (8), Dieterici (9), Grollman and Frazer (lo), and Hacker ( l l ) , permit greater accuracy. A static method was therefore desirable. The method chosen consisted of a modification of the static isoteniscope of Smith and Menzies (12). The partial pressure of hydrogen sulfate or sulfur trioxide over aqueous sulfuric acid solutions is not appreciable below 150°C. According to Thomas and Barker (13) the partial pressure of hydrogen sulfate over a 99 per cent solution is 0.5 mm. a t 180°C. Hence a method measuring the total vapor pressure of a solution will, in effect, give results for the aqueous partial pressure. From a thesis presented by the author t o the Faculty of Princeton University in partial fulfillment of the requirements for the degree of Doctor of Philosophy. 1191
1192
EDWARD hl. COLLINS EXPERIMENTAL
Isoteniscope For details of the static isoteniscope one is referred to the original art,icle (12). The static isoteniscope as described there is applicable in the case of single substances and saturated solutions. It is, however, not
FIQ.1
directly applicable in measuring vapor pressures of liquid pairs or unsaturated solutions. I n these cases the repeated boiling-out to insure the expulsion of all the dissolved or absorbed foreign gases would be followed by a change in the concentration of the solution. It was necessary to devise an isoteniscope which would permit the expulsion of the foreign gases with no accompanying change in concentration. Jacketing the part of the isoteniscope nearest the manometer with a water condenser would prevent loss
PARTIAL PRESSURES OF WATER
1193
of water from the instrument itself, but the solution in the bulb would become more concentrated, while the solution acting as the confining liquid would become more dilute on account of the returned condensed water vapor. This difficulty was obviated by abbreviating the instrument to a J-tube (figure 1). It was made from glass tubing (7 mm. internal diameter). The closed end, A, was approximately 6 cm. long and the length over all was 40 cm. The water jacket, E, was constructed from 20-mm. tubing and was fitted to the isoteniscope with rubber stoppers. The condenser was put in place before sealing on the Kjeldahl trap D. It proved convenient to have both the inlet and outlet tubes of the condenser inserted through the top stopper. The Kjeldahl trap was added as a safety device to protect the pressure measuring system from an influx of acid solution should the pressure in the isoteniscope accidentally become much larger than the applied pressure. The side tube, F, which may be opened or sealed a t will, permits greater facility in washing and charging. Since approximately 5 cc. of solution were needed there arose the question of a change in concentration through the possibility of condensed water in part C at times other than during the process of air removal. On no occasion was any water ever observed laving the inner wallsof the condenser, so that the amount of water must have been small. And yet no vapor escaped past the condenser, which was proven by the subsequent analysis of several samples of acid solutions after having been used in the pressure determinations from 25°C. to the vicinity of 125”C., the concentration agreeing with the initial concentration within the experimental error. T o get an idea how small an amount of water, if any, was on the walls of the condenser the following experiment was performed. The glass tubing between the isoteniscope proper and the condenser was bent so that it was approximately 70” from the vertical and the condenser 20” from the vertical. A piece of glass tubing several centimeters long and having a closed end was sealed to the underside just below the condenser forming an “appendix.” The isoteniscope was charged with an acid solution with a boiling point of approximately 110°C. and connected with the pressure measuring system. The isoteniscope proper was maintained a t 105°C. and the pressure in the system adjusted accordingly. The “appendix” was kept cold by immersion in a bath of water. At the end of two hours the “appendix” was cut off and its water content determined. The water which collected was inappreciable, i.e., less than 0.01 g., an amount having too small an effect upon the concentration to cause concern. This may not seem surprising when we stop to consider that the only way for the water to reach the condenser was by diffusion through a tube 14 cm. long and of 7 mm. diameter against an opposing equal pressure of air, except during the process of air removal.
1194
EDWARD M. COLLINS
Preparation of solutions The acid solutions used in these determinations were prepared from redistilled sulfuric acid of the best reagent grade. An analysis of the acid was made, which showed that the statement made by the manufacturer regarding the maximum limits of impurities was correct. This statement showed that the acid conformed to the standards of purity as given by Murray (14). The various solutions were analyzed in triplicate for sulfate by precipitating and weighing as barium sulfate. The analyses checked within 0.04 of a percentage of composition, which amounts to an error of less than 0.1 per cent for most of the solutions. Temperature regulation For the lower temperatures a water-bath consisting of a 4-liter beaker was used; for the higher temperatures oil was substituted for the water. The bath was jacketed with an inch of asbestos, the variety commonly used for wrapping steam pipes. The temperature was regulated by an electric hot-plate in series with a variable resistance. An additional electric heater of the immersion type was used to facilitate raising the temperature rapidly and then disconnected when the desired temperature was reached. A coil of copper tubing immersed in the bath, through which cold water could be run, afforded the means of cooling when necessary to lower the temperature. The temperature was controlled within 0.01 to 0.02"C. Violent stirring was provided and uniformity of temperature throughout the bath was obtained.
ThermometTy The temperature was measured by means of a platinum resistance thermometer of the Callendar type with compensating leads. The platinum wire was wound on a mica form and was jacketed with a porcelain tube. It was calibrated a t the freezing and boiling points of water and a t the transition point of sodium sulfate decahydrate. Resistances were measured by means of a R'lueller type thermometer bridge (15). The makers supplied calibration data showing no deviations which would affect our results. The sensit'ivity of the galvanometer was such that a change in the resistance of the thermometer of 0.001 ohms, corresponding to O.Ol"C., produced a deflection of one half a scale division. No thermal E.M.F. causing more than 0.01"C. was observed. The fixed points showed no change throughout the period of the measurements, probably because the thermometer was never heated above 140°C.
Pressure The manometer was similar to the one described by Smith and Menzies (12). The mercury was purified by the nitric acid treatment and subse-
PARTIAL PRESSURES O F WATER
1195
quent distillation in vacuo. The manometer levels were measured by a graduated steel bar, the length of which was calibrated by comparison with a standard meter, the latter having been calibrated by the late Dr. E. W. Morley. The graduations were compared by means of dividers for 5 cm. intervals. There were no discrepancies large enough to affect our results. The room temperature variation was so small that the bar was not noticeably affected. The steel bar carried a movable sleeve with a vernier scale. To this sleeve was attached a strip of mirrored glass with a horizontal hair line, in such a way that it was behind the manometer tubes on each side of the steel bar. The hair line could be adjusted a t the mercury level, avoiding parallax by the alignment of the mercury level with its image. The manometer readings were reduced to millimeters of mercury a t 0°C. and to sea level a t 45” N.L. Thevalue of the gravityconstant, g, was taken as 980.3, the calculated value for Williamstown, Mass., according to its latitude, longitude, and elevation. This value agrees with that determined in the Thompson Physical Laboratory of Williams College. For pressures less than one atmosphere the “open end” of the manometer was connected to a system containing a Hyvac pump and McLeod gauge. This system was evacuated to a residual pressure of 0.005 mm., thereby obviating the reading of the barometer a t each determination and also increasing the accuracy of the measurements. Pressures less than 10 mm. were read by means of another McLeod gauge protected from water vapor by a guard tube containing magnesium perchlorate trihydrate. Since an open-end manometer was used for pressures above one atmosphere the barometer had to be read. The barometer was by Henry Green of Brooklyn, N.Y. It was tested for accuracy by comparison with the manometer when the open end was at atmospheric pressure. The Green barometer was found to be accurate within the error of reading the instruments.
Manipulation The essential difference in the manipulation of this isoteniscope and that of Smith and Menzies was the method by which the dissolved gases were removed. Enough solution was placed in the isoteniscope to fill the closed arm about one-half full when the levels of the solution in both arms were equal. About 5 cc. were required. Then the isoteniscope was connected to the system and the latter evacuated. Bubbles of gas were given off from the solution and part of this gas collected in the closed end of the isoteniscope. This was removed from time to time by bringing the instrument to horizontal position. Sharp tapping aided in the removal of the gas from the solution. After several repetitions of gas removal it was necessary to warm gently the closed end of the instrument in order to form a bubble of
1196
EDWARD M. COLLINS
water vapor. If the isoteniscope was swung quickly back and forth through a small arc, as a clock pendulum, a sharp click was heard when the solution struck the closed end. This was an index of the absence of foreign gases. As further proof of the removal of gases a measurement of the aqueous tension was made at a given temperature and then the process of gas removal repeated a t a higher temperature. A redetermination of the pressure a t the first temperature was then repeated and identical values ensured the absence of all foreign gases. Any water which might have deposited in the condenser during the process of air removal was allowed sufficient time to reflux and become thoroughly mixed with the solution before the pressure measurements were made. EXPERIMENTAL RESULTS
The experimental results are given in table 1. FINAL TABULATION O F RESULTS
Since the vapor pressures measured were for solutions of various concentrations and a t various temperatures it was necessary to devise a method whereby the aqueous tensions of solutions of intermediate concentrations could be obtained a t any desired temperature. The experimental data was first smoothed by a highly sensitive graphical method. I n this method the log of the pressure is temporarily assumed to be a linear function of the reciprocal of the absolute temperature. Since log P, however, is not strictly a linear function of 1/T there was a difference between the logs of the observed pressures and those so calculated. These differences were plotted as a function of temperature, and by their aid the pressures a t rounded temperatures were calculated for intervals of 5°C. Since the relative vapor pressure for a solution changes only slowly with temperature, graphs showing this relation allowed us to test the data smoothed as above and to smooth it further where necessary. Extrapolation of these curves gave the values for 20"C., 14OoC., and in some cases 135°C. The relative vapor pressures for solutions of rounded concentrations were next obtained graphically. To accomplish this, the logs of the smoothed relative vapor pressures were plotted as a function of the concentration. Although no experimental data was obtained for concentrations below 29.90 per cent the extrapolation of the isotherms from this concentration to 0 per cent was accomplished without the introduction of appreciable error. The logs of the relative vapor pressures for 10 per cent, 20 per cent, and 25 per cent were read directly from these isotherms. Instead of obtaining the logs of the relative vapor pressures for solutions above and including 30 per cent directly from the graphs, we obtained them by a more accurate method. This method involved calculating them on the temporary assumption of a parabolic relation between the log of the relative vapor pres-
1197
PARTIAL PRESSURES OF WATER
TABLE 1 Vapor pressure for suljuric acid-water solutions Experimental concentrations. Pressure in millimeters bf mercury
degree8 C.
degrees C.
degrees C.
degrees C.
24.87 24.41 30.00 29.90 34.45 34.40
17.8 17.0 24.1 23.7 30.7 31 .O
39.69 44.90 49.71 54.91 59.72 64.59
41.1 54.5 69.6 89.9 112.9 141.7
69.85 74.83 79.67 85.32 89.92 94.78
179.0 221.1 270.7 341.9 408.1 491.0
25.82 31.44 36.20 45.88 50.51 55.44
15.0 21.1 27.5 46.3 59.2 74.8
57.32 63.11 69.69 76.29 82.71 86.09
82.4 108.4 146.5 195.4 255.4 292.1
90.67 96.51 103.13 110.80 123.64 129.25
351.5 440.3 561.1 737.7 1133. 1344.
25.20 29.95 36.53 39.99 44.67 49.72
12.0 15.6 23.2 28.0 36.2 47.0
54.17 59.12 64.75 69.76 74.79 79.98
59.2 75.6 97.7 123.7 154.3 193.0
84.39 90.16 94.86 99.73 104.76 109.92
26.04 32.81 41.61 46.58 51.18 55.29
10.1 14.9 24.9 32.4 41.5 51.0
61.60 69.83 74.53 80.04 84.19 89.41
69.8' 102.4 126.4 160.8 191.8 237.5
24.49 29.69 34.48 39.93 44.50 50.15
7.01 9.60 13.0 17.4 22.4 30.7
54.29 59.40 64,57 70.17 74.42 79.65
37.9 49.4 63.6 83.4
,
102.2
127.0
99.71 104.44 110.06 113.85 119.33 123.06
590.7 699.2 845.1 964.4 1155. 1296.
231.6 293.9 353.1 423.0 511.9 614.8
115.56 119.16 124.25
743.3 844.9 996.4
94.75 99.23 104.81 109.64 114.26 121.27
296.1 349.3 430.3 513.6 602.7 766.1
129.11
990.7
85.05 90.29 94.96 100.43 105.37 109.20
159.8 199.9 239.8 298.5 361.2 412.8
115.53 119.75 125.98
523.0 601.2 750.6
.
1198
EDWARD M. COLLINS
TABLE I-Concluded
I(OBBERVED>(( P
t
degrees
C.
1
1
(OBS&VED)
degrees C.
25.11 33.95 40.58 45.05 49.77 55.19
4.82 8 40 12.4 16.1 20.1 28.0
59.95 65.01 70.13 74.95 80.22 85.16
29.72 34.99 39.59 44,66 49.77 54.91
2.89 4.29 5.37 7.58 10.16 13.74
59.98 64.36 69.86 74.52 79.19 84.72
1
'
I
(OBBERVED) P
18.02 22.7 29.8 37.6 46.9 GO.G
'
J(o.SERVED) P
degrees C.
degrees C.
35.5 45.6 58.3 72.7 92.2 114.3
/I
89.87 95.06 100.15 104.10 109.66 114 99
139.5 172.9 211.1 246.2 303.9 370.1
120.74 125.42 130.20 135.61
90.16 94.71 100.20 104.56 109.96 114.74
77.4 93.9 118.5 141.3 174.9 210.6
119.13 124.41 129.46
453.2 533.8 630.0 752.8
,
I
I
247.8 301.2 361.1
sure and the concentration, a method similar to the one described above where a linear function was temporarily assumed. Table 2 presents the final .tabulation of the relative vapor pressures. The first and second columns, respectively, give the temperature and the vapor pressure of water in millimeters of mercury. The other columns give the relative vapor pressures in per cent for the solutions designated a t the head of each column. Values for intermediate concentrations and temperatures may be obtained by linear interpolation from table 2 without introduction of large errors, since the R.V.P. isotherms are sufficiently rectilinear for intervals of 10 per cent and the relative vapor pressure for a solution of given concentration is a linear function of the temperature. Relative vapor pressures at temperatures not far above 140°C. may be obtained by linear extrapolation. Values in table 2 for 145°C. and 150°C. have been so obhined. ACCURACY OF RESULTS
The maximum absolute error in temperature measurement we estimate a t 0.05"C. For our most dilute solution this could cause an error of 0.1 per cent a t 125OC. and 1 per cent a t 25OC. From a study of the smoothing required of our results we find that the average smoothing required below 10 mm. amounts to 1.5 per cent. Between 10 mm. and 100 mm. the average smoothing amounts to 0.2 per cent. Above 100 mm. it amounts to but 0.1 per cent. In terms of pressure this would correspond ,to an average
1199
PARTIAL P R E S S U R E S O F WATER
better than 0.15 mm. below 10 mm., and 0.2 mm. above 10 mm. This error corresponds to the limit of error in reading the manometer, McLeod gauges, and barometer. Since the pressures for 10 per cent, 20 per cent, and 25 per cent were obtained by graphic interpolation of the logs of the relative vapor pressures, TABLE 2 Relative vapor pressures for suljuric acid-water solutions
1
P
(FOR WATER)
RELATIVE VAPOR PREBSURES I N PER CENT FOR SOLUTIONS CONTAINING FROM 10 TO 70 PER CENT HBOI
20
25
per cent
per cent
30 per cent
---
degrees C.
m m . Hg
20 25 30 35 40 45 50 55 60 65 70 75 80 85 90 95 100 105 110 115 120 125 130 135 140 145 150
17.54 23.76 31.82 42.18 55.32 71.88 92.51 118.0 149.4 187.5 233.7 289.1 355.1 433.6 525.8 633.9 760.0 906.1 1075. 1268. 1489. 1741. 1026. 1347. 1710. 3117. 3570.
88.0 88.0 88.0 88.1 88.1 88.2 88.2 88.3 88.3 88.4 88.4 88.5 88.6 88.6 88.6 88.7 88.7 88.8 88.8 88.8 88. g 88. g 89.0 89.0 89.1 89.1 89.2
82.4 82.6 82.6 82.8 82. g 83. o 83.1 83.2 83.3 83.4 83.6 83.6 83.8 83.9 84. o 84.1 84.2 84.3 84.4 84.6 84.7 84. 8 84. g 85. o 85.1 85.2 85.3
75.0 75.2 75.4 75.6 75.8 76.0 76.2 76.4 76.6 76.8 77.0 77. 2 77.4 77.6 77.8 78. o 78.2 78.4 78.6 78.8 79.0 79.2 79..4 79.6 79.8 80.0 80.2
35 per cent _ .
66.0 66.3 66.6 66. g 67.3 67.6 67. Q 68.1 68.5 68.8 69.1 69.6 69.8 70.1 70.4 70.7 71.0 71.3 71.7 72. o 72.3 72.6 72.9 73.2 73.5 73.8 74.2
40
45
50
55
per cent
per cent
per cent
per cent
60 per cent
56.1 56.5 56. 8 57.3 57.7 58.1 58.6 58. g 59.3 59.7 60.1 60.5 60. g 61.3 61.7 6Z2 62.6 63.1 63.4 63. s 64.2 64.6 65.0 65.4 65.8 66.2 66.8
45.6 46.1 46.6 47., 47.5 48. o 48.6 48.9 49.4 49.9 50.4 50.8 51.3 51.7 52.2 52.7 53.1 53.6 54.1 54.5 55. o 55.5 55. 9 56.4 56.8 57.3 57.8 -
35.2 35.7 36.2 36.8 37.3 37.8 38.3 38. Q 39.4 39.8
25.3 25.8 26.3 26.8 27.4 27. II 28.6 29. o 29.6 30.1 30.6 31.1 31.7 32.2 32.7 33.3 33.8 34.3 34. 9 35.4 35.9 36.6 37.0 37.0 38.1 38.6 39.1
16.1 16.6 17.1 17.6 18.0 18.( 19.0 19.5 20.0 20.4 20. g 21.4 21.9 22.4 22.8 23.3 23.8 24.3 24. 8 25.2 25.1 26.2 26.7 27. I 27.6 28.1 28.6
- -
-
40.5
41. o 41.6 42.1 42.6 43.1 43.6 44.2 44.7 45.2 45.7 46.3 46.8 47.3 47.8 48.3 48.8 --
_ .
85
per cent
9.a 9.7 10.1 10.6 11.0 11.4 11.8 12.3 12.7 13.1 13.6 14.o 14.4 14.8 15.2 15.7 16.1 16.5 17.o 17.4 17.8 18.9 18.7 19.1 19.5 19.9 20.4
70 per cent
3.4 3.7 4.1 4.4 4.8 5. L 5.4 5.8 6.2 6.6 6.8 7.2 7. 6 7.8 8.2 8.6 8. 8 9.3 94 9.9 10.3 10.~ 11.0 11.3 11.7 12.0 12.4
-
their accuracy depends upon the accuracy with which the isotherms of log R.V.P. versus concentration were located when they were extrapolated to zero concentration. We estimate that the probable error involved when reading the log of R.V.P. from these isotherms is not more than 0.005 units, which introduces approximately 1per cent error in the relative and absolute
1200
EDWARD M. COLLINS
vapor pressures. For the concentrations from 30 per cent to 65 per cent whose relative vapor pressures were obtained as described, the probable error is not more than 0.3 per cent. Since the pressures for 70 per cent were obtained by extrapolation, their accuracy is probably not better than 1 per cent. Table 3 compares our values a t 25°C. with those of Hepburn, Wilson, and Greenewalt. A comparison with the values of Wilson and GreeneWalt at higher temperatures seems to show better agreement between our values and those of Greenewalt, although in many instances our values are intermediate between those of Wilson and Greenewalt. TABLE 3 Comparative 86' isotherm Rounded concentrations. Pressures in millimeters of mercury SULFURIC ACID
OUR VALUES
VALUES OF WILSON
VALUES OF QREENEWALT
VALUES OF HEPBURN
22.4 20.8 19.4 17.8 15.8 13.5 10.9 8.45 6.15 3.97 2.24 1.03 0.41 0.12
22.7* 20.9' 19.6* 17.8 15.8 13.4 10.9 8.36 6.05 3.95 2.14 0.88 0.41 0.13
par cent
10 20 25 30 35 40
45 50 55 60 65 70 75 80
22.7 20.9 19.6 17.9 15.8 13.4 11.0 8.48 6.13 3.94 2.30 0.88
22.8 21.0 19.7 17.9 15.8 13.5 11.1 8.7 6.3 4.1 2.3 1.2 0.55 0.19
* Agree with the values obtained from the data of Grollman and Frazer (10). RELATED THERMAL DATA
The heat of vaporization of water from 'an aqueous solution of sulfuric acid may be calculated from our vapor pressure data by the aid of the Clausius-Clapeyron equation in its approximate form d-eInP
dT
Q RT2
(1)
Here Q is the heat absorbed in the evaporation of one mole of water from a large amount of solution so that there is no accompanying change in concentration. This is equal to the heat effects involved in removing one mole of liquid water from the solution and evaporating it. If we assume Q to be constant, we may integrate and obtain the equation:
1201
PARTIAL PRESSURES O F WATER
'
log P = -+ B 2.3 RT
'
Upon plotting log P versus 1/T we obtain a line whose slope is - The 2.3R' heat of vaporization, however, and in consequence the slope, is known to decrease with increasing temperature, attaining the value zero a t the critical temperature. The graphs of our results showed that the slopes of these curves were functions of the concentration as well as the temperature. The slope and the heat of Vaporization increase with the concentration of sulfuric acid, since more work is required to remove water from a solution of greater concentration (cf. Greenewalt 2). The heat of vaporization a t a given temperature could be calculated from the slope at that temperature. Since, however, the slope changes with temperature, it is best that there be chosen a small temperature interval to obtain the slope from the ratio TABLE 4 Heats of vaporization for sulfuric acid-water solutions Q (calories per gram) = A 1°C. B
+
A. . . . . . . . B.. . . . . . .
A log P
The values of Alog P and Al/T, however, can be made neces-
sarily small only a t the sacrifice of the accuracy of the ratio. In order to avoid this difficulty the method of calculating the heats of vaporization was modified. Since Q is a function of the rate of change of log P with 1 / T it follows that for two substances the ratio of the heats of vaporization at a given temperature is equal to the ratio of their rates of change of log P a t that temperature. Furthermore, the interval determining the rate of change of log P need not be necessarily small. This is shown as follows: For the first substance QI
=
2.3 RT2 d log PI dT
and for the second substance Qz =
2.3 RTa d log Pz
dT
1202
EDWARD
M.
COLLINS
Then at the same temperature the ratio of the heats of vaporization log Pi - d log Pi --=2.3 RTa d log P1 d log Pa
&I - 2.3 RT2 d
Q,
A log Pi A log P2
By substituting the values for water for QZ and Alog Pzobtained from the International Critical Tables and our data for Alog PIfor intervals of 5°C.we calculated the heats of vaporization of water from the acid solutions. The values so obtained were next smoothed by plotting them as a function of the temperature. I n drawing these curves it was borne in mind that the heat of vaporization for a solution must decrease a t a rate not less than that for water. If this were not so the log P versus 1/T curve for a solution would intersect that for water a t some temperature. Our smoothed heats were found to be represented by a linear equation whose constants are given in table 4. By means of this equation the heats of vaporization at a certain temperature may be calculated with an accuracy of approximately 1 per cent. Upon calculating the heats of vaporization from Greenewalt’s table we concluded that his table was not sufficiently accurate to show the required decrease with temperature. SUMMARY
A redetermination of the aqueous partial pressures over aqueous sulfuric acid solutions has been made for concentrations up to 70 per cent HzS04from 20°C. to 140°C. For this purpose a n isoteniscope was devised which could measure vapor pressures of solutions with no accompanying change in the concentration of the solution. From the observed data a table of relative vapor pressures has been compiled, from which table it is possible to ascertain the aqueous partial pressure for aqueous solutions of sulfuric acid between 0 per cent and 70 per cent H2S04for any temperature from 20°C. to 140°C. inclusive. The heats of vaporization of water from solutions of the above range of Concentration and temperature have been calculated and tabulated. REFERENCES
(1) WILBON: Ind. Eng. Chem. 13,326 (1921). (2) GREENEWALT: Ind. Eng. Chem. 17,502 (1925). And also International Critical Tables, Vol. 111,p. 302. McGraw-Hill Book Co., New York (1928). Proc. Phys. SOC.London 40, 249 (1928). (3) HEPBURN: (4) BURT: J. Chem. SOC.86, 1339 (1904). (5) BRIGGS:J. SOC.Chem. Ind. 22,1275 (1903). (6) SORIL: J. SOC.Chem. Ind. 9, 175 (1890). (7) REGNAULT: Ann. chim. phys. [31 16, 129 (1845). D : physik. Chem. 68, 707 (1910). (8) B R ~ N B T E 2. (9) DIETERICI:Wied. Ann. 67, 865 (1899). (IO) GROLLMAN AND FRAZER: J. Am. Chem. SOC.47,715 (1925). (11) HACKER: Ann. Physik [4] 39, 1342 (1912).
PARTIAL P R E S S U R E S O F WATER
1203
(12) SMITHANDMENZIEB: J. Am.Chem. SOC.32,1412(1910). (13) THOMAB AND BARKER:J. Chern. SOC.127,2820 (1925). (14) MURRAY:Standards and Tests for Reagent Chemicals. D. Van Nostrand Co., New York (1920). (15) Bur. Standards Bulletin No. 288. (16) BURQESSAND LE CHATIDLIER: Measurement of High Temperature, p. 121. John Wiley and Sons, New York (1912).
