INDUSTRIAL AND ENGINEERING CHEMISTRY
1126
n = number of electrons taking part in reaction per molecule of reactant or product 4 = heat flow, B.t.u. per hour A Q = AG AH, reversible heat liberated, calories per grammole AQ' = total heat liberated atoelectrode, calories per gram-mole 2' = absolute temperature, Kelvin Ai!' = teryerature difference between anode and electrolyte,
-
c.
7 = overvoltage, volts = subscript indicating constant pressure
Vol. 43, No. 5
ACKNOWLEDGMENT
The authors are indebted to Frank B. West and R. W. Moulton for several valuable suggestions which have improved the manuscript. L. H. Clark did some of the preliminary work on the problem. William Antonius and John Sundling constructed part of the equipment. LITERATURE CITED
(1) Beck,
CONCLUSIONS
There is a considerable anode-electrolyte temperature difference in the electrolytic production of ozone from perchloric acid. Other reactions of industrial importance are probably significantly affected by electrode-electrolyte temperature differences. For the eutectic perchloric acid system studies, the electrodeelectrolyte temperature difference waa a linear function of the viscosity of the bulk of the solution a t the temperature of the bulk of the electrolyte. High heat transfer coefficients indicate the considerable effect of agitation by evolved oxygen a t an anode on the rate of heat transfer. Neglecting electrode-electrolyte temperature difference should cause considerable error in the determination of the temperature coefficient of overvoltage a t high current densities.
T. R., "Electrode-Electrolyte Temperature Differences," University of Washington thesis, 1950. (2) Beck, T. R., and Putnam, G. L., Trans. Faraday Soc., in press. (3) Bruz, B., 2. physik. Chem., Abt. A , 145, 283-8 (1930). (4) Clark, L. H., and Putnam, G. L., J . Am. Chem. Soc., 71, 34457 (1949). (5) Fink, C. G., U. S. Patent 1,581,188 (April 20, 1926). ( 6 ) Hornbeck, R. D., Lash, E I., Putnam, G. L., and Boelter, E. D., J . Electrochem. Soc., 98 (April 1951). ( 7 ) Xnobel, M., and Joy, D. B., Zbid., 44, 443 (1923). ( 8 ) McAdams, W. H., "Heat Transmission," 2nd ed., p. 242, New York, hIcGraw-Hill Book Co., 1942. (9) Putnam, G. L., Moulton, R. IF'., Fillmore, W.W., and Clark, L. H., 6.Electrochem. Soc., 93, 211-21 (1948). (10) Richards, J. W., Trans. Faraday SOC.,9, 140 (1913). (11) Solanski, D. N., and Sastry, P. S., J . Indian Chem. SOC., 25, 415 (1948). (12) Tarasov, G. Ya., J . Gen. Giza. (U.S.S.R.), 16, 1753-66 (1946). RECEIVED 3Tay 16, 1950.
Some Physical Properties of Sulfur exaf luoride J
H. 6. MILLER, L. S. VERDELLI, AND J. F. GALL Pennsylvania Salt Manufacturing Co., Wyndmoor, Pa.
T h e developing industrial uses of this chemically inert, high-dielectric material prompted an investigation of those physical characteristics which are inadequately treated in the literature, although of significance in the safe handling of the compound as a liquefied gas. Orthobaric densities vary from 0.14 gram per ml. (gas) and 1.47 gram per ml. (liquid) at 9" C. to 0.727 gram per ml., the critical density, at 45.5" C., the critical temperature, and 2794 cm. of mercury, the critical pressure. Pressure-temperature data for steel cylinders at various filling densities (weight of sulfur hexafluoride charged/capacity weight of water at 20" C.) indicate that pressures of 1000 pounds per square inch gage will result at about 69.7",
63.2', and 55.7" C. with filling densities of 100, 110, and 1209" respectively. Nitrogen was found to be appreciably soluble in liquid sulfur hexafluoride, the Bunsen absorption coefficient being about 2.1 at 26" C., and the effect of up to 0.807 weight % of dissolved nitrogen (0.84c/o total nitrogen) on cylinder pressures was determined. Thus the presence of 0.12, 0.42, 0.62, and 0.84% total nitrogen increased the gage pressure at 26" C. from 356 to 365,383,400, and 420 pounds per square inch, respectively. The data obtained as a result of this investigation should be of help in evaluating this industrially new chemical and promote safety in its handling as a liquefied gas in pressure cylinders.
T
of safe filling for cylinders of liquefied gases. With sulfur hexafluoride, whose critical temperature falls in this range, these data will also yield the critical constants. The sealed-tube method of Lowry and Erickson ( 1 ) for determining orthobaric densities was used in this work.
HE potential industrial use of sulfur hexafluoride as an inert material of high dielectric strength has prompted a rather thorough study of this compound (2). However, data that are of significance in its safe handling as a liquefied gas are inadequately treated in the literature, and, as predicted by Schumb ( 5 ) , the reported critical temperature of 54" C. is in error by 8" or 9'. It was the purpose of this study to correct this error and to supply the missing data, which will help in the evaluation of this industrially new chemical and to promote safety in its handling as a liquefied gas in pressure cylinders. ORTHOBARIC DENSITIES
A knowledge of the variation of vapor and liquid densities with temperatures in the range commonly encountered under normal atmospheric conditions is of importance in determining the limit
PROCEDURE. The cross-sectional areas of three glass capillary tubes approximately 2 mm. in bore and 7 to 15 cm. long were determined by the standard procedure, using known volumes of mercury. Purified sulfur hexafluoride was condensed as a solid in the tubes and sealed from the atmosphere, and the weight of the sulfur hexafluoride was obtained. The bulbs were suspended in a water bath behind a safety barrier, and the volume change of the liquid and vapor was observed with variation in temperature. The observations were made through a cathetometer telescope. With these data the liquid and vapor densities of sulfur hexafluoride under its own vapor pressure were calculated by solving the following simultaneous equations:
INDUSTRIAL AND ENGINEERING CHEMISTRY
May 1951
1127
50 where Do, DL
densities of vapor and liquid weights of sulfur hexafluoride in tubes 1, 2, and 3 Vjo, Vzg,Vag = volumes of vapor in tubes 1, 2, and 3 V,L,VZL,V ~ =L volumes of liquid in tubes 1, 2, and 3 a
mi, mz, ma =
L
The method of Lowry and Erickson does not give values for densities in the neighborhood of the critical points, because the meniscus is not well defined in that region. Therefore, these values were obtained by using the law of rectilinear diameter, and the critical temperature as determined by the sealed glass tube method. In the measurements of critical temperature, care was taken to seal into the tube the quantity of sulfur hexafluoride that kept the meniscus nearly stationary and well away from the ends of the tube as the critical temperature was approached. Values between 45.4' and 45.6" C. were noted as the temperature a t which the meniscus disappeared or appeared as the temperature was raised or lowered. The mean of these observed temperatures (45.5' C,), when read on the straight line drawn through the mean values of the liquid and vapor densities, gave the critical density (the law of rectilinear diameter). A smooth curve then w&s drawn from the last liquid and vapor density points to the critical density. The orthobaric densities are given in Table I and plotted in Figure 1. The accuracy of the measurements of liquid and vapor densities is believed to be within . t 5 and *12%, respectively.
40
30 20 IO
.20
.40
60
DENSITY
Figure 1.
.80
- GRAMS
1.0
1.2
1.4
PER G.G.
Orthobaric Densities of Sulfur Hexafluoride
TABLE111. CRITICAL CONSTANTS Critical Critical Critical Critical
temperature density molar volume pressure
45.50 c. 0.727gram/cc. 201 00. 36.8atm. (540 Ib./sq. inch absolute)
PRESSURE-TEMPERATURE RELATIONS AT VARIOUS CYLINDER FILLING DENSITIES
TABLEI. ORTHOBARIC DENSITIES Temperature, OC.
9.0 12.5 16.5 20.0 26.0 30.0 34.0 38.0 40.0 42.0 43.0 45.5
Liquid Density, G./Ml. 1.47 1.44 1.41 1.37 1.30 1.26 1.21 1.14 1.10 1.07 1.03 0.727
Vapor Density,
G./M1. 0.14 0.16 0.17 0.19 0.25 0.27 0.30 0.34 0.35 0.39 0.43 0.727
CRITICAL CONSTANTS AND VAPOR PRESSURE OF SULFUR HEXAFLUORIDE
The critical pressure was determined by extrapolating the vapor pressure curve, obtained by direct static measurements on the pure liquid, to the critical temperature (Tables I1 and 111). ZI
- TABLE11. VAPORPRESSURE OF SULFUR HEXAFLUORIDE t
(Equation for vapor pressure of SF6 between -50' and 40' C. log Pmm. = 7.308 907.98/T) Vapor Pressure Temperature, 'C. Hg, om. Lb./sq. inch abs. 96.6 18.68 -60 21.25 109.9 58 24.11 124.6 56 27.17 140.5 54 157.6 30.49 -52 32.39 167.5 -51 174.4 33.74 50 181.9 35.19 49 189.2 36.59 48 208.9 40 41 46 49.0 253.3 40 73.0 377.4 30 102.0 527.3 20 713.5 138 0 - 10 961.6 0 186.0 248.0 1282.0 10 315.0 1680.0 20 393.0 2030.0 30 487.0 2520.0 40 540.0 45.6 2792.0
-
-
I
I n order to confirm the maximum safe filling capacity of steel cylinders as determined from the above orthobaric densities of sulfur hexafluoride, the following direct observations 'R. ere made. PROCEDURE. Each of two steel cylinders was equipped with a 1000 pounds per square inch steel Bourdon tube gage and a needle valve. The volumes of the cylinders were determined with water. the sizes of the cylinders were chosen so as to obtain a 2- to 3-fold range of volume. (The capacity of the larger cylinder was 1527 grams of water a t 20" C.) The cylinders were evacuated and filled with sulfur hexafluoride a t dry ice tem eratures and any noncondensable gases (impurities) were pumpefoff a t this low temperature. The cylinders were gradually warmed and vented intermittently to afford an opportunity for more fractionation of noncondensable impurities from the gas. This process was repeated and the net weight of the sulfur hexafluoride adjusted to the desired filling density. All of the apparatus except the gage was submerged in a water bath and a constant temperature was maintained for at least 20 minutes before the pressure reading was recorded. Equilibrium was usually obtained in 10 to 15 minutes. Care was also taken to avoid any leaks in the system by silver-soldering all possible connections. The water bath was placed behind a steel barricade, with bulletproof window and exterior tem erature controls, to avoid possible personal injur in case of a c y k d e r failure. The filling lensity is recorded as weight per cent of the water capacity of the apparatus a t 20" C.
yofilling density
=
weight of SFa (grams) water capacity (grams) of apparatus a t 20" C.
x
100
The gages were standardized against a dead-weight gage before and after these measurements to assure accurate readings. RESULTS. The pressure-temperature readings for various filling densities between 100 and 120%, using the apparatus described above, are shown in Figure 2. The over-all accuracy of these measurements is believed to be within *0.5%. I n Figure 3, filling densities for both cylinders are plotted against temperature a t a constant pressure of 1000 pounds per square inch gage for both experiments. The temperature corresponding to this pressure decreases almost linearly with an increase in filling density in the range of 100 to 120% filling density.
Vol. 43, No. 5
INDUSTRIAL AND ENGINEERING CHEMISTRY
1128
1000
5 800 vi
a:
$ 600 cn
W
(r
400
20
30
40
50
60
102 104 106 I08 110 CENTIGRADE
SOLUBILITY OF NITROGEY IN LIQUID SULFUR HEXAFLUORIDE
I n the work on safe filling density determination it became evident that nitrogen TYSS very soluble in sulfur hexafluoride. This solubility of nitrogen in liquid sulfur hexafluoride was determined.
1000
d - 800
cn a:
600
cn cn W
LT
a 400
fluoride pressure reading. At 27' C., two phases exist in the system-liquid and vapor. The amount of nitrogen in the vapor can be calculated from its partial pressure, which was measured as the difference between the total pressure arid the vapor pressure of pure sulfur hexafluoride at 27" C. (The change in vapor pressure due to the presence of an inert gas is considered negligible a t these low concentrations of nitrogen.) After calculation of the total moles of nitrogen in the system and the moles of nitrogen in the vapor a t 27.0" C., the moles of nitrogen dissolved in liquid sulfur hexafluoride is obtained by difference. The quantity of sulfur hexafluoride in the liquid phase a t 27.0' C. was obtained from orthobaric liquid and vapor densities a t this temperature (DL= 1.30 grams per ml., Dv = 0.24 gram per ml.), the total volume of the system (1527 ml.), and the total amount of sulfur hexafluoride in the system (1832 grams) as follows:
+
30 CENTIGRADE
118
Relation of Filling Density to Temperature at Constant Pressure
1 . 3 0 ~ 0.24 (1527 - 2 ) 20
116
(CONSTANT PRESSURE O f 1000 R S I G )
Figure 3.
5
114
%FILLING DENSITY
Figure 2. Pressure us. Temperature for Various Sulfur Hexafluoride Filling Densities
w
112
TEMPERATURE
40
50
60
TEMPERATURE
Figure 4. Pressure us. Temperature for Sulfur Hexafluoride with Various Amounts of Nitrogen
1832
where
x
= 1383
ml. of liquid sulfur hexafluoride a t 27" C.
The Bunsen absorption coefficient, following equation:
CY,
was calculated from the
volume of Kt (S.T.P.) in liquid sulfur hexafluoride ~ ~ ( 2 C.) 7' = PROCEDURE. The apparatus employed for the previous filling volume of liquid SFs a t 27.0" C. X density measurements was used to determine the solubility of partiaI pressure of Ng (atmosphere) nitrogen in liquid sulfur hexafluoride. The apparatus was filled to 119.5 * 0.5% filling density with sulfur hexafluoride, to which various amounts of nitrogen were added. After each addition of nitrogen, the apparatus was shaken to expedite equilibrium conTABLE I17. PRESSURE RE.4DIR'GS ditions and the pressure at 27" and 50' C. was recorded. The Pounds/Sq. Inch Gage over-all change in filling density due to the largest addition of A t 27' C. A t 50° C. Expt. PJo. 0 (pure SF6) 356 840 nitrogen amounted to approximately 1.0%. At 50" C. the entire 860 I (Nzadded) 366 contents of the apparatus may be assumed to exist in the gaseous I1 (Nzadded) 383 910 I11 (Nz added) 400 943 state; therefore the moles of 980 I V (Nzadded) 420 nitrogen added [to the system can be calculated from its parOF NITROGES IR' SULFUR HEXAFLUORIDE TABLE V. SOLUBILITY tial pressure (assuming ideal S z in Vapor Phase at 27' C. S 1 in Liquid SF6 at 27' C. Total Nz in System Absorption mixtures). The partial pressure % by %. by % by Coefficient %.by G , G. weight weight %by volume G. weight volume a t 2 7 O C. of nitrogen was measured as the difference of the sulfur 2.2 0.12 0.3 1.5 2 0.60 0.1 2.1 0.12 0.86 4.3 2.3 0.30 0.41 2.1 0.42 7.4 7.7 h e x a f l u o r i d e p l u s nitrogen 6.9 2.1 0 49 1.4 0.597 3.04 0.62 10.8 11.3 2.0 9.7 1.9 4.07 0.71 14.6 0.807 pressure readings from that 15.4 0.84 of the pure s u l f u r h e x a -
May 1951
INDUSTRIAL AND ENGINEERING CHEMISTRY
RESULTS. The pressure-temperature curves for sulfur hexafluoride with various amoullts of nitrogen (from 0 to o.s4% nitrogen) are shown in Figure 4. The pressure readings at 27" and 50' C. are as given in Table IV. From these pressure readings the quantities of nitrogen in the svstem were calculated, and are tabulated in Table V along with the Bunsen absorption coefficients. The absorption coefficients 'Onstant in agreement with Henry's law, and a rather are high solubility of nitrogen in sulfur hexafluoride is noted.,*
1129
LITERATURE CITED
(1) Lowry, H. H., and Erickson, W. R., J. Am. Chem. Soc., 49, 272934 (1927). (2) Schumb, W. C., IND.ENO.CHEM.,39, 421-3 (1947). (3) Schumb, W. C., Trump, J. G., and Priest, G. L., Ibid., 41,1348-51
(1949). RECEIVED October 12, 1950. Presented before the Fluorine Ohemistry Subdivision, Division of Industrial and Engineering Chemistry, Symposium on Fluorine Chemistry, a t the 118th Meeting of the AUERICANCHEXICAL SOCIETY, Chicago, Ill.
Gelation Times of Various Silica Sols
4
EFFECT OF LOW TEMPERATURES R. W. SPENCER, A. B. MIDDLETON, AND R. C. MERRILL Philadelphia Quarts Co., Philadelphia, Pa. Activated silica sols are becoming increasingly important as coagulant aids in the various phases of raw water and waste water treatment. A more thorough understanding of the fundamental reactions and the times involved in the formation of these sols is therefore important. Gelation times for mixtures of dilute sulfuric acid or ammonium sulfate with a dilute 3.3 ratio sodium silicate containing 1 to 69'0 silica at 25" and 8" C. are given. The influence of the concentrations of gelling agent and the sodium silicate solutions is investigated, The longer gel
times that are evident with the ammonium sulfate as compared to sulfuric acid are illustrated and discussed. Although the curves follow for the most part the expected relationship of increased gel times at lower temperatures, there are several conditions of pH and silica concentration in which this relationship is reversed. The new data should enable those concerned with the chemical treatment of water to utilize activated silica sols more fully and more efficiently. Improved flocculation and coagulation can be obtained when silica sols are prepared in accordance with the results here presented.
T
Hurd and his associates have done considerable work on gelation times for various concentrations of sodium silicates with hydrochloric, acetic, and sulfuric acids, but these have not been complete within the range in which the authors are interested (4, 6). Hay, in his patent on sodium silicate-ammonium sulfate mixtures, has also given some gel times (3). Because actual appljcation of coagulant aids frequently involves operation under near freezing conditions, the authors have determined the times required for gelation of sols made with sulfuric acid and ammoniumsulfate a t 8" C. These supplement previous data at 25" C. (IO).
HE increasing use of activated silica sols as coagulants and
h
coagulant aids in raw and waste water conditioning has made imperative a more thorough understanding of the time factors involved in the formations of the different sols. The work described in this paper has established the limits of time, temperature, and concentration within which satisfactory silica sols can be made without the attendant fears of gelation taking place. In 1936 Baylis a t Chicago developed a process for making silica sols based on partial neutralization of a sodium silicate (Naz0.3.3Si02)with an acid (1, ,$). The authors' present interest in this application of sodium silicates is an extension of these first successful efforts of Baylis in the use of silica sols as an aid in water coagulation. The term "activated silica" designates a negatively charged colloidal particle formed by the reaction of a dilute sodium silicate solution with a dilute solution of an acidic material. This process generally involves the dilution of a concentrated commercial sodium silicate solution containing about 8.9% sodium oxide and 28.7% silica with a specific gravity of 41.0" Baum6. This silic a b is diluted with water until the silica concentration reaches about 2%. Then the addition of a solution of an acidic material such as sulfuric acid, ammonium sulfate, chlorine, sodium bicarbonate, etc., neutralizes a part of the alkali and further reduces the silica content. Aging for 1.5 to 2 hours a t this concentration permits the growth of the silica micelle. By addition of water to bring the silica concentrations down to about 1%, further growth is then essentially arrested and gelation is prevented with resultant increased stability of the sol. Factors other than gelation which influence the commercial appliqations of activated silica sols have been discussed (8,9).
EXPERIMENTAL PROCEDURE
All gel times were determined by adding a dilute solution of the gel-forming reagent to diluted sodium silicate solutions, mixing well, and allowing the mixture to stand in a &ounce bottle kept a t either 8" or 25" C. until the mixture gelled. The mixing time was a matter of seconds and in all cases was considered to be a negligible part of the gel time. Gelation was determined by 109s of uniform fluid flow, the appearance of breakage planes when the mixture was tilted, and the adherence of solid gel t o the glass wall. The gel times obtained in this manner are reproducible within 2y0 and should correspond fairly closely to the "time of set" determined by the "tilted rod" method used by Hurd (6) and his associates. All acids and the ammonium sulfate used were reagent grade chemicals and all water used in diluting the silicate and acids was distilled. The silicate used in all cases was the N sodium silicate manufactured by the Philadelphia Quartz Go., containing S.9y0 sodium oxide, 28.7% silica, and 62.4% water.
