ilitv of J
LIQUID CA
N DIOXIDE
HOSMER W. STOKE University of California, Los Angeles, Calif.
The solubility of water in liquid carbon dioxide is reported for several temperatures from -29" to f22.6" C. These temperatures involve a vapor pressure range for liquid carbon dioxide of 15 to 60 atrnospheres. The values reported indicate a fairly regular increase in solubility from 0.02 per cent water at the lower temperature to 0.10 per cent at the higher.
OLUBILITY was determined directly by the apparatus shown in Figure 1. Although not indicated in the diagram, the carbon dioxide cylinder was immersed in a thermostat and the sodium-calcium hydrate t,ubes were placed in a water bath at room temperature. T o charge the cylinder, a 25-ml. portion of distilled water was added to the clean cylinder; and after replacing the valve carryin the siphon tube, the cylinder was filled two thirds full with liquicf carbon dioxide. To ensure thorough mixing of the two liquids, the cylinder was placed on the rollers of a ball mill and revolved for 8 hours at about 80 r. p. m. The room temperature during this mixing was above that at which the solubility was to be determined. Next the cylinder was placed upright in a thermostat in which the temperature was controlled within 10.05O C. The water of the thermostat covered the needle valve as well as the cylinder to take care of the cooling resulting from the change of state and the drop in pressure which took place at this point. After allowing 24-48 hours for the establishment of equilibrium a t the temperature of the bath, the cylinder valve was opened fully. The needle valve was then opened gradually so that the gas, when bypassed through the flowmeter to adjust the rate, attained a rate of about 100 ml. per minute. After permitting sufficient flow to purge the siphon tube of any possible droplets of entrapped water, the stopcocks were adjusted t o allow the gas to enter the absorption system. The magnesium perchlorate tubes were weighed filled with carbon dioxide while the sodium-calcium hydrate tubes contained dry air for both t,he initial and the final weighings. The completeness of the absorption of water by magnesium perclilorate and of carbon dioxide by soda lime was assured by the second tube of each of these reagents. The fact that the second tube showed little or no increase in weight during a run was eviaence that neither the rate nor the amount of sample had been too great for the capacity of the absorbents. One charge of the perchlorate was adequate for many runs so that the final weighing of one run was used as the initial weighing for the next. The first sodium-calcium hydrate tube received a fresh charge for each determination, since each run was continued until the appearance of the reagent showed that the Diagram of Apparatus 1284
HE composition of certain phases of the equilibrium mixtures of carbon dioxide and water, such as the solubility of carbon dioxide in liquid water, have been determined with satisfactory precision by a number of workers, especially Weibe (3-6); however, considerable uncertainty exists with regard to the solubility of water in liquid carbon dioxide. Lowry and Erickson ( 2 ) attempted to determine this from the density of the gaseous phase coexisting with the liquid carbon dioxide in the presence and absence of water and then computing the solubility b y Raoult's law. Since the differences they observed were less than the experimental error, they concluded that water could not be soluble to more than 0.05 per cent in liquid carbon dioxide over a temperature range of -5.8' to 22.9" C. The solubility is a matter of practical interest in the manufacture of commercial solid carbon dioxide (dry ice) where too much moisture often causes a freeze-up in the plant due to the accumulation of solid water. On the other hand, a certain amount of moisture has been regarded by some as essential to the production of dry ice with suitable physical properties. Without water the dry ice was thought to shatter easily, particularly after storage. Josephson (1) took out a patent covering the deliberate introduction of water for the purpose of improving the physical
A
KEEDGE V A L V E
Figure 1.
properties of dry ice. Although the desirability of any moisture in dry ice is doubtful, it is a point that is frequently raised in connection with the production of solid carbon dioxide. When carbon dioxide is used in certain types of fire extinguishers, the presence of too large an amount of water may cause serious trouble. Pressure drops in the carbon dioxide system may produce refrigeration resulting in the separation of solid water. This in turn may restrict the flow of the carbon dioxide. Hence the question of the solubility of mater in the liquid carbon dioxide becomes of vital importance.
S
INDUSTRIAL A N D ENGINEERING CHEMISTRY
December, 1943
1285
The results listed as D were obtained by the method of anltlysis OF WATERIN LIQUID CARBON DIOXIDE used for C, but the sample was liquid carbon dioxide taken from TABLE I. PERCENTAGE SATURATED AT VARIOUS TEMPERATURES . the National Dry Ice Plant, Niland, Calif., where operating condiSample, Grams 49.42 72.21 58.34 94.13 71.97 67.73 63.93 67.12 73.58
Water, Grams
mater, Per Cent
A. At 5.08' C. 0.0301 0.0609 0.0476 0.0659 0,0365 0,0626 0 0587 0.0624 0.0465 0.0646 0.0442 0 0653 0.0413 0.0643 0,0402 0,0599 0.0499 0.0678 Average 0 . 0 6 3 9 Average deviation from mean 3 . 2 %
Deviation from Mean - 0 0030 +0.0020 - 0 0013 - 0 0015 + O 0007 $ 0 0014 +O 0004 - 0 0040 +O 0039 0 0020
E.
118.63 99.41 91.37 8 8 . 66 89.22
31.8 14.2 13.3 25.4 11.7 67.7
547.3 199.5 306.0 535.7
At 15.0° C. 0.0838 0 0894 0.0899 0 0817 0.0894 0 0811 0 0914 0 0861 0.0954 Average 0 , 0 9 0 0 Arerage deviation from mean 3.0% 0 0994
C. A t 2 2 . 6 =+=lo C . 0.0316 0.0994 0.0155 0.109 0.0130 0.0980 0.0231 0.0910 0,142 0.121 0.0695 0.103 Average 0 . 1 0 4 +4verage deviation from mean 7.5%
D. A t -29OC. 1121 0,0205 0363 0.0182 0631 0.0207 0991 0.0186 Average 0.0195 Average deviation f r o m mean 5.1% 0 0 0 0
0062 -0 0001 -0
0.0006 - 0 0014
+ O ,0054 0 0027
-0.0046
+ O 0050 - 0 0060 -0.0130 +O 0710 -0.0010
0.0078
+0.0010
-0.0013 +0.0012 - 0 0010
0.0011
L
reaction had progressed nearly t o the exit end of the tube. The third sodium-calcium hydrate tube was included as a protection for the system against the water and carbon dioxide of the atmosphere. A sulfuric acid bubbler following the third soda lime tube showed a slight surge of pressure as the carbon dioxide was turned into the system at the start of a run. No further bubbles were observed during the run, a n indication t h a t practically all of the sample was absorbed by the two reagents. A single determination required a steady flow of gas from several hours to overnight. The sum of t h e u-eights of the water and carbon dioxide absorbed was taken as the weight of the sample. The low percentage of water made difficult what otherwise would have been a simple determination. Each soda lime tube and contents weighed about 600 grams, and was capable of absorbing about 100 grams of carbon dioxide. The heavy-duty analytical balance used t o weigh these tubes had a capacity of 2 kg. and a sensitivity of 0.1 mg. The 0.1 mg. sensitivity was not used with the carbon dioxide because the data for the water included only three significant figures. The magnesium perchlorate tubes and contents weighed approximately 100 grams each and were weighed on a balance with a capacity of 200 grams and a sensitivity of 0.1 mg. The method was developed after many less successful attempts had been made with various modifications.
ABLE I, A and B, records the results with the method just described. The values listed under C and D were obtained in a different manner. I n the case of C the entire experiment from the rolling of the cylinder t o the absorption of the water was carried out in a room t h a t showed a temperature variation of * l o C . The weight of the sample was determined by the loss in the weight of the cylinder. Since the cylinder weighed some 12 kg., the sample was weighed only t o within a few tenths gram. After the water had been absorbed by the magnesium perchlorate, the carbon dioxide was vented t o the atmosphere, instead of being absorbed by soda lime as above.
tions indicated t h a t the liquid carbon dioxide was saturated with water at -29" C. Since the author had no facilities for temperature control at such low temperatures, the values from this analysis are included as an indication of the trend of the solubility.
EVERAL features of the experimental method should be considered. These include the fact t h a t the carbon dioxide liquid and gaseous phases in equilibrium with water a t the same temperature and pressure contain different percentages of water, the Possibility t h a t droplets or particles of water form stable suspensions in the liquid carbon dioxide, the effect of the iron dissolved from the cylinder walls by the water-rich phase, t h e refrigerating effect resulting from the change of state and the expansion at the needle valve, and the disturbance of the equilibrium resulting from the increase in volume of the gaseous phase as liquid carbon dioxide is drawn from the cylinder. The fact t h a t the liquid phase contains a different percentage of water from the gaseous phase made it necessary t o take liquid carbon dioxide for analysis rather than samples from the gaseous phase. Analysis of the equilibrium mixtures of carbon dioxide in the two states with water showed t h a t the gaseous phase contained about half the percentage of water t h a t the liquid phase contained for the particular temperature and pressure. The possibility t h a t droplets of water or solid particles of water at low temperature might remain suspended in t h e liquid carbon dioxide appears t o be very real. The fact t h a t the method gives a somewhat regular decrease in solubility of water in liquid carbon dioxide with a decrease in temperature could hardly be explained if the suspension of water in the liquid carbon dioxide were a t all stable. Hence, it has been assumed from a consideration of the densities of water, ice, and liquid carbon dioxide t h a t , where water is in excess in the liquid state, the carbon dioxide liquid floats on top; at most temperatures a t which the water is in t h e solid state, the water-ice floats on top of the liquid carbon dioxide. The presence of iron in the water-rich phase was discovered when clear water which remained in the cylinder after the carbon dioxide had been exhausted was exposed t o the atmosphere in a n open beaker. On standing for some time open t o the air, a brown gelatinous precipitate characteristic of hydrated basic iron appeared. If the carbon dioxide was swept from the clear solution by a rapid bubbling of air, the precipitate appeared a t once. One of the cylinders was copper-plated on the inside t o avoid this difficulty. Later the mqtter was disregarded when the applicability of the findings of Weibe and Gaddy (4) was recognized. They reported the probable formation of an iron bicarbonate and did quantitative work which indicated t h a t the amount of iron dissolved under these circumstances was small. The error resulting from such small amounts of dissolved iron on the vapor pressure of water and hence its solubility in the liquid carbon dioxide is certainly of an order well outside the experimental limits of the work reported here. H E N liquid carbon dioxide changes t o the gaseous state wat the needle valve with a drop of pressure to that of t h e atmosphere, there is a refrigerating effect which might conceivably retain some of the water near the valve. Because the capacity for water of the gaseous carbon dioxide is much greater a t one atmosphere than t h a t of the liquid carbon dioxide at the relatively high pressure within the cylinder, because the valve was immersed in a thermostatically controlled water bath, and because the rate of flow was kept low (about 100 ml. per minute for the gas), it appears unlikely t h a t this refrigeration effect resulted in incorrectly low values for the percentage of water found. Experimental variations in the rate of gas flow from 50 t o 200 ml.
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I N D U S T R I A L A N D E N G I N E E R I NG
per minute were tried and gave no correlation with the deviation of the results. When the valve was not immersed in the water of the thermostat a n d the rate of gas flow was increased to 300 ml. per minute, moisture was seen to condense in the outlet tube from the valve. No such condensation was ever observed under the conditions a t which the determinations were carried out. The question of the disturbance of the equilibrium between the water-rich liquid phase and the carbon-dioxide-rich liquid phase was raised by the fact that, as liquid is drawn from the cylinder, the volume of the gaseous phase was increased. Since this phase is poorer in water a t the 15-60 atmospheres range of pressure in the cylinder, the gradual loss of a greater percentage of carbon dioxide than of water to the gas phase might result in a supersaturation of the liquid carbon dioxide by water. Shaking the cylinder between determinations was tried with a view to minimizing this difficulty. However, this entailed the removal of the cylinder from the bath which upset the thermal equilibrium and introduced the possibility of entraining droplets of water in the siphon-valve part of the system. Since no noticeable change appeared in the values for the water content of the liquid carbon dioxide when this was clone, much of the work was
CHEMISTRY
Vol. 35, No. 12
carried out without any other agitation than that provided by the vibration of the stirring motor of the bath. Also, the sample entered the siphon a t a point close to the layer of water and at some distance from the interface between the liquid and the gaseous carbon dioxide; this fact. likcwisc mitigated against tlie possibility that the sample of liquid carbon dioxide might become supersaturated with water as a result of the increasing volume of the gaseous phase. It is realized that the values here reported leave much to be desired, but in view of the difficulties involved and the inadequacies of the existing published data, the results appear to be of fundamenhl importance. LITERATURE CITED (1)
Josephson, W. S., U. 8. Patent 1,873,131 (1932): R e f T i g . E n y . , 19
25-6 (1930). ( 2 ) Lowry, H. H., and Erickson, W.E., J . Am. Chem. Soc., 49, 272934 (1927). ( 3 ) Weibe, R., Chem. Res., 29, 475-81 (1941). (4) Weihe, R., and Gaddy, V. L., .J. A m . Chem. Soc., 61, 315-18 (1939). ( 5 ) I b i d . , 62, 815-17 (19403. ( 6 ) Ibid.. 63, 478-7 (1941).
Properties of Detergent Solutions THERMAL pH COEFFICIENTS OF ALKALINE SOLUTIONS LESTER E. IWENTZEL, JAMES W. HENSLEY, AND LESLIE R. BACON Wyandotte Chemicals Corporation, Wyandotte, Mich.
T
HIS paper, the fourth of a series ( 1 , '7, 8) concerned with the properties of detergent solutions, presents detailed p H data at 2 5 O , 40°,and 60" C.on distilled water solutions of nine commercial alkalies sometimes used alone or in combination as laundry soap builders over concentration ranges of interest in laundry practice, Thermal p H coefficients, ApH/ AT, for these builder solutions have been presented which permit the calculation of p H values a t elevated temperatures from experimental data obtained at room temperature. The alkalies examined in the present work are: sodium hydroxide, sodium carbonate, sodium bicarbonate, trisodium phosphate, tetrasodium pyrophosphate, sodium tetraphosphate, sodium metasilicate, sodium sesquisilicate, and sodium orthosilicate. I n general, the technique and equipment mere the same as that described in the first paper ( 1 ) . Only electrometric methods were used, however. Some improvements in technique and apparatus design were made to secure greater accuracy and ease of manipulation. Both hydrogen electrode and glass electrode measurements were made on the same solution in most of the determinations. As a rule there was close agreement between the two measurements except for values obtained from solutions a t elevated temperatures and very high p H where the sodium ion error was considerable. iMETHODS AiVD APPARATUS
Hydrogen electrode potentials were ELECTRICAL EQUIPMENT. measured by a simplified Queen potentiometer with which readings could be made t o the equivalent of less than 0.01 p H unit. Glass electrode potentials were measured by a Leeds & iYorthrup
potentiometer-electrometer 7660-A which was calibrated to 0.002 volt (0.03 p H unit) and could be read accurately t o the equivalent of the nearest 0.01 pH unit. National Technical Laboratories No. 1190E glass electrodes were used. These electrodes are especially adapted for use in alkaline solution and have very low sodium ion errors. All connecting electrical leads were well shielded. PLATINUM ELECTRODES. 5 X 10 mm. electrodes were cut from platinum foil, mounted in the end of sealed glass tubes, and palladinized. A careful comparison of results obtained from both platinized and palladhized electrodes revealed negligible differences as long as the palladium coat was freshIy formed. Upon continued use, a difference became noticeable and it must be assumed that the palladinized electrodes are in error; however, since they were much easier t o clean and replate, they were used and freshly plated before each measurement. To secure uniform and reproducible coatings of palladium black, an apparatus was constructed whereby the cleaning and plating operations could be closely controlled. A current density of 160 milliamperes per sq. cm. applied to the electrodes for one minute while being electrolyzed cathodically in a 3 per cent palladium solution in hydrochloric acid produced a light, satisfactory coating. The palladium solution was prepared by electrolyzing a palladium wire in 1 to 1 hydrochloric acid until the proper weight of palladium had dissolved. The solution was mamtained at approx7mately 3 per cent concentration by using a palladium wire anode for the plating operation. ELECTRODE VESSELS. The vessels housing the electrodes and containing the test solution and saturated potassium chloride solution were made of Pyrex (Figure 1). Tubes K and holding calomel electrodes C and C', wet stopcocks S and S , and 9.11 connecting capillary tubing t o liquid junctions J and J ' , are filled with saturated potassium chloride solution to which has been added a small amount of ca1om:l. Tube LW containing duplicate hydrogen electrodes H and H , and tube AT containing inter-
F',
