hlISLEADIKG EQUILIBRIA ENCOUNTERED IX THE MEASUREhlENT O F DISSOCIATION PRESSURES I N SALT-HYDRATE SYSTEMS BY ALAN W. C. MESZIES AND C. S. HITCHCOCK
It is only too obvious to those interested that, while the number of published experimental results in this field increases year by year, there is a regrettably large degree of discordance between the values found by different workers. It is hoped that the present communication may be of use in tending to remedy this unfortunate situation. The errors caused by the neglect of adsorbed permanent gases which are so difficult t o remove from static tensimeters are now better appreciated than in the days of Frowein, faith in whose otherwise excellent results was the cause of much interesting if nugatory speculation as to the cause of the high values yielded by the gas-current saturation method. This difficulty disappeared when it became apparent' that the tensimetric results of Frowein in question were themselves too low. Apparently insignificant points of experimental technique, such as the use of glass wool in an absorption train, have been shown capable2 of producing considerable discrepancies. Profiting by this and other constructive criticisms offered by Menzies, Partington and Huntingford3 repeated the earlier work of Partington4 on the aqueous pressure in equilibrium with the system CuS04, 5-3Hs0 and obtained new values which were markedly higher. In the present article a rapid method of measurement will be illustrated and shown t o be applicable in difficult cases; and examples will be adduced to show that a genuine equilibrium pressure is sometimes reached which, however, is that of the vapor of water merely adsorbed upon the salts rather than an equilibrium dissociation pressure. Methods. Two methods were employed: ( I ) the well-known gas current saturation method; and ( 2 ) a method whose novelty makes necessary some words of explanation. It has often been reported that equilibrium between water vapor and a salt hydrate pair is reached more quickly if the two solid phases are in contact with a liquid in which water is a t least slightly soluble; and this gain in speed is sought in several of the so-called indirect methods of dissociation pressure measurement which utilize solutions of water in such liquids as ether, ethyl alcohol6 or isoamyl alcohol.7 After the solid phases have been 'Cf. Menzies: J. Am. Chem. SOC., 42, 19j1 (1920). Menzies: J. Am. Chem. SOC.,42, 978 (1920). a J. Chem. SOC.,123, 160 (1923). J. Chem. SOC.,99, 467 (1911). 5Linebarger: Z. physik. Chem., 13, 500 (1894). 6 Foote and Scholes: J. Am. Chem. SOC.,33, 1309 (1911). 7 Wilson, Pl'oyes and Westbrook: J. Am. Chem. SOC., 43, 704, 726 (1921). 2
DISSOCIATION PRESSURES IN SALT-HYDRATE SYSTEMS
1661
brought to equilibrium with the liquid, the water concentration and aqueous partial pressure of the latter are then arrived a t by independent experiments. For the work which follows, we wished to employ such a liquid in the hope of accelerating attainment of equilibrium; but we wished also (a) to be able to follow readily by the eye and without analysis the development of aqueous pressure in the system; and (b) to be able to operate far above room temperatures but without the inconvenience of a cumbersome high temperature tank thermostat. To meet these conditions, we found the method briefly outlined in the preceding article .would serve admirably. With this apparatus, also, low pressures may be directly measured in terms of a column of liquid much less dense than mercury. I n general, dissociation pressure equilibria of such salt hydrate pairs as yield only a low pressure are attained only slowly, and it is especially in such cases that the results of different investigators are apt to show discrepancy. For this reason the case following was suitable for study. I. BaC12, 1.8 H 2 0 i n Chloroform. Equilibrium in the system BaC12 2-1 H 2 0 and vapor is universally reputed to be difficult of attainment.’ Baxter and Cooper, whose work is the more to be trusted because they were well aware of the difficulties, have measured this dissociation pressure at I so, 2 5 ’ and 40°, and the three-constant equation which they fit to these observations yields 56.7 mm. for the pressure a t 60.1’~ which was the boiling point of chloroform under the existing barometric pressure, as used by us for our method ( 2 ) . Our sample of dihydrate, Merck’s “reagent” grade, was effloresced in an oven a t 5s’ to a composition BaC12, 1.8H20. Since the expected pressure would have required a column of chloroform to balance it longer than our apparatus is tall, we employed a few cc. of mercury in the inner tube of the apparatus to serve as the manometric liquid. The manipulation and corrections have already been referred to? The density of chloroform a t its normal boiling point was taken as I .4 I 0. Equilibrium was reached within half an hour, duplicate experiments . ~~ 7 mm. , ~ respectively at 60.1’. Our temperature measureyielding ~ 7 and ment, by a Reichsanstalt certificated thermometer whose zero point had recently been redetermined, is good only to the nearest one-tenth degree. The average of these results, j j . 5 , may well be in error by one millimeter, and is, therefore, in good agreement with the extrapolated value of Baxter and Cooper. A correction, 0.3 mm., due to the lowering of the vapor pressure of chloroform within the inner tube by the water dissolved in it may be computed, assuming Raoult’s law, from the following data: solubility of water in chloroform at 60°, 0.8 percent? vapor pressure of water saturated with chloroform, about 149 mm. Because information is lacking on the effect of small mol I Cf. Baxter and Cooper: J. Am. Chem. Soc., 146, 923 (1924);Schumh: 45, 342 (1923); Partington: J. Chem. SOC., 99, 49 (1911). See preceding article J. Phys. Chem., 35, 1655 (1931). Extrapolated from values in I. C. T. 3, 387.
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ALAN W. C. MENZIES AND C. S. HITCHCOCK
fractions of water on the vapor pressure of chloroform, this correction was not applied, but the calculation serves to show that the value of the true correction is presumably small. It is evident that, in this case, attainment of the true dissociation pressure presents no difficulty under our experimental conditions. We had guarded against the adsorption of water by the salts in handling, and no ambiguity was likely as to the hydrates present. A more difficult case was accvdingly selected. 2. CuS04 1-0 HzO. The measurement of the equilibrium pressure for this system has presented such difficulty to various workers that it is not a t all obvious what the true p-t relation really is. We have arrived a t this in the following manner. The three independent investigations of Bell and Taber,' Footel and Crockford and Warrick8 have reached reasonably concordant values for the concentration of aqueous sulfuric acid with which these two solid phases are simultaneously in equilibrium. If we average these values for 25' we obtain a concentration of sulfuric acid about 87.8percent. If now we suitably graph against concentration the logarithms of the aqueous pressures over sulfuric acid solutions at z 5' as given in International Critical Tables, we find a pressure for 87.8 percent nearly 0.017mm. of mercury. Siggel,4 in Nernst's laboratory, in an effort to improve upon the work of Schottky,&
TABLE I Dissociation Pressures in the System CuSO4H20 Observers
Averaged from three Siggel 1, 1) 1,
!f
Pressure mm. Hg Obs'd. Calculated Siggel M. and H.
Temp. OK.
0.0~7
298. I 372 383.2 403.9 410.2 420.4
6.0 12.0
38.2 53.5 90.4
0.019 6.04 11.92 37'9 52.7
88.0
[o.oI~] 5.96 11.9 38.6 53.9 90.6
Average of preceding five values all taken as positive Algebraic sum of the five differences Schottky 363.6 4.6 3.46 Foote and 298. I 0.8 0.017 Scholes Dover and 298.1 1.3 0.017 Marden Miiller298. I 0.5 0.017 Erzbach
* J. Phys. Chem., 12,
(1908). J. Am. Chem. Soc., 37, 288 (191 j ) 171
J. Phys. Chem., 34, 1064 (1930). 2. Elektrochemie, 19, 340 (1913). 2. physik. Chem., 64, 415 (1908).
CuSO4
+ HzO
Difference in mm. Obs'd minus Calc'd Siggel M. and H. -0.002
0.0
-0.04 +0.08
+o.o4
+0.3
+0.8
-0.4 -0.4
+2.4
-0.2
0.72
+3.54
+O.I
0.23 -0.86 +I.14
$0.78 +I28 +0.48
DISSOCIATION PRESSCRES IN SALT-HYDRATE SYSTEMS
1663
an earlier pupil of Sernst, determined the dissociation pressure of this salt system a t five temperatures from 99’ to 147.4’. Siggel proposes a threeconstant equation to average these results. If, however, using the above datum for 2 5 ’ and an average of Siggel’s data we find the constants in a twoconstant equation (E) of the form log p = A - B ‘T, where A evaluates as 11.040 and B as 3818.2, we are able to average Siggel’s results considerably better than he does himself, as is shown in Table I. We have given greater weight to his two lower temperature values, in which Siggel places most trust. To this table we have appended also the values of certain other observers. The value of Foote and Scholesl was obtained indirectly by equilibration with aqueous ethyl alcohol. That of Dover and 11Iarden2by the transpiration method, purports to give only the apparent pressure under their experimental conditions. (a) Gas Current ;Method. Anhydrous cupric sulfate and the monohydrate were obtained by heating the pentahydrate at about 240‘ and 100’respectively. Approximately equal quantities of the two salts, in granules of an average dimension of four mm., were mixed and charged into two “saturators” tubes each of j cm. diameter and 30 cm. length. These were immersed in a thermostat a t 40’ and connected by ground glass joints with two water absorption tubes containing anhydrous magnesium perchlorate, situated outside the thermostat. The weight of the second of these absorption tubes remained sensibly constant. In all, some I joo liters of air freed from moisture by means of magnesium perchlorate were passed through the train a t this temperature at a rate near 1% liters per hour, causing the copper salts to lose some 4 2 0 mg. of water. The equilibrium pressure expected, according to equation (E), is 0.070 mm. The pressures of water found from the observed gain in weight of the absorption tubes for a metered volume of air ranged from ~ . j mm. r progressively downward to 0.08 mm., at which point the bath temperature was raised to jo.0’. At this temperature the anticipated pressure is 0.167 mm. During the passage of j60 liters, the pressure measured decreased from 0.18 smoothly through the dissociation equilibrium value to 0.11 mm. when the experiment was discontinued. In this instance, the dynamic equilibrium of dissociation is apparently very slow in both directions. During a pause in the work following the experiments at 40°, the saturator tubes had been left closed and inactive at room temperature for three summer months. The aqueous pressure within the tubes, however, remained five-fold higher than the dissociation pressure. At any particular epoch during the course of these experiments, measurements by any static method of the vapor pressure of water as yielded by this material would obviously have given concordant results. But the pressure measured would have been that of adsorbed water and not a true dissociation pressure. In answer to the question how the copper salts in the saturator tubes acquired the water which they later yielded up as described, a t least two J. Am. Chem. Soc., 33, 1309 (1911). J. Am. Chem. Soc., 39, 1609 (1917).
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ALAN W. C . MENZIES .4ND C. S . HITCHCOCK
replies might be made: (I) that this amount of water in the form of a thin film was in equilibrium with the aqueous pressure in the drying oven at IOO', for example; or ( 2 ) that the water was picked up a t room temperature from the air of the room during the process of charging the absorption tubes. I t must be recalled that the copper salts employed resulted respectively from the removal of 4 or j molecules of crystalline water from each "molecule" of pentahydrate of cupric sulfate and that the new specific surface so produced must be very large. A quantity of a few grams of gross analysis CuSO.,, 0.9 H10 was found to gain weight on standing in the balance case containing ordinary room air a t the rate of z percent per hour. In those cases where the specific surface of the salts is large, where the dissociation pressure to be measured is low and where the rate of the chemical reaction is slow, the opportunity for error due to adsorbed water becomes most favorable. (b) CuSO4, 0.9 H 2 0 , in Benzene. The solid was prepared by removal of water from the pentahydrate by prolonged heating in a drying oven a t I 10' and was presumed to be a mixture of CuSO4, H?O and CuS04. Adsorption of water during handling was guarded against. At' So0, which is near the normal boiling point of benzene, the expected dissociation pressure for the system is 1.68 mm. of mercury, according to equation (E). The pressure found by method ( z ) , using benzene, was 1.6 i 0 . 1 mm. This was reached with difficulty within an hour, the speed being increased by shaking and by use of larger quantities of the salts. Otherwise several hours might be consumed. The error due to the effect of the dissolved water upon the vapor pressure of benzene, computed as indicated in the preceding article, would in this case reach only 0.006 mm., and is therefore negligible. 3 . CuS04,8.9 H 2 0 , by Dehydration, zn Chlorojorni. The systems composed of vapor and cupric sulfate with five and three and with one and three molecules of water have been studied by Carpenter and Jette' in the ranges 2 jo-90' and 25'-80" respectively. While an examination of their results, which will be discussed elsewhere, points to their lower temperature values for the second system being too high, due, doubtless, to adsorbed water, their values a t 6 0 . 0 ~which ~ by interpolation we find near 83.4 and 58.8 mm. respectively, are probably not far from the truth. As a desirable antecedent to our next experiment, we made two rough trials, using method ( 2 ) with chloroform, with a solid of analysis corresponding to CuSO4, 4.5 HzO, and ob~ I .j mm. I t may be recalledY tained a dissociation pressure at 60.0' of 8 0 . & that one molecule percent of the isomorphous sulfates of manganese, zinc and magnesium in solid solution lowers the dissociation pressure of cupric sulfate pentahydrate by 8.6, 6 . 2 and 2 . 3 percent respectively. Our copper salts gave negative tests for these and other metals. We then prepared a sample of water content corresponding to 2 . 9 H2O by allowing another portion of the material used in experiment (2b) to take up 2
J. Am. Chem. SOC.,45, 578 (1923). Cf. Hollmsnn: Z. phgsik. Chem., 37, 193 (1901).
1665
DISSOCIATION PRESSCRES IN SALT-HTDRATE SYSTEMS
water from the air of the laboratory over a period of two days. Using the same method, pressures a t 60.0' of 7 3 . 2 and 71.0 mm. were obtained in successive experiments, pointing to a slow disappearance of adsorbed water. What is here being measured is clearly not the true equilibrium pressure of either of the systems mentioned, but that of the less hydrated system as vitiated by the presence of adsorbed water whose effect in this case disappears but slowly.
Summary The method very recently proposed by hlenzies is shown to yield, often within an hour, true results for the dissociation pressure of certain salt hydrates, including the difficult cases of BaCl2zH20 BaC12H20 H20 and CuS04H20 % CuSOl H20. 2. A linear equation is proposed by which to represent the pressuretemperature relation in the last-named system. 3. It is shown that, in salt hydrate systems, equilibrium pressures may be attained that are not true dissociation pressures, but rather pressures due t o adsorbed water. The incidence of this source of error is discussed. I.
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Pniiceton Unwerszty, January 1991.
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