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FRANK R. MEEKS,RAMGOPALAND 0. K. RICE
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times larger for the CO-02 flame. This explains the existence of two distinct quenching limits for the two systems if one assumes that the minimum plate diameter for effective quenching is related to the minimum ignition energy, not the queiichiiig distance. This is probably true since the minimum ignition energy of any particular system is related to the flame thickness and, therefore, to the distance a flame must travel before it settles down to a steady state. This evidence points to the existence of a very thick flame zone in the CO-O2 flame.
velocity of a pure dry stoichiometric CO-02 mixture is extremely low-certainly less than 3 em./ sec. We also found that other limiting values for dry mixtures were : ignition energy, greater than 0.5 joule, and quenching distance, greater than 0.4 cm. The data further indicate that the burning velocity is best correlated to hydrogen atom diffusion ahead of the flame. However, a slight contribution from an hydroxyl radical reaction in, or ahead of the flame has not been eliminated by this study. Ignition energies and quenching distances show no effect of isotopic substitution when comConclusions pared to each other and flame velocity. The data Pure dry stoichiometric carbon monoxide-oxygen indicate that hydrogen atom diffusion and deactivaflames have been studied with added hydrogen or tion at the wall is not important to the quenching deuterium. The driest mixtures contain less than process. Our data indicate that the CO-02 flame 10 p.p.m. of equivalent hydrogen before Hz or D2 is much thicker than a hydrocarbon flame with the addition. The data indicate that the burning same burning velocity.
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CRITICAL PHENOMENA IN THE CYCLOHEXANE-ANILINE SYSTEM : EFFECT OF WATER AT DEFINITE ACTIVITY1 BY FRAXK R. MEEKS,RAMGOPALAKD 0. K. RICE Department of Chemistry, University of North Carolina, Chapel Hill, North Carolinu Received August 18, 1968
Water a t controlled activity has been introduced in the cyclohexane-aniline system by allowing the solution to remain contact with a mixture of Li2S04and LieSOd.He0. The critical temperature is raised about 0.3". The shape of thc coexistence curve is altered, and the flat portion a t the critical temperature has largely, if not completely, disappeared. ill
Introduction Critical phenomena in the cyclohexane-aniline binary liquid system and the shape of the coexistence curve in the vicinity of the critical point have been investigated by Rowdeii and Rice2 and Atack and Rice.3 It was shown that the coexistence curve is flat and horizontal within experimental error over a range of aniline mole fraction from 0.43 to 0.465, in work for which the temperature control was 0.001O. Small amounts of water raise the separation temperature, and careful drying over calcined CaO was necessary to get reproducible results. Atack and Rice4 made preliminary measurements on the effect of water, and suggested that it is negatively adsorbed at the iiiterface. In the present work we have examined the effect of introducing water at a controlled activity, by adding a mixture of approximately equal weights of lithium sulfate monohydrate and anhydrous lithium sulfate to each sample and observing the new transition temperature. To find the actual coiicentration of mater, further experiments on the introduction of small, known amounts of water into samples of varying composition will be necessary, and mill be reported lat'er. (1) Work supported by the Office of Ordnance Research, U. S. Ariiiy.
(2) R . W. Rowden and 0. K. Rice, "Changeiiients de Phases," Societe de Chimie Physique, Paris, 1952, p. 78; J . Chem. Phvs., 19, 1423 (1951). (3) D. Atack and 0. K . Rice, i b i d . , 22, 382 (1954); 0. K. Rice, i b i d . , 23, 164 (1955). (4) D. Atack and 0. K. Rice, Disc. Faradag SOC.,16, 210 (1953).
Experimental Aniline was purified by distillation in air, followed by vacuum distillation using a drying tube and aspirator, recrystallization and drying with CaO on a vacuum rack. Purified samples of approximately 20 cc. bearing no tinge of yellow color were stored in sealed ampoules. National Bureau of Standards cyclohexane, Sample No. 209a-25, was used without further purification in preparation of all mixtures but two. Baker and Adsmson Reagent Grade lithium sulfate monohydrate from a single bottle was used in all the work reported here. The lithium sulfate was insoluble in the cyclohexane-aniline mixtures; evaporation of a sample which had been in contact with the salt left no residue. Mixtures were prepared by measuring out the aniline at a known temperature into a tube with a pipet, freezing with Dry Ice-acetone, then measuring in the cyclohexane6; the opening in the top was sealed immediately with a torch. After thorough evacuation and degassing, the mixture was distilled successively into four tubes each containing several grams of fveshly calcined CaO and allowed to remain on the CaO for several hours at room temperature. After distillation into a sample tube provided with a breakseal, the sample was sealed off from the drying train. A correction (of 0.001 to 0.002 mole fraction aniline) was made for evaporation of cyclohexane into the volume above the liquid. A large, well-insulated thermostat was controyed by a specially designed thermoregulator to &0.001 . Temperatures were measured with a platinum resistance thermometer which had been calibrated against a similar N.B .S.-calibrated thermometer. The scale appears t," differ from that used by Atack and Rice3 by about 0.17 , but temperature differences are of most importance. The (5) Density, at 28O, of cyclohexane and aniline, 0.7710 and 1.0149, respectively. J. Timmermans, "Physico-Chemical Conatants of Pure Organic Compounds," Elsevier Publishing Co., New York, N. Y., 1950.
. D
June, 1959
EFFECT OF WATERA T DEFINITE ACTIVITY IN CYCLOHEXANE-ANILINE
bridge used had a small sensitivity to ambient temperature, but relative temperatyres of different samples were controlled to about 0.001 , as in previous work in this Laboratory, by frequently observing several samples simultaneously. Separation temperatures were observed for the pure systems, as described p r e v i ~ u s l y by , ~ lowering the temperature to a certain point, allowing it t o remain for a half-hour, then raising the temperature t o clear the opalescence in order to see if the phases had separated. The drying process was repeated until the separation temperature for the mixture remained constant from one drying to the next. The mixture was then distilled onto a 0.2-0.3 g. mixture of lithium sulfate monohydrate and anhydrous lithium sulfate, thoroughly degassed, sealed off, and the separation temperature determined. All transfers were done in a vacuum of the order of 10-5 mm., but it was ascertained that a half hour’s pumping on the monohydrate reduced its moisture content very little, and the hydrate was kept chilled as much as possible with liquid air. In the case of the samples with the hydrate, the tubes were allowed to st,and for two hours or longer at the separation temperature, rather than a half hour, before observing the separation, and many careful observations were made on each tube. The more detailed examination was dictated in part by difficulty in observing the appearance of the meniscus in the presence of the hydrate, especially for the cyclohexane-rich samples where the meniscus reappears near or a t the bottom of the tube. It is possible that the longer “settling” time allowed for the monohydrate samples, along with the more detailed investigation of single specimens, would permit the observation of differences in the separation temperatures of samples where such differences would be obscured by the shorter settling time, i.e., half an hour or so. The quest.ion of settling rate is a matter which remains somewhat indeterminate a t present. When tubes are left for severtl hours in a thermostat which can fluctuate by even 0.001 , however, there is no certainty that the temperature has not drift,ed farther downward a t some time during the process than is indicated by intermittent temperature measurements on the bridge. However, the fact that many measurements were made with several tubes observed simultaneously minimizes the likelihood of “drift” error. Sometimes there were erratic results in which either the change in temperature with addition of water was only a few hundredths of a degree, oroelse it was considerably larger than approximately 0.3 , the usual value. I n either case the values followed no discernible pattern but) were scattered. Aberrations of the former type might he attributable to too great a proportion of anhydrous salt in the mixture, and those of the latter to too small a proportion. In either case the sample was redried and a new mixture of hydrate and anhydrate was introduced. Numerous samples were rejected because their pureMystem separation temperatures could not, be brought to s low enough value to agree with the othe1.s. This criterion of rejection was applied equally to tubes of the previous worker and to the freshly prepared tubes of the present worker. (Btark’s 0.390, 0.448 and 0.457 had to be rejected hecause of yellowing; subsequent redistillation removed t8hetinge but they cont,inued to disagree both among t.hemselves and with other tubes. The Atack tuhes of aniline mole fractions, 0.466 und 0.445 were not available.)
Results ‘l’he results are giveii in Fig. 1. X A is the mole fraction of aniline, T is the separation temperat,ure of the pure system, and T’ is the separation temperature in the presence of the hydrate.6 Most (6) T h e temperatures for aniline mole fractions 0.4275 and 0.4435 h a v e been arbitrarily adjusted by subtraction of 0.037 from eacli T’ and T; the measnred values were higir-due presumably to their having been prepared with a sample of cyclohexane other t h a n t h e N.B.S. material. Of course these were not subject to coniparison with t h e others, but t h e adjustment is felt to be justified by t h e fact t h a t their temperature increments are in accord with those for other samples. The values of XA for Atack’s samples are slightly different t h a n previously reported,r because t h e space above t h e liquid was a little greater in our tubes, and the correction for evaporation of cyclohexane was correspondingly increased. Also t h e correction was
- 29.88 ‘
h
993
0
6
29.8;
29.GO
29.59
0
G29.58 29.57 29.56
9
1
1 I
4
9 l
l
0.39
0.40
l
i
i
0.41
A-*.
l
0.42
i
l
0.43
i
l
i l
0.44
Fig. 1.-Coexistence curves (black circles, Atack’s mixtures). The numbers under the circles give the last significant figure in the temperature. Thus, for the lowest point, T = 29.569.
I t I
i, 0.29
0.28
t 1
I
I
I
I
U I
\
0.41 0.42 0.43 0.44 XA. Fig. 2.-Effect of salt hydrate on transition teniperaturc (black circles, Atack’s mixtures). 0.31)
0.40
striking are the shift of the critical region of the upper curve to lower aniline mole fraction and the relatively narrower mole fraction range which it covers. If there remains any flat portion a t the top of tjhe coexistence curve its presence cannot be proved by these data. It should be noted, however, that the meniscus seemed to reappear withiii the tube, for samples with hydrate, from aniline inole fraction 0.417 to 0.436. It seemed to appear at the top for tubes more concentrated in aniline than 0.43G, and a t the bottom (as far as could be discerned, considering the presence of the lithium sulfate monohydrate in the bottom of the tube) for those of aniline mole fraction less than 0.417. The change in shape of the curve produced by a small amount of water (less than 2 X mole mole per mole of mixture, per cc. or 2 X according to Atack and Rice4) reemphasizes the caution which is necesssry in the interpretatioiz of curves of this sort; a flat top could be missed easily on account of impurities. In Fig. 2 we have plotted TI-T against mole inadvertently omitted by Atack and Rice* in the samples reported as mole fractions 0.4656, 0.4376 and 0.4296, and these values shor~ldbe increased by about 0.0010.
994
D. M. ALEXANDER
fraction of aniline; this diagram brings out the fact that a given activity of water has a greater effect on the transition point the richer the system is in the substance, cyclohexane, in which water is less soluble. Since the activity of a given concentration of water is expected to be greater the richer tjhe solution is in cyclohexane, the effect would be enhanced if one considered a given concentrat,ion rather than a given activity of water. A further thermodynamic analysis must await the completion of the experiments mentioned a t the end of the Introduction. The differences shown in Fig. 2 may be more accurate relative to each other than the individual measurements. If this is true, Fig. 2 practically eliminates the possibility of a flat top for the coexistence curve in the presence of the hydrate, for it is seen that almost all the horizonal portion of
Vol. 63
the curve in Fig. 2 lies beyond the end of the critical region of the dry cyclohexane-aniline mixture, which occurs3 a t about 0.425 aniline mole fraction. If the flat portion of the coexistence curve, indicating a range of critical concentrations, is caused, when it exists, by the vanishing of the interfacial tension while the two phases are still of different compo~ition,~ addition of water a t fixed activity can be expected to alter the critical temperature, for a t any given temperature there will be a definite amount of adsorption a t the interface, and hence a definite change in interfacial tension, corresponding to the controlled activity. But a gradual change in the activity of the water could not cause a sudden change, or vanishing, of the range of critical concentrations; this suggests, therefore, that these experiments should be repeated with a salt pair having a considerably lower aqueous tension.
A CALORIMETRIC MEASUREMENT OF THE HEATS OF SOLUTION OF THE INERT GASES IN WATER BY D. M. ALEXANDER’ Contribution f r o m the Department of Chemistry, University of Canterbury, Chri~~church, New Zealand Received September d , 1968
A microcalorimeter is described which has been used t o measure heats of solution of inert gases in water. Calorimetric values in kcal./mole of the standard enthalpies of solution of neon, argon, krypton and xenon are at 25’, -1.4 f 0.4, -2.9 f 0.2,-3.8 f 0.2 and -4.1 i 0.2, respectively.
Introduction The heats of solution of gases in water have in the past been calculated from the temperature variation of solubility.2 I n view of the experimental difficulties associated with accurate solubility measurements and the poor agreement between the results of different workers, it was thought that direct calorimetric measurements should be made of the heats of solution of the inert gases in water. No previous measurements of this sort appear to have been made with any gases apart from those of very high solubility, such as ammonia, hydrogen chloride and acetylene. The design of the calorimeter had to be such that the temperature rises of to could be measured with some precision. Such measurements can be made by using a Paschen galvanometer and a suitable thermopile.a It was necessary also that as much gas as possible should dissolve in the shortest possible time and that the heat produced by mechanical stirring be very small. It was essential that no solution should occur before or after the mixing process, and that the quantity of gas absorbed should be measured. Experimental The Calorimeter .-Twin Pyrex glass calorimeters were used, enclosed in an underwater air jacket. Each com(1) Chemistry Department, University of Queensland, Brishsne, Australia. (2) (a) D. D . Eley. Trans. Faraday Soc., 86, 1281 (1939); (b) H. 5. Frank and M. J. Evans, J . Chem. Phus., 18, 507 (1945). (3) F. T. Gucker, H. B. Picksrd and R. W. Planck, J. A m . Chsm. Sot., 61, 459 (1039).
prised a spherical bulb and cylindrical bottom bulb of equal volumes, the bottom of the top bulb and the top of the bottom bulb being connected by a narrow capillary and a ground glass joint. The top of the to bulb was connected to the top of the bottom bulb by capilfary tubing and a tap. The top bulb of one calorimeter was filled with the solvent, gas-free water and the tap closed. The bottom bulb was filled with gas, saturated with water vapor. The top bulb of the other “blank” calorimeter was filled with a saturated solution of hydrogen in water and the bottom bulb with hydrogen saturated with water vapor. Both taps were opened together so that in both cases the liquid displaced the gas. In the first case solution occurred to the extent of about one third saturation as the water flowed from the capillary. In the “blank” calorimeter no solution occurred. The resulting difference in temperature between the bottom bulbs would then be a measure of the heat change associated with the solution of the gas, if the top bulbs were initially at the same temperature. The top bulbs were of 100-ml. capacity and matched in volume to 0.5%. The capillaries a t the bottom of each top bulb were planned to allow the top bulb to empty in a fixed time. One calorimeter took 11 min. and 25 see. t o drain, the other 11 min. and 35 sec. Three way taps were used, one side being connected t o the t o p bulb and the other side having one stem open and the other connected t o the bottom bulb and t o a gas buret in a subsidiary thermostat. Glam columns in each of the bottom bulbs gave surfaces over which the water could trickle. The bottom bulbs were 5.3 em. in diameter and 4.3 cm. apart. Heaters were sealed into the bottom of these bulbs. Each was of 10 ohms resistance, the element being a curved coil of nichrome wire spot welded t o tungsten wire leads 0.03 cm. in diameter and 3 em. in length. These were welded to copper leads which were connected t o current and potential leads. Thermostated glass capillary leads connected each calorimeter to the gas buret. Bakelite supports held the calorimeters in an air jacket totally enclosed in a thermostat, similar t o that used by Gucker, et aZ.8 The thermostat bath was lagged with kapok enclosed in wood and covered with a metal plate and
