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PAULSCHATZBERG
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ing small amounts, -O.Ol(ro of impurities. Addition of oxides of trivalent metals, e.g., A1203, Ga203,causes ai1 increase, whereas oxides of monovalent metals cause a decrease, in the concentration of charge carriers. I n irradiated zinc oxide both zinc aiid oxygen vacancies, as well as interstitial zinc atoms, are produced. Iiiterstitial oxygen atoms are unlikely in stabilized ZiiO and the displaced oxygen atoms probably go to vacancies or grain boundaries to be freed later. The nuclear reaction
P
Zn6*( n , ~ ) Z 11+ ~ ~Ga69 appears to be responsible for the amount of gallium present. I n the range of temperatures investigated so far, it is believed that the lattice defects introduced by y-ray and neutron irradiation are removed slowly and this could explain the aging period during which the catalysts present a lower activity aiid a lower H2/C2H4 ratio compared with the non-irradiated one. It may be meiitioiied that Kohn, Rloore and Taylor4had showii that irradiation effects are relatively stable in refractory oxides. The decrease in catalytic activity of the chromic oxides as a result of both gamma and neutron ir-
radiation may be attributed to the decrease in the number of conduction-band electrons by the creation of chromium in interstitial positions that are not completely annihilated in the temperature range and in the time these catalysts were used. This work has indicated a method for introducing controlled amounts of impurities into catalysts by neutron irradiation that may have importaiit practical applications. Acknowledgments.-The author wishes to express his appreciation to Dr. G. G. Eichholz for his helpful suggestions during this work aiid to all members of the Physics and Radiotracer Subdivision of the Mineral Sciences Division for advice aiid assistance. Thanks are also due to hfr. R . G. Draper for his help in analysis of the gas mixtures used in the calibration of the chromatographic equipment, to Dr. A. H. Gillieson for the spectrographic analyses of the catalysts, and to Mr. H. J. Rfullingtoii of Eldorado Mining and Refining Ltd. for the surface area determination of the catalysts. This work was done during a two-year tenure of a Postdoctorate Fellowship of the Kational Research Couiicil of Canada, which is gratefully acknowledged’.
SOLUBILITIES OF WATER IK SEVERAL NORMAL ALKANES FROM
C 7
TO ClGl
BY PACLSCHATZBERG U . S. Xaval Engineeriny Experiment Station, Annapolis, Maryland Received August SO, lQ6d
The solubilities of water in several n-alkanes from C? t o CIRhave been determined at 25 and 40’. When expressed on a weight basis the solubilities show an inverse relationship t o solvent molecular weight and, when expressed as mole fraction, a direct relationship. The solubilities have been correlated with the solvent properties, solubility parameter and surface tension.
Knowledge of water solubilities in pure hydrocarbon run petroleum fractions show solubilities expressed on a liquids and the effect of increasing molecular weight 01-1 weight basis which decrease with increasing molecular these solubilities is necessary to understand the chemiweight of the hydrocarbon, a trend directly opposite to cal and physical effects of water in hydrocarbon liquids, that shown by the data of Black, Joris, and Taylor. such as jet fuels and lubricating oils. The most I n an attempt to resolve these disagreements and proextensive determinations of water solubilities in hydrovide data beyond octane, the solubilities of water in carbon liquids are found in the work of Black, Joris, several normal alkanes from C7 to cl6 were determined a t 25 and 40’. and Taylor.2 Using tritium oxide as a tracer, they determined water solubilities a t several temperatures Experimental for various hydrocarbons from C4 to Cs. These iiiApparatus and Procedure.-The hydrocarbons were saturated cluded normal and branched-chain alkanes, alkenes, by storing them over a layer of distilled and deionized water in 4-oz. brown glass bottles without any agitation. The bottles cyclohexane, and benzene. illthough a number of were sealed with tightly fitting serum caps and completely subother investigators using different methods have demerged in a water-bath maintained within f 0 . 0 2 ’ of the satutermined water solubilities in hydrocarbons, no such ration temperature for 7 days. This is referred to as the static data are available beyond octane. I n addition, the water saturation method. The storage period of 7 days was based on experience with kerosene fuels. It was found that the data of Black, Joris, and Taylor for the normal alkanes dissolved water content of these fuels contacting liquid water from C4 to Cs, when expressed on a weight basis, indiwithout agitation became constant after 3 t o 4 days. Thus, a cate increasing water solubilities with increasing molecstorage period ot 7 days was considered more than adequate for ular weights; these solubilities are much higher than this method. At the end of 7 days, each bottle was raised out of the bath sufficiently t o expose the serum cap, which was rinsed those found in kerosene fuels of greater chain length (12 to 13 carbon atoms), as pointed out by D a v i e ~ . ~ with acetone and thoroughly dried. A hypodermic syringe pierced the serum cap and was pushed approximately ”3 of the Further, the data of Griswold aiid Kash4 for straightway down into the hydrocarbon liquid. A small hypodermic (1) The opinions expressed in this paper are those of the not t o be construed as official or reflecting the viexs of the the Navy or naval services a t large. (2) C. Black, G . G. Joris, and H. S. Taylor, ,J. Chem. (1948). (3) P. L. Davies, “Fourth World Petroleum Congress 1955, p. 427. (4) J , Orimvold and J . E. Kawh, Ind, Cnu. C h e m , 34, 804
author and are Department of Phys., 16, 537
Section V/E,”
(19423.
needle also pierced the cap. Air pressure through the small needle quickly forced a 20-ml. sample into the syringe. A dynamic a-ater saturation procedure was used for two of the hydrocarbons in order to compare two different methods of saturation. I n this method saturation was achieved by passing ~ of moist air through the hydrocarbon. Filtered air at B f l o rate
approxirnstely 20 ml./min, w a pmaed ~ through n, water m i u r d o r
SOLUBILITIES OF WATERIY NORMAL ALICANES
April, 1963
and from there through a hydrocarbon saturator. Both saturators were submerged in the same water-bath a t the conditions previously stated. Saturation time was approximately 16 hr. S n Aminco-Dunmore conducting-film sensor, situated in the 1 op of the hydrocarbon saturator, measured the relative humidity of the exit air. Khen a constant reading of 97 to 997c relative humidity was obtained, a 20-ml. sample was quickly removed. This was done by attarhing a syringe to a hypodermic needle situated in the hydrocarbon saturator and closing off the air exit, so that the moist air flow forced the hydrocarbon into the syringe. Sampling and transfer to the water analysis apparatus in both saturation methods were achieved in less than 30 sec. All sampling equipment surfaces were rendered hydrophobic with a silicone compound (Clay-Adams) t o minimize any water loss. Water content in the hydrocarbons was determined by the Karl Fischer method. Stabilized Karl Fischer reagent (Fischer Scientific Co.) diluted to tz titer of 1.0 to 1.3 mg. waterlml. was used t o titrate hydrocarbon water content directly in the presence of methanol to a ‘Ldead-stop”end-point, using a Beckman KF3 automatic titrimeter with a 5-ml. microburet. Reagent titer was accurately determined before each separate hydrocarbon analysis by using distilled water introduced into the titration chamber from a weighed microliter syringe. The solubility of water in benzene at 20’ was determined for comparison of the saturation and analysis methods with those of other investigators. Materials.-Table I lists the hydrocarbons, sources of supply, and grades. All hydrocarbons (except benzene) were passed repeatedly through a 4-ft. column of silica gel (Davison No. 12) until essentially no absorption occurred in the spectral range of 220 to 340 mp as measured by a Beckman DK2 ratio recording spectrophotometer. This assured that contamination by aromatics, which exhibit a much higher affinity for water than the alkanes, was eliminated. The silica gel treatment also served t o remove small amounts of polar materials which exhibit a very high affinity for water. The most probable impurities are branched-chain and cyclic isomers in the same boiling range. The effect of such impurities on the water solubilities was considered negligible. The n-heptane was doubly distilled before treatment with silica gel. The refractive indices of the purified solvents agreed with the literature values6 within 0.7 part per 104 or better.
TABLE I HYDROCARBON SOCRCES AND GRADES~ n-Heptane EOC PP, research grade, 99.69 mole yc n-Nonane n-Decane PPI research grade, 99.43 mole yo n-T7ndecane PP, research grade, 99.33 mole 70 n-Uodecane PPI pure grade, 99+ mole 7c n-Tridecane PP, research grade, 99.73 mole PPI pure grade, 99+ mole 7 0 n-Tetradecane n-Hexadecane HW, ASTM normal cetane Benzene MCB, chromatoquality reagent, 99f mole 70 a EOC, Eastman Organic Chemicals; PP, Phillips Petroleum Co.; HW, Humphrey-Wilkinson, Inc.; MCB, Matheson, Coleman and Bell.
Experimental Results Results of the water solubility determinations a t 25 and 40’ are recorded in Table 11. Each value reported is the mean of two to four determinations. The deviations from the mean ranged from 0 to 6% for the 25’ data and from 0 to 2% for the 40’ data. Data other than the experimental data, also recorded in Table 11,will be discussed subsequently. Results of water solubility determinations for n-nonane and n-decane a t 25’ by the dynamic water saturation method were 81 and Yl p.p.m. by weight, respectively, comparing well with the corresponding results obtained by the static water saturation method (Table 11). ( r ? ) F. D. Rossini, K. S. Piteer, R. L. S m e t t , R. M. Braun, and G. C. Pimentel, “Selected Values of Physical and Thermodynamic Properties of Hvdrocarbonfl and Related Compounda,” Carneaie Press, Pittsbinyh, P a , 1953.
777
It is seen from Table I1 that the water solubilities expressed by weight decrease gradually with increasing alkane chain length. The opposite trend is noted when the solubilities are expressed as mole fraction. The water solubilities found compare favorably with those of kerosene fuels having equivalent chain lengths. Disagreement exists, however, with the water solubilities found by Black, Joris, and Taylor.2 For n-heptane at 25’, they found 151 p.p.m. against 91 in this work; their data from Cd to Ca showed a direct relationship between dissolved water by weight and alkane molecular weight, while in this work the reverse was found for C7 to C,e. Results of the water solubility determinations for benzene a t 20’ are recorded in Table 111,along with the methods and results of other investigators for comparison. Excluding the low result of Black, Joris, and TaylorJ2the average solubility of the values in Table IIJ is 545 ppm, wh.ich differs from the value found in this work by 27,. It can be concluded, therefore, that the saturation and analysis methods used in this iiivestigation are comparable to those used by other iiivestigators. Discussion Hildebrand6 has derived an equation to describe the mixing of components of different internal pressures and volumes
RT In ( L Y Z / X Z ) = V Z ( + I ) ~( ~ 6Zd 2 (1) where LY = activity, V = molal volume, 6 = solubility parameter, + = volume fraction, and x = mole fraction. Subscripts 1 and 2 refer to the h.ydrocarbon solvent and water solute, respectively. The following simplifying assumptions can be applied: a2 = 1, and, since the water concentration is very small compared to the h.ydrocarbon concentration, is very nearly unity. Thus, eq. 1 becomes In x2 Values of relation
=
- v2 ~
RT
(&
- 6J2
(2)
and 61were calculated from the Hildebralid 6 = [(AH,
- RT)/V]”’
(3)
Values of tE.e h.eats of vaporization, AH,.25,for the alkanes were obtained from the l i t e r a t ~ r e while ,~ values of AH,40were calculated from =
AH,”
- ACD(313 - 298)
(4) where AC, is th.e difference between the heat capacity of the liquid and that of the gaseous hydrocarbon. Calculated water solubilities sh.own in Table I1 were obtained from eq. 2 . Figure 1 sh.ows the least squares straight lilies through the experimental plot of log water solubility against alkane molecular weight. It is seen th.at the water solubilities are inversely proportional to molecular weigh.t when expressed on a weight basis and directly proportional when expressed as mole fraction. Figure 2 shows the least squares straight lines through the experimental plot of log x2 against as expected from eq. 2 . With the exception of n-heptane, the experimental and calculated (6) J. H. Hildebrand andvR. L. Scott, “Solubility of Non-electrolytes,” Reinhold Pnbl. Corp., New York, N.Y., 1950.
PAULSCHATZBERG
778
Vol. 67
TABLE I1
--
Temp.. O C .
SOLUBILITIES OF WATERIN NORMAL ALKANES
P.p.rn. by wt.-Exptl.-Calod.a7 25 40 25 40
Heptane 91 74 Nonane 79 72 Decane 72 136 70 111 Undecane 69 130 69 109 Dodecane 65 127 66 104 Tridecane 60 123 64 100 Tetradecane 113 96 Hexadecane 54 104 58 91 a Based on the Hildebrand equation presented
--Mole fraction X 1047Exptl.--Calcd,--
25
40
5.1 5.6 5.7 6.0 6.1 6.1
25
40
4.2 5.2 5.6 6.0 6.2 6.5
10.7 11.3 12.0 12.6 12.6 6.8 13.1 7.3 in the Discussion. 6 is
20.0
Exptl./calod. 25 40
1.2 1.1 1.0 1.0 1.0 0.9
-41~-
25
- 61)~-
-(Se
25
40
7.43 7.65 7.72 7.80 7.54 7.89
40
255 248 246 243 242 241
8.8 1.2 7.60 241 9.4 1.2 7.68 239 9.8 1.2 7.72 238 10.2 1.2 7.77 236 10.6 1.2 7.81 235 11.4 0.9 1.2 8.01 7.89 237 232 the solubility parameter in the Hildebrand equation.
100 0
80.0
c
0
60.0
5 E
10.0
8.0
40.0 e
60
P z
.A W
40
0 20.0
cuU a
L
W -I
0
=
10.0
8.0 ez 6.0
80
t" z 8
60
a
100
4.0
W
5 ALKANE MOLECULAR WEIGHT
Fig. 1.-Alkane
water solubilities and molecular weights.
I.o
230
235
240
245
IS2-
TABLE I11 SOLUBILITY OF WATERIN BENZENE AT 20' Ref.
2 7
8 9 10 11 This work
Saturation method
Water vapor Shaking with water and storage Water vapor Shaking with water and heating Shaking with water and heating Shaking with water, cooling, and standing Storage with water
Analysis method
250
255
wt.
DTO tracer Karl Fischer
430 530
HTO tracer Cloud point
520 570
Cloud point
530
Gasometric with CaHg
573
Karl Fischer
532 i 3
water solubilities a t 25O show remarkably close agreement, as seen from their ratios recorded in Table 11. This must be regarded as fortuitous, since eq. 2 involves several simplifying assumptions and neglects a correction for the entropy of mixing molecules of different size. Inclusion into eq. 2 of a correction for entropy of mixing results in calculated water solubilities which are from three to six times larger than the experimental solubilities. (7) T.I. Berkenheim, Zauodskaya Lab., 10, 592 (1941). (8) B. D.Caddook and P. L. Davies, J. Inst. Petrol., 46,391 (1960). (9) L. A. K. Staveley, J. H E. Jeffes, and J. A. E. Moy, Trans. Faraday Sac., 89,5 (1943). (10) D. N. Tarasenkov and E. N. Polahintaeva, Zh. Obshch. Khim., 1, 71 (1931). (11) C. K. Rosenbaum and J. H. Walton, J. Am. Chem. Soc., 62, 3568 (1930).
260
6,)2
Fig. 2.--Alkane water solubilities us. (a2 Results p.p.m. by
EXPERIMENTAL
0
2.0
-
&)2.
UhligIz describes the energy change in placing a gas molecule of radius R into a solvent as consisting of the work to produce a spherical cavity of radius r in the solvent and the interaction energy between solute and solvent molecules. Based on this description and applying the Maxwell-Roltzmann distributiori theory, he derived the equation 4nr2u logB=--+2.303 L T
E 2.303 k T
(5)
where B = Ostwald solubility coefficient, r = radius of the gas molecule, u = the sdvent surface tension, E = the interaction energy between solute and solvent, and IC = the Boltzmann constant. For those gases which have small interaction with solveiit molecules, E is small. Considering the dissolved water as a gas, E should remain nearly constant in the homologous series of solvents used, since the water molecule would always be in an environment of -CHz and -CHI groups. Consequently, the logarithm of the Ostwald solubility coefficient for water when plotted against n-alkane surface tension should produce a straight line. This is seen in the least squares straight lines shown in Fig. 3. The efficiency of fit, standard deviation, slope, and intercept for each line are recorded in Table IV. Based on eq. 5, the slopes and intercepts were used to calculate the radius of the water molecule and the (12) H. H. Uhfig, J . Phys. Chem., 41, 1215 (1937).
CHARGED CARBOXYLATE BASESIN DILUTEACIDSOLUTIOXS
.4pril, 1963
779
TABLE IV LEASTSQUARES DATA Temp., OC.
Efficiency of fit, %
25 40
96 90
Sbandard dev.
x
Slope
Intercept
- 0.0248
- 0.5734 - 0.3384
104
1.78 1.42
- 0.0232
solute-solvent interaction energy. For the 25 and 40’ data, respectively, the water radii calculated were 1.37 and 1.36 A.,while the interaction energies were -782 and -485 cal./mole. Although literature values for the radius of the water molecule range from 1 to 3 A., depending on the assumptions made, some of these are very close to the values determined in this work: 1.25,13 1.32,13 1.39,14 and 1.44.13 The negative interaction energies can be interpreted, according to Uhlig,12 as an absorption of energy by the system as water dissolves, which is indicative of a low solubility. The negative interaction energies account in part, according to Le Chatelier’s rule, for the increase in solubility of water with increasing temperature.
SOLVENT SURFACE TENSION, d y n e r l c m
Fig. 3.-Alkane
Conclusions Results of water solubilities in normal alkanes from C7 to CI6show a gradual decrease with increasing molec(13) N. E. Dorsey, “Properties of Ordinary Water Substance,” ACS Monograph, Reinhold Publ. Corp.. New York, N. Y.. 1940. (14) C. ,J. F. Bbttcher, Rac. trau. chim. (Rotterdam), 66, 14 (1946).
water solubilities and surface tensions.
ular weight, when the solubilities are expressed on a weight basis; when expressed on a mole fraction basis, the reverse trend occurs. The experimental water solubilities agree well with two different theoretical treatments.
ACID-BASE EQUILIBRIA I N CONCENTRATED SALT SOLUTIOKS. 11. CHARGED CARBOXYLATE BASES I N DILUTE ACID SOLUTIONS1 BY JAMES S. DWYER AND DONALD ROSEXTHAL~ Department of Chemistry, The University of Chicago, Chicago, Illinois Received September 6, 1968 ~HG.E - pMH is shown to be a constant for a given salt solution in slightly acidified 1 to 8 &r LiCI. (pHa E . is the pH measured using a cell with a glass indicator electrode and a saturated calomel reference electrode. pMH = -lag total strong acid concentration.) Values of p H c . ~ . pMH are reported. In dilute acid solutions of acetate or formate ion KBH= ( [B-][H+]/[BH])QBH where KBHis the thermodynamic acid dissociation constant of the acid (BH), and QBR is a constant for a particular base, B-, and a particular salt solution. The values of ~ H G . Ecalculated . using this equation are in good agreement with the eyperimental values. Values of log &BE are reported for acetic acid in 1-8 M LiCl. It is shown that calculated values of log QBHusing the equation log QBH = - ( d Z / l A dZ) BM do not differ significantly from the experimental values. Calculated and experimental QBH values are compared with published results for dilute LiCl solutions.
+
+
Introduction I n a previous study3PHG.E. - pMH was shown to be a constant for dilute acid solutions in 4 and 8 M LiC1, 6 M NaC104,6 M Na,:l‘Oa, and 4M CaCL (PHG.E.is the pH measured using a cell with a glass indicator electrode and a saturated calomel reference electrode and pMH is -logarithm of the total molar concentration of strong acid.) Further, it was found that the quantitative aspects of equilibrium between a weak uncharged base, B, and its conjugate acid, BH+, in such solutions can be accounted for using the equation KBHc =
([B][total strong acid c ~ n c n . ] / [ B H + ] ) & ~ ~ (1)
(1) Taken in p a r t from the Ph.D. research of James 9. Dwyer. This work s u p p o ~ t e dby research grants from the U. S. Public Health SerTiee
\vas
(RG-9583 and RG-8069). (2) To n h o m inquiries should be direrted a t Clarkson College of Technolony, Potsdam. New York. (31 D. Rosenthal and J. S. Dwyer .I. Phrs. Chem., 66. 2687 (1962).
where K B H + is the thermodynamic molar acid dissociation constant and QBH+ depends upon the nature and concentration of salt and the nature of the base, B. Apparently, the various factors4-” which are important in concentrated salt solutions can be satisfactorily incorporated into the QBH+term. (4) H. S.Harned and B. B. Owen, “The Physical Chemistry 01 Elert,rolytic Solutions,” Reinhold Publ. Corp., New York, N. Y., 1958, pp. 509-547. ( 5 ) H . 9. Frank. et at., J . Chem. Phus., 13, 507 (1945); Ann. Res. P h y s . Chem., 6, 43 (1954). R. W. Gurney, “Ionic Processes in Solution.” 410Craw- Hill Book Co.. Inc., New York, N. Y., pp. 248-260. ( 6 ) J. B. Hasted, et al., J . Chem. Phys., 16, 1 (1948); 29, 17 (1959). (7) (a) R . 4. Robinson and R . H . Stokes, “Electrolyte Solutions.” Butterworths Publications Ltd.. London, 1959, p. 62; (b) E. Glueckauf in “The Structure of Electrolyte Solutions,” M7. J. Hamer (editor), John Wiley and Sons, Inc., New York, N. Y . , 1959, pp. 97-112. (8) J. Beck, Physik. Z., 40, 474 (1939); G. W. Brady, J. Chem. rhus., 28, 464 (1958); H. S.Frank and P. T. Thompson in ref. 7b, pp. 113-134. (9) Ref. 4, pp. 614, 634-643. (10) E. Grunwald, G. Baughman, and G. Kohnstam, J . A m . Chem. Soc.. 62, 5801 (1960). (11) J. K. BrSnsted, Trans. Faraday Soe., 2S, 430 (1927); G. Scatchard in rof. 7b, pp. 9-18.
