Environ. Sci. Technol. 1986,20, 830-836
Effects of Solute Concentration and Cosolvents on the Aqueous Activity Coefficient of Halogenated Hydrocarbons Christoph Munz* and Paul V. Roberts Civil Engineering Department, Stanford Unlversity, Stanford, California 94305
The effect of solute concentration and the effect of model cosolvents (methanol and 2-propanol) on the solute’s aqueous activity coefficient at infinite dilution, rim, were investigated at 20 OC. The multiple equilibration technique with direct aqueous-phase analysis was used to measure Henry’s law constant of CHCl,, CC14,and C2C16, which is directly proportional to yIy. It was found that 77 = constant is a valid approximation for compounds with aqueous solubility mole fractions xi,6 Cosolvents reduce the solute’s yi only at cosolvent mole fractions x >5X The more hydrophobic the cosolvent the stronger the effect on a given solute, and the more hydrophobic the solute the stronger the effect of a given cosolvent, in agreement with general thermodynamic principles. The validity of a semiempirical thermodynamic model, UNIFAC, is assessed in the dilute concentration range. UNIFAC predicts cosolvent effects conservatively, while prediction of solute concentration effects agrees with the findings of this work.
Introduction The fate of trace organic chemicals in the environment is customarily modeled on the basis of pure compound properties such as aqueous solubility, octanol-water partition coefficient, and Henry’s law constant. In most cases, however, the contamination of a water stream is caused by more than one substance, and the concentration of any given compound may vary over several orders of magnitude on a case by case basis. Substantial solute-solute or solute-cosolvent interactions could affect the transport of organic contaminants in groundwater or the extent of removal in treatment processes such as activated carbon adgorption and gas-liquid contacting processes. The objectives of this investigation are (1)to establish criteria for defining the dilute concentration range as it applies to hydrophobic compounds in binary aqueous systems and (2) to quantify the degree of influence of cosolvents on the thermodynamic properties of solutes infinitely dilute in solution. Liquid-phase mole fraction concentrations xi 5 ( 1 w t %) are typically considered dilute in the chemical engineering field, or alternatively, phase equilibrium can be represented by a linear relationship (Henry’s law) in this concentration range. Hwang and Dasgupta (1)recently confirmed the linearity of equilibrium of hydrogen peroxide in dilute aqueous solution in the mole fraction concentration range 2.7 X lo* f x < 2.7 X lo9. Since H202 is miscible with water in all proportions (2),this system can be considered fairly ideal compared to the typically more hydrophobic solutes (i.e., slightly soluble) of environmental concern, where deviations in the same mole fraction concentration range, if any, are expected to be more apparent. However, the linearity of equilibrium of slightly soluble compounds has not been consistently tested over a wide concentration range, independently of the methodology used isopiestic measurements (3),headspace N
*Address correspondence to this author at the Swiss Federal Institute for Water Resources and Water Pollution Control (EAWAG), 8600 Dubendorf, Switzerland. 830
Envlron. Sci. Technol., Vol. 20, No. 8, 1986
experiments ( 4 ) ,or in the determination of Henry’s law constant in batch air-stripping experiments (5-9). Moreover, previous studies (3-9) report conflicting results on the effect of solute concentration on the solute’s activity coefficientor Henry’s law constant. A stringent test of the assumption of linearity of equilibrium needs to embrace concentrations from trace amounts up to near the solubility limit, where deviations, if any, are expected to be most evident. In this paper we report such experiments for three solutes of differing hydrophobicity and/or aqueous solubility: chloroform (CHCl,), carbon tetrachloride (CC14), and hexachloroethane (C2C16). Most studies reporting cosolvent effects are based on liquid-liquid, solid-liquid, and solid-solid equilibrium measurements in water (10-17). However, the thermodynamic interpretation of these data is often confounded by the possible nonidealities of the organic phase. Nevertheless, most of these findings suggest that cosolvent effects will only be significant (210%) and quantifiable more precisely if a relatively soluble cosolvent is used that can be added in substantial amounts. Hence, in order to avoid these complications, the approach taken in this study was to measure Henry’s law constant or air-water partition coefficient of the three solutes in a closed system in the absence and presence of varying concentrations of miscible cosolvents: methanol and 2-propanol. Since at low pressure and ambient temperature the vapor phase may be assumed to behave ideally (18) and Henry’s law constant @ directly proportional to the solutes’ liquid-phase activity coefficient, the thermodynamic interpretation is straightforward. And finally the validity of a semiempirical thermodynamic model, UNIFAC, for predicting activity coefficients in the dilute concentration range is assessed.
Thermodynamic Principles The liquid-phase fugacity of compound i can be expressed in several forms, depending on the choice of reference state (19)
f? = xiyfp,
= xiYTHxZc
(1)
where x i is the mole fraction, fiR is the pure component reference fugacity in the liquid state at the system temperature T and pressure P, and yI and y: are the activity coefficients in the symmetric and asymmetric convention, respectively. Hxi,,is Henry’s law constant of compound i in a pure solvent or solvent mixture at the system T and P. It is customary to assume the symmetric convention (Lewis-Randall rule), yi 1as x , 1,such that the pure liquid becomes the reference state. The activity coefficient then characterizes the nonideality of compound i in the liquid solution, analogous to the fugacity coefficient in the gas phase, With 7: 1as xi 0 it follows from eq 1that Henry’s law constant is directly proportional to the activity coefficient HXIC= 78, (P) (2)
- - -
or alternatively, when Henry’s law constant is expressed as the dimensionless ratio of gas to liquid volumetric concentrations, e.g., (mol~m~3),,/(mol~m~3)~i,
0013-936X/86/0920-0830$01.50/0
0 1986 American Chemical Society
(3) where R is the gas constant (= 8.2 X loi atmm3.mol-1.K-1), T i s the absolute temperature, and us is the molar volume of the solution (rn3.mol-l). Linearity of Equilibrium in Dilute Binary Aqueous Systems. Equating fugacities for a liquid solute i in equilibrium with its aqueous solution and considering that water is only slightly soluble in the organic phase, and vice versa, the mole fraction solubility becomes 1 xi, N - = constant (4)
5
For compounds that are solids at the temperature of interest, such as CzC16,eq 4 needs to be multiplied by the ratio of solid fugacity to liquid or reference fugacity (18). At constant temperature the fugacity ratio is also constant. From such theoretical considerations it can be shown that if there were a concentration effect, it would be to lower the activity coefficient (or Henry’s law constant) with increasing solute concentration. Simply consider the Lewis-Randall rule applied to a water (l)/solute (2) system: as x1 1, y1 1, and as x2 -* 0, y; maximum. Thus, if x 2 is slightly larger than zero, then yz < 7;. Cosolvent Effects. According to Yalkowsky and Roseman (12),the effect of a cosolvent on the solubility of a solute is primarily dependent upon the polarity of the solute with respect to the solvent (water) and cosolvent. The octanol-water partition coefficient, is an appropriate polarity scale. For nonpolar solutes the effect of a cosolvent can be mathematically approximated by assuming that a mixed solvent is simply a linear combination of its component solvents (12). By use of this approach in terms of activity coefficients it can be shown that log (yi,m/yi,w) f c log ( ~ i , c / y i , w ) (5) where fc is the cosolvent fraction, subscripts m, w, and c refer to mixed solvent, water, and cosolvent, respectively, and yi is the solute’s activity coefficient. Since for nonpolar compounds yi,c/yL,w C 1, it follows that yi,m/yi,w< 1. Consequently, the solute’s activity coefficient (and Henry’s law constant) will be smaller in the mixed solvent than in pure water, or its solubility will be greater. From eq 5 it can also be inferred that, for a given solute and a given cosolvent concentration, the less polar the cosolvent, i.e., the smaller yi,c,the stronger will be its effect in reducing yi,m. This simple approximation will break down with increasing solute polarity. However, while water has log -1.4 (20), organic solutes of environmental concern generally have log 2 2. This in turn indicates that these solutes can be considered nonpolar or semipolar relative to water. Activity Coefficient Model (UNIFAC). UNIFAC is a semiempirical thermodynamic model that uses the solution-of-groups approach to predict the thermodynamic behavior of compounds in solution. The UNIFAC method was introduced by Fredenslund et al. (21) and has been revised and extended on several occasions (22-24). It is based on the UNIQUAC model for determining activity coefficients for nonideal liquid mixtures (25). The advantage of functional group models is that activity coefficients can be predicted for virtually any compound, provided the parameters are available for the functional groups needed. UNIFAC predictions of parameters such as aqueous solubility, Henry’s law constant, and organicsolvent water partition coefficients have been evaluated by several investigators (26-28). UNIFAC activity coefficients are calculated from the sum of two contributions:
-
+
+
Table I. log K , , of Compounds Studied” compound water methanol 2-propanol
log -1.38 -0.77 0.05
compound
1% K 0 , W
chloroform carbon tetrachloride hexachloroethane
1.95 2.83 3.94
“All values measured (20),except for that of C2C18, which was calculated according to the fragment method (20).
a combinatorial (C) part essentially due to differences in size and shape of the molecules in the mixture and a residual (R) part essentially due to energy interactions. The combinatorial and residual parts correspond to the entropic and enthalpic contributions, respectively. The resulting model is In y i = In yF
+ In yP
(6)
All components in solution must be nonelectrolytes, nonpolymers, and condensable; no distinction can be made for isomers (e.g., dichlorobenzenes). Further details about the UNIFAC method can be found elsewhere (21-26).
Experimental Section Choice of Solutes and Cosolvents. Chloroform and carbon tetrachloride were chosen as solutes because they are frequently identified in drinking water and/or are common groundwater contaminants, and hexachloroethane was chosen because of its strong hydrophobic character. Methanol and 2-propanol were selected as model cosolvents on the basis of the following considerations: (a) expected effect on the solutes on the basis of model predictions (UNIFAC) and previously reported experiments (10-12); (b) difference in hydrophobic character (log (c) solubility in water; (d) noninterference with the analysis of solutes by gas chromatography. Table I lists the log of the compounds considered: the differences in log among the solutes (-2) and the cosolvents (-0.8) appear to be large enough to permit studying the effects of a given cosolvent on solutes of variable hydrophobicity, and vice versa. Approach. Of the several possible approaches to measure Henry’s law constant (6, 29, 30), the multiple equilibration (ME) technique (29) of a closed system with analysis of the aqueous phase was found to yield the best results in a comparative evaluation conducted in our laboratory (5). Keeping in mind that the purpose of this study is to quantify differences in H,, (or yi) that are likely to be very small, minimizing sample handling is of prime importance. Liquid-phase concentrations were determined by direct aqueous injection into a gas chromatograph (CHC13and CClJ and by I4C liquid scintillation counting (CZCl6), thus minimizing experimental artifacts (31, 32). By use of this technique, Henry’s law constant is evaluated by fitting the experimentally determined concentrations, CL,,with a linear least-squares regression according to
I)
[ ( +;
log CLi = log CLo- i log
1
-Hc
(7)
where i = 0-n, n is the total number of equilibrations, and VL and VGare the liquid and gas volumes of the closed system, respectively. Assuming that the errors involved in measuring VL and V G are negligible, the sensitivity of the ME technique can be assessed from the expression:
Environ. Sci. Technol., Vol. 20, No. 8, 1986 831
Table 11. Initial Concentrations Studied, Solubilities, and Ratio VL/Vo for Individual Solutes
CHC13, mg/L 0.14 2.0 20
200 1400 4400
solubility, mg/L VL/ V G
82000
0.5
CCl,, mg/L
C2Cls,mg/L
0.03 0.13 4.0 50 500
0.009 0.073 0.63 7.5
803O
7.7b
I
1.21
$
8
SOLUTE = CHLOROFORM 1.1:
/
4
1
a Recommended values of Horvath (34). Measured in our laboratory (5).
Equation 8 represents a typical (Gaussian) error propagation calculation. Hence, the larger the number of equilibrations n,and the smaller the ratio VL/ VG,i.e., the larger the concentration drop in successive equilibrations, the more accurate the determination of H,. On the other hand, the objective of assessing the validity of linearity of equilibrium calls for minimizing CL,/CL, over an entire experiment. The compromise between these opposing considerations was satisfied by choosing VL/ VG such that the liquid concentration would decrease approximately 20% per equilibration. This was considered sufficiently small to study concentration effects over 4-5 orders of magnitude, while avoiding such effects in a single experiment. Concentration Dependence Experiments. Duplicate experiments for the three so1utes-CHC1,, CC14, and CzC16-were carried out over the largest possible concentration range, i.e., from slightly above detection limit to slightly lower than the solubility limits. Table I1 shows the range of concentrations investigated for each solute, their solubilities, and the ratios VL/VG used. All measurements were conducted at 20 f 0.1 "C. For each solute the experiments were conducted in order of increasing initial concentration, and VL/ VG was held constant. Cosolvent Experiments. On the basis of UNIFAC model prediction5 (23),the interesting range of cosolvent concentration is 10-4-10-2 mole fraction. However, initial experiments at mole fraction resulted in only small effects. Hence, methanol and 2-propanol cosolvent mole fractions of 5 x lo-,, 2X 5X and 10-1 were used with each individual solute at infinite dilution ( x , -130 pg/L). All experiments were carried out in duplicate at 20 f 0.1 "C and the ratios VL/V, used for each solute are those indicated in Table 11. For the system CzCl6-2-propanol-water, three additional duplicate experiments were done with v L / v G = 0.5 at 7 X 9x and 10-1 2-propanol mole fraction. For a given solute the experiments weie conducted in random order. Apparatus and Procedure. The containers used to carry out these experiments were 100-mL gas-tight syringes (Glenco, Spectrum Laboratories, Houston, TX) with 1-mL graduations and a threaded glass tip to which a microvalve could be attached. To the other end of the microvalve a screw-hub nut was attached with two Teflon-faced rubber septa placed in between, providing the means for sampling and assuring a tight seal. The solution in the syringes was spiked with a few microliters of a saturated solute solution through the microvalve, or for the concentration dependence experiments, large amounts were transferred directly by connecting two syringes with a coupling. No solvent was used to predissolve the solutes, with the exception of small amounts of methanol required for l4CZCl6.Equi-
-
832 Environ. Sci. Technol., Vol. 20, No. 8, 1986
0.8'
1 tE
'
""""
10"
'
""""
10'
'
""""
1 is
'
""""
I6 '
' ' ' * * > J
1f 3
MOLE FRACTION SOLUTE u
lo-'
loo
10'
lo2
10'
IO'
CONCEN TRA T ION SOLUTE h g / l I Figure 1. Sensitivity of solute (CHCI,) activity coefficient to solute concentration.
librium between the gas and liquid phase was achieved by placing the syringes first on a rotator (60 rpm) for a minimum of 15 min and thereafter upside down into a water bath for a minimum of 10 additional minutes. During these 10 min, the syringe was withdrawn twice from the water bath and was shaken relatively vigorously for 10-20 s. This procedure was carried out 5 times (6 times for C2C16)(= number of equilibrations), for a total of 6 (7 for C2C16)replicate concentration measurements. Analytical. The chromatographic analysis (CHCl,, CC14) was carried out in a Finnigan Model 9610 gas chromatograph having a 6-ft column, i.d. 4 mm, packed with Chromosorb 101,60/80 mesh. Carrier flow was 40 mL/min methane (5%) in argon. No makeup gas was used. Oven temperature was 130 "C (isothermal); the detector was 63NiECD operated at 330 "C; the injector temperature was 200 "C; the injection volume was 3 pL. The nonlinear detector response at higher concentrations (21 mg/L) was calibrated the day prior to a given experiment. 14C-Labeledsamples (0.7-1.8 mL) were analyzed 2 or 3 times for 20 min with a Packard Tricarb (Model 4530) scintillation counter. The radiolabeled hexa~hloro[U-~~C]ethane (18.2 mCi/mmol; Amersham Corp., Arlington Heights, IL) showed no contamination from gas chromatographic analysis. The unlabeled C2C16(Supelco, Inc., PA), used for the experiments a t higher concentrations, contained -3% of CCl2=CCI2 as a contaminant. The purity of CHC1, (J. T. Baker Chemical Co.) and CC14 (Aldrich Chemical Co.) was >99.5%; methanol and 2propanol were of "pesticide grade" purity (Fisher Scientific). Distilled water was further purified by passing through a "Milli-Q, Reagent-Grade Water System" and vacuum stripping for a minimum of 1 2 h. Results
Concentration Dependence. It is appropriate to recall that the measured Henry's law constant, H,,, is directly proportional to the infinite dilution activity coefficient, rr. The results will be mainly presented in terms of activity coefficients to allow a direct comparison with predictions from the UNIFAC model. Figures 1-3 show the effect of solute concentration on the solute's activity coefficient, expressed as a relative effect on the ordinate, for CHCl,, CC14, and C2C16,respectively. Also shown are the pre-
1.2,
1
Table 111. Henry's Law Constants (20 "C)
compound
c:
2
6
1.
0
4: C
2 L
H,"
CHC13 CC1, C&ls
it
C.V.,* % ne litad
SD
0.125 h0.00222 0.975 f0.0128 0.119 10.00854
ref
11 0.129 8,9,30, 43, 44 10 0.877 43-45 8 NAe
1.8 1.3
7.2
Average values corresponding to Figures 1-3. C.V. = coefficient of variation = (SD X 100)/mean. c n = number of determinations. dAverage value of listed references. "A = no experimentally determined values available. (I
I-
1 UNIFAC PREDICTION
0.9-
1
SOLUBILITY
LIMIT
MEASURED 0
5
0.81 ' ' I tiQ
".""
'
' """'
I tiB
'
' """'
1ti7
'
' """'
'
' ' * . U J
16'
1cis
16O
I
1.1,
MOL E FRACTION SOLUTE u.
16'
I Liz
loo
lo3
lo2
lo1
CONCENTRA TION SOLUTE Cmg/ll
COSOLVENT = METHANOL
Figure 2. Sensitivity of solute (CCI,) actlvity coefficient to solute concentration.
SOLUTES: 0
1.31
0
1
A
CHLOROFORM CARBON TETRACHLORIDE HEXACHLOROETHANE
MEASURED UNIFAC P R E D I C T I Z
SOLUTE = HEXACHLOROETHANE I i5
I
10'
c3
I i2
I E-'
MOLE FRACTION COSOL VENT Flgure 4. Effects of methanol on solutes' activity coefficients. 1.1
UNIFAC PREDICTION
0
2
. MEASURED 0.7 10
'
.a
' """'
'
lCiQ
' """'
'
'
""."
10-
'
1ti7
'
c : -
LIMIT
0
' ' * * . . & J
IO"
0 .7. T
.
,
.6.
1
.
.s-
MOLE FRACTION SOLUTE
COSOLVENT = ISO-PROPANOL
'
I ti3
16'
16'
loo
lo1
G C 4
CONCEN TRA T ION SOL UTE Cmg/ll Figure 3. Sensitivity of solute (C,CI,) activlty coefficient to solute concentration.
9
.4 '
SOLUTES:
.2
0 0 A
.3.
h
,
.I' 8.0'
dictions from UNIFAC represented by a solid line. The ordinate in Figures 1-3 is the following: ratio infinite dilution activity coeff =
(Xi
10-10)
ri" ( X i
= 10-10)
Thus, the reference infinite dilution activity coefficient is at a solute concentration of x i = 10-l" mole fraction ( 1 pg/L). This was used to represent the UNIFAC predictions. The concentrations shown (Figures 1-3) are the initial solute concentrations, CLo. As can be seen from Figures 1-3, the solute's activity coefficient or Henry's law constant exhibits no concentration dependence over the entire concentration range studied. Accordingly, the reference activity coefficient or Henry's law constant was taken as the arithmetic average of the measured values. Therefore, the plotted values also give an indication of the variability of the results around a mean value and/or the precision of the multiple equilibration technique with liquid-phase analysis. The error bars shown represent the 95% confidence intervals of the individual values. Table I11 shows the average values and standard deviation of the N
CHLOROFORM CARBON TETRACHLORIDE HEXACHLOROETHANE
I
E-'
'
" ' ' ' ' '
I
c'
'
'
""*"
,
" " " "
I B3
,.,m.'c4
I E-2
IE-1
MOL E FRACTION COSOL VENT Figure 5. Effects of 2-propanol on solutes' activity coefficients.
data plotted in Figures 1-3. When these values of Henry's law constant are used, the actual concentration drop during an entire experiment, CL,/C~o,can be calculated as 0.33 for CHC1, and CC14 and 0.51 for CpC16. Cosolvent Effects. The relative effect on the three solutes of each cosolvent, methanol and 2-propanol, are shown in Figures 4 and 5, respectively. The data points are the average values of duplicate experiments and are joined by the solid lines. Also shown are the corresponding predictions from UNIFAC (dashed lines). The ordinate in Figures 4 and 5 representing the relative effect is defined as follows: ratio infinite dilution activity coeff =
rl rf (~,o,olv (XC0,OlV
f
0)
= 0)
where 7; ( x ~ ,=~0)~is~evaluated at a solute mole fraction Environ. Sci. Technol., Vol. 20, No. 8, 1986
833
x, = for the UNIFAC predictions and, for the measured values, is taken as the average obtained from the concentration experiments (Table 111) in the case of CHC13 and CC4. For C2C1, the reference value chosen is the average value of the two experiments at the lowest solute concentration investigated (9 X mg/L; Figure 3), as these experiments and the cosolvent experiments were conducted only with labeled C2C&. At the higher cosolvent concentrations used, the molar volume of the water-alcohol solutions, us, is appreciably greater than that of pure water. This correction was considered when activity coefficients were calculated from the measured Henry’s law constant, according to eq 3. Hence, the relative effect based on activity coefficients is stronger than the relative effect based on the dimensionless Henry’s law constant, H,.
Discussion Experimental Error and Data Analysis.. Random errors include the analytical error and the precision of the volume measurements, VL and VG. Potential systematic errors include (a) incomplete equilibrium between the gas and liquid phase, (b) leaks from the syringes (tightness), (c) absorption (diffusion) of the solutes into the Teflon polymer matrix, and (d) adsorption of the solutes onto glass or Teflon. Adsorption experiments indicate that 1-2% of the C2C& mass might have adsorbed onto the walls of the experimental vessels over the duration of an experiment (4-5 h). Similarly, findings from absorption/desorption experiments with CC14suggest that most solutes will absorb or diffuse into Teflon in substantial amounts if given enough time. However, we are confident that none of these potential systematic errors (a-d) had a significant influence on the experimental data, as discussed in detail in ref 5. The 95% confidence intervals of the individual Henry’s law constants are approximately *2.3% for CHC1, and CC4, and f8.0% for C2C16of the estimate in the concentration dependence experiments (Figures 1-3). As the cosolvent mole fraction is increased, the error of prediction increases slightly. Hence, according to eq 8, it can be inferred that the analytical error is the largest source of error in the experimental procedure used: &0.7% for CHC1, and CC14 and f3.6% for C2C16. The cosolvent concentration remains virtually unchanged during the course of an experiment (CLn/Cb 0.999) as the air-water partition coefficients of methanol and 2-propanol are very small (Hc (33)). Comparison with Previous Work. The average values of experimentally determined Henry’s law constants reported in the open literature for the compounds investigated are given in Table 111. While for CHC1, the agreement is excellent, our value for CC14 at 20 “C is -10% larger. For C2C16,on the other hand, no experimentally determined value of Henry’s law constant has been reported to our knowledge. Concentration Effects. The data for CHC1, and CC14 (Figures 1and 2) give no indication of any concentration dependence of the activity coefficient, in accordance with thermodynamic principles. Similarly, the individual measurements for CzC&(Figure 3) indicate no significant differences (CY = 0.05). The apparent variation of the C2C& activity coefficient manifested in the extreme low concentration range (i.e., xcKcl, lo4) of Figure 3 is attributed to experimental (analytical) artifacts (5). Hence, we conclude that infinite dilution activity coefficients are independent of solute concentration up to solute mole fractions of -10-3. The experimental and analytical methodologies usedmultiple equilibrations with direct aqueous analysis-are
-
-
-
834
Environ. Sci. Technol., Vol. 20, No. 8, 1986
clearly superior to the batch air-stripping technique, in which the interaction between mass transfer rate and equilibrium can confound the interpretation of the data (5). Accordingly, the studies reporting higher Henry’s law constants with increasing solute concentration (7,8) should be regarded with caution, as the air-stripping technique was used; furthermore, there is no theoretical basis that could conceivably explain this effect. Platford (3),on the other hand, reports an almost linear decrease of approximately 20% in the aqueous activity coefficient of CC14for the mole fraction concentration range of 2.5 x to 9 X (approximatelysolubility limit). Although the effect is in the direction expected from theory, the experimental error is on the order of 10-15%, and the results cannot be considered conclusive, as discussed by Horvath (34,p 102). Dobbs and Tavener ( 4 ) tested and confirmed the linearity of equilibrium for CHC1, and CC14 at very low concentrations-from -6 down to pg/L. That study (4), coupled with our own findings, confirms the validity of a fundamental thermodynamic principle: 7,” = constant over the full range of miscibility for slightly soluble compounds. Moreover, a slightly soluble compound can be defined for the present purposes as a compound with an aqueous mole fraction solubility, xis 5 lo-,. The corresponding relative effect estimated from UNIFAC calculations at the solubility limit (Figures 1-3) causes a slight decrease of the activity coefficient compared to the value at infinite dilution (xi = 10-lo): 3.5% for CHCl,, 5 X hydrophobic (larger Ko,,J the solute and/or the cosolvent. Accordingly, it is probably unlikely that less soluble cosolvents, i.e., other solutes, will interact significantly with the solute despite their correspondingly stronger hydrophobic character. For example, UNIFAC predictions indicate that the activity coefficient of CCll at infinite dilution will be reduced by only a few percent in a chloro-
form-saturated aqueous solution compared to the CC4H,O system. Hence, interactions in multisolute systems are not expected to deviate significantly from binary solute-water systems. Accordingly, multisolute transport in groundwater can be modeled on the basis of single and pure compound properties, within the dilute solution range defined in this work. Further, treatment efficiencies in air-stripping processes are expected to be the same whether a given compound is the only contaminant present or one of many as usually encountered in practical situations. Exceptions should be anticipated mainly in the presence of high concentrations of miscible cosolvents; criteria are presented for assessing the likely importance of cosolvent effects. Registry No. CHC13, 67-66-3; CC&, 56-23-5; C2Cl6, 67-72-1; MeOH, 67-56-1; i-PrOH, 67-63-0.
Literature Cited Hwang, H.; Dasgupta, P. K. Enuiron. Sci. Technol. 1985, 19, 255-258. Perry, R. H.; Chilton, C. H. Chemical Engineers’ Handbook, 5th ed.; McGraw-Hill: New York, 1973. Platford, R. F. J. Chem. SOC., Faraday Trans. 1 1977, 73, 267-271. Dobbs, A. J.; Tavener, L. J. Chemosphere 1982,11,465-470. Munz, C. Ph.D. Dissertation, Stanford University, Stanford, CA, 1985. Mackay, D.; Shiu, W. Y.; Sutherland, R. P. Enuiron. Sci. Technol. 1979,13, 333-337. Gossett, J. M.; Lincoff, A. H. final report, Air Force Office of Scientific Research, Directorate of Chemical and Atmospheric Sciences, Grant AFOSR-81-0074, 1981. Lalezary, S.; Pirbazari, M.; McGuire, M. J.; Krasner, S. W. J. Am. Water Works Assoc. 1984, 76, 83-87. Nicholson, B. C.; Maguire, B. P.; Bursill, D. B. Enuiron. Sci. Technol. 1984, 18, 518-521. Banerjee, S. Enuiron. Sci. Technol. 1984, 18, 587-591. Leinonen, P. J.; Mackay, D. Can. J. Chem. Eng. 1973,51, 230-233. Yalkowsky, S. H.; Roseman, T. J. In Techniques of Solubilization of Drugs; Yalkowsky, S. H., Ed.; Marcel Dekker: New York, 1981. Herzel, F.; Murty, A. S. Bull. Enuiron. Contam. Toxicol. 1984,32, 53-58. Eganhouse, R. P.; Calder, J. A. Geochim. Cosmochim. Acta 1976,40, 555-561. Tewari, Y. B.; Martire, D. E.; Wasik, S. P.; Miller, M. M. J. Solut. Chem. 1982,11,435-445. Miller, M. M.; Wasik, S. P.; Huang, G. L.; Shiu, W. Y.; Mackay, D. Environ. Sci. Technol. 1985, 19, 522-529. Chiou, C. T.; Schmedding, D. W.; Manes, M. Environ. Sci. Technol. 1982,16, 4-10. Mackay, D.; Shiu, W. Y. J. Phys. Chem. Ref. Data 1981, 10,1175-1199. Smith, J. M.; van Ness, H. C. Introduction to Chemical Engineering Thermodynamics,3rd ed.; McGraw-Hill: New York, 1975. Hansch, C.; Leo, A. J. Substituent Constants for Correlation Analysis in Chemistry and Biology; Wiley: New York, 1979. Fredenslund, A.; Jones, R. L.; Prausnitz, J. M. AIChE J. 1975,21, 1086-1099. Skjold-Jorgensen, S.; Kolbe, B.; Gmehling, J.; Rasmussen, P. Ind. Eng. Chem. Process Des. Dev. 1979,18, 714-722. Gmehling, J.; Rasmussen, P.; Fredenslund, A. Ind. Eng. Chem. Process Des. Deu. 1982,21, 118-127. Almeida-Macedo, E.; Weidlich, U.; Gmehling, J.; Rasmussen, P. Ind. Eng. Chem. Process Des. Deu. 1983, 22, 676-678. Abrams, D. S.; Prausnitz, J. M. AIChE J. 1975,21,116-128. Arbuckle, W. B. Enuiron. Sci. Technol. 1983,17,537-542. Banerjee, S. Enuiron. Sci. Technol. 1985, 19, 369-370. Campbell, J. R.; Luthy, R. G. Enuiron. Sci. Technol. 1985, 19. 980-985. Environ. Sci. Technol., Vol. 20, No. 8 , 1986
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(29) McAuliffe, C. Chem. Technol. 1971,1,46-51. (30) Lincoff, A. H.; Gossett, J. M. In Gas Transfer at Water Surfaces; Brutsaert, W., Jirka, G. H., Eds.; D. Reidel: Boston, MA, 1984. (31) Grob, K.; Habich, A. J. High Resolut. Chromatogr. Chromatogr. Commun. 1983,6, 11-15. (32) Grob, K. J. Chromatogr. 1984, 299, 1-11. (33) Ioffe, B. V.; Vitenberg, A. G. Head-Space Analysis and Related Methods i n Gas Chromatography; Wiley: New York, 1984. (34) Horvath, A. L. Halogenated Hydrocarbons, SolubilityMiscibility with Water; Marcel Dekker: New York, 1982. (35) Nkedi-Kizza, P.; Rao, P. S. C.; Hornsby, A. G. Environ. Sci. Technol. 1985, 19, 975-979. (36) Tokunaga, J. J. Chem. Eng. Data 1975,20,41-46. (37) Alessi, P.; Kikic, I.; Fredenslund, A.; Rasmussen, P. Can. J. Chem. Eng. 1982,60, 300-304. (38) Kikic, I.; Alessi, P.; Rasmussen, P.; Fredenslund, A. Can. J. Chem. Eng. 1980,58, 253-258. (39) Fredenslund, A.; Rasmussen, P. SEP 8419, Instituttet for Kemiteknik, The Technical University of Denmark,
Lyngby, Denmark, 1984. (40) Kehiaian, H. V. Fluid Phase Equilib. 1983, 13, 243-252. (41) Arbuckle, W. B. Environ. Sci. Technol. 1981,15,812-819. (42) Belfort, G. In Chemistry in Water Reuse; Cooper, W. J., Ed.; Ann Arbor Science: Ann Arbor, MI, 1981; Vol. 2, pp 207-241. (43) Leighton, D. T., Jr.; Calo, J. M. J. Chem. Eng. Data 1981, 26, 382-385. (44) McConnell, G.; Ferguson, D. M.; Pearson, C. R. Endeavour 1975,34, 13-18. (45) Hunter-Smith, R. J.; Balls, P. W.; Liss, P. S. Tellus 1983, 25B, 170-176.
Received for review November 18,1985. Accepted April 1,1986. This research was funded in part by the US.Environmental Protection Agency under Assistance Agreement CR-808851 to Stanford University. However, this publication has not been subjected to the Agengy's peer and administrative review and therefore does not necessarily reflect the views of the Agency, and no official endorsement should be inferred.
NOTES Potential Artifacts in the Determination of Metal Partitioning in Sediments by a Sequential Extraction Procedure Franqois Rapln,+Andri Tessier," Peter G. C. Campbell, and Richard Carlgnan
Universit6 du Quebec, INRS-Eau, C.P. 7500, Sainte-Foy, Qugbec, Canada G1V 4C7 The partitioning of trace metals in sediments) as obtained with a sequential extraction procedure, may be affected by (i) the techniques used to preserve the sediments before analysis and (ii) the presencelabsence of atmospheric oxygen during the extraction steps. No storage method tested completely preserved the initial chemical and physical characteristics of the sediments. Drying of the sediment (freeze-drying;oven-drying) should be avoided; acceptable preservation techniques include freezing or short-term wet storage (1-2 "C). Among the different metals (Cd, Co, Cr, Cu, Ni, Pb, Zn, Fe, and Mn), copper, iron, and zinc were particularly sensitive to sample pretreatment. For anoxic sediments) the maintenance of oxygen-free conditions during the extractions is of critical importance. Introduction
In principle, the partitioning of sediment-bound metals could be evaluated both by thermodynamic calculations (provided equilibrium conditions prevail) and by experimental techniques. Although the former approach holds considerable promise in this regard ( I , 2)) it is of limited application at the present time since the thermodynamic data needed for handling the complex sediment-water systems are as yet incomplete. Direct determination of specific sediment-trace metal associations is also difficult, 'Present address: Institut F. A. Forel, Universitg de GenBve, 1290 Versoix, Suisse. 838
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if not impossible, because of the great variety of solid phases that can bind trace metals, their amorphous character, and the low trace metal concentrations involved. As an alternative to metal partitioning models, methods have been suggested for fractionating the sediment chemically (3-8). The partitioning obtained by such procedures is, however, influenced by factors such as the choice of reagents used for the various extractions and the extraction sequence, the time of extraction, the ratio of extractant to sediment and by inherent analytical problems such as incomplete selectivity and readsorption (7,9-11); it follows that the distribution of a metal among various fractions does not necessarily reflect its association with discrete sediment phases, but rather should be considered as operationally defined by the methods of extraction. In addition to these inherent analytical problems, there is also a potential difficulty in preserving sample integrity between the time of sample collection and extraction; sample preservation is of considerable practical importance as there are often unavoidable delays between the two operations. In effect, the importance of maintaining sample integrity has been alluded to in the literature (12,13), but with the exception of the studies reported by Thompson et al. (14) for estuarine sediments) published experimental data are virtually nonexistent. As a result, there exist numerous published reports in which little or no heed has been paid to the possible effects of sample pretreatment on metal partitioning. In this paper we examine the effects of preservation techniques (wet storage, freezing, freeze-drying,oven-drying) on metal partitioning and evaluate the importance of performing the extractions
0013-936X/86/0920-0836$0 1.50/0
0 1986 American Chemical Society
