In the Laboratory
Air–Water Partitioning of Environmentally Important Organic Compounds An Environmental Chemistry or Integrated Laboratory Experiment B.R. Ramachandran, John M. Allen, and Arthur M. Halpern* Department of Chemistry, Indiana State University, Terre Haute, IN 47809 The processes that control chemical transport in the environment are not presently well understood. Of particular current interest is the transport of chemical species across phase boundaries—for example, the air–aqueous interface. Partitioning between the atmosphere and bodies of water is one of the most important processes affecting the transport of many chemical species in the environment, and studies of this process should be incorporated in undergraduate laboratory courses. Partitioning may be quantitatively accounted for by Henry’s law as applied to dilute aqueous solutions. For aqueous solutions that contain many other chemical species, which is ordinarily the case under environmental conditions, the term “air–water distribution ratio” is frequently used and can be approximated by the Henry’s law constant (1). Thus the basic theory governing airwater partitioning is straightforward and is covered in undergraduate general, environmental, and physical chemistry courses. A number of recently published reports have been directed toward improving the experimental methods by which Henry’s law constants are measured. Henry’s law constants have historically been difficult to measure owing to the limitations of available experimental techniques. As a result, many Henry’s law constants reported in the literature have been estimated from other thermodynamic data. However, new approaches, and especially the procedure described here, can make the measurement of Henry’s law constants accessible to virtually any laboratory that has access to a gas chromatograph (2, 3). The motivation for recent studies stems from the need to quantify the partitioning behavior of chemical species in the environment. For example, some investigations have dealt with the partitioning effect of airstripping toxic compounds, such as gasoline and industrial solvents, from contaminated groundwater (4), and of volatile organic compounds (VOC), such as chlorinated solvents, into room or workplace air from tap water (5). Another environmental application is the partitioning behavior of compounds such as hydrogen peroxide, organic peroxides and hydroperoxides, aldehydes, and lowmolecular-weight carboxylic acids between air and the aqueous phase of cloud, fog, and aerosol droplets in the troposphere (6–8). We describe here an elegantly simple yet accurate method of measuring the Henry’s law constant (or partitioning coefficient) of organic compounds in water. The *Corresponding author.
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methodological simplicity of the technique also permits students to determine the temperature dependence of the Henry’s law constant, where desirable. These experiments encourage students to make a cognitive connection between the unfamiliar concept of thermodynamic Henry’s law constants and the more intuitive concepts of solubility and vapor pressure. We and other colleagues have observed that students are more motivated—and even demonstrate considerable enthusiasm—when given the challenge of laboratory experiments that they perceive as being relevant to environmental issues. These experiments are thus well placed in environmental chemistry or physical chemistry laboratory courses and are also appropriate for the integrated or interdisciplinary laboratory curriculum. The Theory Behind the Experiment Close to two hundred years ago, William Henry, an English chemist, observed that the solubility of a gas in a liquid was directly proportional to its pressure in the vapor phase above the liquid (solution). He expressed his results by the following equation, now known as the Henry’s law (9). p i,V = x i,L HP,i
(1)
where pi,V is the partial pressure of component i in the vapor phase, x i,L is its mole fraction in the liquid phase, and HP,i is the proportionality constant, known as the Henry’s law constant, for that component. HP,i and pi,V have the same units. Henry’s law is reasonably well obeyed by gaseous and (volatile) liquid solutes for which (i) there is no reaction with or dissociation in the solvent, (ii) the solubility is low, and (iii) the vapor phase exhibits nearly ideal behavior. If the system (liquid and vapor phases) is in a closed container, eq 1 can be modified as
c i,V =
n i,V H P,i = x i,L V HS RT
(2)
where ni,V and ci,V are the number of moles and molarity, respectively, of component i in the headspace; VHS is the headspace volume; R is the gas constant; and T is the absolute temperature. For dilute solutions of component i (this condition is satisfied for solutions considered here, since we are dealing with gases or organic compounds having low solubilities in water), xi,L can be approximated by c i,L/c s, where ci,L is the molarity of component i in the liquid
Journal of Chemical Education • Vol. 73 No. 11 November 1996
In the Laboratory
phase and cS is the molarity of the pure solvent (here, water). Equation 2 can then be rearranged to read
H P,i c i,V c i,L = c S RT = H i
(3)
The right-hand side of equation 3 is called the dimensionless Henry’s law constant, Hi, which, from the lefthand side of equation 3, can be recognized as the familiar distribution coefficient, K D,i , for a solute i partitioned between two phases (here, the headspace and the liquid phases). It may be noted that Henry’s law constants are also sometimes expressed in units of atm-dm3-mol{1 by multiplying the dimensionless constant H by RT. We will hereafter use the symbol H to denote the dimensionless Henry’s law constant.
Distribution of a Volatile Organic Solute between the Aqueous and Vapor Phases For simplicity we will consider an aqueous solution of a single volatile organic solute contained in a closed vessel. Let VL and VHS represent the volumes of the liquid
We now point out that the equilibrium composition of the VOC in the headspace volume is determined by gas chromatographic analysis, in which the peak area of the analyte, A, is proportional to the vapor concentration of the sample, nV /VHS. Thus, equation 7 shows the central experimental relationship between the dependent variable, A, and the independent variable, the headspace-to-liquid phase ratio, VHS/VL:
AV L =
kn Lo H H(V HS / V L) + 1
VL is separately measured for a given sample, and nLo is the same for all samples, as will be discussed below; k is the proportionality constant between the peak area A and the number of moles of the analyte. Since the magnitude of k depends on the volume of headspace sample injected in the chromatograph, the injection volume must be constant throughout the experiment. Equation 7 can be directly used to determine the dimensionless Henry’s law constant through an appropriate weighted nonlinear regression analysis. The two regression parameters are H and k9 (k9 = knLo), the latter of which is of no physical interest. Equation 7 can be transformed into a linear equation, namely:
1 = 1 (V / V ) + 1 L k′H AVL k′ HS
(aqueous) and vapor (headspace) phases respectively, and let cLo be the bulk molar concentration of the solute before thermal and distribution equilibria are reached. Let cV and cL be its concentrations in the headspace and aqueous phases, respectively, after equilibrium is reached. Also, let nLo represent the bulk amount of the solute in moles and nV and nL the number of moles of the solute in the headspace and in the aqueous phases, respectively. Equation 3 may be rewritten for this twocomponent system as
nV V HS n L =H V L
(4)
A mass balance of the solute in the two phases requires that nLo = n V + n L and thus the following relation emerges:
n V = n Lo
H(V HS / V L) H(V HS / V L) + 1
(5)
Equation 5 forms the basis of a number of analytical methods that have been described for the determination of Henry’s law constants using headspace chromatography (2–4). Our approach represents a useful modification of these techniques, combining experimental simplicity with analytical rigor. It is also a somewhat more sensitive method for determining H values for highly soluble VOCs. Equation 5 may be rearranged as
nV H V = c V V L = n Lo V HS L H(V HS / V L) + 1
(6)
(7)
(8)
in which it can be seen that for a series of solutions with the same bulk amount of solute, nLo , a plot of 1/AVL vs. VHS/VL yields a straight line, and the Henry’s law constant H can be obtained from the slope-to-intercept ratio. The latter values are obtained from an appropriately weighted linear regression analysis (10). 1 In the experimental approach described here, systematically different VHS/VL ratios are developed in four stoppered sample bottles in such a way that all bottles contain the same bulk quantity of the solute. This condition is achieved by injecting equal volumes of a stock solution of the solute into different known volumes of water (solvent) contained in those bottles. After allowing time for the samples to equilibrate, identical vapor volumes are drawn from the headspaces using a gastight, valve-locking syringe. These are analyzed using a gas chromatograph equipped with a flame ionization detector (FID). The FID is a mass flow rate-dependent detector and the signal produced is proportional to the mass of the organic analyte injected.
Effect of Temperature on the Henry’s Law Constant In most cases the solubilities of gases and volatile liquids decrease with increasing temperature (11, 12). Thus, the Henry’s law constant usually increases with temperature and follows a van’t Hoff-type temperature dependence, as given by equation 9. ln H = (A/T) + B
(9)
The constant A is associated with the solvation enthalpy of the gas-phase solute; that is, the formation of the ideally dilute solution from the gas phase. The temperature dependence of the Henry’s law constant can be readily studied as either a part of or an extension of the experiment, since equilibration of the samples takes place in a thermostated bath. These ex-
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In the Laboratory
periments not only permit students to predict the effect of seasonal changes on air–water partitioning of VOCs, but also encourage them to think about the thermodynamics of solvation of these compounds. Experimental Procedure A working solution that contains the organic solute whose value for H is to be determined must first be prepared. Aliquots of this solution are then transferred to several sample bottles, which are prepared so that each has a different liquid-to-headspace ratio. The bottles are allowed to reach thermal equilibrium, and the headspace samples are withdrawn and analyzed by gas chromatography.
Preparation of the Working Solution Since Henry’s law describes the distribution equilibrium between the solution (liquid) phase and the vapor phase (headspace) showing near-ideal behavior, it is necessary that the test solutions used in this experiment are (i) dilute and (ii) below saturation. To this end, two different methods can be used satisfactorily to prepare the working solution in this experiment. In the first method, which we used for toluene, chlorobenzene, and trichloroethylene, a working solution is prepared by injecting 10 µL of the solute into 10 mL of methanol taken in a 10-mL volumetric flask. It is assumed that the trace amounts of methanol that would eventually be involved in the aqueous test solutions prepared from this stock solution do not appreciably affect the thermodynamics of the organic solute–water system. In the second method, which we used for butyronitrile and methylcyclohexane, an excess amount of the solute is added to deionized water taken in a separatory funnel and shaken vigorously for about an hour, and the two-phase system is allowed to stand undisturbed overnight. A portion of the aqueous layer is carefully run into a vial with a septum and one-holed lid just before the preparation of the aqueous test solutions.
Preparation of Samples and Development of Headspaces
that each bottle can hold up to the brim. Additional bottles can be used if higher precision in H is required. The rubber septa should be lined with Teflon tape in order to minimize the adsorption of organic vapors onto the rubber surface. Five, 7.0, 15.0, and 30.0 mL of deionized water are pipetted into the four bottles, and into each of them an appropriate aliquot (e.g., 10.0 µL for method 1 or 100–500 µL for method 2) of the stock solution is injected. The tip of the syringe needle should be kept below the surface of the water during injection. The bottles are then swirled gently for a few seconds; more vigorous agitation should be avoided because such action would strip away the volatile solute from the aqueous phase. The bottles are kept in a constant-temperature bath with occasional momentary swirling for about 30 min, in order for the samples to reach equilibrium. Longer times may be necessary for highly soluble organic solutes. A water bath equipped with a chiller will be required if H values are to be determined at sub-ambient temperatures.
Chromatographic Analysis A Varian model 3300 GC with an FID was used to quantify the organic component in the headspace samples. However, if H values for chlorinated organics are to be determined, an electron capture detector may be more appropriate. In our procedure, a 2 m × 1/8 in. stainless steel OV-101 packed column was used. A capillary column is also suitable. Helium was used as the mobile phase at a flow rate of approximately 30 mL/min. The column oven temperature was set so as to give a retention time of about 2 min for the particular compound being studied. A special 500-µL gas-tight, valvelocking microsyringe (Scientific Glass Engineering model 500R-V-GT) was used to withdraw headspace samples and inject them into the GC. We used 250-µL headspace volume samples in our experiments. Because the removal of the headspace sample may disturb the equilibrium in the sample bottle, sufficient time should be allowed for the reestablishment of equi-
The sample containers used in this experiment are 65-mL narrow-mouthed reagent bottles whose droppers had been removed. Each bottle is equipped with a rubber septum and a one-holed lid. Four such bottles are needed, and their exact capacities should be determined beforehand by determining the mass of deionized water
Table 1. Dimensionless Henry’s Law Constants and Standard Deviations of Several VOC’s Using Head-Space Chromatography at 25.0 °Ca σH H Solute Lit. Value toluene
0.264
0.023
0.263 (2)
chlorobenzene
0.138
0.013
0.155; 0.127 (13)
methylcyclohexane trichloroethylene
2.77 0.484
0.76 0.089
butyronitrile
0.0030
0.0018
5.13 (27.3 °C) (3,4) 0.420 (2); 0.371(20 °C) (13); 0.392 (14) —
a
Analysis performed using weighted linear regression and eq 8. H values can be converted to kPa m3 mol{1 units by multiplying by RT (here, 2.48).
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Figure 1. Plot of 1/AVL vs. VHS / VL for data obtained with trichloroethylene in water at 25.0 °C. See eq 8.
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librium before taking additional samples from the bottle. Calculation of Results Plots of 1/AVL vs. VHS/VL are constructed from the data, preferably using an appropriate computational resource, such as Excel or RS/1. Linear regression analysis of the data can be performed to obtain the slope and intercept, and hence the value of H and its standard deviation. We tested this method of determining Henry’s law constants of five representative organic compounds: toluene, chlorobenzene, methylcyclohexane, trichloroethylene, and butyronitrile. The results are displayed in Table 1, and the data used to determine H for trichloroethylene according to equation 8 are plotted in Figure 1. Note 1. It should be emphasized that in order to obtain correct values of H from either nonlinear or linear regression analysis using eqs 7 and 8, respectively, weighting factors for the dependent variable, A, must be used. In the nonlinear and linear analyses, the weighting factors for the ith area measurement, Ai, are 1/(εAi ) 2 and (Ai / ε)2 , respectively, where ε is the relative error in measuring A (e.g., 4%).
Literature Cited 1. Schwartzenbach, R. P.; Gschwend, P. M.; Imboden, D. M. Environmental Organic Chemistry; Wiley: New York, 1993; pp 109–123. 2. Robbins, G. A.; Wang, S.; Stuart, J. D. Anal. Chem. 1993, 65, 3113–3118. 3. Hansen, K. C.; Zhou, Z.; Yaws, C. L.; Aminabhavi, T. M. J. Chem. Educ. 1995, 72, 93–96. 4. Hansen, K. C.; Zhou, Z.; Yaws, C. L.; Aminabhavi, T. M. J. Chem. Eng. Data 1993, 38, 546–550. 5. Tancrede, M. V.; Yanagisawa, Y. J. Air Waste Manage. Assoc. 1992, 40, 1658– 1662. 6. Lind, J. A.; Kok, G. L. J. Geophys. Res. 1986, 91, 7889–7895. 7. Betterton, E. A.; Hoffmann, M.R. Environ. Sci. Technol. 1988, 22, 1415–1418. 8. Khan, I.; Brimblecombe, P. J. Aerosol. Sci. 1992, 23, 897–900. 9. Carroll, J. J. J. Chem. Educ. 1993, 70, 91–92. 10. Bevington, P. R.; Robinson, D. K. Data Reduction and Error Analysis For the Physical Sciences, 2nd ed.; McGraw–Hill: New York, 1992; pp 58–62. 11. Brown, T. L.; LeMay, H.E., Jr.; Bursten, B. E. Chemistry: The Central Science, 6th ed.; Prentice Hall: Englewood Cliffs, NJ, 1994; pp 464–465. 12. Levine, I. N. Physical Chemistry, 4th ed.; McGraw–Hill: New York, 1995; p 249. 13. Mackay, D.; Shiu, W. Y. J. Phys. Chem. Ref. Data 1981, 10, 1175–1199. 14. Ashworth, R. A.; Howe, G. B.; Mullins, M. E.; Rogers, T. N. J. Hazard. Mater. 1988, 18, 25–36.
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