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Environmental Toxicology and Chemistry, Vol. 18, No. 8, pp. 1637–1642, 1999q 1999 SETAC

Printed in the USA0730-7268/99 $9.00 1 .00

THERMODYNAMIC INVESTIGATION OF TRICHLOROETHYLENE ADSORPTION INWATER-SATURATED MICROPOROUS ADSORBENTS

JAMES FARRELL,* BENJAMIN HAUCK, and MARTIN JONESDepartment of Chemical and Environmental Engineering, University of Arizona, Tucson, Arizona 85721, USA

(Received 14 April 1998; Accepted 24 August 1998)

Abstract—Adsorption of trichloroethylene (TCE) in adsorbents containing hydrophilic and hydrophobic micropores was investigatedin order to determine the mechanisms responsible for TCE adsorption on mineral solids. A high-pressure liquid chromatographymethod was used to measure TCE adsorption isotherms on three microporous adsorbents. Silica gel and zeolite type NaX wereused as hydrophilic model adsorbents, and hexamethyldisilazane (HMDS)-treated silica gel was used as a model hydrophobicadsorbent. Batch uptake and desorption isotherms were also measured on the hydrophilic silica gel. Uptake of TCE by all threeadsorbents was linear over the concentration range investigated. However, the silica gel desorption isotherm was highly nonlinear,as indicated by its Freundlich isotherm exponent of 0.58. Capillary phase separation into hydrophobic micropores was postulatedas being responsible for the isotherm hysteresis. Supporting this hypothesis was the conformance of the TCE adsorption isothermto Dubinin–Radushkevitch volume filling of micropores theory. The enthalpies for TCE adsorption on all three solids were determinedby van’t Hoff analysis of distribution coefficients measured over a temperature range from 5 to 908C. The TCE adsorption enthalpieson the silica gel and HMDS silica gel were exothermic, but on the zeolite adsorption was endothermic. High exothermic adsorptionenthalpies on the silica gel adsorbents indicated that TCE adsorption was occurring in hydrophobic micropores, and that adsorptionon surfaces with large radii of curvature contributed only minimally to the total uptake. This indicates that the predominant mechanismfor TCE adsorption on these mineral solids is not partitioning into the vicinal water layer.

Keywords—Micropore adsorption Trichloroethylene Silica gel Zeolite NaX

INTRODUCTION

The interaction of hydrophobic organic compounds (HOCs)with soils and sediments plays a dominant role in controllingtheir transport and fate in the environment. Several recentinvestigations have reported that micropores, defined as poressmaller than 20 A in diameter, may be responsible for ad-sorption isotherm nonlinearity and slow desorption phenomena[1–3]. These micropores may be associated with either themineral fraction of the adsorbent, or may reside within con-densed organic matter. Although a limited number of inves-tigations have focused on HOC interactions with microporousadsorbents, experimental evidence on HOC–micropore inter-actions at the molecular scale are always indirect, and thus thenature of these interactions is not well understood.

Adsorption of nonpolar organic solutes on soils and sedi-ments is promoted by the hydrophobicity of the solute [4]. Forsmall nonpolar compounds such as trichloroethylene (TCE),the hydrophobic effect can be attributed to reduced entropy ofthe water comprising the solvation envelope around the solute[4]. This conclusion is supported by the exothermic dissolutionof TCE in water, which indicates that TCE–water interactionsare stronger than TCE interactions with itself [5]. IncreasedTCE interactions with water can be attributed to the additionalDebye (dipole-induced dipole) contribution to the total attrac-tive van der Waals interaction energy.

Enthalpy changes accompanying adsorption can be used toexamine the molecular scale interactions associated with the

* To whom correspondence may be addressed(farrell@engr.arizona.edu).

Presented at the 218th Meeting of the American Chemical Society,Las Vegas, NV, USA, September 7–11, 1997.

adsorption process. However, because adsorption involves re-moving the solute from water and displacing water from theadsorption site, the observed adsorption enthalpy is composedof contributions from solute–surface and solvent–surface in-teractions, as well as solute–solvent and solvent–solvent in-teractions [6]. Therefore, the observed enthalpy change uponadsorption (DHobs) can be broken down as

DHobs 5 2 2solute solvent soluteDH DH DHint int diss (1)

where and are the enthalpy changes associatedsolute solventDH DHint int

with formation of solute–surface and solvent–surface inter-actions, respectively, and is the enthalpy change as-soluteDHdiss

sociated with dissolution of the solute in water. For assessmentof sorbate–surface interactions, several previous investigatorshave used a paradigm similar to Equation 1, except that the

term was not included, presumably because of thesolventDHint

assumption that water is not displaced from mineral surfacesby nonpolar HOCs [6,7].

Because of the presence of polar surface functional groups,most mineral surfaces are hydrophilic and are coated withseveral monolayers of structured water [8]. Because nonpolarcompounds cannot displace strongly adsorbed water from min-eral surfaces, some investigators have concluded that adsorp-tion of HOCs occurs by absorption into the vicinal water layer,and there is little or no interaction with the mineral surfaceitself [4]. This hypothesis has been supported by data fromseveral investigations [9–11]. For example, a recent investi-gation of polycyclic aromatic hydrocarbon adsorption on min-eral solids found that the sign and magnitude of the mineralsurface charge did not affect the amount of adsorption [7]. Ifthe HOCs did interact with the mineral surface, the charge-induced dipole interaction should have resulted in greater HOC

1638 Environ. Toxicol. Chem. 18, 1999 J. Farrell et al.

Fig. 1. Conversion of the hydroxyl-terminated silica gel surface to anonpolar trimethylsilyl-terminated surface by reaction with hexame-thyldisilazane.

adsorption with increasing surface charge density. The factthat the surface charge density did not affect the amount ofadsorption supports the contention of minimal HOC interactionwith hydrophilic mineral surfaces. Under this scenario, HOCadsorption into the vicinal water layer occurs with a smallenthalpy change, due to weak HOC–surface interactions, andthe fact that the solute remains surrounded by the solvent [4].

For HOC adsorption on nonpolar surfaces, weak adsorbate–surface interactions lead to small [12]. Additionally,soluteDHint

displacing water from an adsorption site on a nonpolar surfaceis expected to occur with a small , because the watersolventDHint

molecules that were adjacent to the surface before adsorptionare adjacent to a nonpolar HOC after adsorption. These factorshave lead several previous investigators to assert that for weaksolute–surface interactions, is primarily responsible forsoluteDHdiss

DHobs [6,9]. However, this relationship may not hold for HOCadsorption in hydrophobic micropores. Although HOC–ad-sorbent interactions in hydrophobic micropores may also beweak, displacement of water from these pores will be exo-thermic [13], because of the increased electrostatic interactionsof the displaced water molecules with the bulk aqueous phase.

In contrast to surface adsorption, which is promoted byinterfacial energy changes, adsorption in micropores may bepromoted by both surface energy and structural forces. Thiscan be seen from the expression for the internal energy change(dU) accompanying a given process for a fluid confined be-tween two flat surfaces, as given by [14]

dU 5 2p dV 1 T dS 1 m dn 1 2G dA 2 fA dH (2)

where p, V, T, S, n, and m represent pressure, volume, tem-perature, entropy, mole number, and chemical potential, re-spectively. The first three terms on the right hand side ofEquation 2 operate in all systems, whereas the last two termsoperate only where interfaces are involved. The penultimateterm in Equation 2 accounts for changes in internal energyarising from interfacial energies (G) and changes in interfacialarea (dA), whereas the last term arises from structural forces(f), and represents the energy change accompanying a changein separation distance (dH) of interfaces of area A. Interfacialforces promote surface adsorption in pores of any size, whereasstructural forces are significant only where pore surfaces areseparated by distances less than 20 molecular diameters [15].

Structural forces arise from unfavorable packing geometriesbetween surfaces separated by distances that are nonintegralmultiples of the molecular diameter. For adsorption in micro-pores, structural forces may create hydrophobic regions be-tween polar surfaces. For most liquids, structural forces decayin a damped oscillatory manner with increasing distance be-tween pore walls [16]. However, for water and other associ-ating liquids, hydrogen bond effects give rise to an additionalmonotonic component of the structural force that is super-imposed on the oscillatory component [17,18]. For hydro-phillic surfaces, the monotonic component is repulsive, andarises from an increase in the local fluid density; whereasbetween hydrophobic surfaces, the monotonic component isattractive, and arises from a decrease in the local fluid density[16]. Experimental measurements indicate that at separationson the order of 10 A, the monotonic structural force is similarin magnitude to the interfacial tension between water and non-polar hydrocarbons (;30 mJ/m2) [16]. Therefore, adsorptionof hydrophobic solutes may be promoted by the reduction inpotential energy accompanying replacement of water-filled mi-cropores by hydrocarbon-filled micropores. Adsorption of this

type, promoted solely by structural forces, has been exhibitedin Monte Carlo simulations of binary fluid mixtures [19].

Understanding adsorption and transport phenomena ofHOCs in water-saturated microporous materials is essential inpredicting the behavior of organic contaminants in ground-water aquifers. Two previous investigations of TCE adsorptionon silica gel indicated that molecular-sized pores were re-sponsible for TCE isotherm nonlinearity and desorption resis-tance [1,2]. Continuing that work, this research investigatesTCE adsorption in hydrophilic and hydrophobic microporesin order to understand the mechanisms affecting organic soluteinteractions with mineral solids.

MATERIALS AND METHODS

Adsorption of TCE was investigated on three microporousadsorbents: silica gel, hydrophobic silica gel, and the zeolitetype NaX. The silica gel, Davisil, was obtained from Alltech(Deerfield, IL, USA) and had a BET measured surface area of297 m2/g. The zeolite was prepared from sodium aluminate,sodium metasilicate, and the templating agent triethanolamine,according to the method of Charnell [20]. The BET measuredsurface area of the zeolite was 780 m2/g. The hydrophobicsilica gel was prepared by treating the silica gel with hexa-methyldisilazane (HMDS), according to the method of Uhrovaand Malicek [21]. The HMDS treatment converts polar hy-droxyl groups on the silica surface to nonpolar methyl groups,as illustrated in Figure 1. Because the HMDS silica gel isunstable at high temperatures, the surface area of the hydro-phobic silica gel could not be determined by the BET methodused to characterize the other adsorbents. However, Uhrovaand Malicek [21] estimated that the silanizing procedure resultsin a surface area reduction of approximately 50%. The un-treated silica gel contains mesopores up to 150 A in diameter,and has a mesopore volume of 0.76 ml/g. The micropore vol-ume of the untreated silica gel was determined in a previousinvestigation to be 0.11 ml/g [2]. The zeolite is completelymicroporous with a total micropore volume of 0.8 ml/g [22].However, 16.7% of the micropore volume resides within so-dalite cages, which are not accessible to adsorbing species [23].The remainder of the NaX micropore volume is accessible toboth TCE and water, and has pore channel diameters of 7.8A, and cage diameters of 12 A [22].

For the hydrophilic silica gel, TCE adsorption was mea-sured using a batch, long-term equilibration method, and adynamic high-pressure liquid chromatography (HPLC) meth-od. In the batch method, 25-cm-long by 0.9-cm-internal-di-ameter (i.d.) stainless steel columns were packed with ap-proximately 5 g of oven-dried silica gel. The columns werethen water saturated with deionized water, and purged withnitrogen gas to create approximately 30% gas-filled porosity.

Trichloroethylene adsorption in microporous adsorbents Environ. Toxicol. Chem. 18, 1999 1639

Fig. 2. Comparison of batch measured trichloroethylene (TCE) ad-sorption and desorption isotherms on silica gel at 308C.

This treatment resulted in an unsaturated column, packed withwater-saturated silica gel particles. The final water loading inthe columns of more than 1 ml of water per gram of silicawas greater than the internal pore volume of the silica gel,indicating that the adsorbent pores were fully water saturated.The columns were then loaded with TCE by measuring a TCEvapor breakthrough front, as previously described [1]. Thecolumns were then sealed with stainless steel caps and allowedto equilibrate with the TCE vapor for 1 month. After the 1-month equilibration period, the equilibrium vapor phase con-centrations in the columns were determined by purging thecolumns with 10 ml/min of humidified nitrogen gas througha flame ionization detector connected to an SRI Instruments(Torrance, CA, USA) gas chromatograph (GC). The columnswere purged until the signal reached the detection limit on theflame ionization detector. At that point, the adsorbed TCEremaining in the columns was recovered by thermally desorb-ing the columns at 1808C onto a Tenaxt (Chrompack, Bergenop Zoon, The Netherlands) adsorbent trap. The TCE recoveredon the Tenax trap was then quantified by a Hall electrolyticconductivity detector (SRI Instruments, Torrance, CA, USA)attached to the SRI GC. The Henry’s constant and the waterloading on the adsorbent were used to convert the vapor iso-therms to adsorption isotherms from the aqueous phase [1].Desorption isotherms for the silica gel were measured by therepeated purging of two silica gel columns through a flameionization detector. A 1-month equilibration period was al-lowed between successive desorption steps. The desorptionisotherms were then constructed from data obtained duringeach sequential purging, according to methods previously de-scribed [1].

Adsorption isotherms for all three adsorbents were mea-sured using a dynamic HPLC method. For these measurements,the silica gel and zeolite were packed in 30-cm-long by 0.7-cm-i.d. stainless steel columns, and the HMDS silica gel waspacked in a 5-cm-long by 0.4-cm-i.d. PEEKy (Alltech) col-umn. Isotherms were determined from the response peaks of250- to 500-ml injections of aqueous TCE solutions. The ef-fluent concentration profiles were measured with a BeckmanUV absorbence detector (Beckman, Fullerton, CA, USA) op-erated at 210 nm. Pulse concentrations were varied between1 and 140 mg/L to check for isotherm linearity. Gravimetricmeasurements and the response to pulses of the nonadsorbingtracer, sodium nitrate, were used to determine the void volumesof the columns. The method of moments was used to analyzethe measured response profiles to determine the dimensionlessequilibrium partitioning constant, K. The partitioning constantis related to the more familiar distribution coefficient, Kd, byKd 5 K/rg, where rg is the grain density of the adsorbent. Theequilibrium constant is related to the first moment (s) of thepulse response profile by the relation [24]

`

C · t dtE0 L

s 5 5 [e 1 (1 1 e)K] (3)` eu

C dtE0

where C is the aqueous phase solute concentration, L is thecolumn length, e is the column void fraction, and u is theinterstitial water velocity. The equilibrium constant was de-termined by measuring the first moment at several flow ratesand plotting s versus (eu)21. Adsorption enthalpies were de-termined by measuring adsorption isotherms over a range of

temperatures from 278 to 363 K. The columns were maintainedunder isothermal conditions by placement in an insulated waterjacket connected to a circulating water bath.

RESULTS AND DISCUSSION

The mechanism responsible for TCE adsorption to a hy-drophilic microporous adsorbent was investigated by measur-ing batch adsorption and desorption isotherms for TCE onsilica gel. Figure 2 compares the batch measured uptake iso-therm to the batch measured desorption isotherm. For the ad-sorption isotherm, the Freundlich isotherm exponent is equalto 1.02 6 0.02 (695% confidence interval). At the 95% con-fidence level, this number is statistically identical to one, in-dicating linear adsorption. For the desorption isotherm, theFreundlich isotherm exponent is 0.58 6 0.03, indicating non-linear desorption. The linear adsorption isotherm is indicativeof a constant adsorption energy, whereas the desorption iso-therm indicates increasing adsorption energy with decreasingaqueous concentration. Because physical adsorption on opensurfaces is a thermodynamically reversible process [25], thepresence of hysteresis suggests that TCE adsorption on thissilica gel is occurring in micropores, and not on more opensurfaces with large radii of curvature.

Hysteretic adsorption is a well-known phenomenon for ad-sorption of gases on porous adsorbents, and results from theformation of a curved meniscus between the adsorbed fluidand the bulk gas phase [26]. When the adsorbed fluid wets thepore surface, the curved interface leads to a reduction in vaporpressure of the pore fluid in accordance with the well-knownKelvin equation [27]. An analogous situation exists for soluteadsorption in micropores. As the micropores become filled,rearrangement of the solute molecules leads to formation of acurved meniscus between the solute phase in the microporeand the bulk solvent phase. This process is known as capillaryphase separation and leads to hysteretic adsorption [28]. Ad-sorption is hysteretic because the solute–solvent interfacialenergy changes associated with the adsorption process differfrom those associated with the desorption process. For a solutephase that wets the pore surfaces, the increase in interfacialarea associated with solute desorption is greater than the cor-responding decrease associated with adsorption. The surfaceenergy associated with increasing the interfacial area betweenthe solute and solvent phases contributes to a lowered aqueoussolubility (S) for the solute, as given by the Freundlich–Ost-wald equation [28]

22Gv cos uS 5 S exp (4)o [ ]rRT

1640 Environ. Toxicol. Chem. 18, 1999 J. Farrell et al.

Fig. 3. Batch measured trichloroethylene adsorption isotherm on sil-ica gel plotted in Dubinin–Radushkevitch format.

Fig. 4. van’t Hoff analysis of the solid–water distribution coefficientsfor trichloroethylene adsorption on silica gel, hexamethyldisilazane(HMDS) silica gel, and zeolite NaX.

where So is the solubility of bulk TCE, G is the TCE–waterinterfacial tension, v is the molar volume of TCE, r is theradius of curvature, and R is the universal gas constant. Forcontact angles ranging from 0 to 458, G 5 34.5 mN/m2 [29],and v 5 93.2 ml/mol, Equation 4 predicts an aqueous solubilitybetween 6.2 and 7.7% of So in a pore with a 10-A radius ofcurvature. An environmental consequence of this solubilityreduction is a smaller concentration gradient for desorption,and thus slower desorption rates.

To investigate the hypothesis of a capillary phase separationmechanism in the silica gel micropores, the adsorption iso-therm depicted in Figure 2 is replotted in Figure 3 accordingto the Dubinin–Radushkevitch (D-R) volume filling of micro-pores theory. The D-R isotherm is given by [30,31]

2B C

2ln V 5 ln V 2 T ln (5)o 2 [ ]b Csat

where V is the equivalent liquid volume of adsorbed solute,Vo is the adsorbent micropore volume, C is the solute con-centration, Csat is the aqueous solubility of the solute, b is anaffinity parameter that characterizes solute–surface interac-tions, and B is a parameter that is dependent on the porestructure of the adsorbent. The D-R theory is based on thePolanyi adsorption potential concept in which partitioning ofthe solute between the solution and micropore phases followsa Boltzmann distribution. The D-R theory assumes the adsor-bent contains sites with a distribution of adsorption potentials,and the adsorbate in a micropore is considered to be in thecondensed state. The adsorption potential of a particular siteprovides the reduction in chemical potential necessary for con-densing a neat solute phase from an undersaturated solution.For a solute obeying Henry’s law, this reduction in chemicalpotential is given by

(msat 2 m) 5 RT ln(Csat/C) (6)

where msat and m are the chemical potentials of the solute insolutions with concentrations of Csat and C, respectively.Therefore, the net free energy change upon adsorption is equalto the difference between the adsorption potential of a partic-ular site, and the energy required for condensation from theundersaturated solution, as given by Equation 6. When theuptake isotherm is plotted in D-R format, the free energy ofadsorption can be determined from the isotherm slope usingthe relation [31]

2Rb CsatDG 5 2 1 RT ln (7)1 24BT rs

where rs is the solute liquid phase density. For the data inFigure 3, the slope of the D-R isotherm yields an adsorptionenergy of 224.2 6 0.75 kJ/mole. This value is greater thanthe excess Gibbs energy of solution for TCE in water, whichis 21.7 kJ/mole at 308C, and is greater than that normallyassociated with hydrophobic adsorption on open surfaces. Be-cause adsorbed molecules retain some contact with the bulksolution, only ;20 to 33% of the excess Gibbs energy ofsolution is normally recovered upon hydrophobic adsorptionon open surfaces [4,32]. However, in pores of molecular di-mensions, adsorption energies may be increased through agreater shielding of the adsorbed solute from contact with theaqueous phase.

The enthalpy for TCE adsorption on the silica gel wasdetermined by measuring the solid–water distribution coeffi-cient (Kd) for TCE as a function of temperature using the HPLCtechnique. The 308C Kd value of 3.5 ml/g measured with theHPLC technique was close to the batch measured value of 3.6ml/g. This indicates that both techniques were measuring thesame phenomena. The HPLC measured Kd values are plottedas a function of reciprocal temperature in the van’t Hoff plotin Figure 4. The observed enthalpy of adsorption can be de-termined from the slope of the van’t Hoff plot and the tem-perature dependence of the TCE aqueous phase activity co-efficient (g) from

d ln(K /g)dDH 5 2R (8)obs d(1/T )

Application of Equation 8 to the silica gel Kd values yields anexothermic adsorption enthalpy of 222.2 6 1.5 kJ/mole. Forthis calculation, the temperature dependence of g was deter-mined from the temperature dependencies of the TCE Henry’sconstant, H [33], the TCE vapor pressure, Po [34], and themolar volume of water, vw [35], according to g 5 (131.4H)/(Povw). An exothermic adsorption enthalpy of this magnitudeis not consistent with TCE adsorption on hydrophilic or hy-drophobic surfaces, but is consistent with adsorption in hy-drophobic micropores. For adsorption in a hydrophobic mi-cropore, is expected to be small, and is expectedsolute soluteDH DHint diss

to be similar to that associated with transfer of TCE from waterto a bulk TCE phase, which is an endothermic 6.3 kJ/mol at258C. Therefore, the major contributor to DHobs appears to bethe exothermic enthalpy associated with displacement of waterfrom hydrophobic micropores. This phenomenon has been pre-viously observed for displacement of water from nonpolarcavities in dissolved proteins [36]. Assuming a dielectric con-

Trichloroethylene adsorption in microporous adsorbents Environ. Toxicol. Chem. 18, 1999 1641

stant of four for the hydrophobic region, Ernst et al. [37]calculated the enthalpy change associated with transferringwater molecules from hydrophobic protein cavities to bulkwater as 211.3 kJ/mole. The presence of hydrophobic micro-pores on a nominally hydrophilic silica gel may be attributedto small regions with siloxane (Si-O-Si) surface functionalgroups, or to structural forces rendering the pores hydrophobic.

The high exothermic adsorption enthalpy for TCE on silicagel is inconsistent with many previous investigations of hy-drophobic adsorption on mineral surfaces, which reported en-thalpies near zero, after correcting for the dissolution enthal-pies of the HOCs in water [9–11]. The results here are alsoinconsistent with adsorption enthalpies for other HOCs mea-sured on silica gel. For example, Noll [38] reported an enthalpyof 20.11 kJ/mol for benzene adsorption on silica gel. The lowvalue reported by Noll can most likely be attributed to thehigh adsorbed phase concentrations at which the adsorptionenthalpy was measured. The Noll investigation measured ad-sorption from aqueous solutions saturated with benzene, andobtained adsorbed phase concentrations that were approxi-mately two orders of magnitude greater than those measuredhere. Because hydrophobic micropores contain water at lowerconcentrations than in bulk solution [36], the water may becompletely displaced from hydrophobic micropores at low ad-sorbed phase HOC concentrations. Therefore, after the wateris displaced from the hydrophobic micropores, subsequent ad-sorption occurs as a low enthalpy process. This conclusion issupported by results from two recent investigations that re-ported decreasingly exothermic adsorption enthalpies with in-creasing adsorbed phase concentrations. For example, Werthand Reinhard [2] reported isosteric heats of adsorption thatincreased from 29.2 to 245 kJ/mol with decreasing adsorbedphase concentration for TCE desorption from the same silicagel adsorbents used in this study. A second example of thisphenomenon comes from a recent investigation utilizing anHPLC technique to measure PAH adsorption on the pure min-eral solids, hematite and corundum [7]. In that investigation,Mader et al. found that adsorption enthalpies on hematite wereless exothermic the greater the adsorbate Kd value. For ex-ample, naphthalene adsorption was exothermic with an ad-sorption enthalpy of 215 kJ/mol, whereas adsorption of phen-anthrene, anthracene, and pyrene was endothermic, with en-thalpies of 1.2, 2.9, and 8.5 kJ/mol, respectively. The exo-thermic adsorption of naphthalene may be attributed to itslower adsorbed phase concentration, because its Kd value rang-es between 31 and 235 times smaller than those for the othercompounds.

Adsorption of TCE on a hydrophobic adsorbent was in-vestigated using the HMDS silica gel. As indicated by the Kd

values in Figure 4, TCE adsorption on hydrophobic silica gelwas approximately two orders of magnitude greater than thaton the untreated silica gel. However, the heat of adsorption onthe hydrophobic silica gel of 220.8 6 1.7 kJ/mole was similarto that observed for the hydrophilic silica gel. This adsorptionenthalpy is greater than those normally observed for hydro-phobic adsorption on mineral surfaces, and for partitioninginto soil organic matter, for which corrected enthalpy valuesnear zero have been reported [4,6]. Additionally, sorption en-thalpies for partitioning into reversed phase silica gels havebeen reported to be small. For example, Noll [38] reported anenthalpy of only 21.35 kJ/mole for partitioning of benzeneinto a C18 phase bonded onto silica. The difference betweenthis study and that of Noll is that the methyl-terminated sur-

faces of the HMDS silica apparently do not permit bulk phasepartitioning, as did the C18 phase. The high exothermic heatof adsorption found in this study is consistent with adsorptioninto glassy organic matter [12], which has been reported tocontain hydrophobic pores of molecular dimensions [3]. An-other conclusion that may be drawn from the exothermic ad-sorption enthalpy on the HMDS silica gel is that adsorptionof TCE on open hydrophobic surfaces contributes only min-imally to the total uptake, because this process should be en-dothermic, with an enthalpy close to that associated with re-moving TCE from water.

Experiments measuring TCE adsorption on zeolite NaXwere performed to investigate TCE adsorption in completelyhydrophilic micropores. Zeolite NaX is completely hydrophilicbecause the high charge density of the zeolite framework re-sults in charged sites too closely spaced to accommodate TCEadsorption at uncharged areas. Comparison of Kd values inFigure 4 for the three adsorbents indicates that TCE adsorptionon the zeolite is greatly reduced, despite having a specificsurface area more than twice as great as the silica gels, andan accessible micropore volume more than six times as greatas the hydrophilic silica gel. In contrast to the behavior on thesilica gels, TCE adsorption on the zeolite is endothermic, withan adsorption enthalpy of 20.3 6 12.1 kJ/mole. Endothermicheats of adsorption are known to occur for solutes that are lessstrongly adsorbed than the accompanying solvent [39]. Thus,TCE adsorption on the zeolite is endothermic because it re-quires the displacement of water from hydrophilic pores.

CONCLUSION

The high exothermic heats of adsorption observed for TCEon the hydrophilic and hydrophobic silica gels indicate thatadsorption is occurring in hydrophobic micropores, with theconcomitant displacement of water. Hysteretic adsorption onthe silica gel, and the conformance of the uptake isotherm tothe D-R pore-filling model also support the hypothesis of mi-cropore adsorption. Thus, over the concentration range inves-tigated, the predominant mechanism for TCE adsorption onthese mineral solids is not partitioning into the vicinal waterlayer. Adsorption in hydrophobic micropores may explain thelack of correlation between mineral surface area and adsorbateuptake that has often been observed in studies on HOC ad-sorption by mineral solids [7,40].

Acknowledgement—We would like to thank the two anonymous re-viewers and Charlie Werth for their helpful comments. This work wassupported by the U.S. National Science Foundation Chemical andTransport Systems Program through grant CTS-9624724.

REFERENCES

1. Farrell J, Reinhard M. 1994. Desorption of halogenated organicsfrom model solids, sediments, and soil under unsaturated con-ditions. Environ Sci Technol 28:53–72.

2. Werth CJ, Reinhard M. 1997. Effects of temperature on trichlo-roethylene desorption from silica gel and natural sediments. 1.Isotherms. Environ Sci Technol 31:689–696.

3. Xing B, Pignatello JJ. 1997. Dual-mode sorption of low-polaritycompounds in glassy poly(vinyl chloride) and soil organic matter.Environ Sci Technol 31:792–799.

4. Schwarzenbach RP, Gschwend PM, Imboden DM. 1993. Envi-ronmental Organic Chemistry. John Wiley & Sons, New York,NY, USA.

5. Werth CJ. 1997. Analysis of trichloroethylene desorption fromsilica gel and natural sediments using column studies, numericalmodeling, and nuclear magnetic resonance spectroscopy. PhDthesis. Stanford University, Stanford, CA, USA.

1642 Environ. Toxicol. Chem. 18, 1999 J. Farrell et al.

6. Chiou CT, Peters LJ, Freed VH. 1979. A physical concept of soilwater equilibria for nonionic compounds. Science 26:831–832.

7. Mader BT, Uwe-Goss K, Eisenreich SJ. 1997. Sorption of non-ionic, hydrophobic organic chemicals to mineral surfaces. En-viron Sci Technol 31:1079–1086.

8. Drost-Hansen W. 1969. Structure of water near solid interfaces.Ind Eng Chem 61:10–47.

9. Mills AC, Biggar JW. 1969. Adsorption of 1,2,3,4,5,6-hexachlo-rocyclohexane from solution: The differential heat of adsorptionapplied to adsorption from dilute solutions on organic and in-organic surfaces. J Colloid Interface Sci 29:720–731.

10. Boucher FR, Lee GF. 1972. Adsorption of lindane and dieldrinpesticides on unconsolidated aquifer sands. Environ Sci Technol6:538–543.

11. van Bladel R, Moreale A. 1974. Adsorption of fenuron and mon-uron (substituted ureas) by two montmorillonite clays. Soil SciSoc Am Proc 38:244–249.

12. Luthy RG, et al. 1997. Sequestration of hydrophobic organiccontaminants by geosorbents. Environ Sci Technol 31:3341–3347.

13. Matthews BW, Morton AG, Dahlquist FW. 1995. Use of NMRto detect water within nonpolar protein cavities. Science 270:1847–1849.

14. Evans R, Marconi UMB. 1987. Phase equilibria and solvationforces for fluids confined between parallel walls. J Chem Phys86:7138–7148.

15. Christensen HK, Claessen PM. 1988. Cavitation and the inter-action between macroscopic hydrophobic surfaces. Science 239:390–392.

16. Israelachvili JN. 1992. Intermolecular and Surface Forces, 2nded. Academic, New York, NY, USA.

17. Pashley RM. 1981. DLVO and hydration forces between micasurfaces in Li1, Na1, K1, and Cs1 electrolyte solutions: A cor-relation of double-layer and hydration forces with surface cationexchange properties. J Colloid Interface Sci 83:531–546.

18. Lis LJ, McAlister M, Fuller N, Rand RP, Parsegian VA. 1982.Interactions between neutral phospholipid bilayer membranes.Biophys J 37:657–665.

19. MacElroy JMD, Suh SH. 1987. Computer simulation of moder-ately dense hard-sphere fluids and mixtures in microcapillaries.Mol Phys 60:475–501.

20. Charnell JF. 1971. Gel growth of large crystals of sodium A andsodium X zeolites. J Crystal Growth 8:291–294.

21. Uhrova M, Malicek O. 1981. Hydrophobization of supports forextraction chromatography II. Hydrophobization of silica gel withhexamethyldisilazane. Sb Vys Sk Chem Technol Praze Anal ChemH-16:105–115.

22. Chen NY, Degnan TF Jr, Smith CM. 1994. Molecular Transportand Reaction in Zeolites. VCH, New York, NY, USA.

23. Awum F, Narayan S, Ruthven D. 1988. Measurement of intra-

crystalline diffusivities in NaX zeolite by liquid chromatography.Ind Eng Chem Res 27:1510–1515.

24. Ching CB, Ruthven DM. 1988. A liquid phase chromatographicstudy of sorption and diffusion of glucose and fructose in NaXand KX zeolite crystals. Zeolites 8:68–73.

25. Adamson AW. 1982. Physical Chemistry of Surfaces. John Wiley& Sons, New York, NY, USA.

26. Gregg SJ, Singh KSW. 1982. Adsorption Surface Area and Po-rosity. Academic, New York, NY, USA.

27. Alberty RA, Silbey RJ. 1996. Physical Chemistry, 2nd ed. JohnWiley & Sons, New York, NY, USA.

28. Miyahara M, Okazaki M. 1992. An estimation of liquid phaseadsorption isotherms based on the capillary phase separation con-cept. Proceedings, Fundamentals of Adsorption, 4th InternationalConference of Fundamentals of Adsorption, Kyoto, Japan, May17–22, pp 80–85.

29. Weiss G. 1988. Hazardous Chemicals Data Book, 2nd ed. NoyesData, Park Ridge, NJ, USA.

30. Stadnik AM, El’tekov YA. 1978. Calculation of the free energyof adsorption by the theory of volume filling of micropores. RussJ Phys Chem 52:1041–1042.

31. Derylo-Marczewska A, Jaroniec M. 1987. Adsorption of organicsolutes from dilute solutions on solids. In Matijevic E, ed, Surfaceand Colloid Science, Vol 14. Plenum, New York, NY, USA.

32. Melander W, Horvath C. 1980. Thermodynamics of hydrophobicadsorption. In Suffet IH, McGuire MJ, eds, Activated CarbonAdsorption of Organics from the Aqueous Phase, Vol 1. AnnArbor Science, Ann Arbor, MI, USA, pp 65–69.

33. Gossett JM. 1987. Henry’s law constants for C1 and C2 chlori-nated hydrocarbons. Environ Sci Technol 21:202–208.

34. Jordan TE. 1954. Vapor Pressure of Organic Compounds, In-terscience, New York, NY, USA.

35. Streeter VL, Wylie EB. 1979. Fluid Mechanics, 7th ed. McGraw-Hill, New York, NY, USA.

36. Matthews BW, Morton AG, Dahlquist FW. 1995. Use of NMRto detect water within nonpolar protein cavities. Science 270:1847–1849.

37. Ernst JA, Clubb RT, Zhou H, Gronenborn AM, Clore GM. 1995.Demonstration of positionally disordered water within a proteinhydrophobic cavity by NMR. Science 267:1813–1817.

38. Noll LA. 1987. Adsorption of aqueous benzene onto hydrophobicand hydrophilic surfaces. Colloids Surf 28:327–329.

39. Lin YS, Ma YH. 1989. A comparative study of adsorption anddiffusion of vapor alcohols and alcohols from aqueous solutionsin silicalite. In Jacobs PA, van Santen RA, eds, Zeolites: Facts,Figures, Future, Proceedings, 8th International Zeolite Confer-ence, Amsterdam, The Netherlands, July 10–14, pp 877–885.

40. Huang W, Schlautman MA, Weber WJ Jr. 1996. A distributedreactivity model for sorption by soils and sediments. 5. The in-fluence of near-surface characteristics in mineral domains. En-viron Sci Technol 30:2993–3000.