Adsorption of Organic Compounds from Dilute Aqueous Solutions onto the External Surface of Type A Zeolite Anthony C. Savitsky, Brandon H. Wiers, and Richard H. Wendt" The Procter and Gamble Co., Packaged Soap and Detergent Division, Cincinnati, Ohio 45217
Type A zeolite, an insoluble crystalline sodium aluminosilicate, has been developed as a water-hardness controlling agent for laundry detergent products. In order to assess the environmental safety of this material, we have examined the adsorption of organic compounds of varying polarity onto the external surfaces of calcium-exchanged type A zeolite in aqueous medium at pH 8. The compounds included p-chlorophenol, dieldrin, dodecylhexakis(oxyethy1ene) alcohol (ClzE6), and tetradecyltrimethylammonium chloride. In addition, adsorption onto the sodium form of the zeolite was studied with methylene blue (to characterize the zeolite) and with linear alkylbenzenesulfonate (LAS). Type A zeolite provides a relatively weak adsorptive surface for organics of all polarities studied when compared to adsorption of these compounds onto natural clay minerals and sewage sludge. The introduction of type A zeolite into environmental waters through its use in laundry detergents should therefore not alter the normal fate of organic molecules in the environment. Type A zeolite has been developed for use as a builder in laundry detergents. Although insoluble, this synthetic zeolite, Na12[(A102)(Si02)]12-27HzO, is capable of reducing the hardness of wash water by the exchange of calcium ion for sodium ion, thus promoting more effective utilization of surfactants in the detergency process. The use of laundry products containing this material will result in its introduction into sewage at low ppm levels. Most studies of the chemical properties of synthetic zeolites have emphasized their gas-adsorption properties (1-4). Several studies on the environmental considerations of this zeolite have been, or will soon be, published. These studies show that type A zeolite is almost entirely removed from wastewater by conventional treatments and that it does not adversely affect the sewage treatment process (5-9). Other studies show that any zeolite in sewage effluent will not present a hazard to aquatic life (10-12) and will hydrolyze irreversibly to silica, aluminum oxide, and natural aluminosilicate (13,14). In our study we have measured the adsorption properties of type A zeolite for various classes of organic compounds and compared these properties with those of minerals that are naturally present in the environment. This information defines the relative potential of type A zeolite to affect the fate of organic compounds in the envrionment. The tests were conducted mostly on the calcium form of the zeolite because this is the predominant form in which the material is discharged from laundry systems. The chemical classes of organics that were studied include three types of surface-active agents (anionic, nonionic, and cationic) as well as a phenol, a pesticide, and a dye. These materials provided a broad representation of the chemical classes of organic compounds that might be found in sewage or surface waters. Measurements of their adsorption properties onto type A zeolite permit their effects on the fate of organics in wastewater to be assessed. 0013-936X/81/0915-1191$01.25/0
@ 1981 American Chemical Society
Experimental Section Experiments to establish adsorption isotherms require a knowledge of the time needed to achieve equilibrium. A series of kinetic experiments was therefore conducted for each material before the equilibrium experiments were begun. The concentration ranges selected for the equilibrium experiments ranged from the lower limit of analytical capabilities to concentrations greater than any likely to occur in environmental waters. For the surfactants, equilibrium concentrations approached or exceeded critical micelle concentrations in order to more fully understand the nature of their adsorption isotherms. The adsorption studies used a high concentration of zeolite in water and a pH of 8 in order to minimize dissolution and yet be within the pH range of many natural waters. Apparatus. All equilibrium experiments were performed in a constant-temperature room at 20 "C. The test solutions containing surface-active compounds were equilibrated with type A zeolite in vials attached to a rotating wheel and tumbled end over end at 12 rpm. The tumbling ensured complete movement of the solid phase to facilitate mass transfer and avoided the foaming that might have decreased the reproducibility of the equilibrium concentrations. Experiments with non-surface-active compounds were carried out either on the device just described or on a wrist-action shaker (Burrell Model 75). Kinetic experiments were performed at room temperature in 250-mL Erlenmeyer flasks stirred by a magnetic bar. Materials. Commercial type A zeolite, in the sodium form (Na-A) as developed for use in detergents, is a hydrated, crystalline aluminosilicate with a cubic crystal lattice, obtained from the Huber Corp. as Arogen 2000. Na-A was examined by X-ray diffraction and found to contain type A zeolite with only a trace of hydroxysodalite but no type X, Y, or P zeolites. Elemental analyses (anhydrous basis) found: Na, 15.8%; Al, 19.3%; Si, 19.0%. Calcd for Na12A112Si12048: Na, 16.2%;Al, 19.0%;Si, 19.8%.Water content, based on weight loss at 800 "C: 21.1%. Calcd for 27H20: 22.2%. The calcium exchange capacity was 0.97 equiv/mol of Al. The external surface area, determined from the complete adsorptiondesorption Langmuir isotherm with nitrogen, was 1.9 m2/g. The particle-size distribution of the sample, obtained by Stokes' law sedimentation using the Andreasen pipet method (151, is shown in Figure 1;the mass-median diameter, found by interpolation, was 2.85 pm. The Na-A was converted to the calcium form (Ca-A) by repeated ion-exchange treatments as follows. Na-A (100 g) was stirred with 2000 mL of 1.2 M aqueous calcium chloride for 1h. The zeolite was separated by continuous centrifugation at 15 000 rpm on a high-speed centrifuge (Sorvall Model SS-3). This exchange process was repeated two additional times with fresh CaCl2 solution. After the third exchange and centrifugation, the solids were washed free of chloride ion (verified by a silver nitrate test) by seven consecutive washes with distilled water and subsequent centrifugations. The product was rehydrated by standing for 1week in a desiccator Volume 15, Number 10, October 1981 1191
over a saturated solution of ammonium chloride. X-ray diffraction analysis showed that the product retained zeolite A crystallinity. Elemental analyses (anhydrous material) found: Ca, 13.4%;Na, 0.69%; Al, 19.4%.Calcd for Ca6AllzSil@& Ca, 14.4%;Na, 0%;Al, 19.4%.Water content, based on weight loss at 800 "C: 22.3%.Calcd for 28H20: 23.2%. The particle size and the external surface area of Ca-A are expected to be the same as those of Na-A because the calcium exchanges with sodium only from the interior of the crystal lattice. All measurements and calculations on type A zeolite have been reported on the normal hydrated material containing waters of hydration except the elemental analyses reported above. The adsorbate structures are shown in Figure 2. Methylene blue (MB), mol wt 320,91% pure (biological stain quality), was purchased from Eastman Organic Chemidals (Rochester, NY). p-Chlorophenol (p-CP),mol wt 129, was also purchased from Eastman. Dieldrin, mol wt 381, labeled with 14C a t the chlorinated carbons, was obtained from Amersham Searle Corp. (Arlington Heights, IL) and found to be >99% radiochemically
0
1 2 3 4 5 6 EQUIVALENT SPHERICAL DIAMETER (pm)
7
6
Figure 1. Particle-size distribution of sodium type A zeolite (Na-A).
Methylene Blue (ME)
CI-
p-Chlorophenol(pCP)
OH
0 ci
Dieldrin
H
Cl Dodecylhexaoxyethylene alcohol (Ci2Ed
Hz CizHz5(OCzH4)60H
Tetradecyltrimethyl ammonium chloride (TAC)
CH 3
I
Ci4HzsN+CH3 CI-
I
CH3 Linear alkyl benzene sulfonate (LAW
CH~-(CHZ)~-CH(CHZ)~CH~
b
S05Na'
Figure 2. Structures of adsorbates. 1192
Environmental Science & Technology
pure by thin-layer chromatography (TLC). n-Dodecylhexakis(oxyethy1ene) alcohol, a pure, single homologue, mol w t 450, commonly called C12E6, was prepared with a random 14C label on the ethoxyl carbons by the Procter and Gamble Co. and found to be >99% radiochemically pure by TLC. Nonradioactive n-Cl2E6 was also prepared, found to be 9101o isom' erically pure by gas chromatography, and mixed with the 14C12E6in order to achieve appropriate activity for accurate measurements. Tetradecyltrimethylammonium chloride (TAC), mol w t 392, was also prepared by the Procter and Gamble Co. with a 14Clabel on one of the methyl carbons; it was >99% radiochemically pure by TLC. Linear alkylbenzenesulfonate (LAS), av mol wt 344, was commercial-grade material prepared by the Procter and Gamble Co. with an average alkyl chain length of 11.7 and the following homologue distribution: C14,0.1990;unknown, 1.28%.2-phenyl, 25.7%; 3-phenyl, 21.0%; 4-phenyl, 18.6%;5-phenyl, 12.2%; 6and 7-phenyl, 22.6%. All other chemicals used in this study were ACS reagent grade. Analytical Methods. The MB solutions were analyzed by measuring their absorbance a t 668 nm on a Varian Techtron Model 635 spectrophotometer (Palo Alto, CA). p-Chlorophenol was measured similarly a t 279 nm after the solutions were acidified to pH 2.0 with dilute H2S04. The 14C-labeled adsorbates (dieldrin, C&6, and TAC) were measured by liquid-scintillation counting after dilution with toluene containing 2,5-diphenyloxazole (PPO), 1,4-bis[2- (5-phenyloxazoly1)lbenzene(POPOP), and Triton X-100 (Rohm and Haas, Philadelphia, PA). LAS, a t low concentrations, was determined colorimetrically a t 652 nm (Beckman Model B, Fullerton, CA) after reaction with methylene blue and extraction into chloroform. More concentrated solutions of LAS were titrated with Hyamine 1622 (Rohm and Haas). The critical micelle concentrations (cmc's) of C12E6, TAC, and LAS were measured by a technique involving solubilization and fluorescence of perylene as described by Mast and Haynes (16). Kinetic Studies. A slurry was prepared from 10 g of zeolite and a distilled-water solution of each adsorbate. The slurry was adjusted to pH 8.0 with 1 N HzS04 (except solutions of TAC were neutralized with HC1) and then diluted to 100 mL. This neutralization produced 0.042 M sulfate (or 0.084 M chloride) in the slurries. At various intervals, 5.00-mL portions of the rapidly stirred slurry were removed and centrifuged, and the solution phases were analyzed by the methods described above. The experiments were continued until the concentrations in the solution phases remained constant. Equilibrium Studies. Methylene blue was investigated first, not because of its environmental importance, but because of its known strong adsorption to solids and its recognized usefulness in characterizing adsorptive materials. The adsorption isotherm of MB was obtained by mixing 0.050 g of Na-A with 20.0 mL of MB solutions at several concentrations in the range 0.50 X 10-6-17.5 X loe6 M. Except for selected points, the data represent single-point determinations. Equilibrium was established by 1-min sonication followed by mechanical shaking for another 5 min. The solids were then settled by centrifugation, and the concentration of MB in the supernatant was measured. The other adsorption isotherms were generated by preparing a series of 15-mL vials containing 1.00 g of Ca-A and 10.0 mL of solution. The solutions contained the adsorbates to be studied and achieved the equilibrium concentrations indicated in Figures 3 and 4. For all adsorbates, the aqueous slurries were adjusted to pH 8.0 with 1 N H2S04 or 1N HC1 and agitated continuously a t 20 "C. After 500 h of agitation the solids were settled by high-speed centrifugation or by overnight sedimentation (calculation and observation verified that overnight sedimentation was valid).
The concentration of adsorbate in solution was determined by analyzing the aqueous phase, and adsorption isotherms were determined classically by “solution depletion”. Adsorption isotherms were also generated in some cases by separation of equilibrated zeolite solids and direct determination of adsorbate following acid dissolution of the zeolite. This “solids analysis” was accomplished in five steps: First, the zeolite was centrifuged to separate the solids from the equilibrated adsorbate solution. The latter was analyzed to determine the equilibrium adsorbate concentration in solution. Second, the zeolite solids were slurried momentarily with a known small volume of distilled water (20OC) and immediately filtered through a 0.45-pm Millipore filter. The filtrate was analyzed to determine rinse adsorbate concentration, C,, expressed as weight of adsorbate per gram of rinse solution. Third, one weighed portion of the filter cake was dried at 800 “C to determine the weight of anhydrous zeolite per gram of filter cake. ‘This, and the known moisture content of the fully hydrated zeolite, were used to calculate the total weight of filter cake produced per gram of hydrated zeolite, WFC. This allowed determination of the weight of rinse water, W,, clinging to 1 g of hydrated zeolite as
Results and Discussion No formal study of adsorption kinetics on type A zeolite was attempted; instead, the kinetic experiments were designed only to provide the data that were needed for subsequent equilibrium studies. The experiments showed that a 500-h equilibration time was sufficient for equilibrium to be attained with all of the adsorbates investigated and that 5 min was sufficient for MB.
1
w, = WFC - I
CMC
Fourth, a second weighed portion of the filter cake was transferred to a volumetric flask and the weight of hydrated zeolite present in this sample was calculated as described in step 3, above. This filter cake was treated with 2 mL of concentrated HzS04 or HC1. The resulting diluted solution was analyzed for adsorbate, and the results were expressed as apparent weight of adsorbate per gram of hydrated zeolite, W’ads. Finally, the true weight of adsorbate per gram of hydrated zeolite, Wads was obtained by subtracting the amount contained in the clinging liquid phase as
Id -6
-4
v
-6
These experiments were performed in duplicate, and the adsorption-isotherm data presented are the averages of two measurements.
T -5
-
by solution depletion
-
.
-73
by solids analysis
CMC
T
7
.5
-4
-3
-5
-
P z P
H9
-6
LASINaA -7 depletion -8
by solution ,depletion
by solution depletion
by solids analysis
-9 Dieldrin by solids enalysls
-10
CMC of 2-DBS I 1 -6
-1 1
1
-5
?I
I , -4
,
-3
-2
LOG SOLUTION CONCENTRATION -7 -6 -5 -4 -3 -2 LOG SOLUTION CONCENTRATION (rnoles/liter) -8
Figure 3. Adsorption isotherms of methylene blue (MB) onto Na-A, p-chlorophenol (p-CP) onto Ca-A, and dieldrin onto Ca-A zeolite.
(molesiliter) Figure 4. Adsorption isotherms of dodecylhexakis(oxyethy1ene)alcohol (CI2E6). tetradecyltrimethylammonium chloride (TAG), and linear aikyibenzenesulfonate (LAS) onto type A zeolite. Volume 15, Number 10, October 1981
1193
The adsorption isotherms of these six organic compounds were determined over a range of concentrations extending well above the environmental levels in order to characterize fully the nature of the adsorptions. As the figures and the following discussion indicate, tppeA zeolite behaves as an aluminosilicate with predictable isotherms and a low adsorption capacity for organics. The isotherm of MB on Na-A (Figure 3) was Langmuirian over the concentration range studied, reaching constancy for solution concentrations above 2.0 X l o p 6 M, and indicating mol of MB/g of coverage in the plateau region of 2.8 X zeolite. Combining this coverage with the measured surface area gives a calculated area of zeolite per molecule of MB of 113 Az. This is similar to the calculated surface area for a flat orientation of MB, viz., 120 AZ (17) and therefore serves to validate the techniques of measurement and the measured surface area of Na-A. The isotherms for both p -chlorophenol and dieldrin onto Ca-A appear to be of the Freundlich type in agreement with the adsorption of dieldrin and other halogenated aromatics onto various natural sediments, clays, and sands reported by Hargrave and Kranck (18). For surfactants, the critical micelle concentration (crnc) is important in the understanding of the adsorption isotherm. Surfactant ions can form organized aggregates, or micelles, in which the lipophilic hydrocarbon chains are oriented toward the interior of the micelle. The cmc is that concentration above which micelle formation becomes appreciable. If the surfactant concentration in a solution is above the cmc, the addition of more surfactant will increase the number of micelles but will not increase the concentration of surfactant monomer available for adsorption. The adsorption isotherms would therefore be expected to reach a plateau in the region of the cmc. Adsorption isotherms for C1& onto Ca-A, for TAC onto Ca-A, and for LAS onto Na-A are shown in Figure 4. No isotherm of LAS onto Ca-A is reported because of the possible precipitation of Ca(LAS)Z and subsequent cosedimentation with the zeolite. For the adsorption of ClzE6 onto Ca-A, TAC onto Ca-A, and LAS onto Na-A, the desorption of the compounds was sufficiently slow that data could be obtained both by solution depletion and solids analysis. The differences seen in Figure 4 between the two types of measurements probably reflect the slight desorption which occurs during washing of the solids. M (i.e., 39 The cmc of C12E6 in 0.6% Na2S04 was 8.7 X mg/L), as indicated in the figure by an arrow a t log (solution concentration) = -4.06. By measurement of depletion from mol/g solution, C&6 had a limiting adsorption of -1.7 X of zeolite a t a concentration close to its cmc. The coverage of zeolite at the plateau region is calculated to be 2110 A2/molecule of C12E6. This represents -3.5% of one monolayer, based on close-packed (vertical) attachment with 60 A2/molecule cross-sectional area (19).By analysis of solids, the same limit appears to be approached asymptotically, although the break point is less well-defined. The C12E6 adsorption found in this work is much less extensive than for silver iodide powder (20),where adsorption produces multilayer coverage. I t is also less extensive than adsorption onto montmorillonite (21)and graphon (22),where C12E6 adsorption reaches full monolayer coverage. Full monolayer coverage is also observed for C12E6 adsorption onto polystyrene latexes (23).The adsorption of C12E6 onto Ca-A is thus uniquely low among all substrates reported. The reason for the very low coverage on zeolite is unknown, but may reflect a highly site-specific, horizontal molecular configuration. For TAC, the agreement between the two measurements of the isotherm is very good, indicating that this material is 1194
Environmental Science & Technology
more strongly adsorbed than C1&. Stronger adsorption is the expected result because, under conditions similar to those used in the present experiments, electrophoretic mohility measurements show the zeolite particles to have a net negative surface charge. Therefore, in the lowest concentration region, adsorption is probably due to ion pairing at negatively charged zeolite surface sites, rather than ion exchange, as might occur with Na-A. Ion exchange with Ca-A is not likely, based on the preference of zeolite A for divalent ions over monovalent. The low level of adsorption suggests that the surfactant orientation a t the interface is horizontal. In the next region, beginning a t -1 X 10-4 M, surfactant orientation probably changes to vertical, and a new process becomes dominant, that of hemimicelle formation. Hemimicelle formation is a surface aggregation phenomenon in which surfactant ions associate in reverse orientation (head groups out) with respect to previously adsorbed surfactant ions. This process is driven forward by a favorable hydrophobic bonding free energy change. In the third concentration region, adsorption rises less steeply and appears to be approaching a plateau. A true plateau was not attained in these experiments because of termination a t concentrations prior to the cmc. However, a plateau probably exists a t an adsorption level equal to or slightly above the levels of a close-packed vertical monolayer. Based on a trimethylammonium ion head-group cross-sectional area of 34 W2/molecule (24),a close-packed vertically oriented monolayer corresponds to adsorption of 9.0 X mol/g (log = -5.04). The adsorption of TAC onto Ca-A thus described is entirely analogous to previously determined adsorption processes for octyl-, decyl-, dodecyl-, and hexadecyltrimethylammonium surfactants onto negatively charged polystyrene latex particles (24),and for hexadecyltrimethylammonium adsorption onto kaolinite ( 2 5 ) . LAS adsorption onto Na-A is also seen to be a multistage process. The initial adsorption probably arises from electrostatic attraction to positively charged surface sites. This process goes to completion a t a concentration slightly below 1X M. The low level of adsorption in this region is consistent with the hypothesis that the surfactant molecules have a horizontal orientation. This interpretation is similar to that given for dodecylsulfate adsorption onto graphon, which exhibits a similar low-concentration plateau (26). At a concenM, a change to vertical orientation tration of -4 X probably occurs, and adsorption proceeds thereafter by hemimicelle formation. This continues to a point near the cmc of the principal homologue in the LAS mixture. The principal homologue is 2-dodecylbenzenesulfonate (2-DBS) with cmc = 2 X 10-3 M (16). As with TAC adsorption onto Ca-A, LAS adsorption onto Na-A appears ultimately to approach a plateau. This is not fully confirmed because of the termination of experiments a t a concentration only slightly above the 2-DBS cmc. However, a plateau probably occurs a t an adsorption level which corresponds to a close-packed vertical bilayer. This is based on an alkylbenzenesulfonate head-group cross-sectional area of 52 A/molecule (27).The close-packed vertically oriented bilayer corresponds to adsorption of 1.21 X mol/g (log = -4.92). This interpretation of the LAS adsorption process is quite similar to that of Scamehorn for nonyl-, decyl-, and dodecylbenzenesulfonate adsorption onto alumina and kaolinite (27). Bilayer coverage was exhibited on these substrates, both for the individual surfactants and for binary mixtures thereof. Likewise, the present findings agree generally with those of Dick et al., who studied 2-, 3-, 4-, 5 - , and 6-dodecylbenzenesulfonate adsorption onto alumina (28). The reason for finding a low-level plateau in the present work is not entirely understood but may be related to the
presence of significant amounts of and longer-chain alkylbenzenesulfonates in the LAS mixture. The cmc of the longest-chain homologue in a mixture controls the result found by the perylene cmc determination technique (16). Applied to the LAS mixture, this technique gave a cmc of 8.2 X 10-5 M, almost exactly where the first plateau begins. The adsorption isotherm for LAS onto Ca-A suggested more extensive adsorption than onto Na-A. However, this adsorption was due to Ca(LAS)2 precipitation onto the calcium-exchanged zeolite, as confirmed by finding significant amounts of sodium reintroduced into the zeolite. In a realistic situation with sewage or surface water, an excess of calcium would be available so that Ca(LAS)2 could form without disturbing the Ca-A. Overall, the adsorption isotherms of TAC and LAS onto type A zeolite are predictable in shape and show no indication of selective adsorption. In all cases, the isotherms and calculations of surface coverage confirm that these organics adsorb only onto the external surface of the zeolite and are too large to enter the pore structure.
Environmental Considerations A comparison of the adsorption of organics onto type A zeolite with similar adsorption by naturally occurring minerals can only be of a general nature. The available literature indicates that only the adsorption of an alkylbenzenesulfonate has been studied on a significant variety of substrates. In most cases, formal adsorption isotherms have not been reported, pH was allowed to vary or was uncontrolled, and adequate characterization of substrates was not provided. In addition, inconsistencies in the literature are evident, but their cause is not. Table I lists the intensity of adsorption of LAS or ABS (branched alkylbenzenesulfonate) onto type A zeolite and onto naturally occurring minerals as taken from Figure 4 and from the literature (29-32). The equilibrium concentrations of LAS or ABS are 5 mg/L except as noted. In terms of adsorption intensity (in pg/m2), type A zeolite is similar to or lower than all other minerals in the table. The relatively low external surface area of zeolite that is available for adsorption
Table 1. Adsorption of LAS or ABS at 5 mg/L onto Various Minerals a mineral
Na-A zeolite
surface area, m2ig
1.9
29
660 at 4.3 mg/L 650 at 4.3 mg/L
29 29 30
1.6
770
30
0.7
430 371 235
30 30
Mississippian sandstone Pennsylvanian sandstone
30
1.4
260 235 385
21.0
1600
76
30
1.5
375
250 455
30
1.o
1.3
275
590
a0
30 31
240
720 1a0
aooo
560
30
37.5
3000 50
3.0
bentonite (montmorillonite)
29
iaoo at 4.5 mg/L
2340
0.015
glauconite
this work
35
Ottawa sand
illite
ref
14
25000
diatomite silica gel
oolitic limestone
adsorption intensity, Irg/m2
27
natural silt calcium carbonate
high-purity limestone
adsorption capacity, pglg
30 31 30
14.3
40
31
kaolinite
65 300
31
goethite
150
31
gibbsite
150
31
32
Studies reported in ref 29-32 were performed with ABS.
Table It. Adsorption of Surfactants at 5 mg/L onto Type A Zeolite and Sludge surfactant
LAS c12E6 TAC nonylphenol ethoxylate (C9E9.5) Tergitol 15-S-9 (C10-15 secondary alcohol E9 ethoxylate)
type A zeolite
27 24 24
adsorption (pg of adsorbatelg of substrate) primary sewage sludge
2000-5000
activated sludge
2000-10000
600 xa70
ref for sludge data
33
33 33, 34
Volume 15, Number 10, October 1981
1195
means that zeolite also has the lowest capacity (in terms of weight of LAS-ABS adsorbed per gram of substrate) of all of the minerals listed in this table. Table I1 presents data comparing the adsorption of LAS, C12E6, and TAC onto type A zeolite to the adsorption of related compounds onto primary sewage sludge and activated sludge ( 3 3 , 3 4 ) .Nearly all of the zeolit:: which is introduced into sewage will be removed with the sludge. However, the zeolite will be a minor component of sludge (5-7), and adsorption of organics onto zeolite will be lower by 2-3 orders of magnitude than their adsorption onto sludge solids. In addition, any small amount of zeolite not removed during sewage treatment will represent only a minor fraction of the suspended solids normally present in sewage effluent (5-7). Thus, in both sewage and surface waters, no change in the normal fate of organics is expected to result from the use of type A zeolite in laundry detergents. This is consistent with treatability studies in wastewater where the addition of zeolite caused no measurable effects on the removal of organics or on the biodegradation of LAS (6, 7 ) .
Conclusions Type A zeolite, when introduced into natural waters, can be expected to behave in a manner similar to that of a wide variety of commonly occurring minerals with respect to adsorption of organics from solution. No unusual or extraordinary adsorptive capabilities have been observed. In fact, the adsorption of organics onto type A zeolite is lower than onto many common minerals and much lower than onto sewage sludge. In the concentration ranges investigated, none of the organic compounds exhibited more than bilayer coverage. Even a t high concentrations, no evidence was found for penetration by the organics into the zeolite structure, as is understandable from the known pore size of the crystal lattice of zeolite. This overall weak adsorption, combined with the hydrolytic instability of the zeolite in aqueous environments, indicates that type A zeolite should not alter the normal fate of organic molecules in the environment. Acknowledgments We thank Professor R. H. Ottewill, University of Bristol, for helpful comments on the interpretation of the data, Messrs. R. J. Grosse and C. W. Steuver for performing the adsorption experiments, and Mr. Lloyd Williams, J. M. Huber Corp., for the surface-area measurements. Literature Cited ( 1 ) Katzer, J. R., Ed. A C S S y m p . Ser. 1977, No. 40. (2) Breck, D. W. “Zeolite Molecular Sieves”; Wiley: New York,
1974.
1196
Environmental Science & Technology
(3) Rabo, J. A,, Ed. A C S Monogr. 1976, No. 171. (4) Barrer, R. M. “Zeolites and Clay Minerals as Sorbents and Molecular Seives”; Academic Press: New York, 1978. (5) Hopping, W. D. J . W a t e r Pollut. Control Fed. 1978,50,433. (6) King, 3. E.; Hopping, W. D.; Holman, W. F. J . W a t e r Pollut. Control Fed. 1980.52.2875. (7) Holman, W. F.; Hopping, W. D. J . W a t e r Pollut. Control Fed. 1980,52,2887. ( 8 ) Fischer, W. K.; Gerike, P.; Kurzyca, G. Tenside Deterg. 1978,15, 60. (9) Roland, W. A.; Schmid, R. D. Vom Wasser 1980,50,177. (10) Maki, A. W.; Macek, K. J. Enuiron. Sci. Technol. 1978, 12, 573. (11) Payne, A. G.; Hall, R. H. M i t t . I n t . V e r . Limnol. 1978, 21, 507. (12) Fischer, W. K.; Gode, P. Vom Wasser 1977,49,11. (13) Cook, T. E.; Cilley, W. A,; Savitsky, A. C.; Wiers, B. H. “Zeolite A Hydrolysis and Degradation”, submitted to Enuiron. Sci. Technol. (14) Lerman, A.: MacKenzie, F. T.: Bricker, 0. P. E a r t h Plant. Sci. L e t t . 1975,25,82. (15) Allen, T. E. “Particle Size Analysis”, 2nd ed.; Pergamon Press: London,’1975; p 193. (16) Mast, R. C.; Haynes, L. V. J . Colloid Interface Sci. 1975, 53, 35. (17) Allingham, M. M.; Cullen, J. M.; Giles, C. H.; Jain, S. K.; Woods, J. S. J . A p p l . C h e m . 1958,8,108. (18) Hargrave, B. T.; Kranck, K. Paper presented at the NBS Symposium on Transport of Organic Species in Environmental Waters, Gaithersberg, MD, May 1976. (19) Lange, H. Kolloid-Z. 1965,201, 131. (20) Mathai, K. G.; Ottewill, R. H. Trans. Faraday Soc. 1966, 62, 750. (21) Barclay, L. M.; Ottewill, R. H. Spec. Discuss. Faraday Soc. 1970, 1,138. (22) Corkill, J. M.; Goodman, J. F.; Tate, J. R. Trans. Faraday Soc. 1966,62,979. (23) Ottewill, R. H.; Walker, T. Kolloid Z . Z. Polym. 1968, 127, 108. (24) Conner, P.; Ottewill, R. H. J . Colloid Interface Sci. 1971, 37, 642. (25) Callaghan, I. C. Ph.D. Thesis, IJniversity of Bristol, England, 1975. (26) Greenwood, F. G.; Parfitt, 3. D.; Picton, H. N.; Wharton, S. G. A d u . C h e m . Ser. No. 1968, 79,135. (27) Scamehorn, J. F. Ph.D. Thesis, The University of Texas, Austin, TX, 1980. (28) Dick, S. G.; Fuerstenau, D. W.; Healy, T. W. J . Colloid Interface Sci. 1971,37, 595. (29) Renn. C. E.: Barada. M. F. Sewage I n d . W a s t e s 1959,31,850. (30) Suess, M. J. J. W a t e r Pollut. Control Fed. 1964,36, 1393. (31) Wayman, C. H. In “Principles and Applications of Water Chemistry”; Faust, S. D., Hunter, J. V., Eds.; Wiley: New York, 1967; p 152. (32) Barbaro, R. D.; Hunter, J. V. W a t e r Res. 1967,1,157. (33) Swisher, R. D. “Surfactant Biodegradation”, Surfactant Sci. Ser. 1970,3, 121. (34) Conway, R. A,; Vath, C. A,; Renn, C. E. W a t e r W o r k s W a s t e s Eng. 1965,2, 28.
Received for reuiew July 7,1980. Accepted J u n e 8,1981
