J. Phys. Chem. 1082, 86, 3861-3866
3861
AH^',^^, kJ/gatm
KVT= P d P b 3 AGvo(YbPb3)= 881200 - 421.32' (J/mol) W"zss(YbPb3) = -164.0 f 4.6 (kJ/mol)
(second law)
Conclusion As the result of this experimental work, a temperature-composition diagram of the ytterbium-lead system has been revised and includes phase widths for the nonstoichiometric intermetallic compounds. The ability of the technique presented in this work to accurately define composition at elevated temperature makes it superior for determining vertical phase transitions. Results of third-law calculations, determined at each experimental point, for the enthalpy of formation of the intermetallic compounds are as follows:
Yb,Pb Yb,Pb, YbPbI.0, YbPb,
*
-58.2 k2.9 -57.7 k0.8 -57.3 k0.8 -38.1 k 0.4
These data represent an addition to the existing thermodynamic literature. Only the YbPb3 enthalpy has been previously investigated by Palenzona and C i r a f i ~ iwho ,~ reported a value of -37.2 kJ/g-atm. The reasonable standard deviations, agreement between second- and third-law values, and agreement with the previous value for YbPb3 lend support to the validity of the experimental results. The results emphasize the applicability of the simultaneous weight-loss-mass-spectrometric method in investigating these complex systems. The lower sensitivity to temperature variations, as well as the smaller standard deviation, makes the third-law values the preferred results.
Adsorption-Desorption Kinetics of Acetic Acid on Silica-Alumina Particles in Aqueous Suspensions, Using the Pressure-Jump Relaxation Method T. Ikeda, M. Sarakl, K. Hachlya, R. D. Astumlan,' T. Yasunaga,' Department of Chemistry, Faculty of Science, Hiroshima University, Hlroshima 730, Japan
and 2. A. Schelly' Department of Chemlstry, The University of Texas at Arlington, Arlington, Texas, 76019 (Received:February 10, 1982; In Final Form: June 3, 1982)
The pressure-jump technique with electric conductivity detection was used to detect two relaxations on the orders of milliseconds and seconds in suspensions of silica-alumina in aqueous acetic acid. The fast relaxation is attributed to the protonation-deprotonation reaction on the silica-alumina surface, and the slow one to the adsorption-desorption process of the acetate ion, accompanied by the elimination of a water molecule from the surface. The intrinsic values of the protonation and deprotonation rate constants determined are 2.9 X lo4 mol-' dm3 s-l and 4.6 X 10 s-l, and those of the adsorption and desorption rate constants 6.5 mol-' dm3 s-l and 3.8 X s-l, respectively, at I = 5.5 X M and 25 "C.
Introduction The application of relaxation methods to the investigation of rapid adsorption-desorption processes in aqueous suspensions has proven quite successful. For example, the kinetics and mechanism of the adsorption of H+on several different iron oxides: proton3and iodate ion4on TiOz, and Pb2+on y-alumina5 have been elucidated this way. The silica-alumina double oxide that was investigated in the present study is an interesting system because it exhibits stronger acidity than either silica or alumina. It has been revealed by IR studies of the gas adsorption of (1) Japanese Ministry of Education Research Scholar. Permanent a d k Department of Chemistry, The University of Texas at Arlington, Arlington, TX 76019. (2) Astumian, R. D.;Sasaki, M.; Yasunaga, T.; Schelly, Z. A. J.Phys. Chem. 1981,85, 3832-5. (3) Ashida, M.; Sasaki, M.; Kan, H.; Yasunaga, T.;Hachiya, K.; Inoue, T. J. Colloid Interface Sei. 1978.67. 219-25. (4) Hachiya, K..Ashida, M.; Sk&i, M.; Karasuda, M.;Yasunaga, T. J.Phys. Chem. 1980,84, 2292-6. (5) Hachiya, K.; Ashida, M.; Saaaki, M.; Kan, H.; Inoue, T.; Yasunaga, T. J. Phys. Chem. 1979,83, 1866-71. 0022-3654/82/2086-3861$01.25/0
ammonia! pyridine?s etcW that both Bronsted and Lewis acid sites are present, depending on the coordination number of the aluminum atom. The four-coordinated aluminum atom in A acts as a Bronsted acid site, and the -AI-O--SI= H
A
AI:O /H H '+
R
three-coordinated one, =Al, as a Lewis acid site."12 In aqueous suspension of silica-alumina, furthermore, it is (6) Tanabe, K. "Solid Acids and Bases"; Kondasha: Tokyo, 1970. (7) Basila, M. R.; Kantner, T. R.; Rhee, K. H. J. Phys. Chem. 1964, 68, 3197-207. (8) Scokart, P. 0.; Declerck, F. D.; Semples, R. E.; Rouxhet, P. G. Faraday Trans. 1977, 73, 359-71. (9) Semples, R. E.;Rouxhet, P. G. J. Colloid Interface Sci. 1976,55, 263-73. (10) Uytterhoeven, J. B.; Christner, L. G.; Hall, W. K. J.Phys. Chem. 1965,69,2117-26. (11) Fripiat, J. J.; Leonard, A.; Uytterhoeven, J. B. J. Phys. Chem. 1965,69,3274-9. (12) Tamele, M. W.Discuss. Faraday SOC.1950,8, 270-9.
0 1982 Amerlcan Chemical Society
Ikeda et al.
3882 The Journal of Physical Chemistry, Vol. 86, No. 19, 1982
well-known that the Lewis acid site acts as an acceptor of the lone-pair electrons of the water molecule and is thus converted into a Bronsted acid site as in B. These sites, along with the silanol group, d i O H , determine the acidity of silica-alumina in aqueous suspensions. The mechanisms of the emergence of the acidities, however, have not been established because the reactions on the active sites are usually very fast for oridinary kinetic methods. To obtain information about the elementary processes involved in the adsorption-desorption of ions and the emergence of the strong acid sites, we preformed a pressure-jump relaxation method study on silica-alumina suspensions in aqueous acetic acid, and the results are reported in the present paper.
Experimental Section Chemicals and Sample Preparation. Silica-Alumina was prepared from zeolite 4A, TSZ (with the formula Na20.Al2O3.2SiO2.nHz0, supplied by the Toyo Soda Co.), through ion exchange of Nat with H+. It was purified by electrodialysis until the electric conductivity was equal to that of distilled, ion-exchangedwater. The silica-alumina was determined to be amorphous, by X-ray diffraction, and free of Nat by atomic absorption analysis. The diameter of the particles was smaller than 1pm. Acetic acid and sodium nitrate (Wako Chemical Co., reagent grade) were used without further purification. The ionic strength of the silica-alumina-acetic acid system was adjusted to I = 5.5 X M with sodium nitrate, and all measurements were performed on suspensions equilibrated for 1day at a particle concentration of [PI = 30 g/dm3. The amount of adsorbed acetic acid was determined indirectly from the concentration change in the supernatant solution by gas-chromatographic analysis with Shimalite TPA support (Shimazu Co.), and by alkalimetric titration. Prior to both measurements, samples of the silica-alumina suspension containing the acetic acid were centrifuged for 30 min at loooOg in order to effect complete settling. The concentration of H+ was determined with a pH meter, and the temperature was controlled at 25 OC. Since it is known that silica-alumina dissolves in solutions of low pH, the aluminum ion concentration was determined colorimetrically in the liquid phase of the suspensions by the stilbazo method,13 but was found to be negligibly small in the pH range of our investigations. Experiments. The pressure-jump apparatus with conductometric detection has been described previou~ly.~ It has a time constant of 100 ps at a bursting pressure of 200 atm. Results and Discussion Kinetic measurements were carried out in aqueous silica-alumina suspensions containing acetic acid by using the pressure-jump technique with electric conductivity detection. It is assumed that changes in the conductivity observed were due to the change in bulk phase ionic conductivity with no significant contribution to the conductivity by the particles. Two relaxations were observed where the directions of both relaxation signals indicate a decrease in the conductivity of the suspension during relaxation. Relaxations were not observed in aqueous solution of acetic acid, the supernatant solutions of silicaalumina-acetic acid, or silica-alumina-NaN03 systems. Furthermore, neither silica-acetic acid nor alumina-acetic acid systems exhibit relaxation phenomena. This leads to the conclusion that the relaxations must be caused by the presence of a species that is absent in both pure alumina (13)Eigen, M. Angew. Chem., Int. Ed. Engl. 1964, 3, 1-19.
1.5
-1,
‘
0
2
4
6
8
added acetic acid , 10.2moldm’
Flgure 1. Dependence of the reciprocal fast and slow relaxation times, TI-’ (0) and T ~ - I ( o ) , respectively, on the added acetic acid concentration in the silica-alumina-acetic acid system, at 25 ‘C. [PI = 30 g dm-3, I = 5.5 X M. 61
I
I
I
0
2
4
6
CHAcl
+
0
L A C - I , lO-’mol d m 3
Figure 2. Adsorption isotherms of acetic acid, and H+ in the silicaalumina-acetic acid system at [PI = 30 g dm-3, I = 5.5 X M, and 25 ‘C. is the amount of H+ (0)or acetic acid (0)adsorbed.
raa
and pure silica dispersions, i.e., the four-coordinated aluminum found in A. The dependences of both fast and slow reciprocal relaxation times, T ~ and - ~ T ~ -on ~ ,the added concentrations of acetic acid are shown in Figure 1. The values of both decrease with increasing concentration of acetic acid. To explain the above results, of course, one has to distinguish the actual free acetic acid concentration from that added. Therefore, the amounts of acetic acid adsorbed were determined, and the adsorption isotherm obtained is shown in Figure 2. The pH of the silica-alumina-acetic acid suspensions is slightly higher than that of pure acetic acid solutions of the same concentration, indicating that also H+is adsorbed on the surface of the particles. Taking into account the amount of acetic acid adsorbed, as well as its dissociation in the bulk solution, we calculated the concentrations of adsorbed Ht and these are also shown in Figure 2. To obtain information about the sites available for proton adsorption, we titrated aqueous silica-alumina suspensions with HN03. The adsorption isotherm obtained is depicted in Figure 3, and it does not seem to be of the simple Langmuir type. A t saturation ([H+] m), the amount of H+ adsorbed is 1.45 X mol g-l. The nonlinearity of the Langmuir plot (Figure 4) indicates that the equilibrium constant Kl (or reaction 1)varies with the amount of H+ adsorbed.
-
Adsorption Kinetics of CH,COOH on Silica-Alumina
61
I
1
I
I
The Journal of Physical Chemistry, Vol. 86, NO. 19, 1982 3863
L
-ol E
'0, 2
CH'I
, 1 0 - 3 m o l dm'3
Figure 3. Adsorption Isotherm of H+ in the sillca-aiumlna-HN03 system a1 I = 5.5 X lo-, M and 25 O C . 1
--
07
:
*I
52
-.
(1) (1)-(2)-(6) (11) (1)-(3)-(6) (111) (1)-(4)-(6) (IV) (3)-(2)-(6) (V) (3)-(4)-(6) (VI) (3)-(5)-(6) Mechanisms I-VI, of course, include the very f a d 3 dissociation of the acetic acid (reaction 6) in the bulk liquid. If the expressions for the reciprocal fast relaxation time ~ 1 -are l derived with the assumption that the adsorptiondesorption reaction of H+ is much faster than that of HAC or Ac-, mechanisms I-VI are in qualitative agreement with the dependence of the experimental T i 1 on the acetic acid concentration. On the other hand, if the opposite is assumed, the theoretical T ~ for - ~ all six mechanisms are in disagreement with the data. Thus, no easy distinguishing is possible between the mechanisms based on the concentration dependence of T i 1 . Therefore, the selection of the best mechanism must be based on the analysis of the concentration dependence of the slow relaxation time T2-l. Now, we shall show that mechanism I is consistent with both the kinetic and equilibrium data. If the surface potential on the silica-alumina surface is considered constant, the rate equations for this mechanism are -d[~Al-OH]/dt = -k,[H+] [=Al--O] + k-l[=Al-OH] (7) - d [ ~ A l : A ~ ] / d=t -k,[Ac-] [=Al:OH2+] + k-i[=Al:Ac]
L'
(8)
with K1 = kl/k-l = [=Al-OH]/([=Al--O][H+]) 0 CH'
K i = - -k2
IO
5
I" , 10' mol-'dm3
k-2'
Figure 4. Langmulr plot of the adsorption isotherm In the sliica-aiumina-HNO, system.
Now, let us consider the reactions that may be involved in the relaxations. As mentioned earlier, the participation of A is essential. With this in mind, one can consider reactions 1-6. For the sake of brevity in the derivations (a) s u r f a c e r e a c t i o n s EAILO-SiG
+
H+
& =AI-0-SI= H k-I
(1 1
,,+
-
=AI--O-SI~
-
H =AI-O--SiG
t HAC
+
Ac-
t
'H
(3)
IAc i
(4 1
S =Al--O-SiE
Ac SAI--O--SiE
[=Al:Ac] = - KZ [=Al:OH2+] [Ac-] [HzOI
(10)
KA = ([H+l[Ac-l) / [HAC] (11) k-4 = k-Z[H,O] (12) where the brackets indicate concentration, and K,, K2, and KA are the surface charge dependent equilibrium constants of each reaction. In the neighborhood of equilibrium, the linearized relaxation rate equations can be obtained in the usual fashion:14 d6[~Al-OH]/dt = ~116[=Al-OH] ~126[=Al:Ac] (13) -d6[=Al:Ac] /dt with
-\= S A 1 : O NH 7 -AI:OH
-
(9)
+ = ~216[=Al-OH] + U~~~[=AI:AC] (14)
(5)
(b) bulk reaction HAC H+ + Ac(61 that follow, the symbols =AI-OH, ~ A t - 0 =A1:OH2+, , and =MAC will be introduced for the Bronsted acid site consisting of the four-coordinated aluminum atom in A, the same species without the proton, the Bronsted acid site converted from Lewis acid site, and the adsorbed state of the acetate ion (as in eq 2), respectively. The following simplest combinations of reactions 1-6 were examined to explain the two relaxations observed:
a22 =
,
[Ac-]
+ [=A1:OH2+]
KA + [Ac-] KA + [H+] + [Ac-]
+ k-2/ (18)
The symbol 6 refers to a small deviation from the equilibrium concentration, and starting from eq 13, the brackets denote only the equilibrium concentrations.
3864
Ikeda et al.
The Journal of Physical Chemistry, Vol. 86, No. 19, 1982 C 1 , 10-3moldm-3 2 4
C Ac- I-' , 1 O3mol"dm3 0
1 . 1 . I
1
I
1.5}
T
's
0
?
0
I
:
1.01
I
I?.
w
1
0
2
I
C H'l-'
i / /
, 1 O4 mol'ldm'
Figure 5. Langmuir plots of the adsorption isotherms for H+ (0)and acetate ions (0). I
I
potentials wi vary with the amounts of ions adsorbed (see eq A2 and A3 in the Appendix), we adopted the procedure of Davis et al.15 for the calculation of the surface charge dependence of the equilibrium parameters K, and K i . The result is shown in Figure 6; the PK,'~ and the intrinsic values, p K P t and pKiint, are 4.76, -2.80, and -1.40, respectively. The activation potentials, as estimated from the differences between pK and p P t , are different for the two ions, indicating that the activation potentials of the adsorption sites of H+ and Ac- differ and that the adsorption of the two ions takes place in different planes. The kinetic data summarized in Figure 1can be interpreted with eq 20 and 21 if the dependence of the rate constants on the surface potentials is taken into account according to eq A6. In this case the expressions for the reciprocal relaxation times become
-1
2-
1
,
I
4 ;
I
0
6
4
2
C 2 , 10-Zmoldm-3 Figure 7. Reciprocal relaxation times 7,-'and T ~ - ' as a function of the experimental values of C (eq 24) and C 2 (eq 25), respectively. The solid lines are the theoretical curves.
, 10-3mol d m 3
Figure 8. Surface charge dependence of pK, (0)and pK,' (0). uo is the difference of the concentrations ['AI:OH,+] and [eAI--O].
Solving the simultaneous differential eq 13 and 14, we obtain the expressions for the reciprocal relaxation times, 71-l and T ~ - ' ,as follows: 7,?-I -9-
all
=
+ a22
* [(all
f; a22)'
2
1"'
+ a12a21- a11a22
72-l
(19)
If two distinct relaxation processes are observed, one of the conditions all >> a22or a22>> all must be true. Therefore, 71-l and 72-l can be expressed as follows: (9 all >> a22 rl-l = all (20) a12a21
T2-l
[~A1--0]
= a22- -
KA + [H+1 KA + [H+] + [Ac-]
%)(
= k 2int exp( - 2kBT
[=Al: OH2+] KA
[Ac-]
72-1
=
a11 -
-
+
[H+l[Ac-l (KA + [H+] [Ac-])~
[=Al:OH2+] [~ A 1 - - 0 ]
+
[H+] + [ ~ A 1 - - 0 ] \
K2' L, kptC2 (25)
=- a22 ~ a12a21
kptCl (24)
+ [Ac-] + [H+] + [Ac-]
(ii) a22>> all ~
1
KA
a11
T
L, K1
The plots of both (23)
a22
With mechanism I in mind, the Langmuir plots of the adsorption isotherm of H+ and Ac- are shown in Figure 5. At saturation, the amounts of both ions adsorbed were found to be .=1.45 X mol g-l. Since the activation
rl-l
vs. C1 and
72-1
vs. C2 are shown in
(14) Eigen, M.; DeMaeyer, L. In 'Technique of Organic Chemistry";
Part 2, Friess, S.L.;Lewis, E. S.;Wessberger, A., Ed.;Interscience: New
York, 1963; Vol. 8, pp 895-1054. (15) Davis, J. A.; James, R. 0.; Leckie, J. 0. J. Colloid Interface Sci. 1978, 63, 480-99. J. J.; Izatt, R. M.; Hansen, I. D. J. Am. Chem. Soc. (16) Christensen, 1967,89, 213-22.
The Journal of Physlcal Chemistv, Vol. 86, No. 19, 7982 3865
Adsorption Kinetics of CH,COOH on Silica-Alumina
Figure 7. The linearity of the curves and their intersecting the origin support mechanism I. Since the slopes of the lines yield the values of the intrinsic forward rate constants kiint,the reverse rate constants k-i'"t can be calculated from the known values of Klht and KZht. The numerical values of the rate constants obtained are klint = 2.9 X lo4mol-l dm3 s-l, k-lint = 4.6 X 10 s-l, kzint = 6.5 mol-' dm3 s-l, and k-2int = 3.8 X s-l. Next, let use investigate the possibilities of mechanisms 11-V. If r2-lis plotted vs. the experimental values of the concentration terms of the expressions derived for these mechanisms (see Appendix B), negative values are obtained for the rate constants k2, which, naturally, exclude the possibility of the corresponding mechanisms. Mechanism VI is formally identical with I, thus they cannot be distinguished based on the kinetics above. However, both khz are of the same order of magnitude as those for alumina-acetate complex formation in a homogeneous aqueous system;" thus perhaps preference should be given to mechanism I, in which Ac- is attached to a moiety structurally similar to the one involved in the homogeneous case. Acknowledgment. Acknowledgment is made to the donors of the Petroleum Research Fund, administered by the American Chemical Society, for partial support of this work, and to the Robert A. Welch Foundation for additional support. R.D.A. thanks the Japanese Ministry of Education for a research scholarship. The authors thank the Toyo Soda Co. for the supply of zeolite. Appendix A Under the condition of variable surface potential, the rate coefficients khi of the adsorption-desorption reactions are dependent on the electrostatic activation potentials $tii. The relationship can be expressed as3J8
where the superscript "int" means intrinsic and e is the elementary charge, kB the Boltzmann constant, and T the absolute temperature. In this case, the equilibrium parameters expressed by eq 9 and 10 also become functions of as follows: Klint =
[=Al-OH]
-
[~ A 1 - - 0 ][H+]
[rAl-OH] exp( %) kBT = K1 exp( %) kBT (A2) [=Al--O] [H+] =
&tint
[=Al:Ac] [=A1:OH2+][Ac-]
-
exp( - 3) = K i exp( - 3)
[=Al:Ac]
kBT
[=AkOH2+][Ac-]
kBT (A31
with w1 -w2
+ $*-1 = $t2 + $'-2
= $fl
(A4) (A51
where the subscript s refers to the surface, and w1 and w2 are the activation potentials for H+ and Ac- adsorption, respectively. Since the potential wi must be a function of distance only if the surface of the particle is considered as a plane, $ti = $t-1,3J8and eq A1 becomes
khi = khiintex.(
7s)
Appendix B For mechanisms 11-V, the reciprocal slow relaxtion times are given by expressions Bl-B4.
mechanism 11: 72-l
+ [=A1--0]
= k p t exp( -$)[[Ac-]
K1-'
+ [EAl-OH)
KA +
KA + [Ac-] KA + [H+] +[Ac-] [H+l [H+] + [Ac-]
)(
[Ac-]
K1-l + [H+] t' [=Al:OH]
+ [=Al-OH]
KA
+
[Ac-I [H+] [Ac-]
+
KA + [H+1 KA + [H+] + [Ac-]
mechanism 111:
-'KA
+ [H+] + [Ac-]
(17)Hirai~hi,M.; Harada, S.; Uchida, Y.; Kuo, H.L.; Yaaunaga, T. Znt. J. Chem. Kinet. 1980,12, 387-92.
(18)Ashida, M.;Sasaki, K.; Hachiya, K.; Yasunaga, T. J. Colloid Interface Sci. 1980, 74, 572-4.
J. Phys. Chem. 1982, 86,3866-3870
3866
mechanism IV 72-1 =
kzintexp( -
$)I
[Ac-] + [=A1:OH2+]
KA + [Ac-] KA + [H+] + [Ac-]
\
.
.,I:OH]
lH+] KA + [H+] + [Ac-]
)(
[Ac-]
K1-' + [H+] + [=Al:OH]
+ [=Al:OH] KA + [H+] + [Ac-]
KA + [H' KA + [H'] + [Ac-]
mechanism V: [HAC]+ [sA1--0]
[=A1:OH2+] + [Ac-]
)/(
1
[=Al:OH] + [HAC] + KAKl [=AkOH2+] + [Ac-]
))
+ k-2
(B4)
Surfactant-Polyelectrolyte Interactlons. 1. Blndlng of Dodecyltrlmethylammonium Ions by Sodium Dextran Sulfate and Sodium Poly(styrenesulfonate) in Aqueous Solution in the Presence of Sodlum Chkride KatumRu Hayakawa' and Jan C. 1.Kwak' Department of Chemistfy, Dalhousie University, &/ifax, Nova Scotia, B3H A13 (Received:Februaty 16, 1982; In Final Form: June 4, 19821
Isotherms for the binding of dodecyltrimethylammonium (DTA') ions by sodium dextran sulfate (NaDxS) and sodium poly(styrenesulfonate)(NaF'S) in the presence of added NaCl are reported. The binding isotherms were determined by using a potentiometric technique based on surfactant ion selective solid-state electrodes. The solid membranes used in the electrodes consist of poly(viny1chloride) (PVC) plasticized by bis(2-ethylhexyl) phosphate with a DTA-dodecyl sulfate carrier complex. The electrodes exhibit Nernstian response for DTA+ down to concentrations as low as 1 X mol kg-* even in the presence of a large excess of NaC1, allowing for sensitive and accurate free surfactant ion determinations. The binding of DTA+ to both polyanions is shown to be highly cooperative. The cooperativity parameter from the Zimm-Bragg theory may be estimated at 650 f 100 and 200 f 100 for the NaDxS and NaPS cases, respectively, and is independent of the NaCl concentration in both cases. The binding constant K of DTA+ to an isolated site on the polyanion is considerably larger in the PS-DTA system than in the DxS-DTA system, presumably because of differences in the hydrophobic/ hydrophilic properties of the two polymers. K is found to decrease strongly with increasing NaCl concentration; this decrease is similar in magnitude to the decrease in the critical micelle concentration (cmc) of dodecyltrimethylammonium bromide (DTAI3r)with increasing total counterion concentrationin the presence of added NaC1. Introduction The interaction between dissolved surfactants and polymers or colloidal particles is of interest in areas as diverse as polymer solubilization,lJ conformational change in bi0polymers,3-~and mineral flotation and flocculation, including coal flotation6,' and clay flocculation.g10 In the specific case of the binding of ionic surfactants by dissolved polymers, adsorption isotherms exhibiting a marked degree of cooperativity have been This behavior is similar to what is observed in the binding of dyes like acridine orange and proflavine by linear biopolymers.15J6 Surfactants may, in fact, be more suitable for such binding studies because they do not dimerize or associate below the critical micelle concentration (cmc), their hydropho'Permanent address: Department of Chemistry, Kagoshima University, Kagoshima, Japan.
bicity can be estimated, and interaction energies between the hydrophobic parts of the molecule can be compared (1) T. Isemura and A. Imanishi, J. Polym. Sci., 33, 337 (1958). (2) M. N. Jones, J. Colloid Interface Sci., 23, 36 (1967). (3) D. K. Sarker and P. Doty, R o c . Natl. Acad. Sci. U.S.A., 55, 981 (1966). (4) M. J. Grourke and J. H. Gibbs, Biopolymers, 5, 586 (1967). (5) I. Satake and J. T. Yang, Biochem. Biophys. Res. Commun., 54, 930 (1973).
(6) V. L. Basenkova and Yu. N. Zubkova, Khim. Tuerd. Topl. (Moscow), 11, 137 (1977). (7) R. V. Przhegorlinskaya and Yu. N. Zubkova, Khim. Tuer. Topl. (Moscow), 12, 125 (1978). (8)H. S. Hanna and P. Somasundaran,J . Colloid Interface Sci., 70, 181 (1979). (9) J. P. Law, Jr., and G. W. Kunze, Soil Sci. Soc. Am. Proc., 30, 321 (1966). (10) W. F. Howler, Clays Clay Miner., 18, 97 (1970). (11) H. Arai, M. Murata, and K. Shinoda,J . Colloid Interface Sci., 37, 223 (1971).
0022-3654/82/2086-3866$01.25/00 1982 American Chemical Society
