706
Langmuir 1988, 4 , 706-710
these experiments, is shown in Table 111. Along with the ordering, we have included the gas-phase proton affinities for each alcohol as well as the low-coverage heats of adsorption in H-ZSM-5 for the alkane corresponding to the alkyl group for each alcohol. Assuming that sharing of the proton from the A1 site in H-ZSM-5 with the alcohol is a major part of the stabilizing interaction of the 1:l complex, the adsorption affinities should correlate with the gasphase proton affinity values. For water, methanol, and ethanol, the order of the adsorption strength does correlate well with the proton affinities of the molecules. With the larger alcohols, the correlation is less clear. This is probably due to the fact that the proton affinities of these alcohols are close enough so that other factors begin to play an important role in directing the course of the displacement experiments. These other factors may be either kinetic or thermodynamic. Column 3 of Table III shows that, where data exist, there also is a correlation between the preferred adsorbates in these experiments and the isosteric heats of adsorption of the corresponding alkanes. The reported heats of adsorption increased by approximately 2.5 kcal/mol with each additional carbon atom, ranging from 7 kcal/mol for ethane to 14.5 kcal/mol for pentane. It is perhaps not too surprising that both proton affinity and alkane heats of adsorption predict the trend in the displacement experiments since both are known to scale with molecular polari~abilities.~'J~ This may imply that polar and nonpolar (17) Dixon, D. A.; Lias, S. G. In Molecular Structure and Energetics; Liebman, J. F.; Ed.; VCH New York, 1987;Vol. 2,p 269.
interactions between the zeolite and adsorbed alcohol are equally plausible explanations for the adsorption affinity trends. However, it should be noted that proton transfer from the A1 sites must be the major component in the adsorption process since simple alkanes can be easily evacuated from the zeolite at 295 K, whereas each of the alcohols that we have studied cannot be evacuated below a coverage of 1molecule/Al site even after 20 h of evacuation at 295 K.2 Summary By performing a series of sequential adsorption and coadsorption experiments, we have been able to characterize the displacement reactions of adsorption complexes formed by alcohols in H-ZSM-5. In particular, we have shown that there is a clearly defined ordering of adsorption strengths for the alcohols. Though this ordering approximately follows the gas-phase proton affiiity values of the different alcohols, interactions between the carbon backbone and the zeolite channel walls become increasingly important with the larger alcohols and influence both the strength of binding and the chemical properties of the adsorption complexes. Acknowledgment. We are grateful to the Mobil Oil Corporation for supplying us with the ZSM-5 samples. This work was supported by the NSF, MRL Program, under Grant DMR82-16718. Registry No. Methanol, 67-56-1;ethanol, 64-17-5; l-propanol, 71-23-8;2-propanol, 67-63-0; l-butanol, 71-36-3. (18)Richards, R. E.; Rees, L. V. C. Langmuir 1987,3, 335.
Adsorption at Solid/Solution Interfaces. 5. Surface Complexation of Iminodiacetic Acid on Hematite+ Roberto Torres,* Nikola Kallay,s and Egon Matijevie* Department of Chemistry, Clarkson University, Potsdam, New York 13676 Received November 17, 1987. In Final Form: February 10, 1988 Adsorption of iminodiacetic acid (IDA)on uniform spherical colloidal hematite particles was measured
as a function of reactant concentrations and pH. Data were analyzed by using a modified surface com-
plexation model, described earlier, which allowed for identification of the solute species (with respect to their charge) that actually adsorb. The first step in this interpretation considers the adsorption isotherm at the conditions of the isoelectric point (where no electrostatic interaction is present), which is then extended to the entire pH region of interest. In the system studied essentially only singly charged IDA anions, LH-, are adsorbed, occupying an area of 10 A2 with an intrinsic equilibrium constant of 350. A comparison of the calculated surface potentials with the measured electrokineticpotentials led to the estimation of the slipping plane separation of 16 A. Introduction Interactions of metal oxides with complexing molecules are of practical and theoretical significance. For example, chemical decontamination of nuclear power plants by chelating agents is essential to their operation. It has been shown that the dissolution of the corrosion products is directly dependent on the adsorption of such organic additives. The most commonly employed complexing agent Supported by the NSF Grant CHE-8619509. Part of a Ph.D. Thesis. *On leave from the Faculty of Science, University of Zagreb, Zagreb, Yugoslavia. f
has been EDTA, the effects of which with metal (hydrous) oxides have been extensively Since EDTA decomposes at elevated temperatures to yield-among others-iminodiacetic acid and its derivatives, it was of interest to investigate the interactions of such compounds with hematite. (1) Chang, H.-C.; Healy, T. W.; MatijeviE,E. J.Colloid Interface Sci. 1983,92,469. (2)Chang, H.-C.; MatijeviE, E. J. Colloid Interface Sci. 1983,92,479. (3). Blesa, M. A.; Borghi, E. B.; Maroto, A. J. G.; Regazzoni, A. E. J. Colloid Interface Sci. 1984,98, 295. (4)Blesa, M. A.; Maroto, A. J. G. Mater. Sci. Monopr. 1985. 28A (React. Solids, Pt. A), 529.
0743-746318812404-0706$01.50/0 0 1988 American Chemical Society
Langmuir, Vol. 4, No. 3, 1988 707
Adsorption a t SolidISolution Interfaces
From the fundamental point of view, the effects of complexing molecules that are chemically similar (with respect to the bond formation) but differ structurally merit special attention. For example, oxalic and citric acids showed similar adsorption behavior but significantly different ability to dissolve hematite.s~s These effects were explored in terms of surface coverage; Le., the higher surface concentration of the smaller oxalic acid resulted in higher dissolution rate under otherwise identical conditions. The surface interactions with chemisorbed species have been analyzed using various models;' however, even with the simplest adsorbates (neutral electrolyte systems) the conclusions are still questionable? Obviously, an approach is desirable that would lead to a proper explanation of the experimentally observed phenomena. In earlier studies on adsorption of oxalic and citric acids on hematite an interpretation of data was introduced, based on a surface complexation model, that took into consideration all possible adsorbate species.6 The employed model made it possible to evaluate the surface potentials taking into consideration various combinations of these solutes at the interface. The calculated potentials were then compared with the measured electrokinetic potentials over the same pH range, which enabled one to select the actually adsorbed species. This study applies a similar analysis to colloidal spherical hematite particles in contact with iminodiacetic acid (IDA), which is simpler in terms of complexity than EDTA. Adsorption data at the isoelectric point (iep) are especially useful in order to ascertain the degree of protonation of the interacting species and the corresponding intrinsic equilibrium constants, as well as the surface area occupied by the adhered ion(s). A refinement in the analysis quantitatively relates surface to t potentials. This procedure made it possible to test the employed model and to gain additional insight into the structure of the solid/liquid interface. Specifically, the distance of the electrokinetic slipping plane was calculated, which leads to a proper use of the { potential in the interpretation of the double-layer phenomena.
Experimental Section Iminodiacetic acid (IDA; HOOCCH2NHCH2COOH)reagent was used without further purification. Dispersions of colloidal spherical hematite particles of narrow size distribution were prepared by a somewhat modified procedure described earlier? Proper amounts of FeC13stock solutions and concentrated HCl were diluted with doubly distilled filtered water and preheated to 90 O C to obtain a solution 0.02 mol dm4 in FeC13 and 0.001 mol dms in HCI. This solution was then aged at 105 "C for 24 h in a forced-convection oven. After cooling in cold water, the so generated dispersion was diluted with water (1:2) and centrifuged at llOOOg for 30 min, the supernatant solution was discarded, and the precipitate was redispersed with dilute perchloric acid (pH 3) in an ultrasonic bath. The solids were again separated in the ultracentrifuge and redispersed in a NaOH solution of pH 12. After the system was treated with an ultrasonic probe, the particles were allowed to settle and the supernatant solution was decanted. The same washing cycle was repeated many times until no presence of chloride or ferric ions could be detected in the separated liquid. The final dispersion (- 125 g solid/dm3) was kept at pH 9. The average particle diameter as evaluated from electron microscopy was 0.10 0.015 pm, and the
*
(5) Kallay, N.; MatijeviE, E. Langmuir 1986, I, 195. (6) Zhang, Y.;Kallay, N.; MatijeviE, E. Langmuir 1988,1, 201. (7) Adsorption of Inorganics at Solid Liquid Interfaces; Anderaon, M. A., Rubin, A. J., Eds.; Ann Arbor Science: Ann Arbor, MI, 1981. (8) Koop+, L. K.; van Riemedijk, W. H.; Roffey, M. G. J. Colloid Interface Scr. 1987, 118, 117. (9) MatijeviE, E.; Scheiner, P. J . Colloid Interface Sci. 1978,63, 509.
r
"
0
I
Hematite: lOmg cm-' 25°C
,
5
10
lo3c,,(lDA)/mol
15
dm-3
Figure 1. Adsorption isotherms at 25 "C of iminodiacetic acid (IDA) on hematite at different pH values at constant ionic strength, I = 0.01 mol dm-3. specific surface area, measured by the BET multipoints method (Monosorb Surface Analyzer), was 14.4 m2 g-l. Most adsorption data were obtained at 25 "C at a constant ionic mol dm-3 (NaC104) by using 10 mg cm-3 strength of 1 X hematite. The pH was adjusted with perchloric acid and sodium hydroxide. After given equilibration times hematite was separated by centrifugation, the supernatant solution was filtered through a 0.05-pm pore size Nuclepore membrane, and the content in the adsorbate species was determined by a spectrophotometric technique with the ferric-chelate complex.1° Electrophoretic mobility as a function of pH was obtained in a Rank Brothers Mark I1 microelectrophoresisapparatus in a flat cell thermostated at 25 OC. The measurements were conducted with the same systems in adsorption studies (Le., the particles were suspended in the mother liquor). In doing so the adsorption and electrokinetic data refer to analogous conditions.
Results Time-dependent adsorption experiments of iminodiacetic (IDA) acid showed that equilibration was achieved within minutes; any data after the 20-min reaction time remained constant. The adsorption isotherms of IDA for different pH values are shown in Figure 1. The total surface concentration increases with decreasing pH. The experiments were carried out at still higher acid concentrations, and in each case saturation was established, although the maximum adsorbed amount was pH dependent. Reversibility of the adsorption process was confirmed by taking the Fe,O,-IDA system equilibrated at pH 6.6 and acidified to pH 3.4 and vice versa; the isotherms shifted accordingly. The electrophoreticmobilities of hematite in the absence and in the presence of different concentrations of IDA as a function of pH at 25 "C are presented in Figure 2. The isoelectric point (iep) of pure hematite is at pH 7.3, in good agreement with the previously reported value.'l With increasing concentration of the chelating agent the mobility changes appreciably. The shift of the iep to lower pH values indicates the specific adsorption of chelating anions. With data shown in Figures 1and 2 it was possible to determine the surface concentration of IDA on hematite a t the iep for different total concentrations of this acid (Figure 3). Figure 4 illustrates the effect of the addition of acid or base (pH) on the uptake of the complexing agent by he(10) Bhattachayya, S. N.;Kunder, K. P. Talanta 1971, 18, 446. (11) Hesleitner, P.;BabiE, D.; Kallay, N.; MatijeviE, E. Langmuir 1987, 3, 815.
708 Langmuir, Vol. 4, No. 3, 1988
Torres et al.
3
7
2
9
5
'
"E
\
r*
35'
1
I
Q 0
HEMATITE- IDA I = 0 01 mol dm'3 25'C
i
-' J
g
I
I
I
-2
5
-3 2
4
6
8
10
PH Figure 2. Electrophoretic mobilities of hematite particles in the absence (- -) and in the presence of different total concentrations of IDA (solid lines) as a function of pH at 25 OC. I = 0.01 mol dm-3, except for the two highest concentrations of IDA, where I = 0.1 mol dm".
-
HEMATITE-IDA at iep.25"C I
\
PH Figure 4. Bottom: effect of pH on the surface concentration of IDA on hematite at 25 O C at a constant total concentration [IDAIM = 0.005 mol dm-a and I = 0.01 mol dm-3. Top: corresponding equilibrium concentrations of IDA.
used to calculate surface potentials over the entire pH region of interest. By comparison of these data with the measured electrokinetic potentials the separation of the slipping plane can be estimated. The relevant expressions are given below for the hematite-IDA system. Solution Equilibria. For IDA (H2L):
3
+ H+
H2L
H3L+;
YcH&+
K , = -- 66
(1)
CH~L~H+
HL- == L2-
Y3CLzaH+
+ H+;
K2 =
= 1.3 X
(3)
CHL-
where c and a represent the concentration and activity of solute species. The values of equilibrium constants were taken from the literature12 and corrected for the ionic strength effect. The activity coefficient y for 1:l aqueous electrolytes at 25 "C is given by -log y = 0.5091'1/2/(1 (4)
I, B +
OO
0 01
0 02
0 03
c (IDA) /mol dm3 Figure 3. Bottom: adsorption isotherms at 25 "C of IDA on hematite at the conditions of the isoelectric points. Top: corresponding pH(iep) values. I = 0.01 mol dm", except for the two highest concentrations of IDA, where I = 0.1 mol dm-3.
where I' is the ionic strength in mol dm-3. Surface Equilibria. Assumed equilibria are as follows:
S + H3L+
matite at the constant total concentration of 5 x mol dm-3 IDA at the ionic strength I = 1 X mol dm-3. There is a pronounced decrease in r with increasing pH, suggesting that the adsorbate species are negatively charged.
Interpretation of Data The interpretation of data is based on an approach refined from the one previously r e p ~ r t e d .The ~ degree of protonation of the adsorbed acid was evaluated from the isotherm at the isoelectric points (iep). Under these conditions electrostatic interaction need not be considered. The resulting values of the maximum surface concentration, r-, and the intrinsic equilibrium constant are then
S
SLH3+;
+ HL- + SLH-;
Kpint exp(-F$,/RT) =
K,"
exp(F$,/RT) =
~HL-
~SYCHL(7)
S
+ L2-
SL2-;
rLzr&cL2-
K P t exp(2F+,/RT) = - (8)
S is the surface site, r represents the surface concentration (12) SillBn, L. G.;Martell, A. E.Stability Constants of Metal Ion Complexes; Spec. Publ. No. 17, The Chemical Society: London, 1964.
Langmuir, Vol. 4, No. 3, 1988 709
Adsorption ut SolidlSolution Interfaces
HEMATITE- IDA ( H L-I I = 0 01 mol d ~ n ' ~ZS'C ,
0
I
d
0
E
3.
\
v
0.
0.050
c
o
d
h
I
I
I
200 I
HEMATITE- IDA I : 0.01 mol dm-3 25°F
,
400
1
4.0~10~
I
10
PH
I
I
I
2.0 106 I
I
600
8
6
I
Figure 6. Plots of surface potentials (L7 as a function of pH calculated from the adsorption data of IDA on hematite (Figure 4) assuming HL- as the dominant adsorbate species, Klht = 350, and r,, = 1.7 X lod mol dm' (0). Dashed line is the calculated (L,. function from the smoothed ( potential curve (solid line) assuming slipping plane separation of 16 A. Circles represent measured (potentials. Temperature, 25 "C; Z = 0.01 mol dm-3.
Q = [PO + aH+(iep)-'Kgl + ~ ~ + ( i e p ) - ~ K + ~Kzp~ a H + ( i e p ) K ~/J[ 1 + y-'aH+(iep)-'K1 +
nations of adsorbed species. The proper choice of the adsorbed species should lead to a linear relationship between l/r and l/Qc(IDA). Figure 5 displays such plots for all possible combinations of IDA species. The linearity with positive slope as predicted by eq 11-13 was obtained for adsorption of HL- (reaction 7), for combinations HL+ L2- (reactions 7 and 8), and for HL- + H3L+ (reactions 7 and 5 ) . The calculated Qc(IDA) values for these three cases were essentiallythe same due to the negligible contents of H3L+ and L" in the liquid medium over the pH range of interest. In principle these species (especially L2-) could form surface complexes with hematite but were present in concentrations too low to make a detectable contribution to the total adsorbed amount. Consequently, only LHanion need be taken into consideration in the further analysis of the data. Other combinations of adsorbates do not yield a linear relationship and, therefore, can be dismissed. Taking doubly negatively charged anion, L2-, as the only interacting species, one obtains a negative slope of the same plot, which clearly indicates that this assumption is untenable. The calculated value of the equilibrium constant for HLis Klht = 350,and the maximum surface concentration r= 1.7 X mol m-2, which corresponds to an area of 10 A2/molecule. Once the interacting ions are established and the corresponding values of Pnt and rmaare obtained, one can evaluate the surface potential +? from the adsorption isotherm over the entire pH region by means of eq 10. Figure 6 compares the so calculated values for the data of Figure 4, assuming HL- as the only adsorbing anions, with measured { potentials. Both sets of points follow the same trend, and \Ly values are always larger in magnitude than the corresponding { potentials, as expected. Results shown in Figure 6 can also be used to estimate the separation of the slipping plane ({-plane) from the particle surface, or more correctly from the y-plane, representing the location of charged adsorbed species. For this purpose the Gouy-Chapman expression for a 1:l electrolyte was used
Taking p = 1 for strongly interacting adsorbates and assuming p = 0 for all others, one can calculate Q. The values of the later were obtained for all possible combi-
where +x is the potential a t the distance x from the onset
O.5n1O9
l.0n109
1.5=1O9
( Qc / mol dm-3)-' Figure 5. Plot of data shown in Figure 3 according to eq 11-13. Different curves correspond to assumed adsorbed species. a (0): HL-, HL- + L", HL-+ H3L+,HL- + L" + H3L+. Similar resulta
are obtained with combinations H L + HL- and H2L + HL- +
L2-+ HaL+. b (0):H2L,H2L+ Ll-, H2L+ HsL+,H2L+ L2- + H3L'.
c (A): H3L+, H3L+ + L2-. d (V): L2-.
of the adsorbed species, Ptis the intrinsic equilibrium constant, and q7 is the surface potential &e., the electrostatic potential that determines the energy state of the adsorbed ions at the y plane). Mass Balance. The total surface concentration of the adsorbed IDA species, r, is r = r H & + + r H & + rHL- + r L " (9)
Adsorption Isotherms. Equations 1-9 yield
(KOht+ ad+KIKlht&y/RT+ a$+KIK&htem*y/RT
+
UH+KaKpinte-F*y/RT) (10) At the iep
= 0) eq 10 assumes the form r-' = r& + [r,,IQntQc(IDA)]-'
(+?
(11)
which is similar to the Langmuir isotherm. The intrinsic equilibrium constant, IF, refers to predominantly adsorbed species and is related to the individual equilibrium constants by Koht = p p t ; &'Dt = p l p t ; (12) K2'Dt p21Q"t; KPint = P P t For the predominantly adsorbed species the weighting parameter p = 1,while for adsorbates that exhibit a lesser affinity toward the surface 0 < p < 1. The function Q depends on the pH(iep) and on the bulk dissociation equilibrium constants:
yaU~+(iep)-2K& + y-'a~+(iep)K,] (13)
+,
710 Langmuir, Vol. 4, No. 3, 1988
Torres et al.
of the diffuse layer (0-plane), K is the Debye-Htickel reciprocal length defined by K
= (2FI/&NAT)’I2
(15)
and go is related to the potential t,bo (at x = 0)
Accordingly, the distance between the slipping plane ({-plane) and the y-plane is given as xt-x,
=
(17)
Separation distances were calculated in this manner for each experimental adsorption point (Figure 6) and were found to be -16 A and independent of the pH. The dashed line in Figure 6 represents the surface potentials, calculated from the smoothed { potential function assuming the separation of 16 A.
Discussion The interpretation of adsorption data can be facilitated, if in the investigated systems the isoelectric point (iep) and point of zero charge (pzc) are sufficiently different. Such a case was analyzed with aqueous dispersions of colloidal titania, where the uptake of H+ and OH-ions was experimentally determined by potentiometric titration and the difference in the iep and pu: explained in terms of surface ~ontamination.’~ Alternately, one can measure the specific adsorption of inorganic solutes, such as pho~phates,’~ arsenates,15 or organic complexing agental26 (oxalic or citric acids, EDTA, etc.), and use the data to account for the shift in the iep. The surface complexation models that have been employed so far do not yield unique seta of values of parameters describing surface equilibria.* The results depend strongly on the method of optimization of the input data (various numerical and graphical procedures), even for the simplest systems. This deficiency is due to the abundance of parameters and the complexity of their interrelationship, which is further aggravated by the sparsity of the pertinent experimental information. For example, the common potentiometric titration technique yields the difference between the total amount of adsorbed protons and hydroxide ions but does not give the absolute amounts for each ion. In addition, the obtained values include free adsorbed ions as well as the same ions paired up with counterions, i.e., SH+ and SH+.N03-. With more complex systems, such as polyprotic acids, one deals with several possible adsorbate species, which makes it unreasonable to base the interpretation on an a priori assumed set of surface reactions. The treatment developed earlier5 and refined in this study differs from the usual approach in that only two parameters (I’maand Pt) are assessed, yet their values are more reliable. The described procedure also reveals, (13) Kallay, N.; BabiE, D.; MatijeviE, E. Colloids Surf. 1986, 19, 375. (14) Yates, I).E.; Healy, T. W. J. Colloid Interface Sci. 1976,52,222. (15) Anderson, M. A.; Ferguson, J. F.; Gavis, T. J . Colloid Interface Sci. 1976, 54, 391.
in the case of acids, the state of protonation of the actually adsorbed species. In the interpretation of data the adsorption of H+ and OH-ions was considered through their influence on t,bo. It was also assumed that IDA species do not compete with the potential-determining ions for the surface sites. Owing to the low adsorption density of these anions, such an assumption does not affect significantly the calculated t,bo values. The finding in this study that only the HL- ion interacts with the solid surface appears to contradict the previously reported results5 on the adsorption of oxalic and citric acids on a-Fe203. These acids are stronger electrolytes than IDA; i.e., the second dissociation constant (K2)for IDA is - 4 X lo4 smaller. As a consequence L2- species, which could preferentially adsorb, is present in a concentration insignificant to affect the system. This conclusion is also supported by the low value of the occupied area of the molecule (10 A2) for the adsorbed IDA. The L2-ion would bind surface Fe ions with two carboxyl groups, thus occupying a larger area. For e ~ a m p l ethe ,~ doubly charged citric acid anion, which is comparable in size, covers 33 A2. Furthermore, attachment by two bonds would be associated with a higher energy, which can be estimated from the intrinsic equilibrium constants; for the oxalate ion, L2-, Grit = 9800, while for the citrate ion of the same charge, Pt= 7800. Both values are considerably higher than Pt= 350 established for the adsorption of IDA LH- species. The correlation between t,b7 and l potentials further supports the assumption that only the HL- anion is adsorbed at the a-Fe203/IDA interface. The calculated distance of 16 A applies to the separation of the slipping plane from the y plane, where charges of adsorbed HLions are located. Somewhat higher values of the slipping plane separation were suggested by different analyses of simple ~ystems.’~J’ It would seem that the distance of 15-20 is typical for colloidal particles in aqueous environment at room temperature. Thus, one may use electrokinetic data as an additional input in the treatment of the double-layer equilibria evaluated by other means, such as potentiometric titrations. In doing so constrainta are introduced that result in more reliable values of the relevant parameters responsible for interfacial reactions. The rather large separation of the slipping plane, obtained by different approaches, clearly indicates that its usual identification with either the outer Helmholtz (pplane) or the third plane (6-plane) in the triple-layer modeP8J9is not always justified. Slipping plane separation of 15-20 A as obtained from electrokinetic data is in good agreement with the results of this study.20 Knowing the distance of the slipping plane makes it possible to calculate the “surface potential” from the measured { potential. It is precisely the former that is responsible for the stability properties of colloidal dispersions. This potential is characteristic of the onset of the diffuse layer, which governs the interactions between approaching particles. Registry No. IDA, 142-73-4; FezOs, 1309-37-1; hematite, 1317-60-8. (16) Harding, T. H.; Healy, T. W. J . Colloid Interface Sci. 1985, 107, 382. (17) TomiE, M.; Kallay, N. Langmuir, in press. (18) Sprycha, R.; Szcypa, J. J. Colloid Interface Sci. 1984, 102, 288. (19) Sprycha, R. J. Colloid Interface Sci. 1986,110, 278. (20) Kallay, N.; TomiE, M. J. Colloid Interface Sci., submitted.
