Kinetics and mechanism of surface reaction of salicylate on alumina in colloidalaqueous suspension
Z. WANG,1,* C. C. AINSWORTH,2 D. M. FRIEDRICH,1,† P. L. GASSMAN,1 and A. G. JOLY1
1Environmental and Molecular Science Laboratory Pacific Northwest National Laboratory, Richland, WA 993522Interfacial Geochemistry Group, Pacific Northwest National Laboratory, Richland, Washington 99352, USA
(Received January13, 1999;accepted in revised form September13, 1999)
Abstract—The reaction kinetics of salicylate with Al(III) in aqueous solution and at the colloidal alumina–water interface was studied by stopped-flow laser fluorescence spectroscopy. Temporal evolution of thefluorescence spectra suggests that formation of a carboxylate monodentate complex was the reaction inter-mediate that occurs transiently at the beginning of the reaction in aqueous salicylate–Al(III) solution.However, by lowering the pH to 2.0, the formation of such an intermediate can be directly observed as it isthe only species formed. The reaction of salicylate with aqueous Al31 is completed within 10 min at pH 3.3but is significantly slower at pH 2.0. At both pH the aqueous reaction follows a single pseudo-first order ratelaw. In alumina suspension the reaction was initially fast but slowed down after;30 s. Completion of thereaction took up to 12 h, depending on pH and ionic strength. The formation of a carboxylate monodentatesurface complex as a transient species is clearly observed in alumina suspensions at near neutral pH. Theinitial rapid reaction (,30 s), accounting for;70% of the total reaction, can be best described by the Elovichrate equation and the slower reaction, accounting for;30% of the total reaction, obeys pseudo-first orderkinetics. These results are consistent with a sorption reaction mechanism that is controlled by the leavinggroup lability at the surface sites (Al–OH2
1 and Al–OH). The pseudo-first order rate constant varies little withinitial salicylate concentration, ionic strength, or pH. 4, suggesting that the slow reaction pathway involvesligand substitution reactions between salicylate and the hydroxyl groups for which the Al–O binding andactivation energy are affected by site heterogeneity or site density to a lesser degree than Al–OH2
The adsorption of aqueous inorganic oxyanions and organicacid anions to hydrous metal oxides has been described interms of surface complexation reactions involving the forma-tion of mono- and bidentate or binuclear surface complexes.Although the surface speciation resulting from adsorption hasduring the past 20 years been intensely studied, adsorptionkinetics studies are limited and less successful because of theinherent difficulties in data collection and analysis. Kineticadsorption data have often been treated as a pseudo-first orderreaction, a series of pseudo-first order reactions, or a diffusion-controlled reaction (Hingston and Raupach, 1967; Chen et al.,1973; Huang, 1975; Hingston, 1981; Ainsworth et al., 1985;Raven et al., 1998; Westall, 1987). The physical meaning ofsimple rate laws applied to such complex systems is often opento question. A simple rate law may be observed if the param-eters relevant to the kinetic rate constants are not too sensitiveto site or sorbent heterogeneity. However, under the conditionsthat the binding or reaction activation energy varies signifi-cantly with site heterogeneity, the observed reaction rate mayevolve continuously with increasing surface coverage (Lasaga,1981).
Because of their inherently heterogeneous nature, adsorptionreaction pathways or mechanisms are difficult to elucidate.
However, through the use of relaxation methods (Sparks andZhang, 1991) reaction mechanisms for several inorganic oxya-nions (Zhang and Sparks, 1989, 1990) and an organic acidanion (Ikeda et al., 1982) have been deduced. The adsorptionmechanism for these anions is conceptually analogous to theaqueous phase Eigen-Wilkens-Werner mechanism that is thefundamental pathway for strong metal complex formation inaqueous solution (Sposito, 1994). This mechanism, for mono-dentate ligands, involves a two-step process with rapid forma-tion of an outer sphere complex followed by the slower forma-tion of an inner sphere complex with the elimination of a H2Omolecule.
Many adsorption reactions of ligands at a metal oxide–waterinterface have been observed to share common characteristicsof structure, reactivity, and thermodynamic stability with theirhomogeneous aqueous phase counterparts (Kummert andStumm, 1980; Ainsworth et al., 1998; Yost et al., 1990; Tunesiand Anderson, 1992). Cylindrical internal reflection–FTIRstudies indicate that the spectra of ring-substituted benzoicacids at the TiO2 surface are identical to those obtained fromsolution phase Ti(IV) complexes (Tunesi and Anderson, 1992),which reflects the similarity in structure of complexes formedin solution and at solid–water interfaces. Similar spectroscopicresults have been reported for salicylate–Fe(III) complexes insolution and at the goethite–water interface (Yost et al., 1990),as well as the salicylate–Al(III) complex in solution and at thealumina–water interface (Ainsworth et al., 1998).
Using dynamic17O nuclear magnetic resonance spectros-copy, the water exchange rate in aluminum complexes withhydroxide and fluoride (Phillips et al., 1997b, 1998), oxalic
*Author to whom correspondence should be addressed ([email protected]).† Present address: Optical Coating Laboratory, Inc., 2789 NorthpointParkway, Santa Rosa, California 95407-7397, USA. E-mail: [email protected].
acid (Phillips et al., 1997a), and methylmalonate (Casey et al.,1997) was determined. Complexation of Al(III) with theseanions increases the exchange rate of inner hydration spherewater molecules as much as two orders of magnitude or more.Such results are consistent with the interpretation of Secco andVenturini (1975), Perlmutter-Hayman and Tapuhi (1977,1979), and Rakotonarivo et al. (1989) for the faster reactionrate of hydrolysed Al(III) with organic acids than the nonhy-drolysed Al(III) in aqueous solution. In such a reaction removalof the first water molecule from the inner hydration sphere isthe rate determining step. Binding of an anion weakens theAl–OH2 bond of other water molecules and therefore, acceler-ates the rate of ligand complexation. It has also been suggestedthat the increased water exchange rate can be correlated withthe ligand-enhanced dissolution rate of minerals (Pulfer et al.,1984; Zutic and Stumm, 1984; Phillips et al., 1997a,b, 1998;Casey et al., 1998).
In a previous fluorescence spectroscopy investigation of sa-licylate adsorption at thed-Al2O3–aqueous interface we iden-tified four surface complexes (Ainsworth et al., 1998). Theseconsisted of one outer sphere and three inner sphere specieswith a bidentate inner sphere complex being the dominantspecies under the conditions of the study (Fig. 1a). At the lowsurface coverages of this study the bidentate surface speciesaccounted for$90% of the total adsorbed salicylate. At equi-librium a monodentate phenolate surface complex (Fig. 1c) waspresent and could be distinguished from the bidentate species.The monodentate carboxylate surface complex (Fig. 1b) wasnot as unambiguous. At low salicylate surface coverage be-tween pH 4 and 6, the equilibrium concentration of the mono-dentate carboxylate surface complex was beyond the limit of
spectral resolution, although it was observed in pH 2.0 aqueousAl(III) solutions.
Although the above study addressed a number of questionsregarding salicylate surface speciation, issues concerning salic-ylate adsorption kinetics and mechanism arose. First, prelimi-nary kinetic data (unpublished) suggested that, similar to oxya-nions and other organic acids, salicylate time-dependentadsorption is composed of a very rapid adsorption followed bya slower process. Second, investigation of the monodentatecarboxylate surface complex (Ainsworth et al., 1998) suggestedthat the bidentate surface species is formed through a mono-dentate intermediate. Together with the observed presence ofan outer sphere complex (Fig. 1d), this suggested that the pathof formation of the predominate salicylate bidentate surfacespecies is similar to the aqueous phase Eigen-Wilkens-Wernermechanism. This study was undertaken to determine the rateand possible mechanism of formation for the bidentate surfacespecies that were previously shown to dominate the surfacespeciation of sorbed salicylate. Adsorption kinetics of salicylateon d-Al2O3 was studied by stopped-flow with fluorescencedetection. The temporal evolution of the salicylate fluorescencespectra and the growth and decay kinetics of the fluorescence ofsalicylate complexes, over a range of pH, ionic strength, andinitial salicylate concentrations, were used to elucidate theoverall adsorption process leading to the observed time-depen-dent adsorption curve.
2. MATERIALS AND METHODS
2.1. Materials
Aqueous suspensions of aluminum oxide (Aluminum Oxide C,d-Al2O3, Degussa AG.) and solutions of sodium salicylate were pre-pared as described previously (Ainsworth et al., 1998). Briefly, thecommerciald-Al2O3 was washed, suspended, and pH adjusted forseveral days to achieve a stable pH of 6.0. After centrifugation, thegel-like top layer of the sediment was carefully removed and suspendedagain in the supernatant solution. This produced a stable suspensionwith 70 nm average colloidal particle size. Sodium salicylate (2-hydroxybenzoate; 100.1% assay, Mallinckrodt Chemical Co.; sal orHA2) was used as received and made up into pH-adjusted 1027 mol/Land 1025 mol/L salicylate solutions in background electrolyte ionicstrength (IS) of 0.01 mol/L NaClO4. AlCl3 solution (23 1023 mol/L)was prepared by dissolving reagent grade AlCl3 (Aldrich) and usedimmediately after preparation. The pH of the AlCl3 solutions wereadjusted to either pH 3.3 or 2.0 at the time of dissolution using diluteperchloric acid. Deionized water (18 MV) from a Millipore reverseosmosis and ion exchange system with ultraviolet sterilization of thefinished water was used throughout the study. Care was taken to avoidlong periods of water contact with plastic tubing and vessels to elim-inate contamination by residual fluorescing impurities (e.g., phtha-lates).
2.2. Experimental Methods
A schematic diagram of the apparatus used for the kinetics measure-ments is shown in Figure 2. The experiments involved mixing thereactants rapidly, followed by UV excitation, and detection of visiblefluorescence. Mixing of aqueous sodium salicylate solution, typicallybetween 1027 mol/L and 1025 mol/L, with either 1 g/L aluminasuspension at pH 6.0 and IS 0.01 mol/L (sodium perchlorate) or 0.002mol/L fresh AlCl3 solution at pH 2.0 or 3.3 was performed in a Hi-TechScientific SFA-211 stopped-flow kinetics accessory with a deadtime of50 ms. The temperature-controlled sample cuvette has three polishedquartz windows allowing fluorescence detection at 90° relative to theexcitation beam.
To record the temporal evolution of fluorescence spectra, the solu-
tion mixture was excited with a continuous-wave argon laser (CoherentInnova 400) at either 300 or 334 nm. Fluorescence collected by a 2 inchdiameter f/2 lens was focused by an f/4 lens into a 0.22 m focal lengthspectrograph (Spex 270). Wavelength-dispersed emission was detectedby an intensified diode array detector (1024 elements Princeton Instru-ments IPDA with CSMA data acquisition software).
To record accurate growth and decay kinetics of the fluorescencefrom salicylate complexes, a SPEX Flourolog II fluorimeter was usedto record the fluorescence intensity as a function of time at fixedexcitation and emission wavelength. The fluorimeter was equippedwith a 450 W xenon arc lamp, double monochromators (SPEX 1680)for excitation and emission, and a cooled photomultiplier in photoncounting mode. Because of the higher spatial stability of the xenon lightfocused at the sample, fluorescence intensity data recorded with xenonlamp excitation have higher signal-to-noise ratios than with argon laserexcitation. Therefore, data used in the analysis of the kinetic parameterswere all recorded with xenon lamp excitation using narrow bandpass(1–2 nm) excitation and emission wavelengths of the fluorimeter. Tocorrect for minor variation of the lamp intensity, light from the exci-tation monochromator was reflected by a thin quartz plate into aRhodamine dye-cell. Red emission from this “quantum counter” wasdetected by a second photomultiplier tube and recorded by the fluo-rimeter in an analog reference channel. Fluorescence intensity from thesample was normalized at each time to the corresponding excitationintensity in the reference channel.
As described previously (Ainsworth et al., 1998) the bidentate Al–salicylate (Al:A1) complex (Fig. 1a) is the predominant species at lowsalicylate concentration (1027 to 1025 mol/L) at pH 6.0 and IS 0.01mol/L. Both free salicylate and Al:sal have distinct excitation andemission maxima. Because both free and complexed salicylates areexcited at the absorption maximum near 300 nm, both species contrib-ute to the emission intensity at 415 nm (lmax of salicylate). However,at 334 nm the complexed salicylate is predominantly excited, therefore
at 380 nm only the complexed salicylate is detected (Ainsworth et al.,1998). The monodentate complex Alz HA21 shows an excitationmaximum at 311 nm and emission maximum at 412 nm. Therefore, byperforming the kinetics experiments at characteristic excitation andemission wavelengths the complexes Al:A1 and Al z HA21 can bestudied with high selectivity.
Salicylate samples in both homogeneous aqueous solution and het-erogeneous alumina suspension were examined for the effect of pos-sible photobleaching. The results indicated that within 10 min thedecrease in fluorescence intensity at the excitation and emission wave-length for the samples was,1%. Because the time ranges used forkinetics analysis were within even shorter periods, the effect of pho-tobleaching was negligible.
2.3. Data Analysis Procedure
The fluorescence intensity of the aluminum-salicylate complex in-creases rapidly initially and then asymptotically reaches a maximumlimit for both the homogeneous (aqueous) and heterogeneous (colloi-dal) reactions. The asymptotic intensity limit I` can be estimatedaccurately by fitting the longer time portion of the fluorescence increasecurve to smooth functions such as exponentials (omitting data for thefirst 30 s). For aqueous phase experiments I` represents the fluores-cence intensity of the complexes with equilibrium concentrations thatcan be accurately calculated from the stability constants of the com-plexes and the protonation constants of the acid. Thus, the concentra-tion of the complex, [AlA], representing either Al:A1 or Al:HA21,depending on the excitation and emission wavelengths and solution pH,at time t can be calculated using Eqn. 1:
[AlA] 5It
I`[AlA] e (1)
Fig. 2. A schematic drawing of the experimental setup for the study of salicylate adsorption kinetics in aqueous aluminasuspension. (a) Stopped-flow apparatus; (b) sample cell; (c) collection and focusing optics; (d) filters; (e) mirror; (f)monochromators; (g) intensified diode array detector; (h) computer; (i) photomultiplier tube; (j) photon counter andcomputer (see text for detail).
1161Salicylate-alumina reaction kinetics
where [AlA]e is the concentration of the complex of interest at equi-librium. For clarity, charge was omitted in all concentration expres-sions.
For the kinetics measurements involving alumina suspensions thetotal salicylate concentration is much less than the alumina reactionsites. Under similar conditions previous work by Ainsworth et al.(1998) indicated that during 24 h of reaction time free salicylateconcentration was about 10% as determined by14C activity of 14C-labeled salicylate left in the supernatant after centrifugation. However,with careful centrifugation of the suspensions at extended adsorptiontimes (.24 h), under the experimental conditions of the present study,the concentration of free salicylate was,2% as determined from thefluorescence intensity of the supernatant, and therefore the equilibriumadsorption of salicylate-to-alumina is essentially complete. Thus, thelong time asymptotic limit concentration of the surface complex [Al:Al] infin; is assumed to be equal to the initial salicylate concentration[A] tot. Interference from the fluorescence of free salicylate is mini-mized at the excitation and emission wavelengths (lex 5 334 nm,lem 5 380 nm). At these wavelengths the recorded fluorescenceintensity I(t) is proportional to the concentration of the emitting com-plex. The concentration of the complex at time t, [A]t can be calculatedby
[AlA] t 5 I(t) [A] tot/I` (2)
and the unreacted salicylate concentration at time t can be calculated as
[A] t 5 [A] tot 2 [Al:A] t 5 [A] tot (1 2I(t)/I`) (3)
The reaction kinetics data of salicylate with aqueous Al(aq)31 and colloi-
dal alumina at time.30 s were found to be best fit by pseudo-firstorder reaction and therefore, analyzed based on the following treat-ment.
For the reaction
Al31 1 HA2 -|0k
k9AlA 1 1 H1 (4)
The reaction rate
d[AlA]
dt5 2
d[A]
dt
5 k[Al] z [A] 2k9[AlA][H] (5)
Because only Al and A are present at the start,
[Al] 0 2 [Al] 5 [AlA] (6)
[A] 0 2 [A] 5 [AlA] (7)
where subscript 0 denotes reaction at time zero. Therefore,
Fig. 3. Emission spectra of Al31–salicylate(aq) as a function of reaction time at pH 3.3. [Al31] 5 0.002 mol/L,[Sal] 5 2 3 1025 mol/L. lex 5 300 nm. Time range: 0–4 min.
1162 Z. Wang et al.
At equilibrium, 2d[A]
dt5 0
k[Al] 0 z [A] e 5 k9[H][AlA] e 5 k9[H]([A] 0 2 [A] e) (12)
where the subscript e refers to equilibrium concentrations then,
[A] e 5k9[H]
k[Al] 0 1 k9[H][A] 0 (13)
and Eqn. 11 can be rewritten as
ln S[A] 0 2 [A] e
[A] 2 [A] eD 5 ln S [AlA] e
[AlA] e 2 [AlA] D (14)
or
ln([AlA] e 2 [AlA]) 5 2(k[Al] 0 1 k9[H])t 1 ln[AlA] e (15)
A plot of ln([AlA] e 2 [AlA]) or ln([sal] 2 [sal]e) vs. time t will yielda straight line with slope of2(k[Al] 0 1 k9[H]), the pseudo-first orderrate constant kobsd and intercept of ln[AlA]e.
3. RESULTS AND DISCUSSION
3.1. Salicylate Reaction Wth Al(aq)31
The product of the reaction between Al31 and the salicylateanion (HA2) in slightly acidic solutions (pH$ 3) has beenshown to be a 1:1 bidentate complex (Secco and Venturini,1975; Rakotonarivo et al., 1989; Thomas et al., 1993; Ain-sworth et al., 1998). At the ratio of Al to salicylate used (100:1)in the present study and the one by Ainsworth et al. (1998), thereaction is essentially complete and the only equilibrium salic-ylate species observed is the monosalicylate–Al complex.
Al (aq)31 1 HA23 Al:A (aq)
1 1 H1 (16)
The temporal evolution of the fluorescence spectra from theaqueous reaction at pH 3.3 is shown in Figure 3. Because thepKa1of salicylate is 2.95, approximately 31% of salicylate is inthe protonated salicylic acid form. The fluorescence of salicylicacid is very weak and contributes,1% to the observed inten-sity (Ainsworth et al., 1998). With 300 nm excitation both freesalicylate anion (lmax 5 296 nm) and the Al:A1 complex(lmax 5 312 nm) are excited. If present, however, the AlzHA21 monodentate carboxylate complex (lmax 5 306 nm)would also be excited. As shown by Ainsworth et al. (1998),the bidentate Al:A complex (binding aluminum at both itscarboxy and phenoxy oxygens) is characterized by its strongfluorescence maximum at;383 nm. Thus, Figure 3 shows thefluorescence spectrum of the reaction mixture evolving frominitially free salicylate (410 nm) to predominantly bidentatecomplexed salicylate (;383 nm) at longer times. Because ofthe 15-nm red shift in the excitation band maximum of Al:A1
compared to HA2 (Ainsworth et al., 1998), excitation at 334nm only excites the Al:A1 complex, which produces time-evolving fluorescence spectra dominated by the growth of thebidentate Al:A1 complex (Fig. 4).
By performing the same experiments at pH 2.0, time-evolv-ing fluorescence spectra of the monodentate species can beisolated (Fig. 5). The ability to isolate this species arises fromthe observations that (1) in the absence of Al31 90% of salic-ylate is in the weakly emitting protonated salicylic acid form atpH 2.0 and (2) that the fluorescence intensity of Al–salicylate
Fig. 4. Emission spectra of Al31–salicylate(aq) as a function of reaction time at pH 3.3. [Al31] 5 0.002 mol/L,[Sal] 5 2 3 1025 M. lex 5 334 nm. Time range: 0–4 min.
1163Salicylate-alumina reaction kinetics
complexes in such low pH Al31 solutions are many times moreintense than in the absence of Al31 (Ainsworth et al., 1998).Therefore, the growth of fluorescence intensity in Figure 5 isassigned to the formation of an aluminum complex that bindsthe salicylate ion through the carboxylate functional group(nominally monodentate), but still allows an intramolecularH-bond from the phenol group to the carbonyl group of thecarboxylate moiety. As explained in Ainsworth et al. (1998),the evidence for this intramolecular H-bond is the large energyloss between the excitation maximum (;300 nm) and thefluorescence maximum (;410–415 nm). This energy differ-ence is due to proton transfer from the phenol group to thecarbonyl group in the excited electronic state (excited stateintramolecular proton transfer, ESIPT) (Gormin and Kasha,1988; Nagaoka et al., 1988; Barbara et al., 1989). An intramo-lecular H-bond between the phenol (H-donor) and the carbonylof the ortho-carboxylate group (H-acceptor) is required forESIPT to occur.
Monitoring the kinetics at fixed excitation and emissionwavelengths produced Al:A1 and Al z HA21 growth kineticcurves I(t) of high signal-to-noise ratio. For Al:A1 excitation at334 nm and monitoring the emission at 383 nm produces adirect observation of the Al:A1 complex growth (Fig. 4).Similarly, at low pH the formation of the Alz HA21 monoden-tate complex may be followed by excitation at 300 nm andmonitoring of the emission at 412 nm. A plot of log ([sal]t 2[sal]e) at a function of time t yields straight lines with thepseudo-first order rate constants of 0.040 s21 (Al:A 1 at pH 3.3)and 0.009 s21 (Al z HA21 at pH 2) (Fig. 6). An important
feature of the aqueous reaction kinetics at both pH 2 and 3.3 isthat both are pseudo-first order (i.e., second order) over 1.5decades of intensity data. In the aqueous system there is noevidence of competing reaction mechanisms over this timerange.
Although the reaction between aqueous Al and salicylate hasbeen studied previously (Secco and Venturini, 1975; Rako-tonarivo et al., 1989), these studies relied on detection of UVabsorption at 310 nm and calculations of intermediate step ratesrather than direct detection of the monodentate intermediatespecies. They characterized the reaction path for the Al–salicy-late reaction as an Eigen-Wilkens-Werner mechanism followedby the elimination of H3O
1 and ring closure (Secco and Ven-turini, 1975; Rakotonarivo et al., 1989):
(H2O)6Al 31 1 HA2N [(H2O)5 Al(H2O) 2 HA]21 (17)
[(H2O)5Al(H2O) 2 HA]21N [H2O)5AlHA] 21 1 H2O (18)
[(H2O)5AlHA] 21N [(H2O)4AlA] 11 1 H3O1 (19)
The elimination of water from the inner sphere coordinationsphere of Al(III) (Eqn. 18) has been suggested to control thereaction rate for the formation of the bidentate product ratherthan potential steric factors associated with ring closure (Eqn.19) (Secco and Venturini, 1975). The present results comparefavorably with these previously published results on the Al–salicylate reaction (if differences in the Al(aq)
31 concentrationsand the solution acidity are taken into account) and directlyconfirm the formation of a monodentate intermediate. That is,
Fig. 5. Emission spectra of Al31–salicylate(aq) as a function of reaction time at pH 2.0. [Al31] 5 0.002 mol/L,[Sal] 5 2 3 1025 M. lex 5 300 nm. Time range: 0–4 min.
1164 Z. Wang et al.
the results of the current study show that formation of the AlzAH species is the slow step in the aqueous reaction and hencefollows an Eigen-Wilkens-Werner mechanism.
It was noticed that although the formation of the monoden-tate complex is the rate-determining step, at elevated pH (e.g.,pH 3.3 in this study) the overall reaction rate for the formationof the bidentate complex is actually greater than the rate offormation of the monodentate complex. Similar results werealso obtained by Secco and Venturini (1975) and Perlmutter-Hayman and Tapuhi (1977, 1979) for Al(III) reaction withsubstituted salicylates, and Rakotonarivo et al. (1989) for sa-licylate complexation with hydrolysed Al(III). The increasedreaction rate can be explained by reaction between salicylateand hydrolysed Al(III), such as AlOH21:
(H2O)5AlOH21 1 HA2N [(H2O)4AlOH(H2O) 2 HA]11
(20)
[(H2O)4 AlOH(H2O) 2 HA]11N [(H2O)4AlOH 2 HA]11
1 H2O (21)
[(H2O)4AlOH 2 HA]11N [H2O)4AlA] 11 1 H2O (22)
In these reactions removal of H2O from the inner sphere ofAl(III) (Eqn. 21) is the rate-determining step, following theEigen-Wilkens-Werner mechanism, but the hydroxyl group inthe inner sphere of Al(III) reduces the strength of the Al–OH2
1
bond and results in increased reaction rate with salicylate
(Secco and Venturini, 1975; Perlmutter-Hayman and Tapuhi,1977, 1979; Phillips et al., 1997b, 1998). From calculationsbased on the available hydrolysis data of Al(III) (Martell andSmith, 1995), at pH 3.3 the equilibrium concentration of hy-drolyzed Al(III), such as (H2O)5AlOH21, is about 0.3% of thetotal Al(III) concentration, which is comparable to the salicy-late concentration considering the large [Al(III)]:[sal] ratio inthe solutions.
3.2. Reaction With Colloidal Alumina
Figure 7 shows the temporal evolution of the fluorescencespectra (lex 5 334 nm) from the heterogeneous salicylatereaction
where s- represents the solid and i is the stoichiometric numberof nonprotonated hydroxyl sites (0, 1, or 2). The study wasconducted at approximately one order of magnitude lowersalicylate concentration than in the aqueous solution reaction,and the ratio of the total number of surface sites to salicylatewas about 1000:1 (Ainsworth et al., 1998). The spectra ob-tained from excitation at 334 nm evolve in a manner similar tothe aqueous reaction (Fig. 4). As expected only the surfacebidentate species (383 nm) is observed. However, excitation of
Fig. 6. Pseudo-first order kinetics of Al31 complexation with salicylate as a function of reaction time at pH 3.3 (a) andpH 2.0. (b) [Al31] 5 2 3 1023 mol/L, [Sal] 5 2 3 1025 mol/L, IS 5 0.01 mol/L. At pH 2.0,lex 5 310 nm,lem 5 410nm. At pH 3.3,lex 5 334 nm,lem 5 383 nm.
1165Salicylate-alumina reaction kinetics
this reacting system at 300 nm reveals a spectral evolution thatappears to differ from its aqueous analogue in three importantaspects (compare Fig. 8 to Fig. 3): (1) the early phase of thisreaction appears more rapid, (2) the asymptotic approach to theproduct (383 nm) appears slower (note that after about 60 s the383-nm emission has yet to dominate the spectra as in Fig. 3),and (3) the initial spectrum is significantly red-shifted (;423nm) relative to free salicylate (412 nm). The later point sug-gests that an intermediate (;423 nm) is rapidly formed at thishigher pH and evolves toward the bidentate product.
The intermediate is likely to be the monodentate surfacecomplex in which the salicylate phenolic hydrogen is held in anintramolecular hydrogen bond to the salicylate carbonyl oxy-gen, whereas the phenolic oxygen is hydrogen bonded to theproton of a neighboring surface–aluminol or hydronium group(Fig. 1b). This would preserve the intramolecular H bond,which is responsible for the ESIPT emission above 400 nm.The electrostatic field of the phenol-to-surface intermolecularhydrogen bond could stabilize the phenoxide form of the ES-IPT excited state, leading to a red shift of the fluorescencerelative to a complex with a free or unpolarized phenolicoxygen. Importantly, in our earlier equilibrium study of salic-ylate surface speciation, identification of the transient mono-dentate carboxylate surface complex was ambiguous and be-yond the limit of spectral resolution (Ainsworth et al., 1998).Although in the present study we can observe the evolution ofthe monodentate surface species, we cannot adequately isolate
its growth and subsequent disappearance in a manner thatwould allow the calculation of separate rates for the formationof Al z HA21 and Al:A1.
Rakotonarivo et al. (1989) investigated the rate of aqueoussalicylate complexation with the Al31 monomer and Al13 poly-cation and concluded that the reaction of salicylate with thepolycation was faster than that with monomers. However, theirstopped-flow system used UV detection of absorption at 310nm. At this wavelength the growth of the monodentate (ifpresent as an intermediate in their system) and the final biden-tate product would be detected as a single species. Hence, it isunclear as to what was being monitored in that study. However,the current results suggest that the reaction path is similar to theaqueous reaction, in that a monodentate intermediate complexis formed, and supports the conclusion (at least for the initialstage) that the surface complexation reaction is faster than itsaqueous counterpart. In addition, surface formation of theAl z HA21 monodentate intermediate through the eliminationof H2O (Eqn. 25) does not appear to be the rate-limiting stepunder the conditions of this study. Rather the slow step appearsto be ring closure (Eqn. 26).
Fig. 7. Emission spectra of salicylate in 1 g/L alumina suspension as a function of reaction time at pH 6.0. [Sal]5 2 31026 mol/L. IS 5 0.01 mol/L.lex 5 334 nm. Time range: 0–2 min.
1166 Z. Wang et al.
[(s 2 Al(OH) 2 HA] N [(s 2 Al:A] 1 H2O (26)
To simplify the equations (Eqns. 24–26), the number of non-protonated hydroxyl sites (i in Eqn. 23) was assumed to be one.
It should also be pointed out that there are significant differ-ences between the reactions of a ligand with Al(III) in homo-geneous phase and with the solid–water interface. In light of thefact that in ligand substitution reactions the leaving grouplability determines the reaction rate, hydrolysis, or complex-ation of Al(H2O)6
31 will enhance the rate of additional com-plexation because of the weakening of Al(III) binding to theremaining inner hydration sphere water molecules (Secco andVenturini, 1975; Perlmutter-Hatman and Tapuhi, 1977, 1979;Phillips et al., 1997b). However, this effect is less likely forreactions on the surface where there are no more inner spherewater molecules to be replaced except those hydronium orhydroxyl groups, which themselves will be directly involved inthe ligand substitution reactions (Eqn. 23).
Kinetic growth curves of salicylate surface complexes incolloidal alumina suspensions were obtained by excitation at334 nm and viewing emission at 380 nm (Fig. 9). Plots of log([sal]t 2 [sal]e) vs. time for salicylate reaction with aluminasuspensions were pseudo-first order only after approximately30 s, accounting for'30% of the total reaction (Fig. 10),depending on pH and ionic strength. As can be seen in Figure10, the rapid early portion of the reaction accounts for'70%of the total reaction and cannot be fit to a pseudo-first order ratelaw. Indeed, this data cannot be fit reasonably to any simple
combination of first or second order or two- or three-constantrate laws. However, this initial rapid reaction is fit easily to theintegrated form of the Elovich rate equation (Low, 1960;McLintock, 1967; Hingston, 1981; Ungarish and Aharoni,1981; Sposito, 1994; Aharoni et al., 1991) over a wide range ofpH, ionic strengths, and initial salicylate concentrations.
In the Elovich rate law, the rate of adsorption dq/dt, whichwe take to be the rate of formation of surface complexes,exponentially decreases with increasing coverage q:
dq/dt5 a z exp(2aq) (27)
where “a” is the reaction rate at zero coverage, or, in practice,at the beginning of the measurement. The coverage scale factora is the reciprocal of the coverage q1/e at which the adsorptionrate has fallen to 1/e of its initial value.
(dq/dt)l/e 5 a z exp(2aql/e) 5 a/e (28)
The integrated form of the Elovich equation is
q(t) 5 (2.303/a)log(11 t/t0) (29)
where the fitting constant t0 5 (aa)21 is the time at which thesurface coverage q(t0) has increased toa21ln2 5 q1/eln2. Anexample of fitting the early time growth of the surface complex[Al:A] to the integrated Elovich rate law is shown in Figure 11,which plots the concentration of salicylate adsorbed to thesurface [Al:sal]t (equivalent to the surface coverage q(t)) vs.
Fig. 8. Emission spectra of salicylate in 1 g/L alumina suspension as a function of reaction time at pH 6.0. [Sal]5 2 31026 mol/L. IS 5 0.01 mol/L.lex 5 300 nm. Time range: 0–2 min (reaction is not completed yet).
1167Salicylate-alumina reaction kinetics
ln(1 1 t/t0). The fitting parameter t0 is adjusted to achieve thebest linear fit to the data and the slope isa.
The Elovich equation was developed for and is widely usedin describing the kinetics of heterogeneous chemisorption ofgases on solid surfaces (Low, 1960; Adamson, 1990). Theequation assumes a heterogeneous distribution of adsorption oractivation energies that vary continuously with surface cover-age. It has been used successfully to describe the adsorption ordesorption kinetics of phosphate (and other ions) to soils andsoil minerals (Chien and Clayton, 1980; Hingston, 1981; alsosee Sparks, 1989, and references therein). As discussed bySparks (1989) several investigators have used changes or mul-tilinear segments in the Elovich plot as an indication of multi-site binding or shifts for one site type to another. However,concerns have been raised regarding the use of the Elovichequation to conclude mechanistic information concerning mul-tiple sites, changes in adsorption energies with surface cover-age, or site heterogeneity (Sposito, 1994; Aharoni and Sparks,1991; Stumm, 1992). Yet under conditions that the bindingenergy or reaction activation energy varies significantly withsite heterogeneity, the observed reaction rate constant mayevolve continuously with increasing surface coverage, hencethe usefulness of the Elovich equation. In the present study it isextremely useful as a characterization of rate data and as a way
of bounding the extent of the adsorption reaction described bya pseudo-first order rate law.
The variation of the Elovich parameters (a,a) and the pseu-do-first order rate constant kobsdwith respect to [sal]tot, pH, andIS were studied to explore the relationship between theseexperimental conditions and both the fast reaction and slowerreactions. In every case the kinetics could be accurately fit atlong times to a single exponential (pseudo-first order) rateconstant k and at short times to an integrated Elovich rate law.The pseudo-first order rate constant (46 0.3 3 1023 s21) ispractically independent of [sal]tot, and hence the rate of thisreaction increases only slightly with concentration (Table 1).However, the initial rate (or instantaneous rate at t5 0) for theElovichian part of the reaction (defined by the “a” term) in-creases with [sal]tot. As [sal]tot increases from 1 to 16mmol/L,the portion or extent of the reaction described as pseudo-firstorder increases from 16% to 28%. The first-order rate constantis also invariant with pH and IS, yet the extent of its contribu-tion to the overall reaction increases as pH and IS increase(Table 2). The Elovich initial reaction rate, however, is affectedsubstantially by changes in pH and IS.
The pH-induced changes in the rates, rate constants, andextent of reaction for the two parts of the overall reactionappear to reflect changes in the surface charge and speciation at
Fig. 9. Fluorescence intensity of salicylate in 1 g/L alumina suspension as a function of reaction time at pH 6.0. [Sal]58 3 1026 mol/L. IS 5 0.01 mol/L.lex 5 334 nm.lem 5 383 nm. The dashed line is the asymptotic limit (see text fordetails).
1168 Z. Wang et al.
the aqueous–alumina interface with pH. The results are consis-tent with a sorption reaction mechanism that is controlled bythe leaving group lability at the surface sites (Al–OH2
1 andAl–OH). The distribution of surface sites with pH, based on thetriple-layer model (Davis et al., 1978; Davis and Leckie, 1978)was calculated using FITEQL (Westall, 1982). For the purposeof these calculations the total concentration of surface sites anddensity are 1.33 1023 mol/L and 8 sites/nm2 (surface area).Surface acidity constants and zero point charge (ZPC) of thesolid are as reported in Ainsworth et al. (1998). The concen-tration of the total (.Al–OH2
1) sites are calculated to be 2.2331024, 1.053 1024, and 1.953 1025 mol/L at pH 4, 6, and 8,respectively. Likewise, the concentration of (.Al–OH) sitesover this same pH range are calculated to be 1.13 1023,1.233 1023, and 1.323 1023 mol/L. At pH 4 the (.A–OH2
1)site concentration is about 100-fold greater than [sal]tot. How-ever, this excess decreases to slightly greater than equal molarconcentrations at pH 8. Although the protonated sites are ex-pected to be the most active salicylate adsorption sites, clearlythe concentration of these sites affects the reaction rate. On theother hand, the concentration of the neutral (.Al–OH) sitesremain almost constant and well in excess of [sal]tot. Largerconcentrations of hydronium sites, generated by either loweringthe suspension pH or increasing the alumina solid to salicylateratio, will favor the reaction that is expressed by the Elovich
pathway, which is typically faster than reaction following thepseudo-first order pathway.
These results suggest that there are multiple sites active insalicylate adsorption. Conceptually, we view the protonatedaluminol sites as good candidates for the rapid Elovich pathway
s2 Al(OH)(OH2)1 1 HA23 s2 Al:A 1 2H2O (30)
or
s2 Al(OH2)221 1 HA23 s2 Al:A 1 2H2O 1 H1 (31)
because protonation of the hydroxyls will increase both thesubstitutional lability of the surface aluminol bond and thevariation of its binding (and activation energy) at heteroge-neously distributed sites on the alumina surface. The pseudo-first order rate constant varies little with initial salicylate con-centration, ionic strength, or pH. 4. A plausible interpretationfor the slow pseudo-first order reaction pathway involves li-gand exchange between salicylate and the nonprotonated sur-face hydroxyl groups
s2 Al(OH)2 1 HA23 s2 Al:A 1 H2O 1 OH2 (32)
for which the Al–OH binding and activation energy are muchhigher than those of Al–OH2
1. We hypothesize that the muchless labile Al–OH bond is less affected by site heterogeneity orcharge density as compared with the Al–OH2
1 bond. In addi-
Fig. 10. Pseudo-first order plot of salicylate adsorption kinetics from Fig. 9 at reaction time. 30 s and approximately70% of the adsorption.
1169Salicylate-alumina reaction kinetics
tion, as the results of the surface site speciation calculationsshow, the (.Al–OH) sites are in enormous excess under theexperimental conditions of this study. Thus, there appear to betwo parallel (competing) reaction paths for the formation of thebidentate salicylate species at the alumina–water interface. Themajor pathway is rapid and follows Elovich kinetics, and theminor pathway is much slower following second order (pseu-do-first order) kinetics. Such interpretations are supported bythe calculated extent of reaction (Table 2). As pH increasesfrom 4.0 to 8.0, the extent of reaction following the Elovich rateequation decreases from 98% to 68%, which appears to be adirect consequence of the conversion of the Al–OH2
1 sites toAl–OH sites as pH increases.
Although the pseudo-first order rate constant changes little,the initial reaction rate a for the fast Elovich pathway showeda profound change within the pH and ionic strength rangeexplored. In the Elovich equation, the initial reaction rate a is
the rate at zero surface coverage. Results in Table 2 show thatthe initial rate decreases as pH increases from 4.0 to 8.0 and asionic strength increases from 0.01 to 1.0 mol/L. Electrostaticattraction between cations and anions is responsible for theformation of an outer sphere complex, which is the first step inthe Eigen-Wilkens-Werner mechanism (Eqn. 17). Conversionof hydronium sites to hydroxyl sites decreases the density ofthe cation sites on the alumina surface that may cause the initialreaction rate to decrease. At a fixed pH, the density of hydro-nium sites is relatively constant with respect to ionic strength.However, an increase of ionic strength results in (1) a shrinkingof the electric double layer and (2) local charge neutralizationon the solid surface due to competition with perchlorate anion.Both of these effects lead to reduced electrostatic interactionbetween the surface hydronium sites and salicylate anion andreduce the surface concentration of the outer sphere complex,thus leading to the reduction of initial reaction rate.
Fig. 11. Elovich plot of salicylate adsorption kinetics from Fig. 9 at reaction time, 30 s. In the data fit, t0 5 0.16 s;a 5 1.06 z 1026 mol/L s21.
Table 1. Elovichian and pseudo-first order kinetic data for salicylate adsorption on alumina at different initial salicylate concentration.a
a pH 6.0, I 5 0.01 mol/L, 25°C, 1 g/L alumina suspension; the error limits on t0 a and a are estmated to be#10%; k error limits are estimatedto be#4%.
b [(Sal)]E 5 the adsorbed salicylate concentration described by the Elovich equation.
1170 Z. Wang et al.
4. SUMMARY AND CONCLUSIONS
O-salicylate anion adsorption kinetics were investigated bystopped-flow with fluorescence detection over a range in pH,IS, and [sal]tot. The results of these studies indicate that theadsorption reaction path has both similarities to and differencesfrom the path deduced for the homogeneous reaction of aque-ous Al31 and salicylate. The aqueous reaction is characterizedas an Eigen-Wilkens-Werner mechanism (Eqns. 17–19), andthe slow step is the elimination of H2O in the formation of anintermediate monodentate species. Although the surface reac-tion does form a monodentate carboxylate complex at thesurface followed by ring closure, the slow step of this processappears to be ring closure rather than the monodentate speciesformation. This is seen clearly in the transient build up of the423-nm species, and its subsequent conversion to the 383-nmbidentate surface species. Although we did not observe theformation kinetics of the outer sphere complex in aluminasuspension or the aqueous systems, the presence of an outersphere complex has been indicated (albeit at higher salicylateconcentrations) in equilibrated salicylate–alumina systems(Ainsworth et al., 1998). Hence, we conclude that the majorproduct of salicylate adsorption is a bidentate Al:A complexwhose formation closely follows that observed for aqueous Alcomplexation mechanism: (1) formation of an outer spherecomplex, (2) formation of an inner sphere monodentate car-boxylate complex accompanied by the loss of H2O, and (3)followed by ring closure and the formation of bidentate com-plex. However the monodentate complex is stabilized throughthe interaction between the phenolic oxygen and adjacent alu-minol groups, making ring closure the slow step in the adsorp-tion process rather than monodentate formation as in aqueoussolution.
The kinetics of salicylate adsorption is dependent on surfacesite speciation. Conceptually, we view the adsorption of salic-ylate to alumina to consist of two competing (parallel) reac-tions. The rapid pathway is sensitive to surface site heteroge-neity (the protonated sites) producing Elovich kinetic growth.The slower second order reaction is pseudo-first order under theconditions of low surface coverage (surface aluminol in excessover the sorbate concentration). These two pathways are rep-resented by the overall reactions:
Rapid Elovich:
s2 Al(OH)(OH2)1 1 HA23 s 2 Al:A 1 2H2O (33)
or
s2 Al(OH2)221 1 HA23 s2 Al:A 1 2H2O 1 H1 (34)
Slower pseudo first order:
s2 Al(OH)2 1 HA23 s2 Al:A 1 H2O 1 OH2 (35)
In this chemical reaction scheme, transient intermediates mayform (such as the monodentate complex suggested by thespectra excited at 300 nm) but are not shown in the overallreactions.
The proposed mechanism suggests the hypothesis that siteheterogeneity has a significant influence on lability and activa-tion energy of the protonated sites Al–OH2
1, which results inrapid Elovich kinetics over those sites, whereas the lability andactivation energy of the neutral aluminol sites Al–OH are solittle affected by surface site heterogeneity that the ligandsubstitution reaction on the colloidal particles is simply secondorder with a single pseudo-first order rate constant. We arecurrently studying the role of activation heterogeneity in pro-ducing Elovich-like kinetic behavior.
Acknowledgments—The authors thank Dr. Donald M. Camaioni for aloan of the stopped-flow kinetics apparatus and Dr. Stephan Joyce forhelpful comments regarding the interpretation of the Elovich rate law.Also the authors wish to gratefully acknowledge an anonymous re-viewer for his/her helpful comments, careful review of the data anal-ysis, and constructive criticism that greatly contributed to improvingthis manuscript. This research was supported by the Molecular ScienceResearch Initiative at Pacific Northwest National Laboratory. PacificNorthwest National Laboratory is operated for the U.S. Department ofEnergy by Battelle under Contract DE-AC06-76RLO 1830.
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