Environ. Sci. Technol. 1989, 23, 349-356
(7) Prudic, D. E. Ground-Water Hydrology and Subsurface
Migration of Radionuclides at a Commercial RadioactiveWaste Burial Site, West Valley Cattaraugus County, New York. U.S.G.S. Professional Paper 1325, 1986. (8) Gray, D. H.; Weber, W. J. Presented at the Seventh Annual Madison Waste Conference, University of Wisconsin, Madison, WI, 1984. (9) Quigley, R. M.; Ogunbadejo, T. A. In Glacial Till: A n Inter-Disciplinary S t u d y ; Legget, R. F., Ed.; The Royal Society of Canada: Ottawa, Canada, 1976; Special Publication No. 12, pp 336-345. (10) Patterson, R. J.; Frape, S.K.; Dykes, L. S.; McLeod, R. A. Can. J. Earth Sci. 1978, 15, 162-169. (11) Pankow, J. F.; Rosen, M. E. HRC CC, J. High Resolut. Chromtogr. and Chromtogr. Commun. 1984, 7 , 504-508. (12) Johnson, R. L. Ph.D. Dissertation, Oregon Graduate Center, Beaverton, OR, 1984. (13) Johnson, R. L.; Anderson, M. R.; Pankow, J. F., unpublished
material. (14) Hassett, J. J. University of Illinois, Urbana-Champaign,
personal communicationto Michael Anderson of the Oregon Graduate Center, 1985. (15) Blake, G. R. In Methods of Soil Analysis: Black, C. A., Ed.; American Society of Agronomy: Madison, WI, 1965; pp 374-383. (16) Vomocil, J. A. In Methods of Soil Analysis; Black, C. A., Ed.; American Society of Agronomy; Madison, WI, 1965; pp 299-301. (17) Hydrology Consultants Limited, Toronto, Ontario, 1984.
Unpublished report.
(18) Nesbitt, B. Tricil Limited, Sarnia, Ontario, personal communication, 1984. (19) Karickhoff, S.W. J. Hydraul. Eng. 1984, 110, 707-735. (20) Crank, J. The Mathematics of Diffusion. Oxford Unversity Press: Oxford, UK, 1975. (21) Bear, J. Dynamics of Fluids i n Porous Media; American Elsevier: New York, 1972; pp 106-117. (22) Li, Y.; Gregory S.Geochim. Cosmochim. Acta 1974, 9, 703-714. (23) Bonoli, L.; Witherspoon, P. A. J. Phys. Chem. 1968, 72, 2532-2534. (24) Wilke, C. R.; Chang, P. AZChE J. 1955,1, 264-270. (25) Rowe, R. K.; Booker, J. R. University of Western Ontario,
Faculty of Engineering Science, Geotechnical Research Report GEOT-9-83, 1983. (26) Mabey, W. R.; Smith, J. H.; Podol, R. T.; Johnson, H. L.; Mill, T.; Chou, T.-W.; Gates, J.; Partridge, I. W.; Jaber, H.; Vandenberg, D. Aquatic Fate Process Data for Organic Priority Pollutants. EPA Report No. 440-4-81-014, 1982. (27) Karickhoff, S. W.; Brown, D. S.; Scott, J. A. Water Res. 1979, 13, 241.
Received for review December 9, 1986. Revised manuscript received August 9, 1988. Accepted October 27,1988. W e extend our appreciation to Tricil Limited for financial support of this work, and for the cooperation of Blake Nesbitt and other Tricil staff at the field site. Additional financial support was provided by a Strategic Grant from the National Sciences and Engineering Research Council of Canada and by the Northwest Environmental Research Center.
Aluminum Complexation by Catechol as Determined by Ultraviolet Spectrophotometry Frank J. Sikora" and Murray B. McBride Department of Agronomy, Cornell University, Ithaca, New York 14853
I Methods of ultraviolet (UV) spectrophotometry were
used to determine the stoichiometry and association constant for the Al-catechol complex from pH 3.8 to 4.6. Job's method of continuous variation indicated the Al-catechol complex had a 1:l stoichiometry in the p H range studied. Aluminum titrations of catechol and p H titrations of catechol plus A1 resulted in a shift in the UV spectra due to the formation of an Al-catechol complex absorbing UV radiation uniquely different than that of free catechol. General equations were developed for the determination of association constants assuming an organic and Al-organic complex absorb UV radiation. Aluminum titrations with constant catechol concentration yielded a log K6.1of 16.22 for a 1:l Al-catechol complex. Calculated absorbance as a function of pH agreed well with experimental p H titrations of solutions containing catechol plus Al. The fact that A1 can be complexed by catechol at low p H indicates the o-hydroxy group provides a potential source for A1 complexation in soil and surface waters.
Introduction With the advent of increased A1 mobility in the environment due to acid precipitation, interest in A1 chemistry of soils and surface waters has developed. Adverse effects of soluble A1 include potential toxicity to fish ( I ) ,trees (2), and agricultural crops ( 3 ) . The toxicity of Al, however, is
* Address correspondenceto this author. Present address: F242 NFDC, Tennessee Valley Authority, Muscle Shoals, AL 35660. 00 13-936X/89/0923-0349$0 1.50/0
not related to total A1 concentration. Aluminum can form complexes with organic compounds, reducing the activity and subsequent toxicity of A13+ in solution (1, 4, 5 ) . Equilibrium constants for A1 complexation is therefore important for predicting levels of available AP+ in the environment. The o-hydroxy groups of polyphenols are of particular interest because of the production of polyphenols during the biodegradation of lignin (6). Ultraviolet (UV) spectrophotometry can provide valuable information on the composition and stability constant of a metal-ligand complex. Although the UV method for determining composition and association constants is not valid for mixed-ligand systems and humic substances with multifunctional groups (7,8),the method can be used to define a single metal-ligand complex (9-11). The determination of composition and association constants is based on the fact that an ultraviolet or visible spectrum of the metal or ligand alone is different from the spectrum of the complex. Association constants of A13+ complexes with Tiron, 2,3-pyridinediol, 5-sulfosalicylic acid, and salicylic acid have been determined by using information provided by UV spectral shifts (12-14). In flavonoid chemistry, bathochromic shifts upon A1Cl3 addition are used to indicate the presence of o-hydroxy or o-hydroxycarbonyl chelating groups (15). Objectives of the present study were to develop equations for the determination of association constants of Al-ligand complexes by UV spectrophotometry and to use the equations in conjunction with Job's method of continuous variation to determine the composition and asso-
0 1989 American Chemical Socie!ty
Environ. Sci. Technol., Vol. 23, No. 3, 1989
349
ciation constant of the Al-catechol complex. The association constant for the Al-catechol complex has been previously determined by potentiometric pH titration methods (16-18). Ultraviolet spectrophotometric detection is more direct since the absorbance and concentration of individual species are monitored. Catechol is used in this study to test the validity of the derived equations for 1:l meta1:ligand complexation. Ultraviolet spectrophotometric determination of association constants using the derived equations has potential merit for other simple aromatic organics of environmental importance.
AI SPECIES
OH
OH
IOH^
onZ'
0 2
~ 1 3 +
A 1 (OH) 2+
~i
(on)2 +
CATECHOL SPECIES
nL-
HZL
Theoretical Considerations
L2-
AI-CATECHOL SPECIES
Al-Ligand Association Constants. The complexation of AP+ by an organic ligand can be represented as x ~ 1 3 ++ YLZ- + ( ~ 1 , ~ , ) 3 -
(1)
a1
with a conditional stability constant at an ionic strength of 0.1 M (KC,,,)of
K8.i = m c / ( m ~ i N m ~ ) ~
(2)
where mc, mA1, and mL are the molar concentrations of unhydrolyzed A1 complex, unhydrolyzed free Al, and deprotonated ligand, respectively. Since the concentrations of the completely deprotonated organic ligand (mL) and the unhydrolyzed A1 species (mAl,mc) are dependent on pH, side-reaction coefficients (pi)must be considered as follows:
pL = 1 + PA1
+
=
~IK?,)
(3)
+ + [Kl.K2&YAI/
(aH) 3YAl(OH)31
(mQl-i/
i=l
/ (aH) 2YAl(OH),l
P=l
(4) PAIL
=
+
[KlaYAIL/aHYAIL(OH)l +[KlaK2aYA1L/ (aH)2YAlL(OH)21
+
[KlaK2aK3aYA1L/(aH)3YAIL(OH)~l
(5)
where mH and aHrefer to concentration and activity of H+, respectively. The constants K&, KC,21, ...,KE;? in eq 3 refer to consecutive concentration proton dissociation constants of the organic acid where j is the number of protons. The values of K1 through K3 and Kla through KBarepresent activity-based equilibria constants for the hydrolysis of free and complexed A13+, respectively. Concentration-based proton dissociation constants of 10-9.23and 10-13.0 for catechol (19) were used to calculate pL in eq 3. Aluminum hydrolysis constants were obtained from Lindsay (20). Activity coefficients (7) for H+, A13+,and A1 hydrolysis products were determined by the extended Debye-Huckel law. Chemical species that were considered to be present in solution containing A1 and catechol are shown in Figure 1. The geometry of the o-hydroxy oxygens is favorable for A1 chelation. Aluminum hydrolysis in the complex has not been directly proved or disproved. The equilibria constants between AlL+ and the hydrolysis products of AlL(OH), A1L(OH)2-, and A1L(OH)32-(Kia-K3*) may or may not be similar to the hydrolysis constants for free A13+ (K1-K3). Polymeric species with more than one A1 were not represented in Figure 1, but they may also form at higher pH as postulated with A1 and gallic acid (32). Considering all the species that could be present, the total concentrations of the Al-ligand complex (m:), free 350
I
0%
/OH2
\
/A[\OH2
/OH
/A[\OH2
0 2
0 2
A I L( O H ) 0
A~L+
A l L ( OH) 2 -
Figure 1. Chemical species considered in a solution containing AI and catechol. A third water molecule is hydrolyzed in the species AI(0H) and AIL(OH);-. I n accordance with eq 6 , rn; = [AIL'] jAlL(OH)a] -I-[AIL(OH),-] iAIL(OH),2-], = [H,L] [HL-] [L -1, mi: = [AI3+] [Al(OH)' ] [Al(OH),+] [Al(OH)30], rn, = [AIL'], mL = [L*-], and mA,= [AI3+].
+
+
+
dF
+
+
+
+
ligand (mEF),and free A13+ (mxr) can be respectively represented as
m$ = mcPAIL
mzF = ~
L P L m F = ~ A I P A I (6)
Rearrangement and substitution of eq 6 into eq 2 yields
j
[KIYAl/aHYAI(OH)l
[K1K2?'AI
a1 I
OH2 \
Environ. Sci. Technol., Vol. 23, No. 3, 1989
K6.1 = m$(pL)y(pA,)x/(m~p)x(m~F)YpAU, (7) Stoichiometry of Al-Ligand Complex. The values of x and y in eq 7 can be determined from Job's method of continuous variation where mole fractions of A1 and L (designated as 1- q and q, respectively) are continuously varied while keeping the added molar concentration of both components constant. The theory of Job's method is presented by various authors (8-10,21). A brief outline of the theory is given below. Assuming the absorptivity of the uncomplexed metal is zero at a particular wavelength, the absorbance ( A ) of a solution mixture of ligand and metal will be equivalent to A = ccm$ + ELmzF (8) where tC and cL are the molar absorptivities of the total complex and total free ligand at a specific pH, respectively. The difference in A of a metal ligand and ligand solution is equivalent to the A of the complex, AA, and can be represented as AA = A - tLmxF = ccmE (9)
+
Ideally, a wavelength is chosen where t~ mx, where all the ligand in solution can be potentially complexed. Substitution of these values into eq 20, acknowledging t = mZ with y = 1 (eq 17), and taking logarithms yields
m$ = (A”- tL”mZF)/EC”
A plot of log [(A- - A ? / ( A ’ - A,)] vs log [ m f l- rnT(A,i, - A’)/(Amin- Amax)]from an A1 titration should yield a slope of 1. The association constant for the Al-ligand complex can be calculated from the x intercept as
(11)
Substitution of eq 11 into eq 8 for wavelength 1 and eq 10 into eq 8 for wavelength 2, respectively yields
mrF = [(A’/€,’) - at]/$
(12)
m$ = It - ( A ’ / ~ E L ’ ) ~ / $
(13)
and where $ = 1 - [tc’tL’’”L’tC’’]
{ = A”/tC” cy
= €C’/€L’
w = tC”/tL”
(14)
If wavelength 2 is chosen at an isosbestic point where the absorbance is constant with increasing concentration of the complex, the value for w is equivalent to 1 and 4 = 1 - (EC’/tL’).
The definition of t can also be arranged, as follows, to depend only on variables of wavelength 1. Equation 8 for wavelength 2 can be written at the isosbestic point as A ” = tc”[mE - msCy - 111 (15) where m: is the concentration of total ligand and is equivalent to mxF + ym:. Substitution of eq 15 into t yields
t = mE-m$Cy-1)
(16) Further substitution of eq 13 into eq 16 and rearrangement yields
t = [mE$ + [A’(Y - l ) / ~ t ~ ’ I l / ( $+ Y - 1)
(17) If y = 1, { is simply the concentration of total ligand in solution (m:). Considerin the com lexation equilibria in terms of total A1 (mxl = mfy + xmc) yields
P
K8.1 = ~ $ ( P L ) ~ ( P A I ) ~-/ (xm~)x(m~F)yPA,L ~& (18) Assuming x = 1, substitution of eq 12 and 13 into eq 18 and rearrangement yields
An increase or decrease in A’ can be accomplished with an increase in mxl at constant pH (A1 titration) or an increase in pH, changing PL, at constant mxl (pH titration). Equation 19 can be simplified if y is assumed to equal 1 and an isosbestic point exists where w = 1.
The value for l/~$cL’is equivalent to {/(tL’C- t L ’ a t ) . Since t ~ ’ {equals the absorbance of just the organic at a specific pH, EL{is the minimum absorbance of a pH or A1 titration
log K8.1 = log (PMPL/PA~L)- (x intercept)
(22)
where the Pi constants are side-reaction coefficients as shown in eq 3-5. Rearrangement of eq 19 to allow for calculation of expected absorbance as a function of pH and mzI would be instructive. However, an expression for A’ = f(m:,,pH) cannot be obtained. On the other hand, mI1 can be determined as a function of A’and pH upon rearrangement of eq 19:
mxl =
(PLIYPAI
PALK8,1d1-Y
[
t - (~’/weL’) [(A’/EL’)- .{Iy
]
+
l- (~’/weL’) 4 (23)
Absorbance (A)’ as a function of mxl at constant pH can be obtained by reversing the calculated dependence of mzl on A’. Absorbance as a function of pH at constant mxl can be determined through iterative calculations were values for mzl and pH are input into eq 23 and values for A’are estimated until the right-hand side of eq 23 equals the inputed mxl. In the development of the preceding equations, assumptions were made that w = x = y = 1 in eq 21 and x = 1 in eq 23. For the case of Al-catechol complexes at low pH, there is strong evidence in the literature (17,18), and supporting evidence presented in this study, that x and y do equal 1. For determining Kg,l with other stoichiometries, one must solve a polynomial expansion in eq 18 for x # 1or use the logarithmic form of eq 19 for x = 1 and y # 1.
Experimental Conditions Catechol was obtained from Aldrich Chemical Co. (purity of 99+%) and was used without any further purification. Catechol stock solutions were prepared at 10 mM. Aluminum stock solutions were prepared at 0.1 and 1 M using AlCl,. All solutions in the experiments were kept at a constant ionic strength of 0.1 M by using KCl as the background electrolyte. Solution pH was adjusted with dropwise additions of KOH or HC1 as necessary. Job’s plots were obtained at pH 4.2 and 4.6 from solutions containing lo-, total molar concentration of A1 plus catechol with varied mole fractions of each. Eleven solutions were prepared for each pH with the mole fraction of catechol ( 4 ) ranging from 0 to 1.0 by 0.1 increments. The UV absorbance was obtained a t 12 wavelengths ranging from 230 to 302 nm. The UV absorbance spectra of the solutions were determined with a computer-assisted Perkin-Elmer Lambda Environ. Sci. Technol., Vol. 23, No. 3, 1989
351
Table I. Job's Plot Coefficients, K and q , and qmaxValues Obtained by Using eq 24 To Model the Datan wavelength, nm
K
7
Qmar
K
D
4ma.
230 232 235 239 245 249 279 282 287 292 297 302
0.83 f 0.05* 0.98 f 0.07 1.23 f 0.09 1.08 f 0.09 0.62 f 0.07 0.32 f 0.04 0.39 f 0.04 0.58 f 0.04 0.95 f 0.06 0.98 f 0.08 0.63 f 0.06 0.21 f 0.03
1.04 f 0.07 1.10 f 0.08 1.12 f 0.10 1.14 f 0.10 1.16 f 0.13 1.18 f 0.17 0.79 f 0.09 0.95 f 0.08 1.12 f 0.08 1.15 f 0.10 1.16 f 0.11 1.24 f 0.15
0.49 0.48 0.47 0.47 0.46 0.46 0.56 0.51 0.47 0.46 0.46 0.45
1.78 f 0.39 2.10 f 0.43 2.39 f 0.47 2.28 f 0.43 1.30 f 0.24 0.66 f 0.13 0.78 f 0.19 1.24 f 0.29 2.06 f 0.43 2.07 f 0.40 1.32 f 0.26 0.45 f 0.09
1.08 f 0.25 1.05 f 0.23 1.02 f 0.22 1.01 f 0.21 1.00 f 0.20 0.98 f 0.21 1.02 f 0.28 1.08 f 0.26 1.06 f 0.25 1.02 f 0.21 1.01 f 0.21 1.00 f 0.23
0.48 0.49 0.50 0.50 0.50 0.51 0.49 0.48 0.48 0.50 0.50 0.50
pH 4.2
ORegression models and regression coefficients, confidence intervals.
K
pH 4.6
and q, were all significant a t the 5% level. r2 values were greater than 0.90. *95%
4C spectrophotometer. The effective bandwidth, scan speed, and response were set at 1 nm, 300 nm min-l, and 3, respectively. The pH of the solutions was determined with a Fisher Model 805MP pH meter with an Orion Model 91-15 combination pH electrode. The solution cell used for the titrations consisted of a 150-mL beaker placed in a 600-mL beaker that was partially filled with chilled water to maintain a constant temperature of 25 f 1 "C. The solution was stirred with a magnetic stirrer during the course of the titration. A thermometer and pH electrode were inserted into the solution cell to monitor temperature and pH. To obtain UV absorbance changes during the titration, the solution was continuously pumped through a UV-vis 1-cm cell by a Masterflex peristalic pump. The flow rate was set at 75 mL min-l. Aluminum titrations of 0.4 mM catechol were conducted at pH 3.8, 4.0, 4.2,4.4, and 4.6. Fifty milliliters of 0.4 mM catechol was placed in a 150-mL titration cell and adjusted to the desired pH. Solutions of 1, 10,80, and 800 mM A1 were titrated into the cell to obtain increasing concentration of Al. The pH was adjusted with KOH or HC1 to maintain a constant pH. After titrant addition, the absorbance was determined a t 239,249,282, and 297 nm at an equilibrium time ranging between 2 and 6 min. To maintain a constant concentration of catechol during the A1 titration, the A1 titrant solutions contained 0.4 mM catechol. The pH of the titrant solutions were adjusted in order to reduce the need for large amount of acid or base addition during the titration to maintain a constant pH. The pHs of the 1 mM A1 titrant solutions were adjusted to the respective pH of the titration. The 10 mM AI titrant solutions were adjusted to 3.8 for the pH 3.8 titration and 4.0 for all other titrations. The pHs of the 80 and 800 mM A1 titrant solutions were adjusted to 3.5 and 2.0, respectively. Since exact concentration of added KOH was not recorded, accurate OH/Al ratios for the Al titrant solutions cannot be reported. Approximate OH/A1 ratios ranged from 0.06 to 1.72. The titrant solutions were prepared 1 day before the titration. The pH titrations of 0.4 mM catechol alone and 0.4 mM catechol in the presence of A1 were conducted through dropwise addition of 0.1 or 0.01 M KOH to a 50-mL sample solution. Aluminum concentrations used were 0.3,3, and 30 mM. Before base addition, the pH was lowered below 2.5 with dropwise addition of 0.1 M HC1. The UV absorbance was obtained at 239,249,282, and 297 nm during the titration. A t high A1 concentration and high pH the absorbance was observed to decrease slowly and failed to reach a steady state within 15 min. The slower reaction occurring 352
Environ. Sci. Technol., Vol. 23, No. 3, 1989
1
I
PH 4 . 2
U
0.6 0
co 0.4
a 0.2
\
o0.0 ,2&, o"0
0.1
02
0 3
, 04
,
05
9
, 0 6
, 07
08
-
I
0 9
I O
Figure 2. Job's plots for AI plus catechol solutions at three wavelengths with total molar concentration of 1 mM. Symbols represent experimental data and solid line is best fit function obtained by using the appropriate coefflcients from Table I in eq 24.
at high A1 concentration and high pH was concluded not to be simple chelation of monomeric A1 by catechol, which would occur rapidly. Therefore, only the absorbance data with a short reaction time between 2 and 6 min were analyzed.
Results and Discussion Stoichiometry and UV Spectra of Al-Catechol Complex. Job's plots for A1 complexation with catechol, as shown in Figure 2, were modeled with the empirical equation AA = K [ Q ( 1 - Q ) " ]
(24)
Equation 24 modeled theoretical Job's plot data well [calculated from equations presented by Gilbert (23)]. The empirical equation was used to model the Job's plots so qmax could be quantitatively determined from the first derivative of the function. By using an empirical model, subjective determination of qmaxfrom the intersection of extended lines was avoided. The coefficients, K and 7, were determined for each data set by using the nonlinear least-squares analysis package PCNONLIN (24) and are presented in Table I with qmax. The values for qmax were very close to 0.5 for all the wavelengths (Table I). The qmaxvalue of 0.5 confirmed the observation, through potentiometric titrations ( I 7, I B ) , that an Al-catechol complex with a 1:l stoichiometry dominates below pH 5.0. The roton dissociation conare 10-9.23and stants for catechol (K;l1, and Complexation of A1 by catechol at low pH indicates A1 has
9
2.01
,
I
I
0.6'
-6
952 C O K LIMITS
7
Am,
I
-5
-4
-3
-2
-I
0
tl
t2
Log m;, Figure 4. Spectrophotometric AI titrations of 0.4 mM catechol at 282 nm. Solid lines were calculated from eq 23 by using cc determined from the corresponding A , values in Table I1 (e.g., tC = A mx/rnl) and the Bppropriate log values In Table 111.
G,,
ID3
210
230
250
270
2%
310
330
WAVELENGTH ( nrnl Flgure 3. Ultraviolet spectra of 0.4 mM catechol with (A) Increasing AI concentratlon at pH 4.0, (6)increaslng pH with 3 mM AI and (C) increasing pH.
a high affinity for the ortho-hydroxy oxygens since protons with very low dissociation constants are displaced. The UV absorbance spectra of catechol shifted to longer wavelengths with increased A1 a t constant pH and increased pH a t constant A1 (Figure 3, parts A and B, respectively). The bathochromic shift in the spectra was due to formation of AlL+, which has different UV absorption characteristics compared to H,L. The presence of isosbestic points a t 221.5, 259, and 274 nm supported the hypothesis that a fairly simple equilibrium existed in solution that involves the presence of only two dominant species absorbing UV radiation, namely H2L and AlL+. A bathochromic shift in the H2L spectra was also observed upon an increase in pH to 9,Q4(Figure 3C). The spectra for catechol a t low and high pH are similar to reported spectra in the literature (25). Most catechol molecules have lost only one proton a t pH 9.94 since the pK"Ol of the second hydroxyl is high. The resultant species, HL-, can be visualized as a proton chelated by the two 0 groups (26) as shown in Figure 1. The bathochromic shifts of the H2L spectrum upon formation of AlL+ and HL- are due to a shift of electron density from the oxygens into the aromatic ring (27,28). Electron density is released into the ring since AI or a single H has lower electron-attracting capacities than two H. The similarity in the UV spectral shifts due to AIL+ and HL-
formation indicates that H+ and A13+ on the ortho hydroxy oxygens have similar geometry and electron-attracting capacities (as shown in Figure 1). Although A13+ has a greater charge than H+, the small size of H+ results in electron-attracting capacity comparable to one A13+ nucleus. If the spectra of L2- could be obtained, a bathochromic shift even greater than that obtained for AlL+ or HL- would be expected, since no electron-attracting nuclei would be present on the oxygens and an even greater electron density would be present in the ring system. Association Constant for Al-Catechol Complex. The association constant for the AIL+ complex was determined from A13+ titrations of catechol solutions (data for 282 nm shown in Figure 4). Since stable absorbance readings could not be obtained at higher A1 concentration, A,, was not directly obtainable from the data. An extrapolation method was developed, using a modified Langmuir equation,
+
A = [KJm&/(l K m z ) ]
+ I)
(25)
where K, J, and I)are best fit coefficients determined by PCNONLIN (24). From the best fit coefficients, A,,, was calculated as J + $. The modified Langmuir equation was used since it accurately determined A,, for sets of theoretical data calculated from eq 23. The A,, values for the A1 titrations are shown in Table 11. There was a considerable amount of error in the calculated A,, values, as indicated by the large confidence intervals, because data were not available a t higher A1 concentrations. The A,, values were used in eq 21 to determine log KC,., from each A1 titration. Values for Aminwere taken as the absorbance of catechol in solution without added Al. Plots of log [(Amin - A3/(A' - Amax)] vs log [mxl - mz(A,in A')/(A- - A,)] at 282 nm are shown in Figure 5. Linear regression was performed only on the data that were within approximately 1 log unit of the end point of the titration where log [(A,, - A3/(A'- A,=)] = 0. Values for log K;.l were calculated from eq 22 by assuming the complexed A1 hydrolyzed (RAIL = PA]) and are presented in Table 111. The determined values for log K8.1,and calculated eC values from A,,,, were used in eq 23 to calculate functions of absorbance vs log mzl which are shown as solid lines in Figure 4. The average log K:.l of 16.22 for the association constant of the Al-catechol complex (Table 111) agrees well with the values of 16.30 and 16.15 determined by potentiometric titrations a t an ionic strength of 0.1 M (16, 17). A lower Environ. Sci. Technol., Vol. 23, NO. 3. 1989
353
Table 11. Values of A for the A1 Titration Data as Determined from the Best Fit of eq 26 to the Experimental Datao wavelength, nm pH 3.8 239 249 282 297 pH 4.0 239 249 282 297 pH 4.2 239 249 282 297
Am,b
1.725 f 0.298c 1.002 f 0.174 1.378 f 0.118 0.862 f 0.082 1.440 f 0.166 0.846 i 0.100 1.354 f 0.052 0.807 f 0.036
wavelength, nm pH 4.4 239 249 282 297 pH 4.6 239 249 282 297
Amaxb 1.764 f 0.514 1.003 i 0.291 1.524 f 0.274 0.818 f 0.116 1.435 h 0.214 0.464 f 0.056 1.328 f 0.102 0.759 f 0.121
1.518 f 0.310 0.870 f 0.178 1.396 f 0.162 0.804 f 0.099
Regression models and regression coefficients were all significant at the 5% level. r2 values were greater than 0.99. *Amax= J + +. '95% confidence intervals determined from the 95% confidence intervals of J and &. rj
R .
o p H 3.8
o
I
pH 1.1
Table 111. Linear Regression Results of eq 21 Fit to the Linearized A1 Titration Data (282 nm Shown in Figure 4) and Calculated Association Constants for the Al-Catechol Complex Obtained by using eq 22 wavelength, nm pH 3.8 239 249 282 297 pH 4.0 239 249 282 297 pH 4.2 239 249 282 297 pH 4.4 239 249 282 297 pH 4.6 239 249 282 297 total avd
log
av log
Yint"
slopea
Xint
Kc,,,*
1.24 1.24 1.45 1.35
0.950 0.950 1.04 0.980
-1.31 -1.31 -1.39 -1.38
16.11 16.11 16.19 16.18
16.15
2.00 1.96 1.96 1.89
1.05 1.04 1.04 1.01
-1.90 -1.88 -1.88 -1.87
16.30 16.28 16.28 16.27
16.28
2.24 2.24 2.28 2.24
0.998 0.998 1.03 0.999
-2.24 -2.24 -2.21 -2.24
16.24 16.24 16.21 16.24
16.23
2.18 2.19 2.22 2.42
0.873 0.873 0.895 0.917
-2.50 -2.51 -2.48 -2.64
16.10 16.11 16.08 16.24
16.14
3.44 3.37 3.55 3.44
1.12 1.09 1.15
-3.07 -3.09 -3.09 -3.10
16.27 16.29 16.29 16.30
1.11
G.1'
16.29 16.22 f0.08
"Linear models and coefficients were all significant at the 5% level. r2 values were greater than 0.96. Calculated from eq 22 by assuming flAI = PAIL. From nontransformed K$,l values a t the four wavelengths. From the average nontransformed Kc,,, values at each pH. The 95% confidence interval is reported for the five pH observations. 1.4
-1.2'
-4
'
-3
LOG [ ( m
A, -.cT( ( A m , n - R ' I / ( A
-2 m,
I
EXP DATA WITH o 0.3 rrM A I
-I
- A nox 1) I
Flgure 5. Plots of eq 21 using the spectrophotometric AI titration data at 282 nm (shown in Figure 4). Symbols represent experimental data and solid line is the best fit linear regression model.
value of 15.89 was determined a t an ionic strength of 0.6 M (18). The association constants from the literature were calculated from the original data by using the deprotonfor catechol (19). ation constants of 10-9.23and An advantage of the UV method over potentiometric titration is that the spectroscopic method allows direct monitoring of more than one species involved in the equilibrium whereas potentiometric titrations only measure protons released in the solution. Important chemical information on the electron density of the complex is therefore revealed in the UV spectra that could not be obtained from potentiometric titrations. Since UV spectrophotometry is a more direct method, conclusive evidence on whether more than two prominent species exist in solution can be obtained. For example, the similarity in Job's plots data for 12 different wavelengths (Table I and Figure 2), isosbestic points present in the changing catechol spectra with increased A1 and pH (Figure 3A,B), and nearly equivalent log KC,., values obtained at four different wavelengths at each pH (Table 111) indicate that a fairly simple equilibrium existed between free and complexed catechol. The effects of pH on absorbance at constant A1 concentrations are shown in Figure 6 for 282 nm. The other 354
Environ. Sci. Technol., Vol. 23, No. 3, 1989
0.6'
2
3
5
I 6
p4H Flgure 6. Spectrophotometric pH titrations of 0.4 mM catechol solution in the presence of AI at 282 nm.
wavelengths studied yielded similar results. The solid lines in Figure 6 were calculated from eq 23 by using log KE,, of 16.22 and an average EC value of 3490 calculated from A,, values at 282 nm (Table 11) and by assuming PA~L= PA,. Due to the different activity coefficients in eq 4 and 5, the assumption that /3& = PAIdoes not necessarily mean the activity-based hydrolysis constants for free (Kl - K3) and complexed A1 (Klaare equivalent. By setting PAIL= PAl, the hydrolysis constants were assumed to be similar. The experimental data fit the calculated functions well. The broken line in Figure 6 was calculated by assuming the A13+complexed by catechol did not hydrolyze (e.g., pW = 1). Calculated absorbance for PAU, = 1was very close to the calculated values for PAIL = PAI a t the lower
pH range with 3 and 30 mM Al; therefore, the functions are not shown. The greater experimental absorbance values obtained at 0.3 mM as opposed to calculated values assuming pAIL= 1indicated the A13+complexed by catechol was most likely hydrolyzed and the assumption that PAL = PAIin the calculation of log K8, (Table 111) was appropriate. There is disagreement in the literature as to whether complexed A13+ is hydrolyzed to the same extent as free A13+. Many chemists have suggested that complexed A13+ and free A13+will hydrolyze similarly (29-34). Moketaitis and Martell (17),however, dispute the occurrence of A1 hydrolysis in an organic complex. The UV spectrophotometric data presented in Figure 6 indicate the hydrolysis constants for complexed AP+ are similar to free AP'. The extent of complexed-metal hydrolysis is needed information for accurate speciation of metals in aquatic environments. Therefore, careful consideration of this phenomenon is warranted. Complexed-metal hydrolysis may depend on the degree to which the organic ligand reduces the valence of the metal by donation of electrons into metal orbitals. A reduced effective valence lowers the tendency of coordinated water to be hydrolyzed by the weaker electrical field at the metal. Partially vacant 3d orbitals and empty 4s and 4p orbitals of Fe3+ accept electron density from ligands, reducing the effective charge of the metal. Therefore, the degree of Fe hydrolysis, while complexed to an organic ligand, is reduced (17). In contrast, the noble gas configuration of A13+makes the vacant d orbitals less accessible, resulting in Al-ligand bonds that are less covalent than Fe-ligand bonds. The 27Alhyperfine splitting of the ESR spectra of Al-o-semiquinone radicals complexed to AP+ (35) indicates only about 2% of the electron spin population of o-semiquinone is in the s orbital of Al. Also, ultraviolet, infrared, and nuclear magnetic resonance spectroscopic analyses of the Al-catechol complex show greater electron density in the ring system compared to catechol (36)and are consistent with the lack of electron transfer from the ligand orbitals to the metal. Comparison of hydrolysis constants for A13+ and Fe3+ further suggests a basic difference between the effective electrical fields of the two metals in complexes. Single-step hydrolysis constants of AP+ and A10H2+are 10-4.97and 10-4.93,respectively (37). Reduction of the net charge of the metal by coordination to OH- does not significantly reduce the tendency of the metal to hydrolyze. Constants for Fe3+and FeOHZt of 10-2.19and respectively, are an indication that coordination to a single OH- reduces the tendency of the metal to further hydrolyze. A ternary Al-hydroxy-catechol complex [Al(OH) in Figure 11 is chemically reasonable. The occurrence of a 1:2 A1:catechol complex with one catechol bonded to A1 by two oxygens and the other catechol bonded by one oxygen has been postulated from potentiometric titration data (17). Hydrolysis of A1 in a 1:l A1:catechol complex could provide an alternative explanation for the data. Aluminum polymers, as observed in A1 solutions by 27Al nuclear magnetic resonance (38),did not appear to form to any significant extent under the conditions where titration data were recorded. If a side reaction of A1 polymer formation was occurring, the observed absorbance values would have been less than the calculated values in Figures 4 and 6. The observed absence of A1 polymers was most likely due to the short reaction time of approximately 2-6 min after each titrant addition. At high A1 concentration and high pH, a slow decrease in absorbance was observed (data not shown), which could have been due to a slow formation of A1 polymers resulting in a decreasing con-
centration of the newly formed Al-catechol complex.
Conclusions Aluminum complexation by the o-hydroxy group of catechol a t low pH has important implications to A1 chemistry in the environment. Considerable attention has been given to the potential metal-chelating role of phthalate- and salicylate-type groups (5,39-41), whereas little is known on the chelating role of o-hydroxy groups of organic matter. The method of UV spectrophotometry used in this study confirmed the stability constant of A1 complexation by catechol as determined by potentiometric titrations and also verified the existence of ternary Alhydroxy-catechol. The chelation of A13+by catechol indicates that o-hydroxy groups, if prevalent in soils and surface waters, may be an important metal-binding group. The equations and methodology presented in the current study can be adapted for the determination of association constants involving other simple aromatic organic ligands and metals found in soil solutions and surface waters. An increased compilation of currently unavailable association constants can aid the environmental chemist in speciating A1 and other metals in aqueous systems. Registry No. Al, 7429-90-5; catechol, 120-80-9. Literature Cited Driscoll, C. T.; Baker, J. P.; Bisogni, J. J., Jr.; Schofield, C. L. Nature 1980, 284, 161-164. Ulrich, B.; Mayer, R.; Khanna, P. K. Soil Sci. 1980, 130, 193-199. Alva, A. K.; Edwards, D. G.; Asher, C. J.; Blarney, F. P. C. Soil Sci. Soc. Am. J . 1986, 50, 133-137. Young, S. D.; Bache, B. W. J. Soil Sci. 1985, 36, 261-269. Hue, N. V.; Craddock, G. R.; Adams, F. Soil Sci. Sac. Am. J . 1986, 50, 28-34. Stevenson, F. J. Humus Chemistry: Genesis,Composition, Reactions;John Wiley and Sons: New York, 1982; Chapter 8. Crossor, M. L.; Allen, H. E. Soil Sci. 1977, 123, 268-269. MacCarthy, P.; Mark, H. B., Jr. Soil Sci. SOC.A n . J. 1976, 40, 267-276. Gould, R. K.; Vosburgh, W. C. J. Am. Chem. SOC.1942,64, 1630-1634. Vosburgh, W. C.; Gould, R. K. J. Am. Chem. SOC.1941,63, 437-442. Woldbye, F. Acta Chem. Scand. 1955, 9, 299-309. Baiocchi, C.; Menstasti, E. Ann. Chim. 1979,69, 189-200. Perlmutter-Hayman, B.; Tapuhi, E. Inorg. Chem. 1977,16, 2742-2745. Secco, F.; Venturini, M. Inorg. Chem. 1975,14,1978-1981. Harborne, J. B. In Methods in Polyphenol Chemistry; Pridham, J. B., Ed.; The MacMillan Co.: New York, 1964; Chapter 2. Havelkova, L.; Bartusek, M. Collect. Czech. Chem. Commun. 1969,34,3722-3731. Motekaitis, R. J.; Martell, A. E. Inorg. Chem. 1984, 23, 18-23. Ohman, L.; Sjoberg, S. Polyhedron 1983, 2, 1329-1335. Martell, A. E.; Smith, R. M. Critical Stability Constants. Volume 3 Other Organic Ligands; Plenum Press: New York, 1977. Lindsay, W. L. Chemical Equilibria in Soils; John Wiley and Sons: New York, 1979; Chapter 3. Beck, M. T. Chemistry of Complex Equilibria; Van Nostrand Reinhold Co.: London, 1970; Chapter 5. Rossotti, F. J. C.; Rossotti, H. The Determination of Stability Constants and Other Equilibria Constants in Solution; McGraw-Hill: New York, 1961; Chapter 13. Gilbert, J. W. J . Phys. Chem. 1959, 63, 1788-1789. Statistical Consultants, Inc. PCNONLIN and NONLIN84: Software for the Statistical Analysis of Nonlinear Models. Am. Stat. 1986, 40, 52. Environ. Sci. Technol., Vol. 23, No. 3, 1989
355
Environ. Sci. Technol. 1989,23,356-362
Doub, L.; Vandenbelt, J. M. J. Am. Chem. SOC.1949, 71, 2414-2420.
Ishimitsu, T.; Hirose, S.; Sakurai, H. Talanta 1977, 24, 555-560.
Jaffe, H. H.; Orchin, M. Theory and Application of U1traviolet Spectroscopy;John Wiley and Sons: New York, 1962; Chapter 12. Matsen, F. A. J. Am. Chem. SOC.1950, 62, 5243-5248. Aikens, D. A.; Bahbah, F. Anal. Chem. 1967,39,646-649. Bhat, T. R.; Radha, R. D.; Shankar,J. Indian J.Chem. 1967, 5, 324-327.
Dubey, S. N.; Mehrotra, R. C. Indian J. Chem. 1967, 5, 327-332.
Ohman, L.; Sjoberg, S. Acta Chem. Scand. 1981, A35, 201-212.
Patnaik, R. K.; Pani, S. J. Indian Chem. Soc. 1961, 38, 379-384.
(34) Rajan, K. S.; Mainer, S.; Rajan, N. L.; Davis, J. M. J. Inorg. Biochem. 1981,14, 339-350. (35) Eaton, D. R. Inorg. Chem. 1964, 3, 1268-1271. (36) McBride, M. B.; Sikora, F. J.; Wesselink, L. G. Soil Sci. Soc. Am. J . 1988,52, 985-993. (37) Baes, C. F., Jr.; Mesner, R. E. The Hydrolysis of Cations; Wiley-Interscience: New York, 1976. (38) Bertsch, P. M.; Thomas, G. W.; Barnishel, R. I. Soil Sci. Soc. Am. J. 1986,50,825-830. (39) Gamble, D. S.; Schnitzer, M.; Hoffman, I. Can. J. Chem. 1970,48, 3197-3204. (40) Schnitzer, M. Soil Sci. SOC. Am. Proc. 1969, 33, 75-81. (41) Schnitzer, M.; Skinner, S. I. M. Soil Sci. 1965,99,278-284.
Received for review December 30, 1987. Accepted September 15,1988. The project was supported by a grant from the National Science Foundation, Grant EAR-8512226.
A Novel Description of the Acid-Base Properties of an Aquatic Fulvic Acid James H. Ephralm,* Hans Borgn, Catharlna Pettersson, Irina Arsenie, and Bert Allard
Water and Environmental Studies, Linkoping University, $581 83 Linkoping, Sweden The potentiometric properties of an aquatic fulvic acid, Bersbo FA, have been analyzed by the physicochemical approach developed by Marinsky and co-workers. The complicating factors affecting the potentiometric behavior of the fulvic acid molecule have been identified as the heterogeneity of the FA molecule and formation of a separate microphase by the fulvic acid molecule in solution. The insensitivity of potentiometric behavior to ionic strength in the FA molecule has been attributed to a selective exclusion of counterions from the hydrophobic fulvic acid molecule [i.e., preference for H+ (pH) over M+ (pM)]. A combination of carefully designed experiments in nonaqueous medium and protonation enhancement titrations in the presence of heavy metals has facilitated a meaningful assignment of chelating moieties in the fulvic acid molecule.
Introduction The acidic properties of natural organic acids (humic and fulvic acids) are of primary interest to geochemists, soil scientists, and environmental chemists, but despite years of research, these properties have not been adequately described. Interpretation of the acid-base properties of humic materials has been effected by different models (1-8) over the past 30 years, and even though considerable progress has been made during this period, there is still a lot more that needs to be known about these substances. Humic and fulvic acids have been known to be heterogeneous, consisting of numerous oxygen-containing functional groups and fractions with different molecular weights (9). In a recent review, Buffle et al. (10) outlined the models that have been employed in describing ion binding by humic substances and concluded that these organic acids behave as “ideal” heterogeneous complexants, Le., have an infinite number of sites possessing globally remarkably reproducible properties. Additionally, a review by Perdue (11) concluded correctly that humic substances contain a highly complex mixture of nonidentical functional groups with pK, values that span the entire possible range determined by the leveling effect of water on the strengths of the acids. The complexity of humic substances cannot be overemphasized. However, the description of the observed acid356
Environ. Sci. Technol., Vol. 23, No. 3, 1989
base properties by purely statistical continuum models is, we believe, an overestimation of the problem. The vectors leading to the formation of a stable humic substance product will tend to limit the site distribution to narrower boundaries than asserted by the continuum approach. Little chemical significance can be attributed to an infinite number of sites in a stable humic substance. We are of the contention that even though a spectrum of acidic funtionalities exists in these humic substances, the observed acidity may very well be adequately described by consideration of only the most predominant acidic moieties and by taking into account the complicating factors perturbing the system. In this paper, the acid-base properties of an aquatic fulvic acid are described by employing the physicochemical approach developed by Marinsky et al. (12-141, where the complicating factors have been attributed to the heterogeneity of the molecule and the formation of a separate microphase of the fulvic acid molecule in solution. We show in this paper that with carefully designed experiments the heterogeneity in humic substances may be described without resort to extensive mathematical manipulations. Theoretical Background
The potentiometric titration of a polyacid may be described by relating the pH of the solution and the degree of neutralization, a , as follows:
where dGe/az is the work of removal of one proton from a molecule, z times ionized, and Kintis the intrinsic dissociation constant. Equation 1may be rearranged to yield the following expression: pKaPp- pKint = ApK = (0.434/kT)dGe/dZ (2) where pK,, is defined as pH - log [ a / ( l- a ) ] . In typical descriptions (15-17) of synthetic polyacids, a plot of pKapp versus CY gave a quantitative measure of nonideality behavior by measuring any deviation from a straight line of zero slope. In such descriptions, the intrinsic dissociation constant was obtained by extrapolating the pKappversus a curve to a = 0 and reading off the pKappvalue to be the
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