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
0013-936X/89/0923-0356$01.50/0
@ 1989 American Chemical Society
negative logarithm of the intrinsic dissociation constant. The usefulness of pKappversus a plots for humic substances is limited to obtaining insight into ionic strength effects on their potentiometric properties in an attempt to describe the polymeric nature of these substances. Plots of pH - log [a/(l - a ) ]versus a a t different ionic strengths for a monomeric acid yield essentially a unique line with a slope of zero. However, if the molecule exists as a linear polymeric assembly, ionic strength has considerable effect on the slope of the pH - log [ a /(1- a ) ]versus a curve. The pH - log [a/(l - a ) ]value a t higher a deviates considerably from that at lower a for a given ionic strength, with such deviation being largest in the lowest ionic strength. As the value of a is made smaller, the ordinate value of the pH - log [ a / ( l - a ) ]versus a curve is also lowered, with the difference between the experimental points at the different ionic strengths becoming less and less until they coincide on the ordinate axis where a = 0. In the case of a three-dimensional aggregate or gel, the potentiometric curves, pH - log [a/(l - a ) ]versus a, have been observed to be parallel to each other, the uppermost curve corresponding to the lowest ionic strength. Theoretically, the linear displacement of the curves should be proportional to the concentration ratio of the separate ionic media (15-18). Use of the Donnan concept and the Marcus electroneutral cell model has permitted a comprehensive description of the mechanism operative in the gel situation (12). For a cross-linked polyelectrolyte gel, the following relationship is found to hold: pKapp = PKint
+ (PM - PM)
(3)
where pM and pM are the negative logarithms of the activity of the neutral salt cation in the solution and gel phase, respectively, and pKappand pKintcorrespond to the negative logarithms of apparent and intrinsic acidity constants. The above discussions therefore point out that plots of pH - log [ a / ( l - a ) ] versus a for humic substances a t different ionic strengths may be employed to obtain information with respect to the physical state of these substances in solution, i.e., simple monomeric, linear polymeric, or three-dimensional aggregates. Recently, Marinsky and co-workers (13)connected the behavior of the fulvic acid molecule in solution to a hydrophobic, salt impermeable, second phase entity by noting the similarity of its potentiometric characteristics to those reported for the high molecular weight polyelectrolytes, poly(ethy1eneimine) and poly(vinylamine),which have been positively associated with these properties. Experimental Section (A) Chemicals and Their Preparation for Use. The aquatic fulvic acid, Bersbo FA, was extracted by employing a combination of methods developed by Paxeus (19) and Malcom (20). The preconcentration of humic material on diethylaminoethyl (DEAE) cellulose was followed by purification on XAD amberlite resin (20). The natural water was filtered through a diethylaminoethylcellulose column in the H+-form. The brown-pigmented material absorbed on the DEAE-cellulose was then desorbed with 0.30 M NaOH. The humic substance extract was then acidified with concentrated hydrochloric acid. The resulting mixture was separated into two portions, i.e., soluble (fulvic acid) and insoluble (humic acid) fractions. The soluble portion was adsorbed on XAD-8, desalted, and later desorbed with 0.10 M NaOH. Following cation exchange (H+-form), the fulvic acid extract was lyophilized (21).
Sodium perchlorate solutions were prepared from analytical reagent (Merck). Tetrabutylammonium hydroxide and tetraethylammonium chloride were obtained from Aldrich and were used without further purification. For use in the nonaqueous titrations, the tetrabutylammonium hydroxide was dissolved in 2-propanol. The solution wgs blanketed with N2 and stored in a refrigerator. Prior to use, the TBAH solution was standardized against benzoic acid. CuC12.2H20and LaC1,.7H20 (analytical reagents) were used without further purification. Sodium hydroxide was prepared from analytical concentrate and stored in polyethylene bottles in a desiccator with ascarite underneath. Prior to use, the sodium hydroxide solution was standardized against potassium hydrogen phthalate. Distilled deionized water was used in preparing all solutions. (B) Procedure. The acid-base titrations of the Bersbo FA in aqueous medium were performed in three different ionic strengths (0,001, 0.010, and 0.100 M) with two different initial FA concentrations (7.85 X and 1.57 X g of FA/50 mL of solution) with a Radiometer autoburet, Model ABU 80, a Radiometer research pH meter, pHM84, in conjunction with a combination pH glass electrode, GK2401C. The combination pH electrode was calibrated with buffer solutions (9.18, 7.01, and 4.01 at 25 “C) and then against standard HCIOl solutions at the ionic strength of study to ensure accurate response. The titrations were performed in a thermostated beaker a t 25 OC. In the course of a typical titration, a stream of N2 gas was blown over the reaction vessel to ensure a C02-free atmosphere. Nonaqueous titrations were performed with the Bersbo FA dissolved in dimethylformamide, DMF, with phydroxybenzoic acid, pHBA, as an internal reference. Tetrabutylammonium hydroxide, TBAH, dissolved in 2-propanol was the alkali employed in these titrations. A glass electrode was used in conjunction with a calomel electrode whose internal solution of 4.0 M potassium chloride had been replaced by 1.0 M tetraethylammonium chloride solution. Such a modification of the calomel electrode was necessary to prevent diffussion of K+ through the electrode frit into the nonaqueous medium. When not in use, the glass electrode was stored in 0.10 M hydrochloric acid and the calomel electrode was stored in 0.10 M tetraethylammonium chloride solution. A “calibration” curve was obtained for pHBA and then a mixture of pHBA and FA was titrated against standard TBAH. The two inflection points in both situations permitted computation of COOH and OH functionalities contributed by the FA. Titrations of FA in 0.100 M NaClO., with increasing amounts of heavy metal ions, i.e., Cu(I1) and La(III), separately were performed by using a combination pH electrode and determining the extra acidity observable in the presence of such heavy metal ions due to chelation with release of protons. Normal acid-base titration precautions were employed in the series of experiments. Treatment of Data Acid-Base. The degree of neutralization, a , of the FA at each point of alkali addition was computed as follows: degree of neutralization
O(
= A-/HA,
(4)
where HA, is the total titrable acid in aqueous medium at the ionic strength of study determined by a modified Gran relationship ( 5 ) . A- = (mbub+ hV,/y+ - k,/hJ
(5)
where mb is the molarity of base with corresponding volEnviron. Sei. Technol., Vol. 23, No. 3, 1989
357
Table I Elemental Analysis of Bersbo Fulvic Acidn
element
%
C
H
48.83 3.62
N 0
42.25
1.12
element
7 1
1
%
S
0.93 0.19
resid
1.73
c1
"Drying loss, 8.19%; 60 O C ; 1 mbar.
i
x 0.010 M NaC104
I
0
0.100 M NaC104
z i . , , , , , , , , , , , , , , , , , , , , , , , 2
3
4
5
6
7
8
PH Figure 2. Relationship between pK.,, and pH for the Bersbo FA at
different bulk electrolyte concentrations (0.100M, 0.010M and 0.001M NaCIO,). o 0.001 M NaC104 NaC104 NaC104
3 0.1
0.3
m 1.1
0.5 0.7 0.9 Degree of Neutralization
Figure 1. Potentiometricbehavior of Bersbo FA as a function of bulk electrolyte concentration (0.100M, O.OlQM, and 0.001M NaCIO,).
ume, ob, while V , is the total volume of solution and y+ is the single ion activity coefficient for hydrogen at the ionic strength of study obtainable from Kielland (22),and k, is the water dissociation constant. The hydrogen ion activity from pH measurements is denoted by h. The pKa at each point of alkali addition was therefore computed &y using the Henderson-Hasselbalch equation as follows:
,
Nonaqueous Titrations. If the inflection points in the titration of an aliquot of a pHBA solution for the COOH and OH functionalities are represented by a and b, respectively, and in the presence of FA with the same quantity of pHBA, if the inflection points for the two acidic functionalities are c and d , respectively, then the contributions of COOH and OH functionalities by the FA may be computed as follows: COOH contribution by FA = (c - u)NTBAH
OH contribution by FA = (d - c - b
(7)
+ u)NTBAH (8)
where NTBm is the normality of the tetrabutylammonium hydroxide solution.
Results Elemental analyses of the Bersbo FA are shown in Table I. The proportions of the elements H, C, 0, and N in this aquatic FA are not significantly different from results obtained from other FA samples (23). The titratable acid capacity for the fulvic acid sample in aqueous medium was determined to be 4.65 f 0.15 mequiv/g of fulvic acid. Figure 1 shows plots of pH - log [ a / ( l - a ) ]versus a for the Bersbo FA at different ionic strengths (0.100, 0.010, and 0.001 M NaC10,). Each curve at a given bulk electrolyte concentration represents the two different FA concentrations (157 and 314 ppm) employed in the study. The convergence of the pH - log [ a / 1 - a ) ]versus a for 358
Environ. Sci. Technol., Vol. 23, NO. 3, 1989
the different ionic strengths as a approaches zero is apparent, suggesting therefore that the Bersbo FA molecule in solution exhibits polyelectrolyte behavior. The divergence of the pKa,, versus a curves for the different ionic strengths, however, is smaller than expected for a typical high molecular weight polyelectrolyte. At a greater than 0.60, the difference in pKa,, values from 0.10 to 0.01 M NaClO, is 0.4 pK unit. The value obtained from 0.01 to 0.001 M NaClO, is 0.35 pK unit. This apparent insensitivity of the potentiometric behavior to ionic strength has been explained by examining carefully the potentiometric behaviors of high molecular weight polyelectrolytes with salt-impermeable characteristics (24). The potentiometric studies of such high molecular weight linear polyelectrolytes, poly(ethy1eneimine) and poly(vinylamine), show relatively insensitive response of pH to sodium chloride concentrations encountered. This behavior has been attributed to the salt impermeability of the polyelectrolyte molecules resulting in a selective exclusion of counterions from the surface of the hydrophobic polyelectrolyte molecule. In carefully examining the potentiometric properties of such high molecular weight linear polyelectrolytes, it was observed that a plot of pKap versus pH for the different ionic strengths studied yieldes a unique line (12-14). From such an observation, it was deduced that to test for hydrophobicity of a polymeric molecule in solution, it is reasonable to plot pK,,, versus pH for different bulk electrolyte concentrations. Figure 2 is a plot of pKapp versus pH for the Bersbo fulvic acid for experiments performed at three different ionic strengths (0.100, 0.010, and 0.001 M NaC10,). The figure shows that the curves diverge at higher pH values, while at pH of 3.0 and below, there is a convergence. These results suggest that the Bersbo fulvic acid molecule is essentially hydrophobic at lower pH values, i.e., pH below 3, while at higher pH values the hydrophobicity is lost. This conclusion is consistent with the observation during isolation of humic materials where the solution is acidified to pH of 2 before its adsorption on XAD amberlite resins. Background salt concentration, I = 1.00 M NaClO,, was determined as the critical, limiting bulk electrolyte concentration beyond which the observed apparent dissociation constant pKappwas no longer a sensitive function of ionic strength. The computed pKappat this ionic strength was, therefore, considered to be due to the heterogeneity of the fulvic acid molecule. Figure 3 shows the plots for the pKappversus a for the Bersbo FA molecule including
o 1,000 M NaC104 x 0.010 M NaC104 o 0.100 M NaC104
6 -
g
d
$ 5 @
I
i+
c
a
4
3
I D e g r e e of Neutralization
Figure 3. pK,,, versus degree of neutralization, a,including data at 1.OM NaCIO,. 2.0'I 1.8 f Y
a
r-
2
C u 2 + / a n l y 16040 m e q i
---
C
Cu2'/FA
D
Cu2+IFA
E
Cu2+lFA
F
["?+/FA
T
T h e o r e t i c a l Curve
1
0
1.6f
A
i-
084
127
c
211
I
I
05
0.42
10
15
m i (OH')
1.4:
Figure 5. Potentiometric titrations of Bersbo FA in the presence of increasing quantities of Cu(I1) ions.
aJ
; 1.2L
," 1.0.s
a 0.84
%
vr
+ O.1OOM NaC104 x 0.010M NaC104 0.001M NaC104
0.4
0.2
1 1
o.o&-n&
0
0
0
b
i
z
i
0
0
0
0
0
0
0
0
0
1
5
6
i
a
s
b
0 0 0 0 0 0 0 0 Degree of Neutralization. Alpha
0
4
1 9
r n 7
Flgure 4. Separate phase effect, ApK, as a function of degree of neutralization at different bulk electrolyte concentrations (O.lOOM, 0.010M and 0.001M NaCIO,). data at I = 1.00 M. The variation of separate phase effect (ApK) a t the three different ionic strengths (I = 0.100, 0,010, and 0.001 M) as a function of degree of neutralization, CY,was determined by employing the pKappversus a curve at I = 1.00 M as the reference line. Figure 4 presents this variation for the Bersbo FA a t the three ionic strengths. The increase in separate phase effects with decrease in ionic strength is observable after the fulvic acid molecule starts to lose its hydrophobicity (beyond 30% ionization). The enhanced separate phase effects in dilute neutral salt medium are believed to indicate the promotion of head-to-tail interaction between FA molecules upon dilution of background salt concentration (14). Extra Proton Release in the Presence of Heavy Metals. By titrating the FA in aqueous medium in the presence of increasing amounts of Cu(I1) and La(II1) and comparing the measurable acidity to that obtained in aqueous medium of FA alone, additional acidic moieties that are too weak to be detected in normal acid-base titrations can be determined. Figure 5 presents the titration curves of the Bersbo FA molecule in the presence of increasing amounts of Cu(I1) ions. Curve A in the figure represents the titration of Cu2+ only. Between pH 6.4 and 6.8 the system is significantly buffered due to tbe formation of hydroxy complexes of Cu2+precipitating in solution. Curve B is a typical titration curve of FA. Curves C-F represent the situation
5
7
'1
-I. .' 0
,.'
'
02
0.4
06
08
m l (OH-)
Figure 6. Potentiometric titrations of Bersbo FA in the presence of increasing quantities of La(II1) ions. where the ratios of Cu/FA, on a milliequivalent basis, are 0.42,0.84,1.27,and 2.11, respectively. It is observed from the figure that the FA molecule completely solubilizes the Cu2+ions as long as the ratio of Cu/FA is below unity. However, in curve E where the ratio of Cu/FA is 1.27, a buffered shoulder appears around pH 6.0. This buffered shoulder is due to the presence in solution of excess Cu2+ ions reacting with the OH- ions being added to form hydroxy complexes. The more excess Cu2+ions left in solution, the broader the buffered shoulder, as shown in curve F. Curve T is the theoretical shape of the titration curve in the absence of excess Cu2+in the solution, i.e., the shape of the titration curve if no excess Cu2+existed in solution. In Figure 6, the curves obtained when increasing amounts of La3+were added to fulvic acid are presented. When La3+ is titrated alone (curve A), a buffered region Environ. Sci. Technol., Vol. 23, No. 3, 1989 359
Table 11. Functional Group Analysis of Bersbo FA Using p -Hydroxybenzoic Acid as Internal Reference
PHBA, g X lo2 3.853 3.853 3.853 3.853 3.853 3.853
FA, g 0 0 0
3.32 X 6.64 X 1.33 X
"TBAH = 5.316 X bution.
eauiv Dts, mL/TBAH" 1st 2nd 5.43 5.43 5.43 5.73 6.03 6.68
10.86 10.86 10.86 11.25 11.63 12.45 av
caaacitv. mequ&/g of FA COOH OH total 7.24 7.24 7.24 4.80b 4.72* 4.84b 4.78
7.24 7.24 7.24 1.3gb 1.31b 1.34b 1.35
6.19 6.03 6.18 6.13
4
/
i
pHBA+ FA
>
- 400
I
M. bAfter subtraction of pHBA contri-
is observed at around pH 9.10-9.25. This buffered region is attributable to the formation of hydroxy complexes of La3+in solution. In the presence of FA, as long as the ratio of La3+/FA remains less than unity, no buffered region appears. This observation is attributable to complete solubilization of the La3+ ions by the FA due to complexation. When excess La3+ is present, Le., La3+/FA is greater than unity, the buffered region reappears around pH 8.5. Once again, curve T is the theoretical shape of the titration curve if no excess La3+ existed in solution. An extra acidity of 65-70% was determined for the Bersbo FA with increasing amounts of Cu2+ ions, while 25-30 % extra acidity was determined with increasing amounts of La3+. The extra acidity determined in the presence of Cu2+was attributed to the presence of acidic configurations like a salicylic acid like moiety, dihydroxy-like moiety (catechol-type), and acetoacetonatelike moiety in the enol form. In the case of La3+,the extra acidity has been attributed to the presence of a dihydroxy moiety or acetylacetonate-like moieties in the FA molecule. Similar thermodynamic arguments ( I I ) have been employed to arrive at these conclusions. The use of the expressions salicylic acid like, dihydroxy-like, etc., is for configurational description only. No aromaticity or otherwise is implied. Nonaqueous Titrations. Results of the nonaqueous titrations have been presented in Figure 7 and Table 11. These results show that the Bersbo FA molecule has a carboxylic capacity of 4.78 f 0.06 mequiv/g of FA and an OH capacity of 1.35 f 0.04 mequiv/g of FA. The OH capacity is considered to represent "very weakly acidic functionalities" comprised of both phenolic and acidic alcohols with pKs (in aqueous medium) greater than 5 but less than 11. It must be pointed out here that no pK assignments were made to the COOH and OH moieties determined in this phase of the research. Functional Group Heterogeneity. By use of the potentiometric behavior of the Bersbo FA at I = 1.00 M NaC104 as the reference bulk electrolyte concentration where charged separate phase effects are essentially nonexistent, it has been possible to show the two-phase effect (ApK) at the three different ionic strengths (0.10, 0.010, and 0.001 M) (Figure 4). This separate phase effect has facilitated the assignment of acid site heterogeneity in the Bersbo FA. From nonaqueous titrations, approximately 25-30% of the acidic capacity was identified with a weakly acidic functionality of the catechol-type configuration. This assignment was corroborated by the 25-30% extra acidity observed in the titrations with increasing amounts of La3+ and also attributed to a dihydroxy-type configuration. An extra acidity (-60-70%) was detected due to chelation by Cu(I1) with subsequent removal of protons. This 360
- 500
Envlron. Sci. Technol., Vol. 23, No. 3, 1989
- 200 j 0
I
I
I
I
I
I
I
2
4
6
8
10
12
14
I
m i (005136 m l T . B A H
Flgure 7. Nonaqueous titrations of Bersbo FA using an internal reference standard.
extra acidity has been attributed to a combination of protons released due to chelation by two acidic moieties, Le., salicylic-like and catechol-like. Thus, a difference in the extra acidity due to Cu(I1) and that due to La(II1) represents the proportion of the salicylic-like moiety. Therefore of the extra acidity observed in titrations with increasing amounts of Cu2+,30-45 % has been identified with a salicylic acid like moiety while 25-30% has been identified with a catechol-like moiety. Results from the nonaqueous titrations, protonation enhancement experiments in conjunction with the electrostatic deviation (ApK) term has permitted heterogeneity assignment of the Bersbo FA molecule. The experimental degree of neutralization at each pH at the critical, limiting bulk electrolyte concentration (1.00 M NaC104) has been fitted to a model-computed degree of neutralization using five different acid sites with different dissociation constants and abundances. The computed overall degree of neutralization was obtained with the following expression (25):
where ai is the degree of neutralization for the ith acidic site and Ai is the corresponding abundance of the acidic site. The degree of neutralization at the ith site, ai,was computed a t each pH by using the following equation: ai= 1/[1 + ~o(PH+APK-PK~] (10) where pKi is the intrinsic pK of the ith site and ApK is the separate phase effect obtainable from Figure 4 at the given experimental degree of neutralization. Figure 8 shows the plot of the experimental overall degree of neutralization, and the computed overall degree of neutralization, versus pH for the potentiometric titration of Bersbo FA in aqueous medium. The five acid sites that best described the observed potentiometric behavior, presented graphically in Figure 9, have the following pK, and abundances respectively: 1.7 and 20%;3.3 and 25%; 5.0 and 30%;6.5 and 20%; 7.0 and 5%. The acid sites with pKs of 1.7 and 0.70 of 5.0 were considered as two carboxylic acids ortho to each other
of acidic moieties into weak carboxylic acidic functionalities and very weak acidic functionalities which may be considered as a combination of phenolic, alcoholic, or very weak carboxylic acids. The degree of proton release in potentiometric titrations with increasing amounts of heavy metal ions (Cu2+and La3+)has given further insight into the chelating capacity of the FA molecule and has thus facilitated a meaningful assignment of chelating moieties in the fulvic acid molecule.
’‘7 2
0.8-
N I*
5
2
c
0.7-
0.6-
x alpha-expt. o alpha-cornpt. 3
4
5
6
7
PH Flgure 8. Comparison of experimental overall degree of neutralization and computed overall degree of neutralization.
[II] pK=1.70 25.0%
\
Environmental Implications A thorough understanding of the nature and chemical reactivity of humic materials in the environment is essential in order to determine the role played by these materials in the transport of metal ions and inorganic colloids in surface and ground waters. The approach outlined in this paper offers an opportunity to simplify the problem of ion-humate interactions to a stage where already existent programs for speciation determinations of inorganic substances (26) could be applied to these natural organic acids. Registry No. Cu, 7440-50-8; La, 7439-91-0.
Literature Cited .O%
pK=3.30
Sposito, G.; Holtzclaw, K. M. Soil Sci. SOC.M . J . 1977,41,
pK=5.0
Sposito, G.; Holtzclaw, K. M.; Kerch, D. A. Soil Sci. SOC. Am. J . 1977,41, 1119. Eberle, S. H.; Feurstein, W. Naturwissemchaften 1979,66,
330.
5.0%
572.
30.0%
pK=6.50 26.0%
11
1978, 56, 1196. pK=7.0
Figure S. Acidic Heterogeneity of the Bersbo FA.
while 0.3 of the latter (pK = 5 ) exists alone. The acid site with pK of 3.3 was assigned to a carboxylic acid, ortho to an OH with a much higher pK, thus resembling a salicylic acid arrangement. The acid site with a pK of 6.5 was considered as an acidic alcohol next to a more weaker hydroxyl group, thus mimicking a dihydroxyl arrangement, while the acid site with a pK of 7.0 was considered as a hydroxyl next to a carbonyl group, thus mimicking an acetylacetone arrangement. The above configurational assignments are justifiable by consideration of the effects of neighboring group interaction on the numerical values of weak acid constants, pK,’s (11).
Summary and Conclusion The potentiometric properties of an aquatic fulvic acid, Bersbo FA, have been analyzed by using the physicochemical approach developed by Marinsky and co-workers (12). 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+
(PWI.
Gamble, D. S. Can. J . Chem. 1970, 48, 2662. Gamble, D. S. Can. J. Chem. 1972,50, 2680. Burch, R. D.; Langford, C. H.; Gamble, D. S. Can. J. Chem.
Insight into the pK spectrum existent in the FA sample has been attained from nonaqueous titrations employing an internal standard, and this has facilitated the separation
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Received for review M a y 17, 1988. Accepted October 19, 1988. W e are grateful to the Swedish Nuclear Fuel and Waste Management Company for financial assistance and to Lisbeth Samuelsson for graphical assistance.
Acidification of Adirondack Lakes Clyde E. Asbury” Center for Energy and Environment Research, University of Puerto Rico, Rio Piedras, Puerto Rico 00936
Frank A. Vertucci Rocky Mountain Forest and Range Experiment Station, U.S.D.A. Forest Service, Ft. Collins, Colorado 80526
Mark D. Mattson Section of Ecology & Systematics, Corson Hall, Cornell University, Ithaca, New York 14853
Gene E. Likens Institute of Ecosystem Studies, The New York Botanical Garden, Millbrook, New York 12545
The acidification of lakes in the Adirondack Mountain region of New York was estimated directly by comparing data from historic (1929-1934) and modern (1975-1985) regional surveys of lake chemistry. We performed new analyses concerning the quality of the data, rejecting all historic pH data and many modern alkalinity values. When the historic data were corrected for a bias between titration procedures, we found a median loss of 50 pequiv/L alkalinity in 274 lakes with paired data. Eighty percent of the lakes showed a decline in alkalinity. The observed acidification was greatest in the lakes at high elevation and was of the same magnitude as the current precipitation acidity in the region.
Introduction The lakes of the Adirondack Mountain region of New York State have been repeatedly cited as showing the effects of anthropogenic acidification (1-5). Much of the evidence for acidification is indirect, however: e.g., loss of sport fish populations (2, 3), changes in aquatic plant and invertebrate communities (6-8), changes in sediment metal concentrations (9),and results from empirical models of lake acidification (10). The Adirondacks is one of the few areas where historic data are available to measure directly the acidification of lakes through time. While previous researchers have examined smaller data sets from this region, they reached mixed conclusions regarding the changes in pH and alkalinity ( 2 - 4 , I I ) . In this paper we examine the chemical evidence from historic and modern lake surveys to determine if significant lake acidification has occurred in the Adirondack Mountain region. Methods The New York State Conservation Department surveyed hundreds of Adirondack lakes during the summers of 1929-1934. Included in the studies were measurements of alkalinity, pH, and C02acidity (12). Extensive new data from the recent New York State Department of Environmental Conservation (DEC) survey (13) and the U.S. Environmental Protection Agency (EPA) survey (14) along 362 Environ. Sci. Technol., Vol. 23, No. 3, 1989
with existing DEC data (3, 15) were available to form a modern data set for the years 1975-1985. Historic and modern surface alkalinities for each lake were matched on the basis of a unique “pond number” system used by the surveys. In a few cases where pond numbers were not assigned, we used other information, including lake name and location to identify the lakes; all lakes identified ambiguously were removed from the data set. If redundant modern data were available, the most recent survey was used. Both the historic as well as the modern survey results were examined closely to detect any errors or inconsistencies in the data. Historic pH measurements are often unreliable because procedures used to measure pH employed various colorimetric indicator solutions that can alter the pH of the sample being measured (16),especially in the dilute waters of the Adirondacks. In addition, it is well-known that solution pH can vary due to COz exchange with the atmosphere (17). We examined these problems carefully and rejected the use of pH data in favor of a direct comparison of alkalinity values, which can be determined with great precision with procedures that are well documented, and alkalinity is conservative with respect to changes in C 0 2 (17). It has been reported (2,18) that the electrode used by the DEC in measuring both pH and alkalinity during the 1979 survey was malfunctioning. Examination of the laboratory notes kept by the DEC (19) confirmed that the Gran functions used to determine alkalinity were not linear, and the calculated values were biased low. We therefore eliminated 72 lakes from our data set that were included in the 1979 DEC survey. Moreover, 28 lakes that were known to have been treated with lime were not included in our analysis. Our final data set consisted of 274 lakes. In a pair-wise analysis of data such as these, it is important to remove any source of bias or systematic error between the historic and modern measurements. In this case, bias results from differences in the techniques used to measure alkalinities. The historic method employed titration to a fixed pH end point, determined by the
0013-936X/89/0923-0362$01.50/0
0 1989 American Chemical Society
