Environ. Sc/. Technol. 1883, 27, 1182-1 189
Calorimetric Acid-Base Titrations of Aquatic and Peat-Derived Fulvic and Humic Acids Mlchael L. Machesky
Illinois State Water Survey, 2204 Griffith Drive, Champaign, Illinois 61820
Natural organic matter influences the fate and transport of inorganic and organic contaminants. Humic substances constitute a substantial fraction of this organic material, and these substances contain appreciable concentrations of acidic functional groups. Most of these acidic functional groups are carboxylic and phenolic acids (1, 2)) and the degree to which these groups are ionized or protonated (i.e., ambient pH values) and their abundance and distribution significantly influence trace contaminanthumic substance interactions. Humic substances can bind significant quantities of trace metals, and most of this binding is believed to involve acidic functional groups (2). Increasing pH generally leads to increased trace metal binding until favorable electrostatic and steric factors are overcome by metal cation hydrolysis. Proton competition for acidic functional groups also decreases trace metal binding at lower pH values. Trace metal binding can also result in measurable proton release from humic substances to accommodate trace metal-humic substance bonds (3). Organic contaminant-humic substance interactions are also influenced by pH. Humic acids were found to bind more DDT as pH decreased from 9.2 to 6.2 and as Ca2+ increased (4). Similarly, various DDT and PCB congeners became more soluble in the presence of Suwannee River fulvic acid solutions as pH decreased from 8.5 to between 6 and 4 (5). These observations are consistent with the hypothesis that humic substances become more hydrophobic as acidic functional groups are protonated or complexed by metal cations; consequently, hydrophobic trace organic compound sorption increases. Therefore, a thorough understanding of acidic functional group types, concentrations, and distribution is necessary to completely understand
trace metal and trace organic contaminant binding by humic substances. Titration calorimetry is a valuable technique to investigate these acidic functional groups. With this method, temperature changes accompanying titrations are measured, and adding pH electrodes to the basic instrumentation allows simultaneous monitoring of pH. Reaction enthalpies result from these data, and a principal advantage of this technique for humic substance studies is that carboxylic and phenolic groups possess widely different protonation enthalpies (6). Carboxylic acid protonation enthalpies typically range from 6 to -5 kJ/mol (7), and those for simple phenols are -17 to -25 kJ/mol (8). Phenolic hydrogens can also be stabilized by H-bonds to adjacent oxygen-containing groups, which leads to more exothermic protonation enthalpies. For example, the protonation enthalpy for the phenolic hydrogen of salicylic acid is -36 kJ/mol, and that for the most acidic phenolic hydrogen of 1,2-dihydroxybenzene is -34.3 kJ/mol. In addition, reaction enthalpies are less sensitive to electrostatic effects than corresponding free energies. Electrostatic contributions to the ionization enthalpies of polymeric molecules are only about one-third of the corresponding electrostatic free-energy terms near room temperature accordingto available theories (9). Consequently, large enthalpy differences can often be ascribed to other factors such as site heterogeneity or molecular conformation and hydration state changes of the molecules and associated counterions. Titration calorimetry has been used to investigate the protonation behavior of synthetic polyelectrolytes, which in some respects are analogous to humic and fulvic acids. Conformational changes, from expanded to more compact shapes as acidic functional groups are protonated, typically result in exothermic enthalpy contributions because additional bonds form to stabilize the compact structure (10). Other studies have observed that counterion type can have a large influence on observed enthalpy values (11, 12). More endothermic protonation enthalpies are sometimes observed in the presence of larger, more hydrophobic counterions in pH regions where the polymer is appreciably ionized. Hydration-state changes of the polymer and associated counterions are believed to be responsible for these observations. Larger or more hydrophobic ions cannot approach ionized sites as closely as similarly charged smaller ions which are more effective at disrupting electrostricted water molecules near the ionized sites (12). Consequently, upon protonation, more electrostricted water molecules are released in the presence of larger, more hydrophobic counterions with the result that net protonation enthalpies are more endothermic. Moreover, counterions with hydrophobic groups (e.g., tetraalkylammonium cations) can also interact with hydrophobic regions of the molecule. When ionized sites are protonated, the endothermic contribution from the release of electrostricted water can be augmented by hydrophobic hydration of the counterion and polymer (13).
1182 Environ. Sci. Technoi., Vol. 27, No. 6, 1993
0013-936X/93/0927-1182$04.00/0
Titration calorimetry was used to determine the protonation and ionization enthalpies of peat-derived and aquatic humic and fulvic acids from pH 3.6 to 10.4 and for 0.001-0.1 M ionic strength [NaCl and tetraethylammonium chloride (TEAC1)I. Protonation and ionization enthalpies were not completely equivalent, and protonation enthalpies were concluded to more accurately reflect the enthalpies associated with proton binding and release. Below pH 7, protonation enthalpies were similar to simple carboxylic acids (6 to -5 kJ/mol), and no ionic strength effects were apparent. Above pH 9, enthalpies were similar to those of simple phenolic acids (-20 to -36 kJ/mol), and less exothermic values were observed at I = 0.1 M. For NaCI, this was attributed to disruption of H-bonds between phenolic hydrogens and adjacent oxygen-containinggroups by Na+. For TEACI, hydrophobic interactions may be responsible. Thus, protonation enthalpies as a function of pH as well as counterion type and concentration reveal information about the nature and location of acidic functional groups in humic substances.
Introduction
0 1993 American Chemical Society
Therefore, the type of counterion and the location of acidic functional groups in relation to more hydrophobic regions of the molecule can have a large influence on observed enthalpy values. Previous calorimetric studies of humic substances are relatively rare. The most detailed prior studies are those of Choppin and Kullberg (14), Perdue (6, 15), and Machesky (16). In all these studies, protonation enthalpies between pH 6 and 4 ranged from 6 to -6 kJ/mol, which is consistent with those expected for simple carboxylic acids. Between pH 10 and 8, however, Perdue observed endothermic protonation enthalpies of 1.6-32 kJ/mol, while Choppin and Kullberg and Machesky found exothermic protonation enthalpies ranging from -16 to -36 kJ/mol. The latter values are consistent with those expected for phenolic group protonation while the former values could not be adequately explained. Machesky (16) also concluded that protonation enthalpies near -32 kJ/ mol at pH 10 for Suwannee River fulvic acid were consistent with the existence of hydrogen bonds between phenolic hydrogens and adjacent oxygen-containinggroups. Structural models had previously suggested that H-bonds exist between phenolic hydrogens and adjacent ketone groups in this fulvic acid (17). The study by Machesky (16) addressed some of the experimental limitations of previous studies, in particular the need to conduct separate pH and calorimetric titrations (14) and large titration volumes (6,15). In addition, two fulvic acids isolated and purified using widely accepted techniques were studied, and the influence of ionic strength and titration direction on observed enthalpy values was documented. Significant hysteresis (>8kJ/mol) was observed between protonation and ionization enthalpy values at pH values 9. This was attributed to aggregation influencing ionization enthalpy values at low pH and to the small net residual heat values obtained for the ionization (base addition) experiments at high pH. Consequently, it was concluded that protonation enthalpy values more accurately reflected the enthalpies associated with proton binding and release throughout the entire pH 4-10 range investigated. Ionic strength (0.001-0.1 M in NaC1) had no significant effect on the observed protonation enthalpies below pH 9. Above pH 9, however,protonation enthalpies for a Brazilian peat fulvic acid were lower at higher ionic strengths while those for Suwannee River fulvic acid were similar at all ionic strengths. The present study extends that of Machesky (16) to include the influence of titration direction and ionic strength (0.001-0.1 M in NaC1) on the enthalpy values associated with the protonation of two humic acids. It was hypothesized that the type, quantity, and distribution of acidic functional groups in these humic acids might differ enough from the fulvic acids to produce significant enthalpy differences in some pH regions or as a function of ionic strength. In addition, Suwannee River fulvic acid was titrated in the presence of the tetraethylammonium cation to determine if protonation enthalpies are influenced by this larger cation containing nonpolar ethyl groups.
Materials and Methods Four humic substance fractions were utilized for the present study. Suwannee River fulvic acid (SRFA) and peat standard humic acid (PSHA) were purchased from
the International Humic Substances Society (IHSS).These humic substances have been characterized by a variety of techniques, and this information is available from the IHSS and in a recent publication primarily devoted to describing the physical and chemical properties of SRFA (18). Humic and fulvic acids were also isolated from a peat deposit located 30 km south of Rio de Janeiro, Brazil. Isolation and purification procedures for these Brazilian humic (BHA) and fulvic (BFA) acid fractions closely followed those recommended by the IHSS and are detailed in other publications (16, 19). Calorimetric titration and data reduction procedures have also been detailed previously (16,20), so only a brief summary is given here. A TRONAC (Orem, UT) isoperibol/isothermal calorimeter was used for the calorimetric determinations. Experiments were performed in the isoperibol mode, and several modifications were made to the basic instrument. These included adding a second buret and micro pH electrodes (Models MI-405 and MI402 from Microelectrodes, Inc., Londonderry, NH)and automating most instrument control and data acquisition functions. All salts, acids, and bases were ACS reagent grade or better. Freeze-dried humic and fulvic acids were dissolved in distilled-deionized water and passed through H+saturated cation-exchange resin (Bio-Rad AG-50W-X8, 20-50 mesh). Dissolution of the humic acid functions was aided by raising the pH to 7 with NaOH and brief (C5 min) ultrasonic dispersion before the cation-exchangestep. Initial concentrations of these humic substance solutions were 500 mg/L, and pH values ranged from 3.2 to 3.5. Next, 25.0 mL of these solutions was transferred to the calorimeter dewar reaction vessel, and solid NaCl or tetraethylammonium chloride (TEAC1) was added to attain an ionic strength of 0.001 M. Standardized base (0.1 N NaOH or tetraethylammonium hydroxide) was added from one of the previously filled burets to raise the pH to 10.4 f 0.1. After 20 min for thermal equilibration (this was necessary to cool the reaction vessel to near the bath temperature of 25 "C), titrations with standard acid (0.1 N HC1) were initiated. Fifteen to 20 equal volume aliquots (30-40 ,uL) of acid were added to reach the final pH value of 3.5-3.7. Most thermal and pH changes took place within 30 s of an addition, and a total of 90 s was allowed between successive additions for further stabilization. After 20-30 min for thermal equilibration following the acid titration, titration with 0.1 N standard base to pH 10.4 f 0.1 was conducted in identical fashion. This acid-base titration sequence was repeated at 0.01 and 0.1 M ionic strength on the same humic substance solution. All titrations were duplicated, and COZcontamination was minimized by purging the top of the calorimeter insert (the only connection with the laboratory atmosphere) continuously with Nz(g), The basic data reduction strategy was to correct for all extraneous reactions both in terms of proton consumption or release and reaction enthalpies and to ascribe the remainder of these quantities to the protonation (acid addition) or ionization (base addition) enthalpies (kJ/mol H+ bound or released). Nonchemical heat effects were corrected using standard formulas (21).Extraneous proton consumption and release reactions and associated enthalpies were estimated by using a 'theoretical blank' correction procedure detailed previously (16,20), Briefly, a rearranged form of the general titration curve equation Environ. Sci. Technol., Vol. 27,
No. 6, I993 1183
Table I. Acid and Base Consumption (mequiv/g (SRFA Only)
1 SD)for Several pH Intervals in 0.001, 0.01, and 0.1 M NaCl and TEACI
I = 0.001 M humic SRFA (NaCl) SRFA (TEACl) BFA (NaCl) BHA (NaCl) PSHA (NaCl)
Z = 0.01 M
I = 0.1 M
pH range
acid
base
acid
base
acid
base
4-6 8-10 4-10 4-6 8-10 4-10 4-6 8-10 4-10 4-6 8-10 4-10 4-6 8-10 4-10
2.01 (0.03) 0.79 (0.02) 3.72 (0.04) 2.11 (0.02) 0.91 (0.01) 4.01 (0.02) 1.43 (0.08) 0.72 (0.12) 2.82 (0.23) 1.17 (0.02) 0.73 (0.04) 2.85 (0.04) 1.49 (0.05) 0.66 (0.02) 3.18 (0.11)
1.92 (0.01) 0.82 (0.02) 3.58 (0.02) 1.93 (0.02) 0.94 (0.02) 3.79 (0.02) 1.30 (0.05) 0.77 (0.07) 2.71 (0.18) 1.09 (0.02) 0.80 (0.02) 2.81 (0.04) 1.40 (0.05) 0.75 (0.04) 3.07 (0.06)
2.15 (0.02) 0.75 (0.01) 3.80 (0.02) 2.18 (0.10) 0.80 (0.07) 4.06 (0.08) 1.46 (0.05) 0.67 (0.08) 2.81 (0.18) 1.28 (0.04) 0.70 (0.02) 2.85 (0.10) 1.57 (0.02) 0.57 (0.10) 3.06 (0.08)
2.02 (0.02) 0.80 (0.03) 3.65 (0.03) 2.06 (0.02) 0.90 (0.01) 3.86 (0.04) 1.35 (0.04) 0.73 (0.07) 2.70 (0.17) 1.19 (0.04) 0.76 (0.02) 2.79 (0.04) 1.50 (0.05) 0.68 (0.07) 3.01 (0.02)
2.15 (0.02) 0.77 (0.02) 3.67 (0.02) 2.34 (0.02) 0.79 (0.03) 4.18 (0.08) 1.46 (0.02) 0.63 (0.11) 2.68 (0.15) 1.30 (0.04) 0.68 (0.04) 2.87 (0.07) 1.58 (0.09) 0.57 (0.09) 2.94 (0.03)
2.05 (0.02) 0.80 (0.02) 3.54 (0.02) 2.14 (0.03) 0.83 (0.03) 3.85 (0.07) 1.35 (0.06) 0.69 (0.07) 2.58 (0.18) 1.26 (0.02) 0.72 (0.02) 2.69 (0.06) 1.47 (0.13) 0.63 (0.18) 2.76 (0.10)
is used in which the dependent variable is the volume of strong acid or base titrant required to reach an observed pH value during a titration. The difference between this theoretical titrant consumption and the actual quantity of titrant added is assumed to have been bound or released by the humic substances. Correction terms are included for both water formation and dissolved carbonate equilibria, and enthalpy values associated with these equilibria are used to estimate extraneous heat production or consumption. The most significant corrections are associated with water formation-particularly at pH values 9 during acid titrations. This 'theoretical blank' correction procedure has been used to simulate actual titrations of NaCl blank solutions (22). An estimated H+ or OH- consumption of 1 pmol agrees to within 10% (1SD) of the measured consumption between pH 4 and 9. A t pH 10, however, the difference can be 30%. The reason for this increased discrepancy above pH 9 is not clear, but silica dissolution from the glass dewar flask and stirring rod may contribute (22).
16
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Results
1184
Environ. Scl. Technol., Vol. 27, No. 6, 1993
+
BFA SRFA (TEACI)
0
Experimental net titration data (corrected for solution blank effects as outlined above) between pH 3.75 and 10.25 were fit with fifth degree polynomial equations of the form
...+
BHA
SRFA (NaCI)
m.
pmol of H+ bound or released = u ( ~ H ) ~ (+P H ) ~ + f(pH) g (1) for smoothing purposes. These equations estimated the actual pmol of H+ bound or released to within 0.3 pmol at any particular pH between 3.75 and 10.25,and R2values exceeded 0.9992 in all cases. Smaller order polynomials gave less acceptable fits and higher order polynomials resulted in no significant improvement. Table I contains acid and base consumption (expressed as mequiv of H+ bound or released/g of humicsubstance f 1SD)for several pH ranges as calculated using the polynomial equations for duplicate titrations. These duplicate titrations agreed to within 10% between pH 4 and 10, with the largest variability observed for BFA. Moreover, acid and base consumption between pH 4 and 10is similar (within 10% at all ionic strengths, and the total quantity of H+ bound or released between pH 4 and 10 decreases in the order SRFA > PSHA > BHA, BFA.
A b
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Flgure 2. Net (correctedfor Solution blank contributions)buffer capaclty distributions (dpmoi of H+/dpH) at 0.1 M ionic strength.
'Buffer capacity' values (dpmol of H+bound/dpH)were calculated by solving for the first derivatives of the fifth order polynomial regressions for each 0.25 pH unit interval between pH 3.75 and 10.25. Distinct buffer capacity maxima and minima are observed for all the humic substances investigated between pH 4 and 6 and pH 7.5 and 8.5, respectively (Figures 1 and 2). These maxima and minima shift to lower pH values as the ionic strength increases from 0.001 to 0.1 M. Buffer capacity plots generated with the actual experimental data were similar in shape and magnitude but 'noisier' than the smoothed data depicted in Figures 1 and 2.
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Flgure 7. Acid additlon enthalples (kJ/mol of H+ reacted NaCl media as a function of pH for SRFA. -46
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Flgure 8. Acid additlon enthalpies (kJ/mol of H+ reacted f 1 SD) In TEACl media as a function of pH for SRFA.
Enthalpy values (kJ/mol of H+ bound or released) as functions of pH are presented as Figures 3-8. These data represent the average of two separate titrations for each ionic strength and titration direction. Enthalpy values over 0.3 pH unit intervals from 3.6 to 10.5were combined and averaged for inclusion in these figures. Corresponding pH values are the average for each interval. This averaging and smoothing procedure is identical to that followed previously (16)and permits the precision of these enthalpy data to be estimated. Except for Figure 4, only protonation (acid addition) enthalpies are depicted.
Some general trends and features of these data are as follows: (1)Protonation enthalpies are similar below pH 9 for all humic substances and ionic strengths investigated. Values are slightlyendothermic (0-4 kJ/mol of H+ bound) and slightly exothermic (0 to -2 kJ/mol) for fulvic and humic acids, respectively, at pH 4. Between pH 4 and 6.6, enthalpy values become slightly more exothermic (by 2-4 kJ/mol) and above pH 6.5,enthalpy values increase more rapidly and reach values of -18 to -24 kJ/mol at pH 9. (2) Protonation and ionization enthalpies are not completely equivalent (Figure 3 compared with Figure 4) Environ. Scl. Technol., Vol. 27,
No. 6, 1993 1185
as noted previously (16). Ionization enthalpies are generally greater in absolute magnitude, and this difference can exceed 8 kJ/mol, particularly below pH 4.5 and above pH 9.5. (3)Ionic, strength effects are evident in several instances above pH 9. Protonation enthalpies become sharply more endothermic for SRFA in 0.1 M TEACl compared with lower TEACl concentrations (Figure 8) and 0.1 M NaCl (Figure 7). Similarly, BFA, BHA, and PSHA exhibit less exothermic protonation enthalpies as ionic strength increases above pH 9, although these differences are not always significant (e.g., Figure 3).
Discussion Acid and base consumption values (Table I) help to assess the precision and accuracy of the enthalpy data. For a particular ionic strength and pH range, acid and base consumption agree to within 2 SD or 10%. Thus, the significant hysteresis effects observed in several previous investigations (23-25) are largely absent in this studymost probably because titrations were conducted relatively rapidly (9.5 and
