Reduction and Reoxidation of Humic Acid - American Chemical Society

2 Jul 2010 - in natural environments, with electron accepting (EAC) and electron donating capacities (EDC) ranging up to several moles per kg (3-10)...
0 downloads 0 Views 372KB Size
Download PDF
Environ. Sci. Technol. 2010, 44, 5787–5792

Reduction and Reoxidation of Humic Acid: Influence on Spectroscopic Properties and Proton Binding FELIX MAURER, ISO CHRISTL,* AND RUBEN KRETZSCHMAR Institute of Biogeochemistry and Pollutant Dynamics, Department of Environmental Sciences, ETH Zurich, CHN, Universita¨tstrasse 16, 8092 Zu ¨ rich, Switzerland

Received February 22, 2010. Revised manuscript received June 9, 2010. Accepted June 14, 2010.

Previous studies on proton and metal binding to humic substances have not considered a potential influence of reduction and oxidation of functional groups. Therefore, we investigated how proton binding of a purified soil humic acid was affected by reduction. Reduction of the humic acid was carried out using an electrochemical cell that allowed us to measure the amounts of electrons and protons involved in reduction reactions. We further applied spectroscopic methods (UV-vis, fluorescence, FT-IR, C-1s NEXAFS) to detect possible chemical changes in the humic acid induced by reduction and reoxidation. The effect of reduction on proton binding was determined with acid-base titrations in the pH range 4-10 under controlled redox conditions. During reduction, 0.54 mol kg-1 protons and 0.55 mol kg-1 electrons were transferred to humic acid. NICA-Donnan modeling revealed an equivalent increase in proton-reactive sites (0.52 mol kg-1) in the alkaline pH-range. Our results indicate that reduction of humic acid increased the amount of proton-reactive sites by15%comparedtotheuntreatedstate.Spectroscopicdifferences between the untreated and reduced humic acid were minor, apart from a lower UV-vis absorption of the reduced humic acid between 400 and 700 nm.

Introduction Humic substances (HS) strongly affect the solubility and speciation of both nutrients and contaminants in soils as they contain a large number of functional groups able to bind protons and metal cations (1, 2). High organic matter contents are typically found in poorly aerated wetland and forest soils, where decomposition of organic substances occurs under limited O2 supply. In such soils, HS may be present either in an oxidized or reduced state, depending on the prevailing redox conditions. Humic substances can act as an electron acceptor or donor in natural environments, with electron accepting (EAC) and electron donating capacities (EDC) ranging up to several moles per kg (3-10). Formal electrode potentials reported for HS range from 0.4 to 0.8 V (4, 11). This implies that the presence of HS can affect the redox state of various metals. They may act as a reductant, e.g., for Fe(III) and Hg(II) (5, 12), or as an oxidant, e.g., for Cu (0 and I), Sn(II), and U(IV) (4, 13). Within the carbon backbone of HS, redox activity is mainly * Corresponding author phone: +41 44 633 60 01; e-mail: [email protected]. 10.1021/es100594t

 2010 American Chemical Society

Published on Web 07/02/2010

attributed to quinone-like moieties (3, 9, 10, 14, 15), but nitrogen and sulfur containing moieties may also be relevant. The reduction of HS may affect their cation binding properties in various ways. Reduction of quinone moieties to phenolic groups would provide additional cation binding sites. Reduction reactions may further alter the binding affinity of existing functional groups. Redox-couples of simple organic acids like fumaric-succinic acid and pyruvic-lactic acid can serve to illustrate this latter connection. The pKa values of the oxidized acids, fumaric acid and pyruvic acid, are by more than one logarithmic unit lower than the pKa values of the carboxylic groups of the corresponding reduced acids, succinic acid and lactic acid (16). Information on the effect of redox conditions on proton and metal cation binding is almost absent up to now. Usually, HS are extracted and stored under oxic conditions (17). Investigations on proton and metal cation binding of HS are generally performed under exclusion of CO2 to avoid dissolved carbonate species in the solution, but the redox potential is normally not monitored. Considering the experimental conditions, most published data may reflect humic binding properties under oxic conditions (1). Since strongly reducing conditions may significantly alter proton and metal cation binding properties of HS and data measured under strongly reducing conditions are lacking, the effect of HS on proton buffering and metal binding cannot reliably be estimated for reducing environments. In this study, we reduced a well-characterized soil humic acid using an electrochemical method that allowed us to determine the amount of electrons and protons involved in reduction reactions. The reversibility of the electron transfer was investigated using either O2 as a reoxidant or the redoxdye dichlorophenol-indophenol (DCPIP). We applied several spectroscopic methods to detect possible chemical changes in the humic acid induced by reduction and reoxidation. UV-vis and fluorescence spectroscopy were chosen because redox-active quinone-like moieties are thought to play an important role for the optical properties of HS (15, 18). C-1s NEXAFS spectroscopy was selected because of its sensitivity for aromatic ring substitutents such as hydroxyl groups (19). Further, the effect of reduction on proton binding of humic acid was investigated by potentiometric acid-base titrations in the pH range of 4-10 at four different ionic strengths. The reductive changes in the proton binding were quantitatively described with the NICA-Donnan model.

Materials and Methods Solutions were prepared with ultrapure deionized water (Milli-Q, 18 MΩ cm). Water for the preparation of anoxic solutions was boiled for 30 min while purging with N2. All anoxic work was performed in a glovebox (Braun, Germany) with a N2 atmosphere (O2 < 1 ppm). Redox potentials were measured using a combined platinum Ag/AgCl electrode (Hamilton Slimtrode RX) and are reported as Eh values, i.e., as potentials against the standard hydrogen electrode. Humic Acid. A well-characterized, purified humic acid (HA) extracted from a humic Gleysol in northern Switzerland (20) was used in this study as an example of humic materials. Humic acid solutions were prepared by diluting HA stock solution, purging with N2, titrating to the desired pH (pH 7 if not stated otherwise) with 2 M NaOH, and adding NaCl to adjust the ionic strength. These solutions are referred to as untreated. Solutions termed reduced represent solutions that additionally underwent electrochemical reduction until an Eh value of -0.20 V was reached (see below). Reoxidized VOL. 44, NO. 15, 2010 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

9

5787

solutions were purged with air or O2 for 2 h after reduction to reoxidize HA, followed by purging with N2. All HA solutions were stored in the glovebox until further use. Reduction. HA solutions (3.62 g L-1, 0.02 M NaCl) were electrochemically reduced at an electrode potential of -0.59 V in a pH-stat experiment under anoxic conditions using a glassy carbon working electrode (10). Reduction was stopped after reaching an Eh value of -0.20 V. Electron accepting capacity (EAC) was determined as the electron consumption by integrating reductive currents over time. During reduction, pH was kept constant by automated titration (Titrino 751 GPD, Metrohm, equipped with a Mettler Toledo pH-electrode in a Metrohm 6.1420.100 flow through cell) with 0.01 M HCl (Titrisol, Merck). The number of protons involved in reduction reactions (proton consumption, PC) was calculated from acid consumption. Currents observed in blank experiments using 0.02 M NaCl solutions were negligible. For comparison, HA was also reduced with hydrogen gas and Pd powder (Sigma-Aldrich) in a 0.15 M phosphate buffer at pH 7 following the procedure of Ratasuk and Nanny (9). Briefly, ∼5 mg of Pd powder were added to 20 mL of 1.81 g L-1 (0.01 M NaCl) anoxic humic acid solutions in sealed septum beakers. After purging for 10 min with a gas mixture consisting of 8% H2 and 92% N2, the solutions were placed on an overhead shaker for 24 h and subsequently filtered using 0.45 µm nylon filters (Optiflow, Wicom). These samples are referred to as H2/Pd reduced. EDC Determination by Reoxidation with O2. For the determination of the electron donating capacity with respect to oxygen (EDCO2), the procedure of Bauer and Kappler (21) was adapted as follows. HA solutions were diluted to a concentration of 1.81 g L-1 and buffered with 0.15 M phosphate at pH 7 in the glovebox. An oxygen electrode connected to a PA2000 picoammeter (Unisense, Denmark) and a long cannula for pressure equilibration were installed through the rubber septum of a 7 mL glass vial containing 4 mL HA solution. Two mL of oxygen-saturated deionized water were added in 200 µL steps using a gastight syringe (Vici, Baton Rouge, USA). Each addition was followed by gentle stirring by means of a magnetic stirrer for 15 s. After an equilibration time of 60-120 s, the electrode current was recorded. The total reaction time cumulated to 10 min per sample. EDCO2 was calculated as four times the amount of added O2 needed to detect O2 in the HA solution, assuming a complete reduction of oxygen with four electrons transferred per O2. EDC Determination by Reoxidation with DCPIP. The electron donor capacity with respect to 2,6-dichlorophenolindophenol (DCPIP, Fluka & Merck, purity 95%) was determined using a procedure adapted from Aeschbacher et al. (10). Anoxic DCPIP solutions (50 µM DCPIP, 0.33 M phosphate buffer at pH 7) were mixed inside the glovebox with HA solutions to obtain HA concentrations in the range of 40-130 mg L-1. The decrease of absorbance at 603 nm with increasing HA concentration was immediately measured using a UV-vis spectrophotometer (Varian Cary 50 Bio). The data were corrected for the absorbance of HA. From the decrease of absorbance the amount of reduced DCPIP was calculated and further transformed into amounts of electrons transferred from HA to DCPIP (EDCDCPIP). In addition, 3-fold reduction/ reoxidation cycles were conducted to investigate the reversibility of electrochemical reduction and reoxidation by oxygen. After each reduction, EDCDCPIP was determined. Reoxidation was achieved by purging with air for 1.5-2 h followed by purging with N2. Spectroscopic Characterization. UV-vis absorbance of untreated, reduced, and reoxidized HA (180 mg L-1) was measured against deionized water in a sealed quartz cuvette using a Cary 50 Bio spectrophotometer (Varian). A solution of a reduced redox indicator was used to prove the cuvette 5788

9

ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 44, NO. 15, 2010

to be tight with respect to oxygen diffusion for the measuring time required. Additionally, fluorescence emission-excitation matrices (EEMs) of HA solutions (9 mg L-1) were collected on a LS 50 B luminescence spectrophotometer (Perkin-Elmer) varying excitation wavelengths from 200-520 nm in 10 nm steps. Emission was recorded from 310-580 nm in 0.5 nm steps. Data analysis included subtraction of blank (deionized water) spectra and corrections for inner filter effects (22) and for slight variations in HA concentrations by normalizing intensities to measured HA concentrations relative to the concentration of untreated HA. Fluorescence intensities were normalized to the area of the Raman peak (23) and are reported in Raman units (R.U.). C-1s NEXAFS spectra were recorded using the scanning transmission X-ray microscope (STXM) at beamline X-1A, National Synchrotron Light Source (NSLS), Upton, NY. Specimens were prepared by placing 2 µL HA solution onto a X-ray transparent Si3N4 window (Silson Ltd., Northampton, UK). Preparation, drying, and transfer of specimens into gastight glass vials were performed inside the glovebox. Measurements and data analysis followed the procedures in Christl and Kretzschmar (24). Potentiometric Titrations. The pH- and ionic strengthdependent protonation behavior of untreated and reduced HA was studied by potentiometric acid-base titrations at 25 ( 1 °C. The setup was previously described by Christl and Kretzschmar (25). Details can be found in the Supporting Information. All solutions (H2O, 0.05 M HCl, 0.05 M NaOH (both Titrisol, Merck), and 2 M NaCl (Merck, p.a.)) were prepared under nitrogen atmosphere using O2- and CO2free deionized water. To shield the solutions from atmospheric O2 and CO2, titration tubes were jacketed in additional silicone rubber tubes. The tube interspaces, the buret headspaces, and the glass titration vessel were continuously flushed with water-saturated, O2- and CO2-free nitrogen gas. The setup was proven to be tight with respect to O2 using an aqueous solution of a reduced redox indicator (methylene blue). During titration of reduced HA, solution Eh (at constant pH) drifted at maximum 10 mV toward more oxidizing conditions. NICA-Donnan Modeling. Titration data were fitted with the NICA-Donnan model (26) using the computer program FIT (27). For untreated HA, unconstrained fitting gave unreasonable results. Parameters m1 and m2 were set to values obtained by Christl and Kretzschmar (25). Parameter b was ˜ H,1, and K ˜ H,2 remained fixed to 0.49. Parameters Qmax,1, Qmax,2, K unconstrained. For reduced HA, fitting was performed assuming that reduction reactions had primarily changed the number of high affinity sites as well as the proton affinity of both low and high affinity sites. Accordingly, Qmax,1, b, and mi of reduced HA were fixed to the respective values of ˜ H,2, and Qmax,2 were optimized ˜ H,1, K untreated HA, and only K for reduced HA.

Results and Discussion Electron Transfer and Proton Consumption during Electrochemical Reduction. Cumulated electron and proton consumption during electrochemical reduction increased in parallel over the whole reduction period to final values of 0.52-0.57 mol electrons (EAC) and 0.52-0.55 mol protons (PC) per kilogram of HA, yielding a proton to electron ratio (PC/EAC) of 0.96-1.06 (ranges resulting from four reduction experiments; see Figure 1a and Table S1). For the reduction of quinoids, a PC/EAC ratio of 1.0 is expected. Therefore, the experimental PC/EAC ratio is in agreement with the assumption that quinoid moieties dominate reduction reactions. Metal analysis using ICP-OES indicated that redoxactive trace metals present in the HA solution can contribute to the measured electron accepting capacities only by less than 7%.

FIGURE 2. Average electron donating capacity of humic acid exposed to different treatments as determined with DCPIP (EDCDCPIP) and O2 (EDCO2), respectively. The treatments comprise humic acid which was untreated, electrochemically reduced, reduced by H2/Pd, and reoxidized with oxygen after electrochemical reduction. For comparison, electron consumption (EAC) and proton consumption (PC) during electrochemical reduction are shown, additionally. Error bars represent 2 standard errors.

FIGURE 1. A. Electron and proton consumption during electrochemical reduction of humic acid at pH 7. Reduction was stopped at a redox potential of -0.20 V. B. Oxygen-electrode current as a function of added O2 using untreated and electrochemically reduced humic acid, respectively. The dotted line represents the electrode base current (1.3 pA). Dashed lines show regression lines used to calculate oxygen consumption. The measured EAC of our soil HA is in the range of EAC values determined by other authors for different HS. Bauer et al. (8) measured 0.64 molc kg-1 for Pahokee Peat HA using electrochemical reduction. Kappler and Haderlein (7) determined 0.43 molc kg-1 for a soil HA (EAC determined as difference of unreduced and reduced EDC values with respect to K3[Fe(CN)6]). Aeschbacher et al. (10) related the EAC during electrochemical reduction of HA to their carbon-to-hydrogen (C/H) ratios using linear regression. The HA used in this study has a C/H ratio of 0.87 (20). Correspondingly, an EAC of 0.5 molc kg-1 can be expected, which corresponds well to the measured values. The protons consumed during reduction of HA represent only ∼1% of the total proton content (20) and ∼15% of the total proton buffering capacity in the pH range 4-10 of the untreated HA (see below). In a later section of this paper, we will discuss to which extent the consumed protons correspond quantitatively to changes in proton-reactive sites. Electron Transfer Reversibility. Figure 1b shows the oxygen-electrode current resulting from the addition of oxygen to untreated and reduced HA. While the addition of oxygen to untreated HA immediately resulted in an increase of oxygen concentration (detected as a higher electrode current), the electrode current in the reduced sample remained constant at the beginning and increased not until larger oxygen doses. The average EDCO2 of reduced HA resulting from these experiments was about 0.40 molc kg-1 independent of the reduction method. For DCPIP, similar EDC values were obtained (see Figure 2 and Table S1 for detailed values). Differences were observed, however, for the oxygen and the DCPIP method regarding EDC of untreated and reoxidized HA. About 0.09-0.13 molc kg-1 of electrons were transferred from untreated and

reoxidized HA to DCPIP, while no electron transfer was observed when using oxygen as the oxidant. This indicates that DCPIP can acquire electrons from the untreated HA in contrast to oxygen. While the EDC values of HA were very similar for different reduction treatments, the measured Eh values differed between electrochemically reduced samples (-0.195 to -0.199 V) and H2/Pd reduced samples (-0.258 to -0.273 V, see Table S1). We conclude that the number of redox moieties reacting in the Eh range from -0.20 to -0.27 V was small compared to the total EDC. Overall, the two reduction methods appear to reduce HA to a similar extent. As EDC values were smaller than EAC values, reoxidation was only partial using either reoxidation method (75% using the O2 method, 55% using the DCPIP method and subtracting the EDC of the untreated humic acid). The most likely explanation is that both reoxidation methods captured only kinetically fast reoxidation reactions as the reaction time for both methods was 10 min at maximum. Aeschbacher et al. (10) found that within the first minute, only about 50% of the electrons electrochemically transferred to Leonardite humic acid reacted with oxygen. Another pool of ∼35% reacted within 24 h, while the remaining part was recalcitrant to oxidation with oxygen but not to mediated electrochemical oxidation. In the redox cycle experiment (Figure S1), EAC values for the three consecutive reduction steps were 0.53, 0.45, and 0.40 molc kg-1. EDCDCPIP of the reduced solution were 0.36, 0.38, and 0.35 molc kg-1. The untreated and the reoxidized solutions had all EDCDCPIP values around 0.12 molc kg-1 (Figure S1). The ratio of EDC to EAC thus increased with each step from 0.68 to 0.84 and finally to 0.88. The reason for the initially lower EDC-to-EAC ratio may be again a fraction of reduced humic acid moieties that cannot be reoxidized by oxygen within the two hours of reoxidation time. Despite kinetic limitations, the two reoxidation methods provide a valuable tool for the assessment of the redox properties of HA. The DCPIP method, which is experimentally simple, provides a good measure for EDC of humic acid solutions with respect to kinetically fast reoxidation with oxygen. UV-vis Spectroscopy. Figure 3a shows UV-vis spectra of untreated, reduced, and reoxidized HA at pH 7 in the range of 400-800 nm. Between 400 and 700 nm, reduced HA exhibited lower absorbance than untreated and reoxidized HA, respectively. Above 350 nm, UV-vis absorbance of humic VOL. 44, NO. 15, 2010 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

9

5789

FIGURE 3. A. UV-vis absorption spectra of untreated, electrochemically reduced and reoxidized humic acid at pH 7. Humic acid concentration was 180 mg L-1. B. C-1s NEXAFS spectra of untreated, electrochemically reduced, and reoxidized humic acid at pH 5 and pH 7. Photon energy ranges exhibiting spectral features of C-1sfπ*C)C, C-1sfπ*C)C-O, and C-1sfπ*C)O transitions assigned to H and C-substituted aromatic, phenolic, and carboxyl carbon are labeled with characters a, b, and c, respectively. acids typically decreases with increasing wavelength in an exponential-like fashion. It was shown that the absorption in this region cannot be explained by simple linear superposition of the spectra of independent chromophores but may result from intramolecular charge-transfer interactions between hydroxy-aromatic electron donor and quinoid acceptor groups (18). Accordingly, the decreased absorption of reduced HA between 400 and 700 nm may be explained by an elimination of charge transfer contacts due to reduction of humic electron acceptor groups. Correspondingly, reoxidation of these groups may re-establish the charge transfer groups, explaining the restored absorbance of reoxidized HA. Fluorescence Spectroscopy. Figure 4 shows fluorescence excitation-emission matrices (EEMs) of untreated, reduced, and reoxidized HA. All EEMs show a peak at excitation/ emission wavelengths of 320/440 nm, irrespective of the redox treatment. Fluorescence intensities were very low and fairly similar for all samples. Peak intensities amounted to 0.060, 0.072, and 0.070 R.U. for untreated, reduced, and reoxidized HA, respectively. The results indicate that the fluorescence properties of HA were virtually unaffected by electrochemical reduction and reoxidation. Klapper et al. (6) related peak and intensity differences between EEMs of freshly extracted and microbially reduced sediment fulvic acids to differences in redox states of redox5790

9

ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 44, NO. 15, 2010

FIGURE 4. Fluorescence excitation-emission matrices (EEMs) of untreated, electrochemically reduced, and reoxidized humic acid at pH 7. Colors correspond to fluorescence intensities given in Raman units (R.U.) as indicated by the intensity scale bar at the bottom. Humic acid concentration was 9 mg L-1 for all measurements. active moieties. In later studies, parallel factor analysis (PARAFAC) was used to split the EEM spectral information into a linear combination of spectra of individual compound groups, promoting fluorescence spectroscopy as a tool for the inves-

FIGURE 5. Proton binding behavior of untreated and electrochemically reduced humic acid as a function of pH and ionic strength I. Plotted data were derived from potentiometric acid-base titrations conducted in the absence of O2. The data sets of untreated and reduced humic acid are superimposed at pH 7 and I ) 0.01 M for better comparison. tigation of the redox state of HS (15, 28). It was further proposed to express the degree of reduction of a sample as the ratio of reduced to the total (reduced and oxidized) quinone-like fractions as estimated by the PARAFAC analysis (“reducing index”) (29). For unambiguous identification of redox effects on EEMs, it is essential to compare samples that differ only in the redox state of redox-active moieties to minimize potential side effects due to origin of material, pH, or reduction method. So far, published studies did not provide clear evidence for redox-induced changes in peak positions. Fulton et al. (28) compared EEMs of fulvic acids isolated from different depths of a stratified lake before and after microbial reduction in the laboratory. The spectral changes were small, and peak positions and intensities varied depending on the sample origin. Fimmen et al. (30) compared fluorescence properties of native and electrochemically reduced lake organic matter. They found an increase in intensity only for one of three samples at unchanged peak positions with the exception of sharper peaks in the region of 240-260 nm (excitation). Most recently, Macalady and Walton-Day (31) investigated the change of the reducing index after reduction with Zn in the laboratory for several organic matter samples. They did not find a consistent relation between reducing index and redox state for their samples. Likewise, our data do not show any changes in peak positions upon reduction of the humic acid. Overall, our results clearly confirm recent doubts about the suitability of fluorescence measurements for

the determination of redox states of redox-active moieties in natural organic matter (31, 32). C-1s NEXAFS Spectroscopy. The C-1s NEXAFS spectra of HA samples show three pre-edge features: 2 peaks positioned at 285.1 and 288.4 eV and a shoulder at 286.6 eV (Figure 3b). The pre-edge features correspond to C-1sfπ* resonances of H- and C-substituted aromatic carbon (π*C)C, 285.1 eV), O-substituted aromatic carbon (π*C)C-O, 286.6 eV), and carboxyl carbon (π*C)O, 288.4 eV) as labeled in Figure 3b with the characters a, b, and c (24). The two sharper peaks in the post-edge region can be attributed to L-edges of potassium present in the electrolyte solution. The absorbance of the pre-edge region was shown to be very sensitive to aromatic ring substitutions with electronwithdrawing groups like, e.g., phenolic groups (19). In addition, a shift of more than 1 eV toward lower energies was reported for π* resonances of p-benzoquinone compared to its reduced counterpart, hydroquinone, in the gas-phase (33). We therefore expected NEXAFS spectroscopy to be suitable for observing changes due to reduction of quinone moieties. Our spectra of differently treated HA, however, overlap almost perfectly in the pre-edge region at pH 5 and 7. Only around 289 eV, reduced HA showed a slightly reduced absorbance at pH 7 compared to untreated and reoxidized HA. The expected energy shift was not detected implying that only a small fraction of total aromatic oxygen substituents was involved in reduction of HA. Similar to NEXAFS spectroscopy, FT-IR spectroscopy did not reveal chemical differences between untreated and reduced HA (see the SI). Overall, our spectroscopic results indicate that the electrochemical reduction and reoxidation with O2 did not significantly change the chemical structure of HA. Potentiometric Titrations. As shown above, HA consumed on average 0.54 mol kg-1 protons during electrochemical reduction. In the following, we will focus on acid-base properties of untreated and reduced HA to elucidate the effect of reduction on proton-reactivity of HA. Figure 5 shows the protonation behavior of untreated and reduced HA as a function of pH and ionic strength as derived from potentiometric acid-base titration. While proton binding isotherms of untreated and reduced HA are almost parallel at equal ionic strength in the acidic part, they diverge in the alkaline part of the studied pH range. Between pH 4 and 10, reduced HA buffered by ∼0.3 mol kg-1 more protons at a given ionic strength than untreated HA. These two observations indicate a pronounced increase of deprotonable groups with a high proton-affinity (e.g., phenolic-like groups) in reduced HA, while the number of low proton-affinity sites (mainly carboxylic groups) appears to be altered less considerably upon reduction. To obtain a quantitative description of reduction-induced changes in HA protonation, proton binding data of untreated and reduced HA were fitted with the NICA-Donnan model. For ˜ H,1, and K ˜ H,2, were reduced HA, only the parameters Qmax,2, K optimized, following the observation that reduction changed mainly the number of phenolic-like groups and probably also the proton affinity for both carboxylic- and phenolic-like groups, respectively. The obtained set of NICA-Donnan model parameters is reported in Table 1. Calculated proton binding isotherms are shown in Figure S5. Despite the strong fitting constraints chosen, proton binding to reduced HA was described accurately

TABLE 1. NICA-Donnan Model Parameters for Describing Proton Binding of Untreated and Reduced Humic Acid (HA)a

untreated HA reduced HA a

R2

b

Qmax,1

logK˜H,1

m1

Qmax,2

logK˜H,2

m2

ΣQmax

0.9979 0.9992

0.49 0.49

2.23 2.23

3.07 2.97

0.48 0.48

2.28 2.80

8.22 7.94

0.23 0.23

4.51 5.03

Fixed and constrained parameter values are in italics (see Materials and Methods).

VOL. 44, NO. 15, 2010 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

9

5791

with the NICA-Donnan model. The parameter estimates lie within ranges commonly reported for humic acids (1). The differences in parameter estimates between untreated and reduced HA (Table 1) reflect the experimental observations shown in Figure 5. While the fitted proton affinity parameters ˜ H,2 were fairly similar for untreated and reduced HA, ˜ H,1 and K K the fitted total number of phenolic-like sites (Qmax,2) was by 0.52 mol kg-1 higher for reduced HA. This calculated increase in proton-reactive sites is in excellent agreement with the amount of protons consumed during reduction (0.54 mol kg-1). We conclude that for the most part, proton consumption during reduction resulted from the formation of additional binding sites exhibiting a high proton-affinity. This finding corroborates that quinoids having a high proton-affinity in their reduced state may play a dominant role as redox-active moieties in humic acids. Compared to natural systems, experimental conditions during acid-base titration of reduced HA correspond to strongly reducing, anoxic environments. Our results show that the total proton buffering capacity of HA increased by 15% upon reduction. This increase may have little effect on the actual proton buffering of HS in reduced soils or sediments because the additional sites deprotonated to a considerable extent not until pH values markedly exceed the pH range of anoxic environments (pH 5-7). However, the binding sites generated during reduction are expected to be significant for the fate of contaminants having a high affinity for phenolic-like groups (2).

Acknowledgments We gratefully acknowledge Michael Aeschbacher and Michael Sander for valuable discussions and help with electrochemical reduction of humic acids as well as Kurt Barmettler and Sue Wirick for support in the laboratory and at the STXM, respectively. This research was financially supported by the Swiss National Science Foundation under grant No. 200021117933.

Supporting Information Available Tabulated EAC and EDC values, data for consecutive redox cycles, UV-vis difference spectra and fractional loss of UV-vis absorbance, UV-vis absorbance of solutions used for fluorescence measurements, FT-IR spectra, and additional details on acid-base titrations and NICA-Donnan model fits. This material is available free of charge via the Internet at http://pubs.acs.org.

Literature Cited (1) Milne, C. J.; Kinniburgh, D. G.; Tipping, E. Generic NICA-Donnan model parameters for proton binding by humic substances. Environ. Sci. Technol. 2001, 35, 2049–2059. (2) Milne, C. J.; Kinniburgh, D. G.; van Riemsdijk, W. H.; Tipping, E. Generic NICA-Donnan model parameters for metal-ion binding by humic substances. Environ. Sci. Technol. 2003, 37, 958–971. (3) Scott, D. T.; McKnight, D. M.; Blunt-Harris, E. L.; Kolesar, S. E.; Lovley, D. R. Quinone moieties act as electron acceptors in the reduction of humic substances by humics-reducing microorganisms. Environ. Sci. Technol. 1998, 32, 2984–2989. (4) Struyk, Z.; Sposito, G. Redox properties of standard humic acids. Geoderma 2001, 102, 329–346. (5) Chen, J.; Gu, B.; Royer, R. A.; Burgos, W. D. The roles of natural organic matter in chemical and microbial reduction of ferric iron. Sci. Total Environ. 2003, 307, 167–178. (6) Klapper, L.; McKnight, D. M.; Fulton, J. R.; Blunt-Harris, E. L.; Nevin, K. P.; Lovley, D. R.; Hatcher, P. G. Fulvic acid oxidation state detection using fluorescence spectroscopy. Environ. Sci. Technol. 2002, 36, 3170–3175. (7) Kappler, A.; Haderlein, S. B. Natural organic matter as reductant for chlorinated aliphatic pollutants. Environ. Sci. Technol. 2003, 37, 2714–2719. (8) Bauer, M.; Heitmann, T.; Macalady, D. L.; Blodau, C. Electron transfer capacities and reaction kinetics of peat dissolved organic matter. Environ. Sci. Technol. 2007, 41, 139–145.

5792

9

ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 44, NO. 15, 2010

(9) Ratasuk, N.; Nanny, M. A. Characterization and quantification of reversible redox sites in humic substances. Environ. Sci. Technol. 2007, 41, 7844–7850. (10) Aeschbacher, M.; Sander, M.; Schwarzenbach, R. P. Novel electrochemical approach to assess the redox properties of humic substances. Environ. Sci. Technol. 2010, 44, 87–93. (11) Skogerboe, R. K.; Wilson, S. A. Reduction of ionic species by fulvic acid. Anal. Chem. 1981, 53, 228–232. (12) Alberts, J. J.; Schindler, J. E.; Miller, R. W.; Nutter, D. E. Elemental mercury evolution mediated by humic acid. Science 1974, 184, 895–896. (13) Szila´gyi, M. The redox properties and the determination of the normal potential of the peat-water system. Soil Sci. 1973, 115, 434–437. (14) Nurmi, J. T.; Tratnyek, P. G. Electrochemical properties of natural organic matter (NOM), fractions of NOM, and model biogeochemical electron shuttles. Environ. Sci. Technol. 2002, 36, 617–624. (15) Cory, R. M.; McKnight, D. M. Fluorescence spectroscopy reveals ubiquitous presence of oxidized and reduced quinones in dissolved organic matter. Environ. Sci. Technol. 2005, 39, 8142–8149. (16) Martell, A. E.; Smith, R. M.; Motekaitis, R. J. NIST Critically Selected Stability Constants of Metal Complexes; National Institute of Standards and Technology: Gaithersburg, MD, 2004. (17) Stevenson, F. J. Humus Chemistry: Genesis, Composition, Reactions; Wiley: New York, 1994. (18) Del Vecchio, R.; Blough, N. V. On the origin of the optical properties of humic substances. Environ. Sci. Technol. 2004, 38, 3885–3891. (19) Baˆldea, I.; Schimmelpfennig, B.; Plaschke, M.; Rothe, J.; Schirmer, J.; Trofimov, A. B.; Fangha¨nel, T. C 1s near edge X-ray absorption fine structure (NEXAFS) of substituted benzoic acidssA theoretical and experimental study. J. Electron Spectrosc. Relat. Phenom. 2007, 154, 109–118. (20) Christl, I.; Knicker, H.; Ko¨gel-Knabner, I.; Kretzschmar, R. Chemical heterogeneity of humic substances: characterization of size fractions obtained by hollow-fibre ultrafiltration. Eur. J. Soil Sci. 2000, 51, 617–625. (21) Bauer, I.; Kappler, A. Rates and extent of reduction of Fe(III) compounds and O2 by humic substances. Environ. Sci. Technol. 2009, 43, 4902–4908. (22) Ohno, T. Fluorescence inner-filtering correction for determining the humification index of dissolved organic matter. Environ. Sci. Technol. 2002, 36, 742–746. (23) Lawaetz, A. J.; Stedmon, C. A. Fluorescence intensity calibration using the Raman scatter peak of water. Appl. Spectrosc. 2009, 63, 936–940. (24) Christl, I.; Kretzschmar, R. C-1s NEXAFS spectroscopy reveals chemical fractionation of humic acid by cation-induced coagulation. Environ. Sci. Technol. 2007, 41, 1915–1920. (25) Christl, I.; Kretzschmar, R. Relating ion binding by fulvic and humic acids to chemical composition and molecular size. 1. Proton binding. Environ. Sci. Technol. 2001, 35, 2505–2511. (26) Koopal, L. K.; Saito, T.; Pinheiro, J. P.; van Riemsdijk, W. H. Ion binding to natural organic matter: General considerations and the NICA-Donnan model. Colloids Surf. A 2005, 265, 40–54. (27) Kinniburgh, D. G.; Tang, C. K. FIT; British Geological Survey: Wallingford, UK, 1998. (28) Fulton, J. R.; McKnight, D. M.; Foreman, C. M.; Cory, R. M.; Stedmon, C.; Blunt, E. Changes in fulvic acid redox state through the oxycline of a permanently ice-covered Antarctic lake. Aquat. Sci. 2004, 66, 27–46. (29) Miller, M.; McKnight, D.; Cory, R.; Williams, M.; Runkel, R. Hyporheic exchange and fulvic acid redox reactions in an alpine stream/wetland ecosystem, Colorado Front Range. Environ. Sci. Technol. 2006, 40, 5943–5949. (30) Fimmen, R. L.; Cory, R. M.; Chin, Y.-P.; Trouts, T. D.; McKnight, D. M. Probing the oxidation-reduction properties of terrestrially and microbially derived dissolved organic matter. Geochim. Cosmochim. Acta 2007, 71, 3003–3015. (31) Macalady, D.; Walton-Day, K. New light on a dark subject: On the use of fluorescence data to deduce redox states of natural organic matter (NOM). Aquat. Sci. 2009, 71, 135–143. (32) Boyle, E. S.; Guerriero, N.; Thiallet, A.; Del Vecchio, R.; Blough, N. V. Optical properties of humic substances and CDOM: Relation to structure. Environ. Sci. Technol. 2009, 43, 2262–2268. (33) Francis, J. T.; Hitchcock, A. P. Inner-shell spectroscopy of p-benzoquinone, hydroquinone, and phenol: Distinguishing quinoid and benzenoid structures. J. Phys. Chem. 1992, 96, 6598–6610.

ES100594T