How to overcome inter-electrode variability and instability to quantify dissolved oxygen, Fe(II), mn(II), and S(−II) in undisturbed soils and sediments using voltammetry
© Slowey and Marvin-DiPasquale; licensee BioMed Central Ltd. 2012
Received: 20 April 2011
Accepted: 26 May 2012
Published: 25 June 2012
Although uniquely capable of measuring multiple redox constituents nearly simultaneously with no or minimal sample pretreatment, voltammetry is currently underutilized in characterizing redox conditions in aquatic and terrestrial systems. Investigation of undisturbed media such as pore water requires a solid-state electrode, and such electrodes can be difficult to fabricate reproducibly. An approach to determine the concentrations of electroactive constituents using indirectly calibrated electrodes has been developed, but the protocol for and accuracy of this approach—the pilot ion method—has not been documented in detail.
A detailed procedure for testing electrode quality is provided, and the application and limitations of the pilot ion method have been documented. To quantify Fe(II) and Mn(II), subtraction of non-linear baseline functions from voltammetric signals produced better calibration curves than did linear baselines, enabled lower detection limits and reliable deconvolution of overlapping signals, and was successfully applied to sediment pore water signals. We observed that electrode sensitivities often vary by tens of percent, and that the sensitivity declines over time. The ratio of calibration slopes of Mn(II) to Fe(II) varied by no more than 11% from one Hg/Au electrode to another and Fe(II) concentrations predicted by the Mn(II) pilot ion were, on average, 13% different from their actual values. However, concentration predictions by the pilot ion method were worse for less than 15 μM Fe(II) (46% different on average). The ratio of calibration slopes of Mn(II) to S(−II) varied by almost 20% from one Hg/Au electrode to another, and S(−II) predicted concentrations were as much as 58% different from their actual values. These predictions of Fe(II) and S(−II) concentrations indicate that the accuracy of the pilot ion method depends on how independent calibration slope ratios are from the electrode used. At medium-to-high concentration for the ocean, naturally derived dissolved organic carbon did not significantly affect the baseline-corrected electrode response of Mn(II) and Fe(II), but did significantly affect the response of S(−II).
Despite their intrinsic variability, Hg/Au electrodes fabricated by hand can be used to quantify O2, S(−II), Fe(II), and Mn(II) without calibrating every electrode for every constituent of interest. The pilot ion method can achieve accuracies to within 20% or less, provided that the underlying principle—the independence of slope ratios—is demonstrated for all voltammetric techniques used, and effects of the physicochemical properties of the system on voltammetric signals are addressed through baseline subtraction.
Reduction-oxidation (redox)-active species are integral to microbially mediated contaminant, nutrient, and carbon transformations in aquatic and terrestrial systems [1–6]. Electrochemical sensors have been routinely used over the past two decades to study redox processes and trace element speciation [7, 8]. Among the electrochemical methods used for this purpose, voltammetry is able to analyze multiple dissolved redox constituents with no or minimum sample manipulation [9, 10]. The pioneering work of Brendel and Luther  and Tercier and Buffle [10, 12] in the development and use of microelectrodes has enabled in situ voltammetric analysis of redox constituents and trace metals in natural systems [13–17]. While progress has been made and researchers continue to voltammetrically analyze aquatic systems with microelectrodes, there is much room for growth in the number of users and deployments . One impediment to growth is poor understanding of how to convert voltammetric signals to analyte concentrations. Although numerous papers show the quantitative results of such conversions, the process is not intuitive and could be explained in more detail and with specific application to natural samples. To help more environmental scientists use voltammetry to decipher microbial processes that mediate redox conditions in undisturbed soils and sediments, this paper shows how to quantify the voltammetric signals of O2, Mn(II), Fe(II), and S(−II) and overcome challenges imposed by the use of solid-state electrodes.
To achieve quantitative in situ measurements, one needs to overcome the imprecision of solid-state electrodes when fabricated and possible alteration of the sensors by chemical or biological agents, which cause initial differences in and instability of their analytical sensitivity. For reasons explained by Buffle and Tercier-Waeber , mercury (Hg) is the sensor material of choice for redox analysis of environmental systems. Hg-amalgam [11, 19] or Hg-film [12, 20–22] sensors are used to construct electrodes that meet additional requirements for in situ measurements such as insertion into porous media. The focus of this paper is on a 1 mm-diameter glass electrode equipped with a 100 μm Hg/Au amalgam sensor.
Once a Hg/Au amalgam electrode is made, there are at least three criteria by which to evaluate the quality of the Hg/Au amalgam sensor. First, in oxygenated solution, a Hg-based electrode should yield (a) an elongated, S-shaped current-potential relationship following O2 and H2O2 reduction and (b) alkali metal reduction at −1.7 V vs. Ag/AgCl or lower. When Au is plated well by Hg, alkali metal (typically Na+) reduction shifts to potentials well below −1.55 V such that a sufficient overpotential can be applied to measure Mn(II) [10, 19, 23]. The overpotential is defined as “the additional potential (beyond the thermodynamic requirement) needed to drive a reaction at a certain rate” .
The second criterion pertains to the current signal measured in a deaerated solution containing no electroactive constituents within the range over which Hg can be electrically polarized. At circumneutral pH, this range extends from about −0.05 V vs. Ag/AgCl, above which Hg is oxidized, down to about −1.7 V, below which alkali metals are reduced . This range encompasses the redox potentials of O2, S(−II), Fe(II), and Mn(II). The electrode-water interface acts experimentally like a capacitor . When polarized, an electrode surface accumulates charge and electrostatically retains an excess of aqueous cations or anions (the point of zero charge of a Hg electrode is around −0.5 V vs. Ag/AgCl [25, 26]). A current flows during this process and is called the charging or capacitative current [9, 24], which is measured continuously as the structure of the electrode-water interface evolves with the change in potential during a voltammetric scan. This capacitative current is also sometimes described as being nonfaradaic, in that it does not arise from electron transfer to or from an electroactive constituent. To quantify electroactive constituents, the faradaic signal needs to be isolated from the nonfaradaic signal. After each potential step or modulation in the voltammetric scan, both the capacitative and faradaic currents spike, but the capacitative current decays faster . If scan rates are lower, the potentiostat can wait longer to sample the current after a potential modulation, thereby allowing the capacitative current to decline, enhancing the proportion of faradaic current. Although nonfaradaic signal can be minimized instrumentally, it cannot be eliminated entirely. An artificial function is typically fit to the nonfaradaic signal and then subtracted. A quality Hg/Au sensor will have a small enough capacitative current relative to the faradaic signal to allow for this subtraction.
The third criterion for Hg/Au sensor quality is whether the faradaic current response of the electrode associated with electron transfer induced by oxidation or reduction of aqueous constituents is within a normal range. Using Mn(II) to evaluate this third criterion also helps check the first criterion because its reduction occurs near the polarization limit of Hg.
Unless polarized at a reducing potential, Hg-based electrodes will oxidize and lose sensitivity after approximately ten hours , although this time can vary depending on the sensor material and size [12, 21]. Regardless, the time between electrode calibration and quantitative measurements is finite. Analytical sensitivities need to be adjusted if measurements occur beyond this point. Several electrodes are most likely needed to measure redox constituents at multiple locations in situ, posing an additional burden of knowing and adjusting several constituents’ calibrations on each electrode.
A typical use of the pilot ion method is described in the Additional file 1.
Explicit means by which voltammetric signals can be quantified using the pilot ion method have not been provided in the literature. In several publications [11, 13, 16, 29–33], graphical representations of how O2, Fe(II), Mn(II), and S(−II) voltammetric signals are distinguished from background are not shown or explained. And for the pilot ion method, these works reference Meites  or Brendel and Luther , which in turn cites Meites . The only example provided by Meites  to substantiate the pilot ion method involved metal analysis at environmentally exaggerated concentrations using a mercury drop (not solid-state) electrode.
In the present study, we show in detail how to quantify the voltammetric signals of O2, Fe(II), Mn(II), and S(−II) and evaluate the accuracy of the pilot ion method. By comparing nine replicate Fe(II) and Mn(II) and twelve replicate S(−II) calibration curves on each of three electrodes, the accuracy of calibrations using the pilot ion method are provided for the analysis of seawater with and without naturally derived dissolved organic carbon (DOC) at two DOC concentrations. The approach is then demonstrated with multi-constituent aqueous solutions and sediment pore waters. In addition to providing a reproducible protocol, this paper is the first to evaluate the accuracy of the pilot ion method under realistic environmental conditions.
Materials and methods
Artificial seawater  was made with reagent-grade chemicals and deionized water and stored at 4°C. Na2S stock solutions were prepared by dissolving rinsed Na2S·9H2O crystals in deoxygenated water, followed by iodimetric standardization. Water was deoxygenated by boiling, purging with ultrahigh-purity (UHP) N2 and storing in a 98% N2, 2% H2 anaerobic chamber (Coy Labs; industrial grade chamber gases were deoxygenated using a Pd catalyst). Fe(ClO4)2·6H2O (Alfa Aesar; desiccated under N2) was dissolved in deoxygenated deionized water. To remove any Fe(III) originating from the dry reagent, the iron stock solution was filtered once daily three times through a 0.02 μm Anotop membrane. Iron stock solutions (pH 2.8) were stored in the glove box and were re-filtered and colorimetrically standardized  on the day of each experiment to remove Fe(III) that inevitably forms despite the low pH of the solution and trace (few ppm) levels of oxygen that can persist inside the anaerobic chamber. Stock solutions made with newly purchased ferrous chloride were found to contain much higher levels of Fe(III) contaminant compared to Fe(ClO4)2·6H2O. While others favor using Mohr’s salt to prepare stock solutions , Fe(ClO4)2·6H2O is also a reasonable choice  and one that we found lost only 3% Fe(II) or less over 1 d presumably due to filtration of oxidized precipitates. MnSO4 was weighed on an analytical balance, dissolved in deionized water, and stored at 4°C. Pony Lake fulvic acid (PLfa) was purchased from the International Humic Substances Society. Fulvic acid from Pacific Ocean water collected from 100 m depth, 170 km southwest of Honolulu, Hawaii was isolated as described by Aiken et al. . Both aquatic organic isolates were weighed using a microbalance in a desiccated chamber, dissolved in artificial seawater, and stored at 4°C. The Pacific Ocean isolate was dissolved to 0.850 (mg C)/L, which is within the medium-to-high range of organic C in the ocean , whereas PLfa, derived mostly from algal biomass , was dissolved at 26.3 (mg C)/L to resemble conditions in biologically productive estuarine sediment.
Electrodes were fabricated by a procedure modified from Brendel and Luther . 100 μm Au wire (Surepure Chemetals) was soldered to shielded copper cable and sealed with Epo-tek® 360 in a glass tube with a tip pulled from 5 mm to approximately 0.8 mm diameter. The tip was flattened by sanding with 400-grit and successively polished by hand using Buehler Metadi 15, 6, 1, and 0.25 μm diamond polishing compounds suspended in AB Metadi lubricant on Buehler TexMet® (15 and 6 μm) and MicroCloth® (1 and 0.25 μm) affixed to a rotating pedestal mounted on a DC motor equipped with an HY152A AC/DC converter. All electrodes were inspected with a 100x microscope throughout the polishing process, verifying a mirror finish before electroplating with Hg. Without delay, the polished electrode tip was rinsed and placed in a 0.1 M Hg(NO3)2/0.05 M HNO3 solution deaerated with UHP N2. Hg was electroplated onto the Au disc at −0.1 V vs. Ag/AgCl/[saturated KCl] for 4 min. The plated electrode was then polarized at −9 V in 1 M NaOH for 90 s , rinsed with water and 0.01 M HClO4 to remove hydroxide, and stored overnight in deionized water. Unless stated otherwise, electrodes were (re)plated, polarized, and stored in deionized water one day prior to performing calibration measurements.
Voltammetric analysis of model seawaters
All analyses were performed at 21 ± 2°C. An Autolab PGSTAT12 potentiostat was used for all voltammetric measurements (Metrohm-Autolab B.V. (formerly Eco Chemie), Utrecht, The Netherlands). By convention, reduction currents are negative while oxidation currents are positive, which is the opposite of some systems . Data acquisition, first-derivative calculations, linear and 4th-order polynomial fitting, baseline subtractions, and peak analyses were performed with Nova versions 1.6 and 1.7 (Metrohm-Autolab). Non-linear baseline functions were fit to the background signal by choosing three data points on the volammogram on the electropostive side of the region of interest and two points on the negative side. A 4th-order polynomial that intersects those five points is then fit. Regression of calibration data, peak deconvolution, statistical analyses (analysis of variance (ANOVA) and chi-squared calculations), and figure data plotting were performed with Igor Pro (Wavemetrics, Portland, Oregon, USA). The electrochemical cell consisted of a 100 μm Hg/Au amalgam working electrode, Ag/AgCl/[saturated KCl] reference electrode, and Pt auxiliary electrode placed in a three-port, 250 mL glass flask that resided inside a polyethylene glove bag (Glas-Col) mounted within an electrically grounded copper screen (Faraday cage) to shield the system from electrical noise. The glove bag was left open to the air to perform cyclic voltammetric analysis of aerated seawater. The system was then deaerated by purging the glove bag under positive pressure with UHP N2 for at least 1 h. Unless stated otherwise, S(−II), Fe(II), and Mn(II) stock solutions were added individually to the flask inside the glove bag under continuous UHP N2 flow under positive pressure. The pH of artificial seawater was 8.2 ± 0.1 throughout MnSO4 addition. To slow Fe(II) oxidation, seawater pH was adjusted prior to and maintained during the Fe(ClO4)2 additions at 6.0 ± 0.5 with CO2 and N2 flow through the glove bag. Acidification is permissible because pH does not affect the electron transfer kinetics from the electrode to Fe(II) [10, 11]. However, the solution should not be acidified too much, because a H+ reduction signal may obscure any Fe(II) signal . Solutions were analyzed after each standard addition with three Hg/Au electrodes in series using a multiplexer (two electrodes were disconnected while the third was used to measure).
S(−II) was measured using normal pulse voltammetry, where the potential (E) was scanned from −0.8 to −0.4 V in 0.5 s, 0.005 V steps, with a −0.8 V base potential and 0.05 s pulse time (for further explanation, see Turner et al.  and section 7.3.2 in Bard and Faulkner ). O2 was measured with cyclic voltammetry (five scans from −0.1 to −1.7 V and back at 0.1 V/s in −7 mV steps, 0.07 s interval time). Fe(II) and Mn(II) were measured by three techniques: first, cyclic voltammetry (five scans from −0.1 to −1.7 V and back at 0.1 V/s in −7 mV steps), then linear sweep voltammetry (−0.8 to −1.7 V at −0.1 V/s in −7 mV steps), and finally square-wave voltammetry from −0.8 to −1.7 V in −5 mV steps and 25 mV amplitude pulses at 8 Hz frequency (0.04 V/s; 0.125 s interval time). The electrode was held at −0.4 V for 10 s prior to cyclic voltammetry or −0.8 V for 10 s prior to linear sweep and square-wave voltammetry. The uppermost potential (−0.1 V) was chosen to prevent oxidation of the Hg/Au sensor. We chose a scan rate of 0.1 V/s and a step potential of 7 mV for cyclic and linear sweep scans and 8 Hz for square-wave scans to ensure that the measurement interval time was long enough for the potentiostat to have sufficient bandwidth (100 Hz) in the low current range to accurately measure the 1 to 10 nA currents typically sensed by the 100 μm-diameter Hg/Au electrodes.
Voltammetric analysis of sediment
To test the electrodes in natural sediment and evaluate data processing approaches with signals measured in complex media, several sediment manipulations were performed in the laboratory. Surface sediments (top 0–2 cm) from a salt pond in South San Francisco Bay, CA (SPA8S3) and a coastal marsh in Mississippi (MSM2) were collected and stored in completely filled mason jars, chilled on wet ice, brought to the laboratory for further sub-sampling under controlled anoxic conditions , and refrigerated at 4°C to retard microbial processes and prevent sample degradation until use. MSM2 sediment contained 57% solids, consisting of 70% silt and clay (<64 μm), 5.6% organic matter (LOI), and 15 μmol chromium-reducible sulfur (CRS) per gram of wet sediment. SPA8S3 sediment contained 47% solids, of which 70% were silts and clays, 13% organic matter, 61 μmol/g CRS, and 70 mg/L dissolved organic carbon in the pore water. Sediments were mixed with deaerated seawater in a three-port, 250 mL glass flask. The electrode tip was placed within the pore space of the sediment. A set of voltammetric scans was performed in the anoxic pore water. The sediment was then oxygenated by injecting aerated seawater followed by periodic measurements.
Results and discussion
Preliminary electrode evaluation
Quantification of O2, S(−II), Fe(II), and mn(II) in model seawater solutions and sediments
Hg/Au electrode sensitivities (calibration slopes) measured in artificial seawater
Sensitivity (10-2 nA/μM)
Artificial seawater (no organic carbon)
Artificial seawater with 0.85 (mg C)/L Pacific Ocean fulvic acid
Artificial seawater with 26 (mg C)/L Pony Lake fulvic acid
Comparison of calibration slope ratios for Mn(II) and Fe(II) for different electrodes
K = sMn/sFe
K = sMn/sS(-II)
K = sMn/sOxygen
average (all K values)
95% confidence interval
Analyses of Fe(II)- or Mn(II)-bearing seawater indicate that subtracting a linear baseline to obtain ip for Fe(II) or Mn(II) is error-prone and of limited use, while 4th-order polynomial subtractions provide analytical sensitivity similar to what is determined in simpler, one-component systems. Results, of which examples are provided by the red symbols in Figure 6d and 4d, consistently indicated that linear baselines produced appreciably offset Fe(II) and Mn(II) calibration curves and precluded detection of less than approximately 50 μM Fe(II) and 20 μM Mn(II). By contrast, Fe(II) and Mn(II) signals as low as 6 μM were detectable, and 10 μM quantifiable, upon subtraction of 4th-order polynomials, as shown by the blue symbols in Figure 6d and 4d. These results suggest that linear baselines may be useable for solutions containing at least 20 to 50 μM Mn(II) or Fe(II), but the extra effort to fit a 4th-order polynomial may be justified to quantify Fe(II) and Mn(II) below those concentrations.
On three separate occasions, Fe(II) and Mn(II) calibration curves were measured with three electrodes in succession. The following comparisons refer to the analytical sensitivities (calibration slopes) provided in Table 1 and do not denote statistical variability, as three electrodes do not constitute a statistical population. On a given day, sensitivities varied from electrode to electrode (relative to the average) by 13 to 76% for Fe(II) (40% on average with a median of 43%), and 4 to 76% for Mn(II) (24% on average with a median of 11%). On four separate occasions, S(−II) calibration curves were measured with three electrodes in succession. On a given day, sensitivities varied from electrode to electrode by 0.3 to 82% for S(−II) (37% on average with a median of 26%). Since all of these electrodes met the three quality criteria explained earlier, this degree of analytical variability reflects the difficulty of reproducing electrode surfaces by hand.
To restore sensitivity, electrodes were periodically replated with Hg and polarized at −9 V for 90 s. In one experiment, Mn(II) and S(−II) were measured separately with three electrodes at several concentrations in seawater, after which the electrodes were placed in deionized water in open circuit. A day later, these electrodes were replated, polarized, stored overnight in deionized water, and then used again to separately measure Mn(II) and S(−II) in seawater. The sensitivity of each electrode to Mn(II) changed by 10, 27, and 58% from the values measured two days before. The sensitivity of each electrode to S(−II) changed by 5, 31, and 76%. In most cases, the reconditioning procedure did not restore the electrode to its original sensitivity.
If possible, sensitivity can be better maintained by continually polarizing an electrode to prevent oxidation of the Hg/Au sensor. After measuring aerated seawater without added dissolved organic carbon, three electrodes were polarized at −0.8 V ten minutes at a time in series (using a multiplexer) for 16 hours. They were then used to measure cyclic voltammograms in aerated seawater amended with 0.85 mg C/L as Pacific Ocean fulvic acid. For two electrodes, Δi for O2 decreased very little: 30.7 to 30.2 nA and 21.3 to 20.2 nA; the Δi for the third electrode dropped more, from 48.3 to 29.0 nA (some small difference can be attributed to the addition of dissolved organic carbon). Two out of three of these results are consistent with an experiment reported by Luther et al. , in which Hg/Au electrodes were stable for as long as two months when polarized at potentials sufficiently negative to prevent Hg/Au oxidation. It is important to point out that Luther et al. continuously polarized their electrodes, whereas electrodes were continually polarized 10 minutes out of 30 in this study. Polarizing electrodes even continually may be unfeasible if a potentiostat is used for other purposes or under logistical constraints like traveling to a field site. In addition to correcting electrode-to-electrode variability caused by hand fabrication, the pilot ion method would be useful to account for changes in sensitivity due to reconditioning or oxidation.
Accuracy of the pilot ion method
To determine O2, S(−II), Fe(II), and Mn(II) concentrations with uncalibrated electrodes, we first verified that the current response for each of these constituents is directly proportional to their concentration (Figure 3, Figure 6, Figure 4, and Figure 11). Next, we needed to determine if the quotient of analytical sensitivities between the pilot ion and constituent of interest (K in equation 1) is, in fact, independent of the electrode used. Contemporaneously measured calibration slopes for Mn(II) and Fe(II) using square-wave voltammetry indicate that averaged 2.4 ± 0.3 (95% confidence interval); that is, at 95% confidence, the difference in response to Mn(II) and Fe(II) varies by no more than 11% from one Hg/Au electrode to another (Table 2). We conclude on the basis of this result that is independent of the Hg/Au electrode used, and so Mn(II) should be a reasonably accurate pilot ion for Fe(II). Mn(II) is a convenient pilot ion because it does not oxidize appreciably over experimental time scales and therefore is easier to measure.
Using current response data for O2 in air-saturated media and Mn(II) calibration slopes, averaged 0.36 ± 0.05 (95% confidence interval; Table 2); that is, the difference in response to Mn(II) and O2 varies by no more than 14% from one Hg/Au electrode to another. Although slightly more variable than , is arguably independent of the Hg/Au electrode used, and so Mn(II) should be a good pilot ion for O2. However, as mentioned earlier, most researchers calibrate each electrode directly for O2 by measuring a single air-saturated current response.
The slope ratio K is more variable for Mn(II) and S(−II) measured by square-wave (SW) and normal pulse voltammetry (NPV), respectively: averaged 0.28 ± 0.05 (95% confidence interval; Table 2); that is, at 95% confidence, the difference in electrode response to Mn(II) and S(−II) varies by almost 20% from one Hg/Au electrode to another. The higher variability in K for reflects differences in Hg/Au oxidation by S(-II) to form HgS and Mn(II) reduction and Mn(0) sorption/amalgamation with Hg/Au. For example, comparing one electrode to another, a 10% difference in the efficiency of Mn(II) reduction to Mn(0) apparently does not imply a 10% difference in the efficiency of Hg/Au oxidation by S(-II). As a result, the K values for each electrode can appreciably differ. In summary, depending on the demands of the user, Mn(II) should be a good pilot ion for Fe(II), somewhat less good for O2, and perhaps unacceptable for S(-II) quantification.
Comparison of actual Fe(II) concentrations to those predicted by the pilot ion method
K = 2.43
August 19, 2010
August 20, 2010
Comparison of actual S(-II) concentrations to those predicted by the pilot ion method
K = 0.284
August 19, 2010
August 20, 2010
Implications of different matrices
Temperature affects the diffusion of chemical species at the electrode-water interface, which can change the current by up to 10% per°C in the range 4–20°C [10, 21]. If temperatures of lab or field systems are expected to differ from the temperature at which calibration data were obtained, the voltammetric figure of merit needs to be determined at a given species concentration as a function of temperature. The voltammetric figure of merit is typically significantly correlated to temperature for electrochemically reversible (HgS/Hg0) and irreversible (O2/H2O2, Fe2+/Fe0, and Mn2+/Mn0) systems and therefore can be used to correct in situ data [54, 55]. The possible compensatory effect of lowered O2 and H2S solubility and higher rate of diffusion (and therefore larger diffusion-limited current) at higher temperatures should also be considered .
Guidelines for baseline fitting
With standard solution measurements, we fit baselines (see the Methods section), subtract them from the total signal, obtain ip for the constituent of interest, plot ip versus concentration, perform a linear regression, and evaluate the quality of the regression. Specifically, how well correlated is ip to concentration? There should be a high degree of correlation (R2 > 0.99) if baseline fits were done correctly. The linear regressions should include computation of the confidence interval of the ip-axis intercept at zero concentration. This offset should be essentially zero if baseline removal was done correctly. But if we do not know the concentration, how do we know if the baseline removal was done correctly?
To start, we compare a sample voltammogram with that measured in standards to get a rough idea of the concentration. Points where the baseline function must intersect (anchors) are then chosen at potentials similar to those chosen with a standard voltammogram. If the overall polynomial fit looks unreasonable, the points are shifted slightly until the fit looks reasonable. By reasonable, we mean that the baseline follows the background signal on either side of the faradaic current region and does not exhibit strange topology, such as inflections within the faradaic region (region A in Figure 19). Furthermore, after subtracting the baseline, the symmetry and coincidence of peaks to the redox potential of Fe(II) or Mn(II) provide further indication of the quality of the baseline fit (e.g., Figure 6, Figure 4, and Figure 11).
Regions of voltammograms at potentials more positive than where Fe(II) and Mn(II) reduction occurs tended to overlap (region B in Figure 19; see also Figure 6 and Figure 4). Therefore, regardless of concentration, we chose three common baseline points in region B (Figure 19) and still achieved good correlations. However, the two remaining necessary anchors, chosen on the more negative side of the reduction signal (region C in Figure 19), need to be shifted by about +/−15 mV per halving/doubling of Fe(II) or Mn(II) concentration to follow the negative shift of the reduction signal as concentration increases (Figure 8). If anchors are chosen too far from the faradaic region, ip will be overestimated. Good calibration curves were obtained when adjacent anchors were separated by about 20 mV. The distance ΔE1 in Figure 19 between the anchor nearest the more positive side of the reduction signal and the inflection point in the voltammogram indicating the onset of Fe(II) or Mn(II) reduction was typically 100 mV, while 40 mV typically separated the inflection point after the reduction signal and the nearest anchor on the negative side (ΔE2 in Figure 19).
As a sensitivity analysis, we adjusted the positions of anchors by 10 mV while maintaining what appeared to be a reasonable fit to the background signal. Such adjustment changed ip by ±30% at the 1 nA level, which translates for our electrodes to uncertainties of 25 to 40 μM Fe(II) and approximately 20 μM Mn(II). At higher concentrations, it is easier to visually optimize the baseline fit, and we found that the uncertainty could be limited to 10% or less.
Implications for redox characterization of aquatic systems
Voltammetric i – E relationships contain more information than potentiometric or amperometric signals, which enables the (nearly) simultaneous quantification of multiple redox constituents in minimally disturbed systems. In eliminating the need to repeatedly gather and process calibration data for every electrode used, the pilot ion method is useful for accurate redox characterization of aquatic systems with voltammetric solid-state electrodes. Aspects including electrode quality and data processing presented here will help new users more efficiently obtain quantitative results and understand their uncertainty. We hope this work will improve understanding of the exceptional capability of voltammetry to fill data gaps in understanding the economic and ecological function of terrestrial and aquatic systems.
Experimental research and initial drafting and revision of the manuscript were funded by a Venture Capital Grant to AJS from the U.S. Geological Survey (USGS). Further manuscript revisions were performed with support from the Subsurface Science Scientific Focus Area funded by the U.S. Department of Energy (DoE), Office of Science, Office of Biological and Environmental Research under Award Number DE-AC02-05CH11231. Associate Editor Prof. Gregory K. Druschel (Indiana University-Perdue University Indianapolis), Dr. Laurie S. Balistrieri (USGS), and anonymous reviewers are acknowledged for constructive comments that helped ensure that the material was presented clearly. The authors would like to thank the graduate students who read the manuscript to ensure that this paper will help future generations of geo-electrochemists. Use of trade names in this paper is for identification purposes only and does not constitute endorsement by the USGS or DoE.
- Windham-Myers L, Marvin-DiPasquale M, Krabbenhoft DP, Agee JL, Cox MH, Heredia-Middleton P, Coates C, Kakouros E: Experimental removal of wetland emergent vegetation leads to decreased methylmercury production in surface sediment. J Geophys Res-Biogeo. 2009, 114: G00C05-View ArticleGoogle Scholar
- Marvin-DiPasquale M, Lutz MA, Brigham ME, Krabbenhoft DP, Aiken GR, Orem WH, Hall BD: Mercury cycling in stream ecosystems. 2. Benthic methylmercury production and bed sediment pore water partitioning. Environ Sci Technol. 2009, 43: 2726-2732.View ArticleGoogle Scholar
- Klump JV, Fitzgerald SA, Waples JT: Benthic biogeochemical cycling, nutrient stoichiometry, and carbon and nitrogen mass balances in a eutrophic freshwater bay. Limnol Oceanogr. 2009, 54: 692-712.View ArticleGoogle Scholar
- Dale AW, Van Cappellen P, Aguilera DR, Regnier P: Methane efflux from marine sediments in passive and active margins: Estimations from bioenergetic reaction-transport simulations. Earth Planet Sci Lett. 2008, 265: 329-344.View ArticleGoogle Scholar
- Kostka JE, Roychoudhury A, Van Cappellen P: Rates and controls of anaerobic microbial respiration across spatial and temporal gradients in saltmarsh sediments. Biogeochemistry. 2002, 60: 49-View ArticleGoogle Scholar
- Thullner M, Regnier P, Van Cappellen P: Modeling microbially induced carbon degradation in redox-stratified subsurface environments: Concepts and open questions. Geomicrobiol J. 2007, 24: 139-155.View ArticleGoogle Scholar
- Reimers CE: Applications of microelectrodes to problems in chemical oceanography. Chem Rev. 2007, 107: 590-600.View ArticleGoogle Scholar
- Buffle J, Horvai G: In Situ monitoring of Aquatic Systems. 2000, John Wiley & Sons, New YorkGoogle Scholar
- Buffle J: Complexation Reactions in Aquatic Systems: An Analytical Approach. 1988, Ellis Horwood Limited, ChicesterGoogle Scholar
- Buffle J, Tercier-Waebe ML: In Situ Monitoring of Aquatic Systems Chemical Analysis and Speciation. Edited by: Buffle J, Horvai G. 2000, John Wiley & Sons, Chichester, 279-405. trade paper edn., 2000, vol. 6, ch. 9Google Scholar
- Brendel PJ, Luther GW: Development of a gold amalgam voltammetric microelectrode for the determination of dissolved Fe, Mn, O2, and S(−II) in porewaters of marine and freshwater sediments. Environ Sci Technol. 1995, 29: 751-761.View ArticleGoogle Scholar
- Tercier ML, Parthasarathy N, Buffle J: Reproducible, reliable and rugged Hg-Plated Ir-based microelectrode for in situ measurements in natural waters. Electroanalysis. 1995, 7: 55-63.View ArticleGoogle Scholar
- Luther GW, Glazer BT, Ma SF, Trouwborst RE, Moore TS, Metzger E, Kraiya C, Waite TJ, Druschel G, Sundby B, Taillefert M, Nuzzio DB, Shank TM, Lewis BL, Brendel PJ: Use of voltammetric solid-state (micro)electrodes for studying biogeochemical processes: Laboratory measurements to real time measurements with an in situ electrochemical analyzer (ISEA). Mar Chem. 2008, 108: 221-235.View ArticleGoogle Scholar
- Druschel G, Emerson D, Sutka R, Suchecki P, Luther GWI: Low-oxygen and chemical kinetic constraints on the geochemical niche of neutrophilic iron(II) oxidizing microorganisms. Geochim Cosmochim Acta. 2008, 72: 3358-3370.View ArticleGoogle Scholar
- Tercier-Waeber ML, Hezard T, Masson M, Schafer J: In situ monitoring of the diurnal cycling of dynamic metal species in a stream under contrasting photobenthic biofilm activity and hydrological conditions. Environ Sci Technol. 2009, 43: 7237-7244.View ArticleGoogle Scholar
- Taillefert M, Neuhuber S, Bristow G: The effect of tidal forcing on biogeochemical processes in intertidal salt marsh sediments. Geochem Trans. 2007, 8: http://www.geochemicaltransactions.com/content/8/1/6,Google Scholar
- Taillefert M, Hover VC, Rozan TF, Theberge SM, Luther GW: The influence of sulfides on soluble organic-Fe(III) in anoxic sediment porewaters. Estuaries. 2002, 25: 1088-1096.View ArticleGoogle Scholar
- Tercier-Waeber ML, Confalonieri F, Koudelka-Hep M, Dessureault-Rompré J, Graziottin F, Buffle J: Gel-Integrated Voltammetric Microsensors and Submersible Probes as Reliable Tools for Environmental Trace Metal Analysis and Speciation. Electroanalysis. 2008, 20: 240-258.View ArticleGoogle Scholar
- Billon G, van den Berg CMG: Gold and silver micro-wire electrodes for trace analysis of metals. Electroanalysis. 2004, 16: 1583-1591.View ArticleGoogle Scholar
- Belmont C, Tercier ML, Buffle J, Fiaccabrino GC, Koudelka-Hep M: Mercury-plated iridium-based microelectrode arrays for trace metals detection by voltammetry: optimum conditions and reliability. Anal Chim Acta. 1996, 329: 203-214.View ArticleGoogle Scholar
- Belmont-Hébert C, Tercier ML, Buffle J, Fiaccabrino GC, de Rooij NF, Koudelka-Hep M: Gel-Integrated Microelectrode Arrays for Direct Voltammetric Measurements of Heavy Metals in Natural Waters and Other Complex Media. Anal Chem. 1998, 70: 2949-2956.View ArticleGoogle Scholar
- Tercier ML, Buffle J: Antifouling membrane-covered voltammetric microsensor for in situ measurements in natural waters. Anal Chim Acta. 1996, 68: 3670-3678.View ArticleGoogle Scholar
- Bristow G, Taillefert M: VOLTINT: A Matlab-based program for semi-automated processing of geochemical data acquired by voltammetry. Comput Geosci. 2008, 34: 153-162.View ArticleGoogle Scholar
- Bard AJ, Faulkner LR: Electrochemical Methods, Fundamentals and Applications. 2004, John Wiley & Sons, New YorkGoogle Scholar
- Cominoli A, Buffle J, Haerdi W: Voltammetric study of humic and fulvic substances: Part III. Comparison of the capabilities of the various polarographic techniques for the analysis of humic and fulvic substances. J Electroanal Chem. 1980, 110: 259-275.View ArticleGoogle Scholar
- Buffle J, Cominoli A: Voltammetric study of humic and fulvic substances Part IV. Behaviour of fulvic substances at the mercury-water interface. J Electroanal Chem. 1981, 121: 273-299.View ArticleGoogle Scholar
- Kounaves SP, Buffle J: An iridium-based mercury-film electrode: Part I. Selection of substrate and preparation. J Electroanal Chem. 1987, 216: 53-69.View ArticleGoogle Scholar
- Meites L: Polarographic Techniques. 1965, John Wiley & Sons, New YorkGoogle Scholar
- Bull D, Taillefert M: Seasonal and topographic variations in porewaters of a southeastern USA salt marsh as revealed by voltammetric profiling. Geochem Trans. 2001, 2: 104-View ArticleGoogle Scholar
- Lewis BL, Glazer BT, Montbriand PJ, Luther Iii GW, Nuzzio DB, Deering T, Ma S, Theberge S: Short-term and interannual variability of redox-sensitive chemical parameters in hypoxic/anoxic bottom waters of the Chesapeake Bay. Mar Chem. 2007, 105: 296-308.View ArticleGoogle Scholar
- Croteau M-N, Misra SK, Luoma SN, Valsami-Jones E: Silver Bioaccumulation Dynamics in a Freshwater Invertebrate after Aqueous and Dietary Exposures to Nanosized and Ionic Ag. Environ Sci Technol. 2011, 45: 6600-6607.View ArticleGoogle Scholar
- Bura-Nakic E, Krznaric D, Helz GR, Ciglenecki I: Characterization of Iron Sulfide Species in Model Solutions by Cyclic Voltammetry. Revisiting an Old Problem. Electroanalysis. 2011, 23: 1376-1382.View ArticleGoogle Scholar
- Sundby B, Caetano M, Vale C, Gobeil C, Luther GW, Nuzzio DB: Root-induced cycling of lead in salt marsh sediments. Environ Sci Technol. 2005, 39: 2080-2086.View ArticleGoogle Scholar
- Kester D, Duedall I, Connors D, Pytkowicz R: Preparation of artificial seawater. Limnol Oceanogr. 1967, 12: 176-179.View ArticleGoogle Scholar
- Stookey LL: Ferrozine-A new spectrophotometric reagent for iron. Anal Chem. 1970, 42: 779-781.View ArticleGoogle Scholar
- Rickard D: Kinetics of FeS precipitation: Part 1. Competing reaction mechanisms. Geochim Cosmochim Acta. 1995, 59: 4367-4379.View ArticleGoogle Scholar
- Stumm W, Lee GF: Oxygenation of ferrous iron. Ind Eng Chem. 1961, 53: 143-146.View ArticleGoogle Scholar
- Aiken GR, McKnight DM, Thorn KA, Thurman EM: Isolation of hydrophilic organic acids from water using nonionic macroporous resins. Org Geochem. 1992, 18: 567-573.View ArticleGoogle Scholar
- Hansell DA, Carlson CA, Repeta DJ, Schlitzer R: Dissolved organic matter in the ocean-A controversy stimulates new insights. Oceanography. 2009, 22: 202-211.View ArticleGoogle Scholar
- Brown A, McKnight DM, Chin Y, Roberts EC, Uhle M: Chemical characterization of dissolved organic material in Pony Lake, a saline coastal pond in Antarctica. Mar Chem. 2004, 89: 327-337.View ArticleGoogle Scholar
- Stojek Z, Kublik Z: Silver based mercury film electrode I. General characteristics and stability of the electrode. J Electroanal Chem. 1975, 60: 349-358.View ArticleGoogle Scholar
- Turner JA, Abel RH, Osteryoung RA: Normal pulse polarographic analysis based on mercury anodization: Sulfide and iodide. Anal Chem. 1975, 47: 1343-1347.View ArticleGoogle Scholar
- Lutz MA, Brigham ME, Marvin-DiPasquale M: Procedures for collecting and processing streambed sediment and pore water for analysis of mercury as part of the National Water-Quality Assessment Program.U.S. Geological Survey Open-File Report. 2008, , , 1279-2008.Google Scholar
- Glazer BT, Marsh AG, Stierhoff K, Luther GW: The dynamic response of optical oxygen sensors and voltammetric electrodes to temporal changes in dissolved oxygen concentrations. Anal Chim Acta. 2004, 518: 93-100.View ArticleGoogle Scholar
- Guld RN, Gundersen JK, Ramsing NB: In Situ Monitoring of Aquatic Systems Chemical Analysis and Speciation. Edited by: Buffle J, Horvai G. 2000, John Wiley & Sons, Chichester, 19-74. vol. 6, ch. 2Google Scholar
- Williams KH, Long PE, Davis JA, Wilkins MJ, N'Guessan AL, Steefel CI, Yang L, Newcomer D, Spane FA, Kerkhof LJ, McGuinness L, Dayvault R, Lovley DR: Acetate Availability and its Influence on Sustainable Bioremediation of Uranium-Contaminated Groundwater. Geomicrobiol J. 2011, 28: 519-539.View ArticleGoogle Scholar
- Davison W, Gabbutt CD: Polarographic methods for measuring uncomplexed sulfide ions in natural waters. J Electroanal Chem. 1979, 99: 311-320.View ArticleGoogle Scholar
- Krznarić D, Ciglenećki-Jušić I: Electrochemical processes of sulfide in NaCl electrolyte solutions on mercury electrode. Electroanalysis. 2005, 17: 1317-1324.View ArticleGoogle Scholar
- Luther GW, Brendel PJ, Lewis BL, Sundby B, Lefrancois L, Silverberg N, Nuzzio DB: Simultaneous Measurement of O2, Mn, Fe, I-, and S(−II) in Marine Pore Waters with a Solid-State Voltammetric Microelectrode. Limnol Oceanogr. 1998, 43: 325-333.View ArticleGoogle Scholar
- Ferdelman TG, Church TM, Luther GWI: Sulfur enrichment of humic substances in a Delaware salt marsh sediment core. Geochim Cosmochim Acta. 1991, 55: 979-988.View ArticleGoogle Scholar
- Chanudet V, Filella M, Quentel F: Application of a simple voltammetric method to the determination of refractory organic substances in freshwaters. Anal Chim Acta. 2006, 569: 244-249.View ArticleGoogle Scholar
- NIST/SEMATECH: Bartlett's Test. http://www.itl.nist.gov/div898/handbook/eda/section3/eda357.htm, Accessed March 14, 2012, 2012
- Slowey AJ: Rate of formation and dissolution of mercury sulfide nanoparticles: The dual role of natural organic matter. Geochim Cosmochim Acta. 2010, 74: 4693-4708.View ArticleGoogle Scholar
- Belmont-Hebert C, Tercier ML, Buffle J, Fiaccabrino GC, de Rooij NF, Koudelka-Hep M: Gel-integrated microelectrode arrays for direct voltammetric measurements of heavy metals in natural waters and other complex media. Anal Chem. 1998, 70: 2949-2956.View ArticleGoogle Scholar
- Boudreau BP: A method-of-lines code for carbon and nutrient diagenesis in aquatic sediments. Comput Geosci. 1996, 22: 479-496.View ArticleGoogle Scholar
This article is published under license to BioMed Central Ltd. This is an Open Access article distributed under the terms of the Creative Commons Attribution License (http://creativecommons.org/licenses/by/2.0), which permits unrestricted use, distribution, and reproduction in any medium, provided the original work is properly cited.