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Abiotic Reduction of Chlorate by Fe(II) Minerals: Implications for Occurrence and Transformation of Oxy-Chlorine Species on Earth and Mars Maeghan Brundrett, Weile Yan, Maria C Velazquez, Balaji Rao, and William Andrew Jackson ACS Earth Space Chem., Just Accepted Manuscript • DOI: 10.1021/ acsearthspacechem.8b00206 • Publication Date (Web): 20 Feb 2019 Downloaded from http://pubs.acs.org on February 21, 2019

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ACS Earth and Space Chemistry

Abiotic Reduction of Chlorate by Fe(II) Minerals: Implications for Occurrence and Transformation of Oxy-Chlorine Species on Earth and Mars

Maeghan Brundrett†, Weile Yan†, Maria C. Velazquez†, Balaji Rao† and W. Andrew Jackson*† †Department of Civil, Environmental and Construction Engineering, Texas Tech University, Lubbock, Texas 79409, United States *Corresponding

author phone number: (806)834-6575; email: [email protected]

ACS Paragon Plus Environment

ACS Earth and Space Chemistry 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Abstract Recent investigations have reported a widespread occurrence of chlorate (ClO3-) and perchlorate (ClO4-) throughout the solar system including terrestrial arid environments. ClO3- and ClO4- are deposited/accumulated at an approximate equal molar ratio with some exceptions such as the Antarctica dry valley soils (MDV) and perhaps Martian surface material, where ClO4- is the dominate ClOx- species. All known ClO4- production mechanisms produce molar ratios of ClO3- : ClO4- equal to or much greater than 1, suggesting that reduced ratios may be due to post depositional mechanism(s). The objective of this study was to investigate potential iron-mediated abiotic reduction of ClO3-, similar to transformation mechanisms reported for nitrate (NO3-) by Fe (II) minerals. Three types of Fe (II)-containing minerals wustite (FeO), siderite (FeCO3), and sulfate green rust (GRSO42-) were investigated in completely mixed batch reactors as potential ClO3- reductants at a range of pH (4 - 9) and iron mineral concentrations (1 to 10 g/L). ClO3- was stoichiometrically reduced to chloride (Cl-) by wustite, siderite, and green rust, but no transformation occurred by dissolved Fe (II). Wustite and green rust reduced NO3- but not by siderite. When both NO3- and ClO3- are reduced simultaneously, ClO3- is reduced preferentially to NO3-, although the effect is somewhat concentration dependent. Increased background salt concentration (NaCl) increased ClO3- reduction but decreased NO3-. The stability of ClO3- and subsequent impacts on the ratio of ClO3- : ClO4- in the environment have implications for understanding the cycling of oxyanions and stability of iron minerals and related to this, the ratio of ClO4- and ClO3- may be an indicator of the past availability of free water. On Mars, these reactions may help to explain the unusually high concentrations of ClO4- compared to ClO3- and NO3-.

Keywords: (per)chlorate, wustite, green rust, siderite, nitrate, reduction, Martian soil chemistry, Antarctica Dry Valleys


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Introduction Chlorate (ClO3-) and perchlorate (ClO4-) are ubiquitously present in arid and semi-arid terrestrial environments 1, 2, 3, extraterrestrial systems (i.e. Mars and lunar regolith) 4, 5, 6, 7 and in meteorite samples 7, 8.

On Earth, chloride (Cl-) is the dominant chlorine species, while the minor ClO4- and ClO3- components

play an important role in the biogeochemical cycling of chlorine in the environment. Terrestrial mechanisms for the production of oxy-chlorine anions (ClOx-) are predominantly controlled by stratospheric oxidation reactions involving ozone (O3), based on 36Cl and Δ17O content, and possibly UV mediated photo-oxidation 2, 9, 10, 11, and are deposited by wet and dry deposition 1, 12. In the absence of liquid water, regardless of the starting material or production mechanism (whether photochemical or ozone driven, atmospheric or surface dominated oxidation), ClO3- and ClO4- are produced on a near equal molar basis. In aqueous solutions, the yield of produced ClO3- is orders of magnitude higher than ClO4- 9, 13- 19.

Terrestrially, ClO4- and ClO3- occur ubiquitously at roughly equimolar quantities, with

concentrations above background only occurring in semi-arid or arid environments where accumulation of deposition occurs 1, 11, 14. However, notable exceptions to this relationship exist, such as in the Antarctica McMurdo Dry Valley (MDV), an arid, cold environment previously studied for its suitability as a Martian analogue where the reported ClO3- : ClO4- ratios in surface soils are 1 but similar to terrestrial arid areas 7. Comparatively, ClO3- is less stable than ClO4- and unlike ClO4-, the post-depositional fate of ClO3is not well understood. ClO3- can be used as an electron acceptor by bacteria under anoxic conditions 25, 26, and transformed by plants 27, 28, 29 or undergo acid decomposition 30. Preliminary data in our lab demonstrates that a likely abiotic transformation mechanism for ClO3- may be iron mediated reduction similar to that observed for NO3- in Don Juan Pond 31. Iron is an abundant element within terrestrial and extraterrestrial systems and is found ubiquitously throughout aquatic environments 32, 33, where potential interactions between iron minerals and ClO3- are likely to occur. ClO3- reduction was reported during extraction of an H type chondrite, potentially due to the presence of native metallic iron (Feo) and/or other reactive iron oxides present 7. The ability of metallic( iron to reduce ClO3- has been applied in packed bed columns for water treatment 34, but there is a paucity of studies examining the reduction of ClO3- by natural iron minerals (Fe2+, Fe3+, or mixed). In comparison, reduction of nitrate (NO3-) by iron bearing minerals has been extensively investigated. In the presence of green rust, NO3- reduction to ammonium (NH4+) occurs readily 35, 36, 37, while the reduction of NO3- by siderite, an iron carbonate, does not readily occur 38. Further investigations have also reported the reduction of NO3- by wustite (FeO) 39. NO3- and ClO3- exhibit similar chemical properties and share common microbial transformation pathways with NO3- respiring bacteria capable of reducing ClO3- 40, 41, but no studies have examined the reduction of ClO3- by Fe(II) bearing minerals. In this study, the reactivity of ClO3- with three Fe (II) bearing minerals was investigated. The minerals were selected in order to represent a broad range of common types of iron minerals found in the terrestrial and extraterrestrial systems and planetary materials (i.e. meteorites) including iron oxides 42-45, iron carbonates 42, 46 and mixed-valent iron minerals. The overall objective of this paper is to report a survey of the reactivity of ClO3- by prevalent Fe(II)-bearing minerals to infer potential geochemical reactions. The reactions of ClO3- with mineral phases may constitute an important part of post


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depositional transformation of oxy-chlorine compounds and contribute to the large variations in the relative occurrence of ClO3-, ClO4- , and NO3- in different terrestrial and extraterrestrial environments. These reactions may bear direct implications to the current and past geochemical surface reactions that produced and evolved the fossil relic of oxy-chlorine compounds. In addition, elucidation of the stability of ClO3- has implications for understanding the past and present presence of water in hyper-arid systems (e.g., Mars). Materials and Methods Overview To prevent air oxidation, all synthesis and handling of minerals and experiments were conducted in an anaerobic chamber (mixture of Ar-H2 and N2) equipped with dual palladium-carbon catalyst wafers (Coy Laboratory Products, Grass Lake, MI). All solutions, including the distilled deionized water used to make batch solutions and synthesize minerals, were purged overnight (~ 12 hours) in the chamber to remove trace levels of O2. All chemical reagents used to prepare Fe(II), ClO3-, and NO3- solutions were of ACS grade. Stock solutions of ClO3- and NO3- were made using NaClO3 and NaNO3. Fe(II) solutions were prepared from ferrous sulfate (FeSO4 ▪ 8 H2O). All experiments were carried out under dark conditions to exclude potential photochemical reactions. Two buffers where prepared and used to maintain a constant pH: (1) a 0.05 M TAPS ([tris (hydroxymethyl) methylamino] propanesulfonic acid) and (2) 0.05 M MOPS (3-(N-morpholino) propanesulfonic acid) 47, 48. No buffer was used for experiments conducted at pH ~ 4.5 and pH was adjusted with H2SO4. The pH was monitored during experiments but little variation (± 0.1) in pH was observed. Mineral Synthesis and Characterization Wustite (FeO) was commercially purchased from Alfa Aesar (Ward Hill, MA) and was used as received. Magnetite (Fe3O4) (laboratory grade iron oxide, Fisher Scientific, Catalog No. I119-500) was used as received. Siderite (FeCO3) was synthesized following methods detailed in Rakshit (38) and

Thamdrup (49) in the anaerobic glove box and is further described here. Sodium carbonate (Na2CO3)



(held at pH 8.5 with TAPS buffer) was added to a solution of FeCl2 to give an equimolar concentration of 0.5 M for each and the mixture was stirred for ca. 10 minutes until a pale grey precipitate formed. The grey precipitate was centrifuged at 10,000 rpm for 10 minutes and excess liquid was decanted from the solid. The solid was rinsed at least 3 times with Ar-purged deionized water until the concentration of Clion in the supernatant was comparable to the background level (< 0.1 mg/L). The solid was re-suspended in the appropriate pH solution for the required experiment. The precipitate was used immediately for experimentation and was not stored for periods longer than 24 hours. Preparation of sulfate green rust (GRSO42-) was adapted from Hansen (35). Specifically, it was synthesized by addition of 100 g of ferrous sulfate (FeSO4 ▪ 8 H2O) to a stirred 100 mL solution of 0.05 M TAPS buffer (pH 9) for ca. 10 minutes. Particles were then collected via centrifugation and rinsed as described above. The precipitate was used immediately within 24 hours. The surface areas of the minerals were analyzed using the BrunauerEmmett-Teller (BET) N2 adsorption method with a surface area analyzer (Nova 4200e, Quantachrome Instrument Corp) at 77 K. Prior to N2 adsorption, the particles were degassed at 150 oC for 8 hours. A relatively low degassing temperature was used as oxide minerals (e.g. ferrous and aluminum) are susceptible to oxidation and phase change at elevated temperatures 50, 51. In spite of the precautions, siderite experienced significant color change from greenish grey to dark brown upon BET analysis. No noticeable changes in visual appearance were observed for the wustite and sulfate green rust particles. Chlorate or Nitrate Reduction Experiments All batch experiments were conducted at 25°C in triplicate, unless otherwise noted, using 50 mL glass serum bottles sealed with butyl-rubber septa and aluminum crimp seals. The specific conditions employed for the reaction of NO3- and ClO3- with different iron materials are described below. Bottles were kept well mixed using a wrist action shaker. Each reactor bottle was sampled sequentially over the time interval for each experiment (sample size: 50 -500 µL). Control experiments were conducted in identical media but without the amendment of iron minerals to evaluate possible removal of ClO3- by background processes. Sampled aliquots were filtered through either 0.2 μm syringe filters or disk filters


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to remove solids before analysis for ClO3-, NO3- and/or chloride (Cl-). The concentration of NO3- was reported as NO3--N. Concentrations of ClO3- were quantified by IC-MS/MS using the method detailed in Rao (1) and further detailed in Jackson (7, 14). NO3- (as NO3--N) and Cl- were analyzed by ion chromatography following EPA Method 300.0 54.

Dissolved Iron (Fe(II)) – To evaluate the impact of dissolved Fe(II), triplicate experiments at pH 4.5 and 6.5 were conducted with a 10 or 100 mg/L solution of ferrous sulfate and 1 mg/L ClO3-. A similar set of triplicate experiments were evaluated for NO3- (NO3--N = 1 mg/L) transformation in the presence of 100 mg/L Fe(II).

Siderite (FeCO3) – Experiments with FeCO3 were conducted with either 1 or 10g/L of FeCO3 in bottles containing an initial concentration of 1 mg/L of ClO3- at pH 6.5 or 8.5. Transformation of NO3- by FeCO3 was examined using similar conditions at an initial NO3- (as NO3--N) concentration of 1 mg/L.

Sulfate Green Rust (GRSO42-) – GRSO42 experiments were conducted with 10 mg/L of ClO3- at pH 6.5 or 8.5 with 1 or 10 g/L GRSO42. A higher initial concentration of ClO3- was used due to its rapid reaction with green rust particles. Similar experiments were performed to assess transformation of NO3- at an initial concentration of 10 mg/L.

Wustite (FeO) – FeO experiments were conducted over a pH range from 4.5 – 8.5 with 1 or 10 g/L FeO at an initial ClO3- concentration of 1 mg/L. NO3- experiments (initial concentration of NO3--N = 1 mg/L) were conducted similarly to ClO3- experiments but at pH 6.5 only. Additional experiments with variable concentrations of FeO (1, 10, 20, or 50 g/L FeO) and ClO3- (1, 10, or 50 mg/L) were conducted to study the effects of concentration on reaction rates. To understand the effects of brine matrixes on ClO3- and NO3- reduction, duplicate experiments were conducted with 10 mg/L of ClO3- or NO3--N and 20 g/L FeO



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in the presence of varying levels (10, 100 or 1000 mM) of sodium chloride (NaCl). All brine experiments were conducted at a pH of 6.5.

Simultaneous Reduction Experiments – Simultaneous transformation of NO3- and ClO3- was evaluated at pH 6.5 for each iron mineral (10 g/L) to evaluate competitive effects. The experimental procedure follows that of the single reductate experiments described above except that both NO3- and ClO3- were added to the reaction mixture at a concentration of 20 mg/L each at the beginning of the experiments.

Kinetic Analysis The reduction rates of ClO3- by various Fe (II) bearing minerals were assessed considering the following overall rate law: ―

𝑑[𝐶𝑙𝑂₃⁻] 𝑑𝑡

= 𝑘[𝐹𝑒(𝐼𝐼) 𝑚𝑖𝑛𝑒𝑟𝑎𝑙]𝑥 [𝐶𝑙𝑂₃⁻]𝑦 (Eq. 1)

where – (d[ClO3-]/dt is rate of reduction of ClO3- at any given time, t; k is the overall rate coefficient; and x and y are the reaction order with respect to the Fe(II) mineral and ClO3-, respectively. Surface reactions can assume complex kinetic behavior and the apparent reaction order may change with the ratio of Fe(II) mineral to ClO3- or NO3. In this study, our experiments were performed with a large stoichiometric excess of Fe(II) minerals relative to ClO3- or NO3-, therefore, the solid concentration was assumed to be constant and Eq. 1 may be simplified to: ―

𝑑[𝐶𝑙𝑂₃⁻] 𝑑𝑡

= 𝑘𝑜𝑏𝑠[𝐶𝑙𝑂₃⁻]𝑦

(Eq. 2)

Where kobs is the apparent reaction rate constant with respect to ClO3-. The experimental data was fit with pseudo-first order rate law with respect to ClO3- (i.e., y= 1), and the apparent reaction rate constants from Eq. 2 were determined using non-linear regression.


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Results We evaluated the reactions between ClO3-, aqueous Fe2+, and three types of solid phase minerals (FeCO3, GRSO42- and FeO) at pHs ranging from 4.5 to 8.5. Similar experiments were conducted with NO3in the presence of aqueous Fe2+or the three minerals at pH 6.5, while competitive studies between NO3and ClO3- were also conducted in the presence of both anions and the minerals at pH 6.5. Further studies evaluated the impact of a background salt (i.e., NaCl) on ClO3- or NO3- reduction. Samples were periodically taken to monitor for the concentrations of ClO3- and the product Cl-, while NO3- experiments were monitored for NO3- loss only. Kinetic rate constants were calculated using Eq. 2. There were negligible reactions of ClO3- with dissolved Fe(II) at pH 4.5 and 6.5 over a 28 day period (data not shown), demonstrating reduction of ClO3- by dissolved Fe2+ does not readily occur. No transformation of NO3- at pH 6.5 by dissolved Fe(II) was observed within our experiments either (data not shown). This agrees with prior studies that showed a lack of reactions between NO3- (and/or NO2-) and dissolved Fe(II) 35, 37- 39, 52. The results show that, aqueous phase reduction by Fe(II) ion from dissolution of minerals should not impact the results of solid phase experiments. Reduction of ClO3- by FeCO3 is relatively slow compared to GRSO42- at the same solid loading and solution pH (Figure 1). In comparison to ClO3-, NO3- was not reduced in the presence of FeCO3 at pH 6.5 over a 45-hour time period. Greater than 95% of ClO3- was transformed after 72 and 30 hours, respectively, for 1 and 10 g/L FeCO3 at pH 6.5, whereas no significant loss of ClO3- was observed in control experiments (Figure 2). Comparatively at pH 8.5, approximately 75% and 90% of ClO3- was lost after 72 hours in the presence of 1 and 10 g/L FeCO3. In all experiments, ClO3- was reduced to Clquantitatively based on the ratios of ClO3- loss to Cl- production (Δ[ClO3-]/Δ[Cl-] ≅ 1.0 - 1.2). In general, the pseudo first-order rate model fits the data well (r2 >0.94) and the apparent rate constants were summarized in Table 1. Transformation of ClO3- was more rapid at pH 6.5 than pH 8.5 (kobs = 2.7 ± 0.3) x 10-2 and (1.5 ± 0.3) x 10-2 hr-1, respectively for 1 g/L FeCO3 (Table 1). Increasing the loading of FeCO3 from 1 g/L to 10 g/L increased transformation rates by twofold. The non-linear response of reaction rate to solid loading is interpreted as the solid surface being in excess at the higher loading of 10 g/L.



The transformation rate of ClO3- by GRSO42- was the highest among the minerals examined (Figure 2). ClO3- was rapidly reduced in the presence of 1 or 10 g/L GRSO42-, with > 90% loss of ClO3- in less than 10 hours occurring at pH 6.5 and 8.5 (Figure 3). At all pH values and GRSO42- concentrations evaluated, ClO3- was reduced to Cl- at an approximate 1:1 stoichiometric conversion. Similar to FeCO3, the reduction rate approximately doubled with a ten-fold increase in the concentration of GRSO42- at each pH. The pH had a prominent influence on the reduction rates, and kobs at a pH of 8.5 was 5 times greater than the rate constant at pH 6.5 (Table 1). In contrast to siderite, NO3- was readily reduced by GRSO42- (pH = 6.5) and the rate exceeded that of ClO3- (Figure 1). NO3- reduction increased by 50% with an increase in GRSO42- concentration from 1 g/L to 10 g/L. Similar to FeCO3, the transformation rate increased by 1.5 and 2.1 times, for NO3- and ClO3-, respectively, due to increased mineral mass from 1 g/L and 10 g/L. The non-linear response of reactions rates to the mass loading of iron minerals again, suggests that GRSO42 is in excess relative to the oxyanions at the higher loading. In aqueous suspensions of FeO, ClO3- was reduced at all pHs between 4.5 and 8.5, leading to near equal molar production Cl- (Δ[ClO3-]/Δ[Cl-] ≅ 0.95 -1.0) (Figure 3). ClO3- reduction was strongly influenced by pH. The rate of reduction was the greatest at pH 6.5, approximately 6 times larger than the rates at pH 8.5 or 4.5, respectively (Table 1) regardless of the loading of FeO. At all pHs, ClO3- reduction rates tripled as the FeO dose increased from 1 g/L to 10 g/L (Table 1). Additional experiments were conducted in which the ratio of ClO3- concentration with respect to that of FeO was systematically adjusted. As depicted in Figure 5, for a given initial ClO3- concentration, increasing the mass of FeO increased the reaction rate, while increasing the ClO3- concentration for a given dose of FeO reduced the reaction rate. This inverse dependence of the rate on ClO3- concentration likely reflects a surface saturation effect, which can arise in mineral-water interactions when the aqueous reactants are in abundance relative to the available surface sites 53. Because of this effect, the highest rate of ClO3transformation by FeO (kobs = ~ (1.7 ± 0.1) x 10-1 h-1) was attained with the largest concentration ratio of FeO to ClO3- (50 g/L : 1 mg/L), while the lowest mineral to ClO3- ratio (1 g/L : 50 mg/L) gave rise to the slowest reduction process (kobs = ~ (2.0 ± 0.2) x 10-2 h-1) (Figure 6). NO3- was reduced by FeO at a pH of


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6.5 (Figure 1). kobs of NO3- reduction increased from 0.059 ± 0.05 to 0.116 ± 0.1 h-1 with an increase in FeO concentration from 1 to 10 g/L (Table 1). Notably, in the FeO suspensions, NO3- was more rapidly reduced than ClO3-, with the kobs of NO3- about 2 to 2.5 times that of ClO3- (Table 1). The effect of brine matrix on ClO3- and NO3- reduction by ferrous materials was examined by performing experiments in concentrated NaCl solutions. Transformation of ClO3- in brine solutions was moderately accelerated ((4.0 ± 0.1) x 10-2 to (6.0 ± 0.3) x 10-2 h-1) with increasing NaCl concentration (Figure 7). On the contrary, NO3- reduction rates decreased significantly with increasing brine concentration, with kobs declining from (6.0 ± 0.1) x 10-2 to (2.0 ± 0.1) x 10-2 h-1 as the brine concentration increased from 0 to 1000 mM NaCl. The natural abundance of ClO3- in the terrestrial environment is relatively low and ClO3- often coexists with other reducible oxyanions, such as NO3-. To evaluate the impact of the co-presence of ClO3and NO3- on ClO3- reduction, reactors were amended with both NO3- and ClO3-, each at 20 mg/L and 10 g/L iron minerals (Figure 6). When FeCO3 was used as the reductant, the presence of NO3- at an equal concentration moderately reduced the observed rate constant for ClO3- transformation from (5.8 ± 0.1) x 10-2 to (4.8 ± 0.3) x 10-2 h-1, when ClO3- was the sole electron acceptor (Table 1, 2). Similar to the results seen in simple NO3- reduction experiments, there was no loss of NO3- within 45 hours when both anions are present. In the case of GRSO42-, when NO3- and ClO3- were both present, NO3- reduction decreased by more than 50%, while ClO3- reduction was minimally impacted (kobs = (6.0 ± 0.4) x 10-2 and (4.1 ± 0.6) x 10-1 h-1, respectively, in simple and competitive reduction experiments). It was also noted that the concentration profile of NO3- with time in the presence of ClO3- deviates from the typical 1st-order behavior. Specifically, there was an initial period of slow reaction, which transitioned to a latter period exhibiting 1st-order decay. This bi-phasic behavior gave rise to a distinctive sigmoidal shape for the NO3concentration profile (Figure 6). This significant impedance in NO3- reduction at the early stage of the competitive reduction and the minor impact experienced by ClO3- during the process indicate preferential reduction of ClO3- over NO3- by sulfate green rust. This observation agrees with preference exhibited in siderite experiments.



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When NO3- and ClO3- were simultaneously reduced by FeO, both anions appear to react more slowly than their reduction in the non-competitive systems. The concentration profiles appear to deviate from the pseudo-first order rate law (Figure 6). Furthermore, while there seems to be a modest effect on ClO3- transformation when NO3- was present (kobs = (6.0 ± 0.4) x 10-2 and (4.5 ± 0.5) x 10-2 h-1, respectively, in simple and competitive reduction experiments), NO3- reaction was severely retarded in the presence of ClO3- (kobs drops from (1.2 ± 0.1) x 10-1 to (1.8 ± 0.9) x 10-2 h-1) (Table 1, 2). The relative extents of rate inhibition portray ClO3- as a more competitive electron acceptor, however, compared to siderite and green rust, NO3- clearly poses a stronger competition on ClO3- in the FeO system. Discussion The final product of ClO3- experiments was predominantly Cl-. Intermediates such as chlorite (ClO2-) and hypochlorite (OCl-) were not detected during sample analysis and their presence can be considered as negligible based on the Cl mass balance monitored during the experimental periods. Therefore, the limiting step of ClO3- reduction is likely the first two-electron transfer to form ClO2- (R.1) or the four-electron transfer to form OCl- (R.2). Figure 8 shows the reduction potential (EH) of various redox pairs relevant to study. As shown, the reduction potential of ClO3-/ClO2- at pH 7 is 0.62 V. The reduction potential of ClO2- to ClO- or Cl- is more positive than that of ClO3-/ClO2- and lies outside the boundary of water stability zone, indicating the oxy-chlorine intermediates generated during ClO3reduction readily transform to Cl-, which agrees with our experimental observations.

ClO3- + 2H+ + 2e- = ClO2- + H2O

(R.1)

ClO3- + 4H+ + 4e- = OCl- + 2 H2O

(R.2)

NO3- + 2H+ + 2e- = NO2- + H2O (note pKa of NO2- = 3.4)

(R.3)

NO3- is the most oxidized form of nitrogen and in the previous studies, the reduction of NO3- by Fe2+ bearing minerals yielded products of NO2-, N2O and NH4+ 31, 35-39, 52. Prior results have shown that NO2- reduction is more rapid than that of NO3- 35- 39, therefore, NO3- to NO2- conversion (i.e., R.3) can be


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regarded as the rate-limiting step of NO3- reduction (R.3). The reduction potential of NO3-/NO2- at pH 7 is 0.43 V, whereas the reduction potential of NO2- to N2O, NO, and NH4+ are 0.77, 0.36, and 0.35 V, respectively, at pH 7 (Figure 8). In the diagram, the potentials of FeOOH(s)/FeO(s) and FeOOH(s)/FeCO3(s) are more negative than all N- or Cl-containing redox pairs. This indicates that the reduction of ClO3- or NO3- by wustite and siderite is energetically favorable. Across the pH range of 5 - 9, the reduction potential of ClO3- is more positive than that of NO3- (assuming NO2- is the intermediate). If the reaction rates were controlled by the free energy changes of the dissolved anions and their immediate products, ClO3- is expected to be reduced at a higher rate than NO3-. However, our experimental observations do not follow this prediction, and this suggests that surface interactions are more likely to control the activation energies and the reduction rates in the systems examined here. Similar notions have been made by Rakshit (39) and Hansen and Koch (36) in studies of NO3- reduction by FeO and GRSO42-. Among the three minerals, GRSO42- demonstrated the fastest reduction of ClO3- at any given pH or mineral loading (Figure 3), followed by FeCO3 and FeO. NO3- reduction was demonstrated to only occur in the presence of GRSO42- and FeO, and the rates by GRSO42- was approximately an order of magnitude higher than that of FeO at either solid loading. No prior data is available on abiotic ClO3- reduction by reduced iron minerals, while studies by Rakshit (39) and Hansen (35) have demonstrated that NO3- is reduced stoichiometrically to NH4+ in aqueous suspension of GRSO42- and FeO 35, 39. It has been reported that siderite reaction with NO3- is slow at circumneutral pH, with negligible NO3- loss in 30 days 38, 55. The large differences across mineral types can be attributed to the differences in the surface areas and/or the intrinsic properties of the surface reaction sites. Sulfate green rust has a layered structure with external and internal sites, and the sulfate anion intercalates the layers acting as a charge balance 37. Siderite is structurally a rhombohedral with only external reactive sites 56. Wustite is also considered as having exclusively external sites 39, 57. In the present study, the experimentally measured BET surface areas of GRSO42-, FeO and FeCO3 are summarized in Table 1. FeO has the smallest specific surface area (0.074 m2/g), and the value is consistent with wustite having external surface only. The relatively low abundance of reactive sites may partially account for the sluggish reactions observed in the presence FeO. In the case



of FeCO3, the surface area measurement was impacted by the sample oxidation issue and the BET results cannot be compared against other iron minerals. Nonetheless, the observations that NO3—N transformed in the presence of FeO and GRSO42- but not FeCO3 and the reactivity series of materials were different for ClO3- and NO3—N imply that the specific interactions between the oxyanions and solid surface play a more important role than the surface area effect. Unlike GRSO42- or FeO, whose surfaces are terminated with hydroxyl groups, the surface of siderite is predominantly covered with carbonate groups 38. The carbonate sites carry negative charges over a wide range of pH, and this can create electrostatic repulsion hindering access to surface reaction sites by the anions 55. Unfavorable charge interactions may explain for the lack of reaction between siderite and NO3-, however, ClO3- is reduced by siderite at rates comparable to or slightly higher than FeO. The difference may be rationalized by the molecular structures of the anions. NO3- has a molecular geometry that is trigonal planar with bond angles ~ 120o, and the N atom does not have an electron lone pair. ClO3- has a molecular geometry that is trigonal pyramidal with bond angles < 109o. ClO3- also has a lone pair of electrons surrounding the central Cl atom. The structure of ClO3- may allow it to coordinate with surface Lewis sites on siderite for electron transfer while such interaction is restricted for NO3-. While there are limited studies on the coordination chemistry of ClO3-, ClO3- adsorption to surfaces such as calcium carbonate (CaCO3) has been reported 58. This may account for, among other factors, the observation of ClO3- reduction by FeCO3, but not NO3-. Though the pH effect was not explicitly considered in the overall rate law equation, our experiments demonstrated a dependence of reduction rates on pH. In general, reactions rates were the fastest at a neutral pH (~ 6.5), with the exception of GRSO42-, which reported a faster reduction constant (kobs = 2.92 ± 0.13 hr-1) for 10 g/L at pH of 8.5 (in contrast to kobs = 0.655 ± 0.21 hr-1 at pH 6.5). A potential explanation for the observations with GRSO42- could be greater stability of green rust particles at an alkaline pH as the GRSO42- was prepared at approximately pH 9. The effect of pH on reaction rates can be related to the influence of pH on the relative distribution and speciation of surface sites. Protonation/deprotonation of surface sites in water and the formation of surface complexes are regulated


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by pH 59. Anion adsorption is preferred at a lower pH, however, protonation of surface sites makes oxidation of Fe(II) significantly slower 60. The net effect of these processes predicts the highest reaction rates should occur at an intermediate pH, which is consistent with the observations of siderite and wusite. Similarly, Rakshit (38) demonstrated the complexity of pH influence on the reduction of NO2- by FeCO3 and predicated a higher fraction of active sites at more neutral pH, as did Dhakal (52) with Fe3O4. We observed competitive effect on NO3- reduction imposed by the presence of ClO3- (Figure 6 and Table 2) with 50 to 85 % decreases in rate constants for NO3- reduction by GRSO42- and FeO, respectively, when ClO3- was present at an equal concentration. ClO3- reduction rate constants, however, were only slightly reduced (by 7- 25 %) in competitive reduction experiments compared to rate constants obtained with ClO3- only. The observations can be attributed to thermodynamic considerations, the molecular structural differences of the two anions as well as the potential products formed during our reactions. ClO3- is a stronger oxidant than NO3-, as shown by the preceding discussion of the reduction potentials of the two anions at neutral pH. Additionally, the difference in molecular geometry may permit ClO3- to be more affinitive to surface reaction sites, which, as noted earlier, may also contribute to the lack of reduction of NO3- by FeCO3. Implications Overall, results from this study demonstrate a previously unknown abiotic reaction between ClO3and Fe (II) bearing minerals, which has significant implications for terrestrial and extraterrestrial systems. Reported ClO3- : ClO4- ratios for the Antarctica MDV soils are orders of magnitudes lower than similar arid environments ratios (i.e. Atacama) 1 and ratios from Lake Vida and the ice-covered lakes of the MDVs have reported ratios that are lower than atmospheric depositional values, presumably due to biodegradation of ClO3- 14, 61. However, other research suggests that biotransformation of any oxychlorine species in Antarctica would be unlikely due to insignificant microbial metabolic rates owing to constant, extreme temperatures (< 0° C) and reported cell replication time of over 120 years 62, 63. As biotic reactions rates may be potentially insignificant, there is a strong interest in understanding the



contribution of abiotic mechanisms in the cycling of Cl and Fe through these environments. Additionally, understanding the soil chemistry of Antarctica MDVs and lakes systems helps to establish a framework for interpreting in-situ measurements from Mars and identifying abiotic or possibly biotic processes permissible under the relevant conditions. Studies investigating the presence of oxy-chlorine species in the Martian soil system have hypothesized various explanations for the abundance of ClOx- in comparison to co-occurring anions. Although only one sampled location has reported a predominance of ClO4- in comparison to other anions4, ClO4- is hypothesized to be the dominating species with ClO3- and NO3- having concentrations below ClO4-. As previously stated, all known dry production mechanisms of ClO4- produce ClO3- and ClO4- at a ratio ≥ 1, therefore the predominance of ClO4- suggests post-depositional processing. Due to the high stability of ClO4-, abiotic iron reactions could play a part in the consumption of ClO3- (and by extension NO3-) without the depletion of ClO4- potentially explaining the hypothetical prevalence of ClO4. This explanation would also indirectly imply the (past) availability of water on Mars. ClO3- also could potentially impact sampling and/or participate in the destruction of organics in Martian soils. NavarroGonzález (21) noted that the presence of ClO4- is important in Martian systems, because it interferes with the detection of organic compounds during pyrolysis-MS, but ClO3- is less stable than ClO4- and would be more likely to participate in the destruction of organics in Martian soils. As heterogeneous reactions between oxidants (e.g., ClO3-) and reductants (e.g. organic substances) are facilitated in the presence of Fe (II)- and Fe (III)-bearing minerals due to the catalytic properties of iron 60, implications into potential reactions for the oxidizing nature of soils can be inferred. The reactions of ClO3- with iron minerals observed in this study point to the need for considerations of potential sample alteration during the sampling and handling of non-planetary and extraterrestrial samples. Jackson (7) reported an observed instability of ClO3- during extraction of H type chondrite meteorites. As reported, concentrations of ClO3- determined from extractions were noted to be taken as minimums, as reduction by potential iron particles within the samples could not be accounted for


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7.

It is recommended that potential interactions between ClO3- or NO3- and iron bearing minerals should be

carefully evaluated prior to interpreting the sampling data.

Acknowledgements The authors would thank Drs. Juliusz Warzywoda for assistance in BET analysis. The authors gratefully acknowledge funding from the NASA Earth and Space Science Fellowship (NNX16AP45H). W.L.Y. acknowledges partial support by NSF (CHE-1308726).

References 1. Rao, B.; Hatzinger, P. B.; Böhlke, J. K.; Sturchio, N. C.; Andraski, B. J.; Eckardt, F. D.; Jackson, W. A., Natural chlorate in the environment: Application of a new IC-ESI/MS/MS method with a Cl18O3-internal standard. Environ. Sci. Technol. 2010, 44 (22), 8429-8434. 2. Jackson, W. A.; Böhlke, J. K.; Gu, B.; Hatzinger, P. B.; Sturchio, N. C., Isotopic composition and origin of indigenous natural perchlorate and co-occurring nitrate in the southwestern United States. Environ. Sci. Technol. 2010, 44 (13), 4869-4876. 3. Jackson, A.; Davila, A. F.; Böhlke, J. K.; Sturchio, N. C.; Sevanthi, R.; Estrada, N.; Brundrett, M.; Lacelle, D.; McKay, C. P.; Poghosyan, A., Deposition, accumulation, and alteration of Cl−, NO 3−, ClO 4− and ClO 3− salts in a hyper-arid polar environment: Mass balance and isotopic constraints. Geochim. Cosmochim. Acta 2016, 182, 197-215. 4. Hecht, M. H.; Kounaves, S. P.; Quinn, R. C.; West, S. J.; Young, S. M. M.; Ming, D. W.; Catling, D. C.; Clark, B. C.; Boynton, W. V.; Hoffman, J., Detection of perchlorate and the soluble chemistry of Martian soil at the Phoenix lander site. Science 2009, 325 (5936), 64-67. 5. Ming, D. W.; Archer, P. D.; Glavin, D. P.; Eigenbrode, J. L.; Franz, H. B.; Sutter, B.; Brunner, A. E.; Stern, J. C.; Freissinet, C.; McAdam, A. C., Volatile and organic compositions of sedimentary rocks in Yellowknife Bay, Gale Crater, Mars. Science 2014, 343 (6169), 1245267. 6. Kounaves, S. P.; Hecht, M. H.; Kapit, J.; Gospodinova, K.; DeFlores, L.; Quinn, R. C.; Boynton, W. V.; Clark, B. C.; Catling, D. C.; Hredzak, P., Wet Chemistry experiments on the 2007 Phoenix Mars Scout Lander mission: Data analysis and results. J. Geophys. Res., [Planets](1991–2012) 2010, 115, E00E10. 7. Jackson, W. A.; Davila, A. F.; Sears, D. W. G.; Coates, J. D.; McKay, C. P.; Brundrett, M.; Estrada, N.; Böhlke, J. K., Widespread occurrence of (per) chlorate in the Solar System. Earth Planet. Sci. Lett. 2015, 430, 470-476. 8. Kounaves, S. P.; Carrier, B. L.; O’Neil, G. D.; Stroble, S. T.; Claire, M. W., Evidence of martian perchlorate, chlorate, and nitrate in Mars meteorite EETA79001: Implications for oxidants and organics. Icarus 2014, 229, 206-213. 9. Rao, B.; Anderson, T. A.; Redder, A.; Jackson, W. A., Perchlorate formation by ozone oxidation of aqueous chlorine/oxy-chlorine species: Role of Cl x O y radicals. Environ. Sci. Technol. 2010, 44 (8), 2961-2967. 10. Sturchio, N. C.; Caffee, M.; Beloso Jr, A. D.; Heraty, L. J.; Böhlke, J. K.; Hatzinger, P. B.; Jackson, W. A.; Gu, B.; Heikoop, J. M.; Dale, M., Chlorine-36 as a tracer of perchlorate origin. Environ. Sci. Technol. 2009, 43 (18), 6934-6938.



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48. Moroney, J. V.; Andreo, C. S.; Vallejos, R. H.; McCarty, R. E., Uncoupling and energy transfer inhibition of photophosphorylation by sulfhydryl reagents. J. Biol. Chem.1980, 255 (14), 66706674. 49. Thamdrup, B.; Finster, K.; Hansen, J. W.; Bak, F., Bacterial disproportionation of elemental sulfur coupled to chemical reduction of iron or manganese. Appl. Environ. Microbiol. 1993, 59 (1), 101-108. 50. Clausen, L.; Fabricius, I., BET Measurements: Outgassing of Minerals. J. Collid Interface Sci.2000, 227 (1), 7-15. 51. Heister, K., How accessible is the specific surface area of minerals? A comparative study with Al-containing minerals as model substances. Geoderma 2016, 263, 8-15. 52. Dhakal, P.; Matocha, C. J.; Huggins, F. E.; Vandiviere, M. M., Nitrite reactivity with magnetite. Environ. Sci. Technol. 2013, 47 (12), 6206-6213. 53. Deng, B.; Stone, A., Surface-catalyzed chromium(VI) reduction: Reactivity comparisons of different organic reductants and different oxide surfaces. Environ. Sci. Technol. 1996, 30 (8), 2484. 54. Pfaff, J.D. Method 300.0 Determination of inorganic anions by ion chromatography; US Environmental Protection Agency, Office of Research and Development, Environmental Monitoring Systems Laboratory 28, 1993. 55. Matocha, C.J.; Dhakal, P.; Pyzola, S.M., The role of abiotic and coupled biotic/abiotic mineral controlled redox processes in nitrate reduction, in Advances in agronomy, Elsevier: 2012; Vol. 115, pp 181-214. 56. Sharp, W. E., The cell constants of artificial siderite. Am. Mineral. 1960, 45 (1), 241-243. 57. Bernal, J. D.; D. R. Dasgupta; Mackay A.L. The oxides and hydroxides of iron and their structural inter-relationships. Clay Miner. Bull. 1959, 4 (21), 15-30. 58. Li, Y.; Xu, T.; Lui, C.; Li, Y. The adsorption of chlorite and chlorate by calcium carbonate in a drinking water pipe network. Desalination and Water Treatment, 2015, 53 (7), 1881-1887. 59. Van Cappellen, P., A surface complexation model of the carbonate mineral-aqueous solution interface. Geochim. Cosmochim. Acta 1993, 57 (15), 3505-3518. 60. Stumm, W.; Sulzberger, B., The cycling of iron in natural environments: Considerations based on laboratory studies of heterogeneous redox processes. Geochim. Cosmochim. Acta 1992, 56 (8), 3233-3257. 61. Kenig, F.; Chou, L.; McKay, C. P.; Jackson, W. A.; Doran, P. T.; Murray, A. E.; Fritsen, C. H., Perchlorate and volatiles of the brine of Lake Vida (Antarctica): Implication for the in situ analysis of Mars sediments. J. Geophys. Res., [Planets] 2016, 121 (7), 1190-1203. 62. Murray, A. E.; Kenig, F.; Fritsen, C. H.; McKay, C. P.; Cawley, K. M.; Edwards, R.; Kuhn, E.; McKnight, D. M.; Ostrom, N. E.; Peng, V.; Ponce, A.; Priscu, J. C.; Samarkin, V.; Townsend, A. T.; Wagh, P.; Young, S. A.; Yung, P. T.; Doran, P. T., Microbial life at -13 °C in the brine of an ice-sealed Antarctic lake. Proc. Natl. Acad. Sci. U.S.A 2012, 109 (50), 2062620631. 63. Ostrom, N. E.; Gandhi, H.; Trubl, G.; Murray, A. E., Chemodenitrification in the cryoecosystem of Lake Vida, Victoria Valley, Antarctica. Geobiology 2016, 14 (6), 575-587. 64. Vanysek P., Electrochemical Series, in CRC Handbook of Chemistry and Physics, 91th Edition. Taylor and Francis: Boca Raton, FL, 2010, pp 8. 65. Stumm, W.; Morgan, J. J., Aquatic Chemistry Chemical Equilibria and Rates in Natural Waters. 3rd ed., Wiley Interscience, 1995, pp 990.


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1 2 3 Table 1. Apparent Rate Constants (kobs) for Reduction of ClO3- or NO3- by Fe (II)-Bearing Minerals and 4 BET analyses. UD = Undetermined. 5 6 7 kobs for ClO₃⁻ (hr ⁻¹) kobs for NO₃⁻ (hr ⁻¹) 8 Specific 9 Surface Area pH 1 g/L 10 g/L 1 g/L 10 g/L 10 Iron Mineral m2/g 11 12 (3.1 ± 0.4) x 10-1 (6.6 ± 0.3) x 10-1 6.5 (5.8 ± 0.2) x 10-1 (8.5 ± 0.1) x 10-1 4.5 13 GR(SO₄²⁻) 1.4 ± 0.2 2.9 ± 0.2 8.5 14 15 (9.0 ± 0.8) x 10-3 4.5 (3.0 ± 0.2) x 10-3 16 17 (2.3 ± 0.3) x 10-2 (6.0 ± 0.4) x 10-2 (5.9 ± 0.5) x 10-2 (1.2 ± 0.1) x 10-1 FeO 0.074 6.5 18 (5.0 ± 0.7) x 10-3 (1.3 ± 0.4) x 10-2 8.5 19 20 (2.7 ± 0.3) x 10-2 (5.8 ± 0.1) x 10-2 UD UD 6.5 21 190 FeCO₃ 22 -2 -2 (1.5 ± 0.3) x 10 (3.2 ± 0.2) x 10 8.5 23 24 25 26 27 Table 2. Apparent Rate Constants for Competitive Reduction of ClO3- and NO3- by Fe (II)-Bearing 28 Minerals. UD=Undetermined. 29 30 31 kobs for ClO₃⁻ (hr ⁻¹) kobs for NO₃⁻ (hr ⁻¹) 32 33 Iron Mineral pH 10 g/L 10 g/L 34 35 36 (6.2 ± 0.5) x 10-1 (4.1 ± 0.6) x 10-1 GR(SO₄²⁻) 6.5 37 38 39 FeO (4.5 ± 0.5) x 10-2 (1.8 ± 0.9) x 10-2 6.5 40 41 (4.8 ± 0.3) x 10-2 UD FeCO₃ 6.5 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 ACS Paragon Plus Environment 60


(A)ClO ClO33-(A)

1.0

C/Co

0.8

0.6

0.4

0.2

(B) NO3- - N

1.0

0.8

C/Co

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FeO Gr(SO42-)

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FeCO3

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0.2

10

20

30

40

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Time, hours Figure 1. Reduction of (A) ClO3- and (B) NO3- (as NO3--N) at pH 6.5 by FeO, FeCO3, and GR(SO42-). Initial ClO3- or NO3- (as NO3--N) concentration was 1 mg/L, except for NO3--N reduction with GR(SO42-) using a starting concentration of 10 mg/L. Mineral loadings were 10 g/L. The solid lines represent kinetic model fits by non-linear regression. Error bars represent the standard deviation determined from triplicate experiments.


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pH = 6.5

(A)

1.0 FeCO3 (1 g/L) - ClO3-

0.8

C/Co (moles of Cl/ moles of initial Cl)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


FeCO3 (10 g/L) - ClO3 FeCO3 (1 g/L) - Cl FeCO3 (10 g/L) - Cl Control - ClO3-

0.6 0.4

-

0.2 0.0 pH = 8.5

(B)

1.0 0.8 0.6 0.4 0.2 0.0 0

20

40

60

80

100

120

Time, hours

Figure 2. Reduction of ClO3- at (A) pH 6.5 and (B) 8.5 by siderite (FeCO3). Initial ClO3- concentration was 1 mg/L and FeCO3 loading was 1 or 10 g/L. The solid lines represent kinetic model fits by non-linear regression. Values of ClO3- concentrations at any time were normalized to the initial ClO3- concentration, while Cl- concentrations were normalized to the theoretical final concentration of Cl- assuming total conversion of ClO3- (the same applies to subsequent figures). Error bars represent the standard deviation determined from triplicate experiments.



pH = 6.5

(A)

pH = 8.5

(B)

1.0 0.8

C/Co (molesof Cl/moles of initial Cl)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 24 of 30

0.6 0.4 0.2 0.0

1.0 0.8

GR(SO4 -) (1 g/L) - ClO3 2 GR(SO4 -) (10 g/L) - ClO3 2 GR(SO4 -) (1 g/L) - Cl 2 GR(SO4 -) (10 g/L) - Cl 2 Control - ClO3

0.6 0.4 0.2 0.0 0

50

100

150

200

250

Time, mins Figure 3. Reduction of ClO3- at (A) pH 6.5 and (B) 8.5 by sulfate green rust (GR(SO42-)). Initial ClO3concentration was 10 mg/L and GR(SO42-) loading was 1 or 10 g/L. The solid lines represent kinetic model fits. Error bars represent the standard deviation determined from triplicate experiments.


Page 25 of 30

pH = 4.5

(A)

1.0 0.8 0.6 0.4 0.2 0.0

C/Co (moles of Cl/initial moles of Cl)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


pH = 6.5

(B)

1.0 0.8

FeO (1g/L) - ClO3

-

FeO (10g/L) - ClO3 FeO (1g/L) - Cl FeO (10g/L) - Cl Control - ClO3

0.6 0.4

-

0.2 0.0 pH = 8.5

(C)

1.0 0.8 0.6 0.4 0.2 0.0 0

50

100

150

200

Time, hours

Figure 4. Reduction of ClO3- at (A) pH 4.5, (B) pH 6.5 and (C) 8.5 by wustite (FeO). Initial ClO3concentration was 1 mg/L and FeO loading was 1 or 10 g/L. The solid red lines represent kinetic model fits. Error bars represent the standard deviation determined from triplicate experiments.



0.20

Rate Coefficient, k (hr -1)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 26 of 30

1 g/L FeO 10 g/L FeO 50 g/L FeO 20 g/L FeO

0.15

0.10

0.05

0.00

1 mg/L

10 mg/L

50 mg/L

-

ClO3 Concentration Figure 5. Effect of initial ClO3- concentration and FeO loading on apparent rates of ClO3- reduction by FeO at pH 6.5. Error bars represent the standard deviation determined from triplicate experiments.


Page 27 of 30

1.2

ClO3- (Only) ClO3- (Competitive) NO3- - N (Competitive) NO3- - N (Only)

1.0 0.8 C/Co

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47


0.6 0.4 0.2 0.0 (A) FeO

(C) GrSO42-

(B) FeCO3

20

40 Time, hours

60

20

40

60

Time, hours

80

100

200

300

Time, mins

Figure 6. Reduction of ClO3- and NO3- (as NO3--N) in simple and competitive reduction systems at pH 6.5 by (A) FeO, (B) FeCO3, and (c) GR(SO42-). Initial NO3--N and ClO3- concentrations were 20 mg/L, while initial concentrations for FeO, FeCO3, and GR (SO42-) were 10 g/L. The solid red lines represent kinetic model fits. Error bars represent the standard deviation determined from triplicate experiments.



0.10 ClO3

-

-

NO3 - N 0.08

Rate Coefficient, k (hr -1)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 28 of 30

0.06

0.04

0.02

0.00

No Brine

10 mM

100 mM

1000 mM

Brine Concentration Figure 7. Effect of NaCl concentrations on apparent rates of NO3- (as NO3--N) and ClO3- reduction by FeO at pH 6.5. Initial NO3--N and ClO3- concentrations were 10 mg/L, while initial concentrations for FeO were 20 g/L. Error bars represent the standard deviation determined from duplicate experiments.


Page 29 of 30 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Figure 8. pH-EH diagram of redox couples relevant to this study. Potentials were calculated assuming half conversion (equal concentration of oxidized and reduced species) and major ions at 10-3 M. The oxidation product of Fe(II) minerals is assumed to be amorphous ferric oxyhydroxide (FeOOH). Thermodynamic data were obtained from sources 64, 65.



For TOC Only

Mars or Earth + Fe

2+


Page 30 of 30

Abiotic Reduction of Chlorate by Fe(II) Minerals - American Chemical

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