A Review of in Situ Chemical Oxidation and Heterogeneity

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Environ Sci Technol. 2014 Sep 2; 48(17): 10330–10336.

In Situ Chemical Oxidation of Contaminated Groundwater past Persulfate: Decomposition by Fe(Three)- and Mn(IV)-Containing Oxides and Aquifer Materials

Haizhou Liu

^†Department of Civil and Environmental Engineering science and ^§Department of Material Science and Engineering, University of California at Berkeley, Berkeley, California 94720, U.s.

Thomas A. Bruton

^†Department of Civil and Environmental Engineering and ^§Department of Material Scientific discipline and Applied science, University of California at Berkeley, Berkeley, California 94720, United States

Fiona Thou. Doyle

^†Department of Civil and Ecology Engineering and ^§Department of Textile Science and Applied science, University of California at Berkeley, Berkeley, California 94720, United States

David Fifty. Sedlak

^†Department of Civil and Environmental Technology and ^§Department of Material Science and Engineering, University of California at Berkeley, Berkeley, California 94720, United States

Received 2014 Apr 26; Revised 2014 Jul thirty; Accepted 2014 Aug 1.

Abstruse

Persulfate (S₂O_eight ^2–) is being used increasingly for in situ chemic oxidation (ISCO) of organic contaminants in groundwater, despite an incomplete understanding of the mechanism through which it is converted into reactive species. In particular, the decomposition of persulfate by naturally occurring mineral surfaces has not been studied in detail. To gain insight into the reaction rates and machinery of persulfate decomposition in the subsurface, and to identify possible approaches for improving its efficacy, the decomposition of persulfate was investigated in the presence of pure metal oxides, clays, and representative aquifer solids collected from field sites in the presence and absence of benzene. Nether weather typical of groundwater, Atomic number 26(III)- and Mn(4)-oxides catalytically converted persulfate into sulfate radical (Then₄ ^•–) and hydroxyl radical (HO^•) over time scales of several weeks at rates that were 2–twenty times faster than those observed in metal-free systems. Baggy ferrihydrite was the about reactive atomic number 26 mineral with respect to persulfate decomposition, with reaction rates proportional to solid mass and surface expanse. As a result of radical concatenation reactions, the rate of persulfate decomposition increased by every bit much as 100 times when benzene concentrations exceeded 0.1 mM. Due to its relatively wearisome rate of decomposition in the subsurface, it tin can be advantageous to inject persulfate into groundwater, assuasive it to migrate to zones of low hydraulic conductivity where clays, metal oxides, and contaminants volition accelerate its conversion into reactive oxidants.

Introduction

Despite more than three decades of effort, contamination of groundwater with organic pollutants remains a significant threat to drinking h2o supplies.¹ Attempts to employ ex situ treatment are frequently expensive and require decades to complete, while bioremediation and physical isolation of contaminants are difficult to use at many complex sites. As a result, engineers are increasingly using in situ chemical oxidation (ISCO) equally an expedient approach for contaminant remediation.²⁻⁵

ISCO has been used for the treatment of chlorinated organic solvents and petroleum hydrocarbons for over three decades.^vi−eight Many of the early efforts to utilise ISCO used permanganate (MnO₄ ^–) or hydrogen peroxide (H₂O₂). MnO₄ ^– reacts quickly with sure contaminants,⁹⁻¹¹ but the germination of manganese oxides (e.g., MnO_2(south)) can clog soil pores and inhibit oxidant transport in the subsurface. H₂O₂ reacts with minerals in the subsurface to produce hydroxyl radicals (HO^•) via Fenton-like reactions.¹² Because HO^• reacts rapidly with H₂O₂ to produce O₂ and H₂O¹³ and Fenton-like reactions typically exhibit depression yields at circumneutral pH values, the process is relatively inefficient.¹⁴⁻¹⁷ In addition, catalase enzyme activity in aquifer materials tin decompose H_iiO_ii through a nonradical pathway.^xiii As a result of its high reactivity, H_twoO_ii decomposes chop-chop in the subsurface, which can make it difficult to remediate contaminants that are located far from the reagent injection bespeak.

While permanganate and hydrogen peroxide are withal used frequently, persulfate (South₂O_viii ^two–) has recently become pop as an ISCO oxidant.^5,18−20 The mechanism through which Due south_iiO₈ ^2– oxidizes contaminants is similar to the way in which H_twoO_two is converted to HO^• by ultraviolet light, with thermolytic cleavage of the peroxide bail slowly producing sulfate radical (SO₄ ^•–) instead of HO^•:

ane

The formation of radicals from persulfate is commonly referred to as persulfate activation. However, in many cases it is difficult to measure sulfate radical concentrations and a decrease in persulfate loss is often used as a proxy for radical production. In such cases it may be more appropriate to refer to persulfate decomposition instead of activation, and this is the convention that is followed throughout this manuscript.

At 25 °C, the half-life of South₂O₈ ^two– is approximately 600 days, which ways that groundwater must be heated if this mechanism is to be applied for remediation.²¹⁻²⁵ Other approaches for accelerating the rate of S_twoO₈ ^ii– decomposition include UV irradiation,²⁶⁻²⁸ electrochemical systems,²⁹ zerovalent iron,³⁰ ferrous³¹ or cobalt salts,³² and base of operations activation.³³ Base activation is used most unremarkably in field applications.³³ Each of these mechanisms has its limitations. For example, use of ultraviolet low-cal in subsurface remediation is impractical, cobalt is a toxic metallic, Fe(II) salts are rapidly oxidized to insoluble Fe(Iii)-oxides in the presence of persulfate, and thermal activation is energy intensive nether the weather condition encountered in soils and aquifers.

Compared to the mechanisms listed to a higher place, persulfate decomposition past reactions at mineral surfaces are not also-known or understood. Results from recent studies propose that aquifer materials and natural organic affair (NOM) can increase persulfate decomposition rates.³⁴⁻³⁶ The charge per unit of decomposition was positively correlated with the content of amorphous atomic number 26 in aquifer solids.³⁴ Empirical studies besides demonstrated that the decomposition rate was enhanced by approximately an order of magnitude by the presence of aquifer solids.^21,34 Although the disappearance of persulfate during ISCO is slow compared to that of H₂O₂, the depression reactivity of persulfate in aquifer solids can show advantageous because information technology may allow the oxidant to reach source zones. It also may be useful in remediation of dilute plumes. All the same, there is a lack of key understanding of the factors that control the rate of persulfate decomposition by aquifer materials. In add-on to Fe(III) and Mn(Iv) content, other factors that may bear on the rate of persulfate decomposition include pH and the concentration of radical scavengers.

To provide insight into the role of naturally occurring minerals in persulfate decomposition and to develop a means of predicting the rate of persulfate loss during ISCO treatment, persulfate decomposition was studied under well-controlled conditions. Experiments were conducted using suspensions of pure minerals and aquifer solids to decompose persulfate in the presence of a pH buffer. The production of SO₄ ^•– or HO^• was followed using benzene equally a probe compound. Concentrations of persulfate and phenol (i.due east., a benzene oxidation production) were measured nether conditions expected in groundwater.

Materials and Methods

All chemicals used in this study were reagent grade or college. All solutions were prepared using deionized water (resistivity >18.ii MΩ, Millipore system). Ferrihydrite (Fe(OH)_three(s)) and pyrolusite (β-MnO_2(s)) were obtained from Sigma-Aldrich. Crystallized silica (SiO_2(s), i.e., pure sand, ACROS Organics, Inc.) was cleaned to remove trace amounts of metals co-ordinate to the procedure described in the Supporting Data (SI) section. Goethite (α-FeOOH_(s)) was synthesized by mixing 1 M Atomic number 26(NO₃)_three solution with 5 One thousand KOH and and so aging the precipitates at 70 °C for 60 h.³⁷ The synthesized goethite particles were rinsed with deionized water and dried using a Labconco Freezone iv.5 freeze-dry system (Labconco Corp., Kansas Urban center, MO). The synthesized goethite was characterized by Ten-ray diffraction (XRD) to confirm its purity (SI Figure S1).

Two clay materials were obtained from the Clay Minerals Society: Australian nontronite (26.0% Iron by weight) and Wyoming montmorillonite (2.59% Atomic number 26 by weight). In addition, aquifer materials were obtained from five sites in California and Arizona. Details of the physicochemical characteristics of pure minerals, clays, and aquifer materials are provided in SI Table S1. Additional information was provided previously.¹⁷ Prior to use, each fabric was crushed with a mortar and pestle and sieved to collect particles with sizes between 38 and 150 μm (sieves with mesh 400 and 100, respectively). The sieved particles were autoclaved at 120 °C and 20 psi for 45 min and stored under sterile conditions.

Persulfate decomposition experiments were carried out under chemical conditions typical of groundwater. Prior to initiation of each experiment, suspensions were prepared with solids concentrations between 12.5 and 125 grand/50 using air-saturated deionized water or constructed groundwater (SI Table S2). In some experiments, the suspension was purged with N₂ gas to achieve O_ii-free atmospheric condition and tests were conducted in a glovebox. In about experiments, the solution was maintained at pH eight.0 with l mM borate buffer prepared by adding 15 mL of a solution of 1 M of boric acid (H₂BO_three) and 62.5 mM of sodium tetraborate (Na_iiB₄O₇). The use of loftier buffer concentration was necessary to maintain a constant pH throughout the experiment. For some experiments, boosted NaHCO_three was added along with the borate buffer to reach a bicarbonate concentration between 0.ane and 10 mM.

Several experiments were conducted with benzene at concentrations ranging from 0.ane to 1 mM. Benzene solutions were prepared by diluting from a 1.2 mM stock solution that was freshly prepared by dissolving 10 μL of anhydrous benzene (purity ≥99.8%, Sigma-Aldrich Inc.) in a 1 L glass volumetric flask filled with deionized water without headspace. No benzene was volatilized during this preparation step.

To beginning a decomposition experiment, persulfate was added to yield initial concentrations ranging from one to 50 mM, using aliquots of a freshly prepared 100 mM Thou_twoSouthward₂O₈ solution. Later on mixing, the interruption was immediately transferred to multiple sealed glass tubes with no headspace and placed on a rotating mixer (Labquake Tube Rotators, Thermo Scientific Inc.). Each tube was treated as a sacrificial reactor and discarded afterwards sampling. Nearly weather condition were tested in triplicate, forth with a homogeneous control consisting of only persulfate. For experiments with benzene, 2 additional controls consisting of just benzene and benzene with solids were included. All experiments were conducted at room temperature of 23 ± ii °C.

At predetermined sampling intervals, each sealed sacrificial tube reactor was centrifuged at 3000 m for iii min to dissever the solids from the solution. The supernatant was immediately removed and filtered through a 0.22-μm nylon filter prior to analysis for persulfate, benzene and oxidation products. The solids were resuspended in acetonitrile, agitated vigorously with a vortex mixer, and so placed on a rotating mixer for 3 h to extract benzene and phenol from the solids. The acetonitrile was separated from the solids using the same centrifugation and filtration procedure and analyzed for adsorbed benzene and phenol.

Persulfate was measured using the KI colorimetric method³⁸ with a Lambda-14 UV spectrophotometer (PerkinElmer Inc., Waltham, MA). Benzene and phenol were analyzed on a Waters Brotherhood 2695 HPLC (Waters Corp., Milford, MA) equipped with a diode array detector. A Waters Symmetry-C18 cavalcade was used with a mobile phase of 40% acetonitrile and 60% 10 mM formic acid at a flow charge per unit of 1 mL/min. Dissolved oxygen concentrations were determined using a YSI Model 58 oxygen probe (YSI Inc., Yellow Springs, OH).

Results and Discussion

Decomposition of Persulfate by Minerals

As expected, the rate of persulfate loss in the absence of solids (i.eastward., homogeneous command) at room temperature was very boring (Effigy ​ 1A). The small loss of persulfate in the homogeneous command was consequent with the predicted thermal decomposition of S₂O₈ ^two– to Then₄ ^•– (reaction one). Using the activation energy measured previously,²¹ a loss of only 3% of persulfate was predicted over 30 days, which was similar to the measured loss of 2% in our experiments. The heterogeneous arrangement with silica exhibited less than 3% persulfate loss over 32 days, which was indistinguishable from the homogeneous control. The addition of fe- and manganese-containing oxides accelerated the decomposition charge per unit significantly, with an average of 12%, 65%, and 45% of persulfate loss over 32 days for 100 one thousand/L suspensions of goethite, ferrihydrite, and pyrolusite, respectively (Figure ​ 1A). The decomposition charge per unit was afflicted by the mass of solids in the suspension. For case, reducing the mass of pyrolusite from 100 to 50 k/L decreased persulfate loss from 45% to x% over 32 days, while increasing its mass from 100 to 200 and 500 g/L increased the loss to 60% and xc% over the same menstruation, respectively (Figure ​ 1B). Increasing the mass of silica did non affect the rate of persulfate loss (SI Figure S2).

Stability and decomposition of persulfate in different systems. Initial persulfate concentration one mM, pH eight.0, ionic strength l mM. (A) Impact of unlike solids along with a homogeneous control, solid mass loading 100 g/L. (B) affect of pyrolusite mass loading on persulfate decomposition.

To examine the effect of suspended mineral mass on persulfate decomposition, information were converted to pseudo-first-order kinetics with the following expression:

In this expression, thou _obs represents the observed pseudo-outset-order rate constant (due south^–i), k _SA is the surface area normalized pseudo-first-order rate constant (Fifty·1000^–2·s^–1), and C _S is the BET surface area of the suspended solids (grand²·L^–one).

The surface area normalized decomposition charge per unit constants for different solids varied by approximately 6 orders of magnitude (Figure ​ 2). The relative reactivity of the minerals followed the guild pyrolusite > ferrihydrite > goethite > silica. Nontronite, an iron-rich dirt and montmorillonite, a clay with a relatively high manganese content, were both capable of persulfate decomposition, with montmorillonite decomposing persulfate at a rate that was approximately an guild of magnitude higher than that of nontronite (Figure ​ ii). The aquifer materials exhibited decomposition rates that were similar to those observed for the iron oxides (Figure ​ 2). These aquifer materials contained between 0.7 and ii.5% Iron and between 0.01 and 0.12% Mn. Iron- and manganese-containing minerals likewise decomposed persulfate at initial persulfate concentrations ranging from 5 to 50 mM with 50 g/L of each mineral (SI Effigy S3).

Comparison of rate constants of persulfate decomposition past different solids at varying concentrations. Rate constants are normalized by BET surface areas and corrected for thermal decomposition at room temperature. Initial persulfate i mM, pH 8.0. Error bar represents ane standard divergence.

Reaction Mechanism for Heterogeneous Persulfate Decomposition

Past illustration to the Fenton system, nosotros propose that the rate-determining pace for persulfate decomposition by an iron-containing mineral can be described past the following reaction:

2

where ≡Iron(III) represents the redox-active iron surface site. Ane-electron reduction of ≡Fe(3) would effect in the formation of persulfate radical (S_twoO₈ ^•–), which could then initiate radical chain reactions. The existence of Southward₂O₈ ^•– is supported by previous studies of the reaction between S₂O₈ ^2– and So_four ^•–, in which S_iiO₈ ^ii– was oxidized to S_iiO₈ ^•– via a one-electron transfer reaction.^39,40

Similar to the Fenton reaction, the overall reaction rate of persulfate decomposition is likely afflicted past the number of reactive surface sites and electron transfer rates.⁴¹ Every bit in the Fenton system, the surface reactivity of each blazon of mineral is likely affected by crystallinity and metallic coordination. For example, in the Fenton organisation, the H_twoO₂ decomposition charge per unit was much higher in the presence of amorphous ferrihydrite than crystalline goethite.¹⁴ A similar trend was too observed in our experiments, with ferrihydrite exhibiting college charge per unit constants than goethite. Similarly, our observation of higher reactivity for manganese oxides was consistent with observations from the Fenton organisation.¹⁷

Previous research has demonstrated persulfate decomposition by dissolved transition metals, including Fe^two+, Mn²⁺, Ag⁺, Coⁱⁱ⁺, and Ni²⁺, under acidic conditions (i.e., pH < three).^31,42 Acidic weather condition are potentially problematic in groundwater because they tin can release toxic metals and alter minerals. Under circumneutral pH weather, transition metals associated with surfaces can undergo similar reactions. Later on surface metals are reduced via reaction 2, ≡Fe(II) should be oxidized quickly past persulfate to regenerate ≡Fe(III) with the production of Then₄ ^•– and sulfate:

3

Therefore, the metallic catalyzed decomposition of persulfate proceeds as follows:

4

The fact that persulfate tin be both oxidized and reduced is like to the machinery by which H₂O₂ acts as both an oxidant and reductant in the Fenton arrangement. In addition, S₂O_eight ^•– can be produced through a radical chain propagation reaction:^40,43,44

v

The fate of South_iiO₈ ^•– is unclear. Previous studies on the fate of the radical did not examine its decomposition machinery.^45,46 1 possible fate of S₂O_eight ^•– is oxidation of water, which would occur through a multistep procedure with the following stoichiometry:

half dozen

OH^– could besides serve as a source of HO^•46 and And so₄ ^•– can also oxidize water,⁴⁵ and nether element of group i weather condition

7

8

nine

In the absence of organic solutes, Reactions half-dozen and nine serve as concatenation termination steps.

The overall reaction of persulfate decomposition in water is

10

Results from selected experiments in which sulfate production was measured along with persulfate loss showed skilful agreement with the overall stoichiometry of reaction 10 (SI Effigy S4A). For Mn(4)-containing materials, a similar reaction pathway tin can take identify through a redox cycle involving Mn(IV)/Mn(3) surface species to generate sulfate radical and initiate chain reactions.

To assess the possible office of dissolved O₂ in persulfate decomposition, experiments were conducted in the absence of oxygen with a pyrolusite suspension. Results showed that the presence of dissolved O_two had no effect on the rate of persulfate loss (SI Effigy S4B). In add-on, dissolved O_two was produced equally S₂O_viii ^ii– decomposed. The observed stoichiometric ratio of O₂ produced per S_twoO₈ ^ii– lost was approximately 0.v (SI Figure S5), which was consistent with reaction 10.

Impact of Organic Compounds on Persulfate Decomposition

In the presence of organic contaminants, the radical chain decomposition of persulfate changes. Using benzene as a representative organic contaminant, the stability of persulfate was examined (Figures ​ 3 and ​ 4). The rate of persulfate decomposition increased equally benzene concentrations increased from 0 to thou μM. For example, in the presence of 50 g/Fifty ferrihydrite, approximately 15% of persulfate was decomposed over thirty days in the absenteeism of benzene. When 100 μM of benzene was added, persulfate decomposition accelerated significantly, with 45% loss over 30 days. At a benzene concentration of 1000 μM, approximately 75% of persulfate was decomposed over the same catamenia (Effigy ​ 3). These results indicated that the half-life of persulfate would decrease from 170 days in the absence of benzene to 45 and 20 days in the presence of 100 and thou μM of benzene, respectively. A like trend was observed for other pure minerals, with the exception of silica, which was not capable of catalyzing persulfate decomposition. For the aquifer materials, benzene appeared to have an effect, but it was harder to notice due to higher variability amidst replicates due to heterogeneity of the aquifer materials.

Bear upon of benzene concentration on the decomposition of persulfate in the presence of 50 g/Fifty ferrihydrite, pH viii.0. No mineral was nowadays in the homogeneous control.

Bear on of benzene concentration on the decomposition of persulfate in the presence of different solids. Solids mass loading 50 chiliad/L, pH 8.0.

The acceleration of persulfate decomposition in the presence of benzene may be owing to radical concatenation reactions initiated by oxidation of benzene. Previous enquiry using HO^• has shown that the intermediate products of benzene oxidation include organic radicals that can propagate the chain.^28,48 For case, the reaction of HO^• with benzene results in formation of a cationic radical that is subsequently hydrolyzed to hydroxyl-cyclohexadienyl (HCHD) radical.⁴⁸ HCHD radical farther reacts with oxygen to produce an organic peroxy radical that generates HO_two ^• and phenol. Under circumneutral pH weather condition, HO_ii ^• will deprotonate to superoxide (O_ii ^•–), which can undergo iron-catalyzed dismutation:

11

12

During this process, the faster rates of reduction of Fe(III) relative to oxidation of Atomic number 26(2) favor the presence of reduced Fe(II),⁴⁹ thereby contributing to oxidant decomposition. H₂O₂ produced past superoxide dismutation reacts quickly with iron and manganese through Fenton-like reactions on mineral surfaces.^13,14,17

Presumably a similar radical concatenation machinery occurs during the reaction of benzene with SO₄ ^•–. A second-order rate constant of 3 × 10^nine Grand^–ane·s^–one has been reported for the initial step of this reaction:⁵⁰

13

Because the branching ratio of reaction thirteen is much bigger than those of reactions 5 and 7 under experimental weather condition used in this study (i.due east., yard ₁₃[benzene] ≫ 1000 _v[S₂O₈ ^ii–]; thou ₁₃[benzene] ≫ thou ₇[H₂O], reaction rate constants provided in SI Table S4), essentially all SO₄ ^•– reacted with benzene instead of persulfate or water when benzene was nowadays.

The acceleration of persulfate decomposition by benzene was besides observed in experiments using synthetic groundwater that contained chloride, bromide, carbonate and natural organic matter (SI Table S2). In these experiments, the rates of persulfate decomposition by iron and manganese-containing minerals were consistent with those measured in benzene solutions prepared with borate-buffered water (SI Table S3).

Under typical groundwater atmospheric condition, boosted reactions between SO₄ ^•– and chloride or bromide tin can take place:^44,51

14

15

Cl^• and Br^• radicals can further react with water to produce other brusque-lived radicals.⁵²⁻⁵⁴ However, in the presence of 1 mM benzene and under typical groundwater weather condition, chloride and bromide are pocket-sized sinks for sulfate radical (i.due east., yard ₁₃[benzene] > k ₁₄[Cl^–]; k _thirteen[benzene] > one thousand ₁₅[Br^–]). Therefore, groundwater constituents only had a small impact on the scavenging of Then_iv ^•– radicals in the presence of benzene.

In addition, under groundwater atmospheric condition, And then₄ ^•– generated through reaction 3 can react with bicarbonate:⁴⁷

16

CO_iii ^•– can then react with persulfate to cause farther decomposition of persulfate:

17

All the same, for the range of concentrations studied in these experiments ([HCO_three ^–] = ane–10 mM and [Southward₂O₈ ^2–] = 1 mM) and in the presence of 1 mM benzene, bicarbonate is non a major sink for sulfate radical (i.e., yard ₁₆[HCO₃ ^–] < k ₁₃[benzene]) and reaction 17 will not be important to persulfate decomposition. This is consistent with observations from experiments in which persulfate decomposition was measured in the presence of benzene at varying concentrations of bicarbonate (SI Figure S6).

The presence of And then₄ ^•– or HO^• during persulfate decomposition in heterogeneous systems was confirmed by measurements of phenol production. In heterogeneous systems with Fe(III)- and Mn(IV)-containing minerals, an initial persulfate concentration of ane mM and an initial benzene concentration of 100 μM, phenol was formed in the presence of different solids (Figure ​ 5). For example, approximately 25 μM of phenol was generated in the presence of 50 g/50 goethite, nontronite, and montmorillonite afterward 30 days. Fifteen micromolar of phenol was produced in pyrolusite, and 5 μM, in ferrihydrite. Benzene loss ranged from twenty to 40 μM in these experiments.

Formation of phenol as oxidation product of benzene by sulfate radical in the presence of fifty g/L dissimilar Fe- or Mn-containing oxides and aquifer materials. Initial benzene concentration 100 μM, initial persulfate concentration 1 mM, pH viii.0, ionic strength 55 mM.

Environmental Implications

The findings from this study accept significant implications for the employ of persulfate for ISCO applications. First, the observation that fe- and manganese-containing minerals decompose persulfate under conditions encountered in the subsurface allow for prediction of the rate at which SO_iv ^•– or HO^• is produced in the absence of added activators, such as heat or base. In aquifers dominated by pure sand, persulfate is very stable with one-half-lives of more ii years. If Fe(III)- or Mn(IV)-containing oxides comprise greater than approximately 2% of aquifer solids, the rate of persulfate decomposition will increase, with half-lives decreasing to less than a twelvemonth. When persulfate encounters a zone that is rich in iron- or manganese-oxides or dirt, its half-life tin can decrease essentially.

Second, one time persulfate is injected into the subsurface, its treatment efficacy may exist affected past the distance that it travels in the surface. In many situations, information technology may be preferable for persulfate to persist in the aquifer for many months to allow it to reach contaminants afar from injection wells. Results from this study prove that persulfate compares favorably to H₂O_ii-based ISCO systems in this respect. Persulfate decomposition was much slower than hydrogen peroxide decomposition in the Fenton organisation, with a half-life that was more than than 2 orders of magnitude longer (SI Effigy S7). Hydrogen peroxide was decomposed over several hours (SI Figure S8A), whereas persulfate decomposition took several weeks (SI Figure S8B). Considering different hydraulic conductivities in the subsurface and uncertainties associated with the rate of oxidant transport in the aquifer, the high stability of persulfate in uncontaminated aquifer zones indicates more effective ship and delivery of oxidants.

Finally, the rate of persulfate decomposition accelerates when the injected oxidant reaches a contaminant feather. The half-life of persulfate in the aquifer can be more one twelvemonth in the absence of an organic contaminant. Once persulfate encounters a contaminant such as benzene, radical chain reactions increase the persulfate decomposition rate and its half-life can decrease to as piffling equally a few days. Its half-life tin can likewise decrease in the presence of iron- or manganese-containing oxides or clays, which are frequently associated with low hydraulic conductivity regions where contaminants frequently accrue.

Acknowledgments

This written report was supported by the U.Due south. National Constitute for Environmental Health Sciences (NIEHS) Superfund Inquiry Program (Grant P42 ES004705) and the Superfund Research Center at Academy of California, Berkeley. Nosotros thank our grouping member Anh Pham for stimulating discussions and Dr. Urs Jans for his aid in conducting initial experiments that helped us to develop a means of studying these slow reactions.

Funding Statement

National Institutes of Health, The states

Supporting Information Available

Additional description of fe and manganese mineral preparation and viii figures on the rate of persulfate decomposition and product formation in different atmospheric condition. This fabric is available costless of charge via the Internet at http://pubs.acs.org.

Author Present Address

^‡ Department of Chemical and Environmental Applied science, Academy of California at Riverside, Riverside, CA 92521, The states.

Notes

The authors declare no competing financial interest.

Supplementary Fabric

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Source: https://www.ncbi.nlm.nih.gov/pmc/articles/PMC4151705/

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