Influences of carbonate and chloride ions on persulfate oxidation of trichloroethylene at 20 °C

Abstract

Application of in situ chemical oxidation (ISCO) involves application of oxidants to contaminants such as trichloroethylene (TCE) in soil or groundwater in place. Successful application of ISCO at a hazardous waste site requires understanding the scavenging reactions that could take place at the site to better optimize the oxidation of target contaminants and identification of site conditions where ISCO using persulfate may not be applicable. Additionally, estimation of the oxidant dose at a site would need identification of groundwater constituents such as alkalinity and chlorides that may scavenge radicals and therefore use up the oxidant that is targeted for the contaminant(s). The objective of this study was to investigate the influence of various levels of chloride and carbonates on persulfate oxidation of TCE at 20 °C under controlled conditions in a laboratory.

Based on the results of the laboratory experiments, both chloride and alkalinity were shown to have scavenging effects on the rate of oxidation of TCE. It was found that at a neutral pH, persulfate oxidation of TCE was not affected by the presence of bicarbonate/carbonate concentrations within the range of 0–9.20 mM. However, the TCE degradation rate was seen to reduce with an increase in the level of carbonate species and at elevated pHs. TCE degradation in the presence of chlorides revealed no effect on the degradation rate especially at chloride levels below 0.2 M. However, at chloride levels greater than 0.2 M, TCE degradation rate was seen to reduce with an increase in the chloride ion concentration. Prior to application of persulfate as an oxidant, a site should be screened for the presence of scavengers to evaluate the potential of meeting target cleanup goals within a desirable timeframe at the site.

Introduction

Trichloroethylene, also referred to as TCE, is a dense non-aqueous phase liquid (DNAPL) commonly used as a cleaning solvent in many industries. The Agency for Toxic Substances and Disease Registry (ATSDR, 2003) reports that TCE is the most frequently reported organic contaminant in groundwater. Because TCE has low water solubility near 1100 mg/L, the environmental health risks related to TCE have been a cause of concern. Although TCE has a high vapor pressure, TCE does not evaporate easily from the soil and can persist in subsurface for long periods of time. In situ chemical oxidation (ISCO) by strong oxidants is one of the promising remedial alternatives for destroying organic contaminants such as TCE in a subsurface medium (Amarante, 2000, Interstate Technology Regulation Cooperation (ITRC), 2005). One of the recently used oxidants for ISCO application is sodium persulfate (Na₂S₂O₈). Some properties of persulfate anion that make it attractive for the ISCO application include: high water solubility (saturated solution: 2.5 M Na₂S₂O₈ at 20 °C) (Behrman and Dean, 1999), no odor, effectiveness of oxidation (redox potential (E°) = 2.01 V) over a wide range of pH (Latimer, 1952), and lower affinity for soil organics (e.g., at 15 °C in subsurface) (Brown et al., 2001). The sulfate free radical (SO₄⁻·) with E°(SO₄⁻·SO₄²⁻) = 2.40 V (Huie et al., 1991) has been extensively studied as an intermediate radical oxidant in the thermal-, metal-, and photochemical-activated decomposition of persulfate anion.

Thermal activation (House, 1962):

S₂O₈²⁻ + heat → 2SO₄⁻·

Metal activation (Kolthoff and Miller, 1951):

S₂O₈²⁻ + Fe²⁺ → SO₄⁻· + Fe³⁺ + SO₄²⁻

Photochemical activation (Neta et al., 1977):

S₂O₈²⁻ + e⁻ → SO₄⁻· + SO₄²⁻

When persulfate is used for ISCO application at relatively low temperatures (e.g., < 20 °C), the oxidation reactions are usually less aggressive due to a slow generation rate of SO₄⁻· (see Eq. (1)). The oxidation of a target contaminant may be accelerated by activation of persulfate, which increases the rate of persulfate decomposition and thereby increases the rate of sulfate free radical formation (Liang et al., 2003). However, in the application of persulfate to total organic carbon (TOC) analysis (e.g., 100 °C), it has been suggested (Peyton, 1993) that at a high reaction temperature the mineralization efficiency of TOC is reduced by a quick release of the sulfate free radicals. Based on Peyton's research, it was hypothesized that by reducing the activation temperature, sulfate radicals may be effectively utilized for the oxidation of target contaminants in the presence of radical scavengers typically found in groundwater. However, reducing the activation temperature is expected to slow the rates of the reactions. Nevertheless, when the sulfate free radical serves as an oxidant it accepts a single electron resulting in the formation of the sulfate anion (SO₄²⁻) as follows (Huie et al., 1991):

SO₄⁻· + e⁻ → SO₄²⁻   E^o = 2.40 V

A high redox potential of the sulfate free radical makes it very reactive in destroying organic contaminants. However, competing side reactions with various species in groundwater system other than the target contaminant (e.g., TCE) can result in scavenging of SO₄⁻· and could possibly limit its oxidation efficiency. Competition for SO₄⁻· could be from reactions with groundwater constituents such as chloride ions (also a byproduct of TCE degradation) and carbonate species. The chemical mechanism and rate constants of persulfate with carbonates or chlorides in the aqueous phase include (but not limited to) the following reactions:

$${\mathrm{SO}}_4^{-}·+{\mathrm{SO}}_4^{-}·\stackrel{k_5}{\to }{\mathrm{S}}_2{\mathrm{O}}_8^{2-},k_5=4\times {10}^8{\text{M}}^{-1}{\text{s}}^{-1}$$

(Huie and Clifton, 1989)

$${\mathrm{SO}}_4^{-}·+{\mathrm{S}}_2{\mathrm{O}}_8^{2-}\stackrel{k_6}{\to }{\mathrm{SO}}_4^{2-}+{\mathrm{S}}_2{\mathrm{O}}_8^{-}·,k_6=6.1\times {10}^5{\text{M}}^{-1}{\text{s}}^{-1}$$

(Buxton et al., 1999)

$${\mathrm{SO}}_4^{-}·+{\mathrm{Cl}}^{-}\underset{k_{7r}}{\overset{k_{7f}}{\leftrightarrow }}{\mathrm{SO}}_4^{2-}+\mathrm{Cl}·,k_{7f}=4.7\times {10}^8;k_{7r}=2.5\times {10}^8{\text{M}}^{-1}{\text{s}}^{-1}$$

$$\mathrm{Cl}·+{\mathrm{Cl}}^{-}\underset{k_{8r}}{\overset{k_{8f}}{\leftrightarrow }}{\mathrm{Cl}}_2^{-}·,k_{8f}=8\times {10}^9{\text{M}}^{-1}{\text{s}}^{-1};k_{8r}=4.2\times {10}^4{\text{s}}^{-1}$$

$${\mathrm{Cl}}_2^{-}·+{\mathrm{Cl}}_2^{-}·\stackrel{k_9}{\to }2{\mathrm{Cl}}^{-}+{\mathrm{Cl}}_2,k_9=1.3\times {10}^9{\text{M}}^{-1}{\text{s}}^{-1}$$

(Huie et al., 1991)

$$\mathrm{Cl}·+{\mathrm{H}}_2\mathrm{O}\underset{k_{10r}}{\overset{k_{10f}}{\leftrightarrow }}{\mathrm{ClHO}}^{-}·+{\mathrm{H}}^{+},k_{10f}=1.3\times {10}^3;k_{10r}=2.1\times {10}^{10}{\text{M}}^{-1}{\text{s}}^{-1}$$

$${\mathrm{ClHO}}^{-}·\underset{k_{11r}}{\overset{k_{11f}}{\leftrightarrow }}\mathrm{HO}·+{\mathrm{Cl}}^{-},k_{11f}=6.1\times {10}^9{\text{s}}^{-1};k_{11r}=4.3\times {10}^9{\text{M}}^{-1}{\text{s}}^{-1}$$

(Kiwi et al., 2000)

$${\mathrm{Cl}}_2^{-}·+{\mathrm{H}}_2\mathrm{O}\stackrel{k_{12}}{\to }\mathrm{Cl}{\mathrm{HO}}^{-}·+{\mathrm{H}}^{+}+{\mathrm{Cl}}^{-},k_{12}\left[{\mathrm{H}}_2\mathrm{O}\right]<100{\text{s}}^{-1}$$

(Yu and Baker, 2003)

$${\mathrm{SO}}_4^{-}·+{\mathrm{HCO}}_3^{-}\stackrel{k_{13}}{\to }S{O}_4^{2-}+{\mathrm{HCO}}_3·,k_{13}=(1.6\pm 0.2)\times {10}^6{\text{M}}^{-1}{\text{s}}^{-1}\left(\text{at}\mathrm{pH}8.4\right)$$

$${\mathrm{SO}}_4^{-}·+{\mathrm{CO}}_3^{2-}\stackrel{k_{14}}{\to }{\mathrm{SO}}_4^{2-}+{\mathrm{CO}}_3^{-}·,k_{14}=(6.1\pm 0.4)\times {10}^6{\text{M}}^{-1}{\text{s}}^{-1}(\text{at}\mathrm{pH}>11)$$

$${\mathrm{HCO}}_3·\stackrel{k_{15}}{\to }{\mathrm{H}}^{+}+{\mathrm{CO}}_3^{-}·,p{K}_a=9.5\pm 0.2$$

(Zuo et al., 1999)

Moreover, in aqueous solution water and hydroxyl ion (OH⁻) could also consume the sulfate radical formed. The reaction with SO₄⁻· could result in the formation of hydroxyl radical (OH·) according to Eqs. (16), (17) (Hayon et al., 1972):

$${\mathrm{SO}}_4^{-}·+{\mathrm{H}}_2\mathrm{O}\stackrel{k_{16}}{\to }\mathrm{OH}·+{\mathrm{SO}}_4^{2-}+{\mathrm{H}}^{+},k_{16}\left[{\mathrm{H}}_2\mathrm{O}\right]<3\times {10}^3{\text{M}}^{-1}{\text{s}}^{-1}$$

$${\mathrm{SO}}_4^{-}·+{\mathrm{OH}}^{-}\stackrel{k_{17}}{\to }\mathrm{OH}·+{\mathrm{SO}}_4^{2-},k_{17}=6.5\pm 1.0\times {10}^7{\text{M}}^{-1}{\text{s}}^{-1}$$

The above Eqs. (5), (6), (7), (8), (9), (10), (11), (12), (13), (14), (15), (16), (17) represent some of the simultaneous chemical reactions that occur immediately after SO₄⁻· is formed by either one of Eqs. (1), (2), (3). Eqs. (16), (17) show the possible coexistence of SO₄⁻· and OH·. The coexistence of both sulfate and hydroxyl radicals has been shown by electron spin resonance (ESR). However, Norman et al. (1970) reported that as the pH is increased above 7, the conversion of SO₄⁻· into OH· becomes increasingly important via Eq. (17). Additionally, a study using ESR to investigate the reactions of the sulfate radical with organic compounds by Dogliotti and Hayon (1967) revealed that as the pH is increased above 8.5, SO₄⁻· decays rapidly in aqueous solution by reacting with OH⁻ to generate OH· (Eq. (17)). Hence, it can be concluded that under acidic to neutral conditions (i.e., pH < 7) the sulfate free radical would be the predominant radical oxidant species.

Domenico and Schwartz (1990) have categorized groundwater constituents such as bicarbonate and chloride as the major dissolved constituents at concentrations greater than 5 mg/L in potable groundwater according to their relative abundance. The reactivity of SO₄⁻· in groundwater system might be affected by the presence of background ions (e.g., Cl⁻, HCO₃⁻/CO₃²⁻). For the ISCO applications, transition metal activators (e.g., Fe²⁺) (see Eq. (2)) or heat (see Eq. (1)) are usually added to induce the formation of sulfate free radials. However, as described earlier, a quick release of sulfate free radicals may not necessarily increase the rate of oxidation of the target contaminant. The SO₄⁻· formed may be scavenged due to cannibalization of SO₄⁻· by Eqs. (5), (6). Comparing the rate constants of Eqs. (5), (6), it may be noted that Eq. (6) is a minor sink for the radical compared with its self-reaction (Buxton et al, 1999). One of the ways to control the generation rate of sulfate free radicals and the efficiency of its use is to manage the activation temperature.

The objective of this study was to experimentally investigate the influence of various levels of chloride and carbonate concentrations on the thermally activated persulfate oxidation of TCE at 20 °C.