LIF study of the formation and consumption of mercury (I) chloride: Kinetics of mercury chlorination

Chemosphere 61 (2005) 685–692 www.elsevier.com/locate/chemosphere

LP/LIF study of the formation and consumption of mercury (I) chloride: Kinetics of mercury chlorination Philip H. Taylor *, Rajesh Mallipeddi, Takahiro Yamada Research Institute, Environmental Engineering Group, University of Dayton, 300 College Park, Dayton, OH 45469-0114, USA Received 18 January 2005; received in revised form 10 March 2005; accepted 15 March 2005 Available online 12 May 2005

Abstract The laser photolysis/laser induced fluorescence (LP/LIF) technique has been applied to studies of gas-phase mercury (Hg) chlorination. Mercury (I) chloride (HgCl) was been detected via LIF at 272 nm from reactions of elemental Hg with Cl atoms generated from the 193 nm photolysis of carbon tetrachloride. While the formation of HgCl was too fast to be observed on millisecond time scales, the kinetics of the consumption of HgCl have been determined at temperatures characteristic of post-combustion conditions. Rate coefficients and Arrhenius parameters for the reaction of HgCl with Cl2, HCl and Cl atoms were determined. The reaction of HgCl with Cl2 was the fastest reaction studied, while the reaction of HgCl with HCl was the only reaction to show any measurable temperature dependence. Estimates of the rate coefficient for the reaction Hg + Cl ! HgCl were determined using a modeling approach. Comparisons of these new measurements with model predictions are discussed.
2005 Elsevier Ltd. All rights reserved. Keywords: Mercury; Chlorination; Kinetics; Combustion

1. Introduction Coal-fired electric generation stations have been identified by the US–EPA as a significant source of anthropogenic mercury emissions to the environment. In 1999, EPA initiated an Information Collection Request (ICR) regarding Hg associated with coal-fired power-generation facilities across the US. The results from this ICR indicate that US coal-fired power plants emit 45 ton Hg/yr and thus pose a potentially significant risk to the environment (Electric Power Research Institute, 2000). As a result of these findings, EPA plans to regulate Hg emissions from coal-fired power plants under * Corresponding author. Tel.: +1 937 229 3604; fax: +1 937 229 2503. E-mail address: [email protected] (P.H. Taylor).

provisions of Title III of the Clean Air Act Amendments of 1990 (US Environmental Protection Agency, 1990) with regulatory compliance scheduled for December 2007 (Brown et al., 2000). Hg is common in coal at a reasonably low level at an average of
0.1 mg/kg. However, at a total annual US coal consumption rate of almost 800 million tons,
75 tons of Hg are burned and
60% of the Hg is emitted to the environment (Electric Power Research Institute, 2000). Hg can be emitted in either the oxidized form (Hg2+) or in the elemental form (Hg0). Either form will eventually reach a waterway where some fraction is methylated into the more toxic methylmercury, exposing the human population through consumption of contaminated fish. Because of the serious health effects of methylated mercury, assessing the contribution of coal-fired power

0045-6535/$ - see front matter
2005 Elsevier Ltd. All rights reserved. doi:10.1016/j.chemosphere.2005.03.089

686

P.H. Taylor et al. / Chemosphere 61 (2005) 685–692

plant Hg to the environmental is critical. Power plant emissions are typically reduced with existing environmental controls, including flue gas desulfurization (FGD), selective catalytic reduction (SCR), and electrostatic precipitators (ESPs). Based on EPAÕs ICR database, coal-fired power plants in the US demonstrate an average 40% Hg capture, but the amount of capture varies widely, from 0% to 99% (Kilgroe et al., 2000). The reasons for this variability are poorly understood, but appear to relate to the oxidation state of the mercury, the properties of the mineral matter associated with the coal, and the type of existing air pollution control equipment installed on the furnace. The essence of the problem is that the fundamental pathways governing the fate of mercury in the post-combustion environment are not known. These fundamental issues will, however, ultimately determine if mercury is retained with the ash or emitted with the stack gas. Oxidized mercury from coal combustion is generally thought to be HgCl2. Relative to Hg0, HgCl2 is slightly less volatile at stack temperatures and at lower ambient temperatures. A key differentiation aspect of these two types of mercury is that HgCl2 is water-soluble and that it tends to interact with the mineral matter and char, and with cold-end air pollution control equipment. Elemental mercury is much less water-soluble and has a much lower tendency to interact with fly ash constituents and air pollution control equipment. This is believed to be the source of the positive correlation between the fraction of mercury in the oxidized state and the removal of mercury from the flue gas (Meij, 1991). The factors that control the division of mercury between the elemental and oxidized states are thought to be of critical importance in understanding mercury emissions. Emissions of trace elements and mercury can be predicted with some success based on their occurrence in the fuel, transformation in the furnace, and ability to penetrate air pollution control devices. Since mercury in a flame is completely converted to volatile elemental mercury, questions regarding its transformations are concerned with cooled, condensed phases (Sarofim et al., 1998). Equilibrium calculations suggest that mercury may condense as HgO, HgCl2, or HgSO4, but nitrates and other forms may also occur as intermediates. The equilibrium conversion to Hg+2 would be essentially complete upon cooling to about 400 C, but measurements of boiler emissions indicate only 35– 95% in the oxidized form, suggesting that kinetic limitations are controlling. Gas-phase kinetic studies at low to moderate temperatures should provide improved insight for predicting the relative concentration of elemental and ionic mercury in the post-combustion zones of both coal-fired power plants and thermal disposal facilities. Recently, room temperature measurements of the reaction of Hg

with Cl and Cl2 have been reported in the literature in relation to its atmospheric transformation (Ariya et al., 2002). These results indicate that the reaction of Hg with Cl atoms is fast (k298 = 1.0 · 1011 cm3 molecule1 s1) while the reaction of Hg with molecular chlorine (Cl2) is negligibly slow (k298 = 2.6 · 1018 cm3 molecule1 s1). The rates of these reactions, particularly Hg + Cl, at elevated temperatures characteristic of post-combustion systems are not known. This manuscript addresses the kinetics of Hg chlorination under post-combustion conditions. Specific objectives of this work were to address the following questions: (1) What is the reactivity of Hg0 with post-combustion chlorine-containing constituents (Cl, Cl2, and HCl) for a range of temperatures that are characteristic of post-combustion conditions? (2) What are the rates of reaction for conversion of Hg0 to HgCl and then to HgCl2?

2. Experimental approach The gas-phase kinetics of Hg chlorination under post-combustion conditions was investigated using state-of-the art kinetic techniques. A laser photolysis/ laser induced fluorescence (LP/LIF) technique was used to study the kinetics of Hg and HgCl chlorination. The following reactions were investigated: Hg þ ðCl; Cl2 ; HClÞ ! reaction products HgCl þ ðCl; Cl2 ; HClÞ ! reaction products The general approach is based on recently published studies of OH radical reactions with chlorinated hydrocarbons and oxygenates (Yamada et al., 2001, 2003). A fused-silica reactor equipped with orthogonal optical access for pump and probe lasers and LIF detection was used for the rate determinations (see Fig. 1). The reactor was wrapped with Nichrome wire with an operational temperature capability of 295–1323 K. The uncertainty in the reactor temperature, based on experimental measurements with a retractable thermocouple, was ±2 K between 295 and 700 K. Gas-phase mercury was introduced into the reactor using concentrationcertified permeation tubes (VICI, Inc.). For HgCl production the following method was employed: photolytic generation of Cl atoms followed by reaction of Cl atoms with vapor-phase Hg. The following Cl atom sources were tested: (1) 193 nm photolysis of carbon tetrachloride (CCl4) and (2) 351 nm photolysis of Cl2. The photolysis laser was a pulsed (10 Hz) excimer laser (Lambda Physik, Compex Model 102). Laser-induced fluorescence was used to monitor the formation of HgCl in the Hg chlorination studies and the consumption of HgCl in studies of its chlorination. A pulsed

P.H. Taylor et al. / Chemosphere 61 (2005) 685–692

687

Photomultiplier Tube (PMT) Narrow Bandpass Filters

f/6 Condenser Lens Exhaust

Nichrome Wire Pump Laser

Window Holder Fused Silica Window

O-ring Seal Thermocouple

Probe Laser Sample Inlet Probe

Sample carrier(Argon) flow Water Bath

Cl2/HCl CCl4 Flow

Electric Heater Hg Permeation Tubes

Main Carrier Flow Argon

Fig. 1. Schematic of fused silica reactor as modified for elemental mercury introduction.

Nd:YAG pumped dye laser (Continuum, Model YG682-20/TDL-51,
1 W cm2 pulse1 @280 nm) was used as the excitation source. Fluorescence was isolated with a narrow bandpass filter/PMT combination. A scanning gate boxcar averager (SRS Model 255) coupled to electronic time delay of the pump and probe laser was used for data collection. Signal averaging consisted of 30 laser pulses at each time delay, reducing the effects of laser and PMT instability. Computerized data collection, storage, reduction, and statistical error analysis was performed after each set of experiments. Argon was used as diluent and carrier gas for the elemental Hg and chlorinated substrates. Each gas was separately metered using mass flow controllers (Porter Instruments, Inc.). Sample lines to the reactor consisted of Teflon, an inert material for quantitative transport of Hg. The vapor pressure of Hg was sufficiently large that sufficient quantities (0.07–2.0 ppm) were delivered to the reactor by passing a stream of Ar carrier gas through a series of Hg permeation tubes held at constant temperature (373–403 K). Cl2 and HCl were introduced using

calibrated gas standards. Cl atoms were generated in the reactor via laser photolysis of suitable precursors. Experiments were conducted under slow-flow conditions, with total gas flow rate of 550–560 cm3/min at an inlet temperature and pressure of 298 K and 745 torr, respectively. Under these conditions, it was expected that secondary reactions will be negligible in their contribution to the measured rate because of the low initial chlorinated substrate concentrations and the single reaction conditions. The single reaction assumption was based on the dimensional thickness of the pump laser pulse in the direction of the flow (ca. 1 cm), the laser repetition rate (10 Hz), and an estimated linear gas velocity of P10 cm s1. Thus, reactive fragments or molecules potentially produced by a variety of processes were swept out of the reaction volume before having an opportunity to react and cause a change in the measured rate. Temperature-dependent reaction kinetic data were collected at reactor temperatures ranging from 398 to 673 K, consistent with the regions downstream of the economizer and upstream of the wet scrubber of a

688

P.H. Taylor et al. / Chemosphere 61 (2005) 685–692

coal-fired boiler. All data were collected at atmospheric pressure, i.e., 735 ± 10 torr.

SeveralPabsorption lines were available from the 2 ðC X Þ transition of HgCl ranging from P1=2 271 to 281 nm (Horne et al., 1968). Fluorescence was detected at 278 nm with a narrow bandpass filter (Omega Optical, Inc.)/PMT combination. The excitation laser wavelength was 272 nm. Experiments at an excitation laser wavelength of 270 nm, where HgCl absorption does not occur, did not produce a fluorescence signal. Reduction in the Hg concentration (from 2.0 to 0.07 ppm) resulted in a dramatic decrease in the signal. Experiments using CCl4 as the Cl atom source (193 nm photolysis) produced the optimum conditions for HgCl LIF. No signal was observed using Cl2 as the source (351 nm photolysis). This was attributed to a lower efficiency for Cl atom production and the rapid reaction of HgCl with excess Cl2 (see results for HgCl + Cl2). The fluorescence signal was a function of Cl atom concentration below a threshold of ca. 10 ppm. At higher Cl atom concentrations, more rapid decays were observed, but the initial signal was approximately constant (see HgCl + Cl experiments). 2

3.2. HgCl + Cl2 ! products The reactions between HgCl and chlorine gas (Cl2) were performed under pseudo-first-order conditions (Cl2 in excess). Initial Hg and Cl atom concentrations were 2 and 20 ppm, respectively. The initial HgCl concentration was limited by the initial Hg concentration. The experiments were performed between 423 and 673 K. Fig. 2 presents semi-logarithmic decay plots of HgCl as a function of Cl2 concentration, providing evidence of the pseudo-first-order conditions. Fig. 3 illustrates HgCl decay rates as a function of Cl2 concentration for various reaction temperatures. Reaction of HgCl with Cl2 was fast with a rate coefficient of 1.2 ± 0.1 · 1011 cm3 molecule1 s1, independent of temperature. Table 1 presents a summary of the rate measurements.

HgCl Fluorescence Signal

3.1. Verification of HgCl fluorescence

[Cl2] = 60 ppm

0.1

0

100

200

300

400

500

Time (µs)

Fig. 2. HgCl decay profiles in the presence of excess molecular chlorine. P = 735 ± 10 torr, T = 523 K.

30000

25000 623 K

523 K

20000

573 K 673 K

15000

10000

5000

0

0

2

4

6

8

10

12

Cl2 concentration (1014 molecule/cm3)

Fig. 3. First-order HgCl decay rates in the presence of excess molecular chlorine at various reaction temperatures. P = 735 ± 10 torr.

Table 1 Experimental results—HgCl + Cl2 [Cl2] ppm

Reactor kbi (1011 cm3 ±2r Error limitsa temperature, molecule1 s1) (1012 cm3 K molecule1 s1)

25–75 25–75 25–75 25–75 25–75 25–75

423 473 523 573 623 673

3.3. HgCl + HCl ! products The reaction between HgCl and hydrochloric acid (HCl) was performed under pseudo-first-order conditions with HCl in excess. Initial Hg, Cl atom and HgCl concentrations were the same as for the Hg + Cl2 experiments. HCl concentrations were varied between 25 and 500 ppm. A measurable increase in the background HgCl decay rate was only observed at HCl concentra-

[Cl2] = 40 ppm

0.01

k' (sec-1)

3. Results

[Cl2] = 25 ppm

1

1.18 1.28 1.59 1.22 1.49 1.08

1.64 1.14 1.94 0.49 3.02 1.90

a Error limits represent random errors and do not include the 5–10% uncertainty for possible systematic errors.

P.H. Taylor et al. / Chemosphere 61 (2005) 685–692

tions above 100 ppm. Kinetic determinations were performed between 398 and 673 K. Reaction of HgCl with HCl was considerably slower than with Cl2. A rate coefficient of 6.3 · 1013 cm3 molecule1 s1 was measured at 398 K, increasing to 1.1 · 1012 cm3 molecule1 s1 at 673 K. The reaction exhibited non-Arrhenius behavior (see Fig. 4) and was best fit by the following modified Arrhenius equation (in units of cm3 molecule1 s1): kð398  673 KÞ ¼ 9.9  1019 T 2 expð548ðKÞ=T Þ. Table 2 presents a summary of the rate measurements. 3.4. HgCl + Cl ! products We have also derived rate measurements for the reaction of HgCl + Cl atoms. Cl atoms were generated by photolysis of CCl4. Variation of the CCl4 concentration at a fixed photolysis rate produced a wide range of Cl atom concentrations. We derived the Cl atom concentra1.3

-12

k(10

tion based on the measured photolysis laser intensity (6 mJ cm2), absorption cross-section for CCl4 at 193 nm (8.7 · 1019 cm2 molecule1, DeMore et al., 1997) and the quantum yield for the reaction (DeMore et al., 1997) and derived values ranging from 10 to 190 ppm for our experimental conditions. The Cl atom concentration was verified in separate experiments where the CCl4 photolysis products (C2Cl4 and C2Cl6) were quantified. The C2Cl4 and C2Cl6 photolysis products were assumed to form via self-recombination of the CCl2 and CCl3 intermediates that were generated from the photolysis of CCl4. The Cl atom concentrations were sufficiently larger than the estimated HgCl concentration (limited by the Hg concentration to 6 2 ppm) to permit pseudo-first-order kinetic measurements. Fig. 5 presents a plot of HgCl decay rates as a function of Cl atom concentration at various reaction temperatures. The experimental rate measurements did not show a clear temperature dependence. Although a slight positive temperature dependence could be inferred from the measurements, a temperature-independent rate coefficient of ca. 4 ± 1 · 1012 cm3 molecule1 s1 between 398 and 573 K was considered representative of the data and the presumed radical combination reaction mechanism. Table 3 presents a summary of the rate measurements.

1.1

3.5. Hg + Cl ! products 1

For our reaction conditions, HgCl LIF reached maximum levels in <50 ls, very near the limit of the temporal resolution of our LP/LIF system (25 ls). Thus, we were unable to clearly establish conditions where the formation of HgCl could be accurately determined. In

0.9

3

cm /molecules-s)

1.2

689

0.8 0.7

18000

0.6 0.5 1.4

16000 1.6

1.8

2.0

2.2

2.4

2.6

-1

1000/T (K )

Table 2 Experimental results—HgCl + HCl 12

12000

10000

8000 3

a

[HCl] ppm

Reactor temperature, K

kbi (10 cm molecule1 s1)

±2r Error limits (1013 cm3 molecule1 s1)

25–500 25–500 25–500 25–500

398 473 573 673

0.63 0.70 0.83 1.13

0.43 1.10 0.42 1.63

a

k' (sec-1)

14000

Fig. 4. Arrhenius plot of HgCl + HCl data. P = 735 ± 10 torr. Error bars denote ±2r random deviations from the measured rate coefficient.

Error limits represent random errors and do not include the 5–10% uncertainty for possible systematic errors.

398 K 473 K 573 K

6000

4000 0

5

10

15

20

25

30

Cl concentration (1014 molecules/cm3)

Fig. 5. First-order HgCl decay rates in the presence of excess chlorine atoms at various reaction temperatures. P = 735 ± 10 torr.

690

P.H. Taylor et al. / Chemosphere 61 (2005) 685–692

Table 3 Experimental results—HgCl + Cl [Cl] ppm

Temperature, kbi (1012 cm3 ±2r Error limitsa K molecule1 s1) (1012 cm3 molecule1 s1)

20–156 10–193 40–156

398 473 573

3.14 3.07 4.93

0.82 0.46 0.59

a

Error limits represent random errors and do not include the 5–10% uncertainty for possible systematic errors.

lieu of experimental results, modeling calculations were performed to provide an estimate of the rate coefficient for this reaction. Using the model shown in Table 4, our new experimental rate measurements for HgCl + Cl2 (R4) and HgCl + Cl (R3), and literature values for HgCl + HgCl (R2) and Cl + Cl + M (R5), we equated the predicted HgCl decay rate to the experimental observed HgCl decay rate by varying the rate coefficient for Hg + Cl. The rate coefficients for the other reactions were not adjusted. Initial Hg and Cl atom concentrations in the model were 2 and 20 ppm, respectively. An HgCl decay rate of 5340 ± 1110 s1 was used for these calculations. This was the average background HgCl decay rate determined from the HgCl + Cl2, HgCl + HCl, and HgCl + Cl kinetic measurements discussed above. The derived rate coefficient for Hg + Cl + M, R1, in third-order units, was 3.5 · 1030 cm6 molecule2 s2. The magnitude of the rate coefficient at atmospheric pressure (6.0 · 1011 cm3 molecule1 s1 in second-order units) indicates this reaction is very fast and consistent with the rapid formation of HgCl on our experimental time scales of ca. 25–50 ls. It is also within uncertainties of the measurements of Horne et al. (1968) (1.5 ± 4.5 · 1011 cm3 molecule1 s1). A recent study by Ariya et al. (2002) reported a rate of 1.0 ± 0.2 · 1011 cm3 molecule1 s1 at atmospheric pressure and 298 K. Ariya et al. measurements were performed in a Table 4 Modeling results molecule1 s1

at

423 K,

1 atm,

in

units

of

No.

Rate coefficient

Reaction mechanism

R1 R2a R3b R4b R5c

Variable 5.2 · 1010 6.0 · 1012 1.2 · 1011 5.5 · 1033

Hg + Cl + M ! HgCl + Md HgCl + HgCl ! Hg2Cl2 HgCl + Cl ! HgCl2 HgCl + Cl2 ! HgCl2 + Cl Cl + Cl + M ! Cl2 + Md

cm3

a The rate coefficient for R2 was taken from Horne et al. (1968). b R3 and R4 were based on the LP/LIF experimental measurements. c R5 is taken from the NIST Chemical Kinetics Database (2000). d Units of cm6 molecule2 s2.

flow tube at atmospheric pressure for reaction times on the order of several seconds. This contrasts with the measurements reported here that were measured at sub-millisecond time scales where surface reactions are practically non-existent. There is one previous theoretical analysis of this reaction. Khalizov et al. (2003) recently reported a variational transition state theory (high-pressure limit) result for this reaction of 3.9 · 1010 cm3 molecule1 s1 at 298 K. 3.6. Hg + Cl2, HCl ! products There was no evidence for formation of HgCl from these reactions, indicating the rate coefficients for these reactions were <1013 cm3 molecule1 s1 for our experimental conditions.

4. Discussion There are very limited prior measurements of gasphase Hg chlorination at temperatures relevant to postcombustion, coal-fired utility systems. Horne et al. (1968) reported a rate measurement for Hg + Cl ! HgCl at temperatures between ca. 383 and 443 K. The observed rate coefficient was very fast, however no rate data were reported for the subsequent reaction of HgCl + Cl ! HgCl2. The only other reported rate measurement for these reactions was by Ariya et al. (2002) at room temperature. The major motivation for these measurements was to verify prior estimates of the rates of these reactions considered relevant to the Hg chlorination reaction system. A comparison of our direct measurements with predicted values for these reactions are presented in Table 5. The direct experimental measurements (except for Hg + Cl) presented here indicate that substantial revision of prior estimates is needed to improve the predictive capability of these models. For example, for the HgCl + HCl reaction, our results indicate this reaction may be important at elevated temperatures. Sliger et al. (2000) reported that this reaction was of negligible importance to Hg chlorination for all temperatures, due primarily to a large activation energy of ca. 80 kJ/mol. The estimates of Senior et al. (2000) were much more consistent with our direct measurements, but overestimated both the absolute rate and the temperature dependence of the reaction. For the HgCl + Cl2 reaction, the estimates of Widmer et al. (1998) overestimated the rate coefficient by a factor of nearly 7 at 423 K, increasing to nearly a factor of 10 at 673 K. The prediction of a near zero activation energy was consistent with our measurements. The consecutive reactions Hg + Cl ! HgCl and HgCl + Cl ! HgCl2 have been modeled by Niksa et al. (2001). They assumed the reactions are pressure dependent at 1 atm. Our results, when com-

P.H. Taylor et al. / Chemosphere 61 (2005) 685–692

691

Table 5 Comparison with model predictions at atmospheric pressure Reaction

Source

Temperature, K

Rate coefficient (cm3 molecule1 s1)

This work (cm3 molecule1 s1)

Hg + Cl + M

Niksa et al.

398 473 573 673

9.1 · 1012 8.4 · 1012 7.6 · 1012 7.0 · 1012

6.0 · 1011 6.0 · 1011 6.0 · 1011 6.0 · 1011

HgCl + Cl + M

Niksa et al.

398 473 573 673

1.2 · 1012 1.1 · 1012 9.8 · 1013 9.1 · 1013

4.0 · 1012 4.0 · 1012 4.0 · 1012 4.0 · 1012

HgCl + Cl2

Widmer et al.

423 473 523 573 623 673

7.0 · 1011 8.0 · 1011 8.8 · 1011 9.6 · 1011 1.0 · 1010 1.1 · 1010

1.2 · 1011 1.2 · 1011 1.2 · 1011 1.2 · 1011 1.2 · 1011 1.2 · 1011

HgCl + HCl

Senior et al.

398 473 573 673

4.6 · 1012 9.5 · 1012 1.9 · 1011 3.0 · 1011

6.3 · 1013 7.0 · 1013 8.3 · 1013 1.1 · 1012

HgCl + HCl

Sliger et al.

398 473 573 673

7.9 · 1025 5.6 · 1023 3.1 · 1021 5.7 · 1020

6.3 · 1013 7.0 · 1013 8.3 · 1013 1.1 · 1012

pared to their predictions at 1 atm pressure, indicate a significantly faster rate of Hg chlorination than predicted. For Hg + Cl, the deviation is ca. a factor of 6– 10, with the deviation increasing at higher temperatures. For HgCl + Cl, the deviation is ca. a factor of 4, independent of temperature.

acknowledge the support of C.W. Lee, EPA Project Officer, Bob Brown of OCDO, and Jarek Karwowski of Arcadis. The authors also acknowledge Pavan Mopidevi and Rajendar Brahman of the University of Dayton Research Institute for assistance with the Cl atom measurements.

5. Conclusions

References

New rate measurements for the chlorination of Hg are reported. The gas-phase reactions of Hg with Cl and HgCl with Cl and Cl2 are fast as expected with a near-zero activation energy. The reaction of HgCl with HCl is much slower with a small activation energy. Comparison of previous predictions with these new measurements indicate revisions are needed to place these models on a more scientifically defensible basis and to more accurately predict the gas-phase component of Hg chlorination in the post-combustion zones of coalfired power plants.

Ariya, P.A., Khalizov, A., Gidas, A., 2002. Reactions of gaseous mercury with atomic and molecular halogens: kinetics, product studies, and atmospheric implications. J. Phys. Chem. A 106, 7310–7320. Brown, T.D., Smith, D.N., OÕDowd, W.J., Hargis, R.A., 2000. Control of mercury emissions from coal-fired power plants: a preliminary cost assessment and the next steps for accurately assessing control costs. Fuel Process. Technol. 65–66, 311–341. DeMore, W.B., Sander, S.P., Golden, D.M., Hampson, R.F., Kurylo, M.J., Howard, C.J., Ravishankara, A.R., Kolb, C.E., Molina, M.J., 1997. Chemical Kinetics and Photochemical Data for use in Stratospheric Modeling: Evaluation Number 12, JPL Publication 97-4, Jet Propulsion Laboratory. Electric Power Research Institute, 2000. An Assessment of Mercury Emissions from US. Coal-Fired Power Plants; EPRI Report TR-1000608; EPRI: Palo Alto, CA. Horne, D.G., Gosavi, R., Strausz, O.P., 1968. Reactions of metal atoms. I. The combination of mercury and chlorine

Acknowledgments This work was supported in part by contracts from the Environmental Protection Agency (EPA) and the Ohio Coal Development Office (OCDO). The authors

692

P.H. Taylor et al. / Chemosphere 61 (2005) 685–692

atoms and the dimerization of HgCl. J. Chem. Phys. 48, 4758–4764. Khalizov, A.F., Viswanathan, B., Larregaray, P., Ariya, P.A., 2003. A theoretical study on the reactions of Hg with halogens: atmospheric implications. J. Phys. Chem. A 107, 6360–6365. Kilgroe, J.D., Srivastava, R.K., Sedman, C.B., Thorneloe, S.A., 2000. Control of Mercury Emissions from Coal-Fired Electric Utility Boilers; Technical Memorandum; US Environmental Protection Agency, NRMRL/APPCD: Research Triangle Park, NC, October. Meij, R., 1991. The fate of mercury in coal-fired power plants and the influence of wet flue-gas desulphurization. Water, Air, Soil Pollut. 56, 21–33. NIST Chemical Kinetics Database, 2000. Standard Reference Database 17, Version 7.10 (Web Version), Release 1.2: http://kinetics.nist.gov/index.php. Niksa, S., Helble, J.J., Fujiwara, N., 2001. Kinetic modeling of homogeneous mercury oxidation: the importance of NO and H2O in predicting oxidation of coal-derived systems. Environ. Sci. Technol. 35, 3701–3706. Sarofim, A.F., Senior, C.L., Helble, J.J., 1998. Emissions of mercury, trace elements, and fine particles from stationary combustion sources. In: Proceedings of Conference on Air

Quality: Mercury, Trace Elements, and Particulate Matter, McLean, VA. Senior, C.L., Sarofim, A.F., Zeng, T., Helble, J.J., MamaniPaco, R., 2000. Gas-phase transformations of mercury in coal-fired power plants. Fuel Process. Technol. 63, 197–213. Sliger, R.N., Kramlich, J.C., Marinov, N.M., 2000. Towards the development of a chemical kinetic model for the homogeneous oxidation of mercury by chlorine species. Fuel Process. Technol. 65–66, 423–438. US Environmental Protection Agency, 1990. The Clean Air Act Amendments of 1990, Section 114; Government Printing Office: Washington, DC. Widmer, N.C., West, J., Cole, J.A., 1998. Practical limitation of mercury speciation in simulated municipal waste incinerator flue gas. Combust. Sci. Technol. 134, 315–326. Yamada, T., Siraj, M., Taylor, P.H., Peng, J., Hu, X., Marshall, P., 2001. Rate coefficients and mechanistic analysis for reaction of OH with vinyl chloride between 293 and 73 K. J. Phys. Chem. A 105, 9436–9444. Yamada, T., Taylor, P.H., Goumri, A., Marshall, P., 2003. The reaction of OH with acetone and acetone-d6 from 298 to 832 K: rate coefficients and mechanism. J. Chem. Phys. 119, 10600–10606.