Degradation of Environmental Contaminants with Water-Soluble

In the Laboratory

Degradation of Environmental Contaminants with Water-Soluble Cobalt Catalysts: An Integrative Inorganic Chemistry Investigation Alexandra L. Evans, Reid E. Messersmith, David B. Green, and Joseph M. Fritsch* Natural Science Division, Pepperdine University, Malibu, California 90263, United States *[email protected]

We present an integrative laboratory investigation incorporating skills from inorganic chemistry, analytical instrumentation, and physical chemistry applied to a laboratory-scale model of the environmental problem of chlorinated ethylenes in groundwater. Here, the kinetics of perchloroethylene (C2Cl4, PCE) degradation by transition-metal catalysis are monitored with gas chromatography to study the stepwise reductive dechlorination of PCE to less chlorinated ethylenes in aqueous media. Although not demonstrated here, the full dechlorination of PCE to a mixture of hydrocarbons (ethylene and acetylene) takes a carcinogenic compound to environmentally benign ones. This laboratory integrates concepts and skills gathered from different chemistry courses to investigate the kinetics of PCE degradation and can be adapted to a number of different courses including inorganic chemistry, analytical chemistry, or physical chemistry. The inorganic chemistry student will find the incorporation of transition-metal oxidation states, oxygen-sensitive compounds, and catalysis valuable to their laboratory experiences. Analytical chemistry courses will benefit from the incorporation of gas chromatography with either mass spectrometry or flame ionization detection into the curriculum. Finally, the determination of first-order rate constants from kinetic data obtained from an interesting and authentic system will be a useful teaching tool for the physical chemist as well. Chlorinated ethylenes represent an important class of environmental pollutants resulting from the improper disposal of dry cleaning solvent (PCE) or metal-degreasing solvent (trichloroethylene C2HCl3, TCE). At its peak, PCE was manufactured at an annual rate of 800 million pounds per year and TCE was manufactured at an annual rate of 650 million pounds per year (1, 2). Although the great majority was disposed of properly, PCE releases result in groundwater contamination that are generally localized to urban centers (3, 4). The presence of PCE and TCE in groundwater is a concern because of nerve, kidney, and liver damage in humans from prolonged exposure and the elevated cancer risks observed in animal studies (5-8). As a result, PCE is classified as a “potential” carcinogen and TCE is considered a “probable” carcinogen (5, 6). Covalent bonds between carbon and chlorine atoms are unusual in nature and there are few natural processes to break these compounds down and draw their elements back into environmental cycles. Methanogenic bacteria under anaerobic conditions were discovered to metabolize chlorinated ethylenes and the active degradative agent was determined to be cobalamin (vitamin B12, 1), a naturally occurring, cobalt-containing macrocyclic complex (9). In biological systems, the cobalt center is 204

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reduced from cobalt(III) to cobalt(I) followed by the catalytic degradation of the chlorinated ethylene.

In nonbiological systems, vitamin B12 has been shown to participate in the reductive dechlorination of chlorinated ethylenes in aqueous media (10, 11). One of the key features of the in vitro studies was the presence of titanium(III) citrate (the bulk reducing agent), which was able to produce the necessary cobalt(I) oxidation state for vitamin B12: vitamin B12 ðCoII Þ þ TiIII f vitamin B12 ðCoI Þ þ TiIV ð1Þ In the reductive dechlorination process, sequential replacement of Cl with H was observed:

Cobalt(I) contains the unusual combination of properties that it is both a strong reducing agent and a “super nucleophile”; that is, ∼104 times more nucleophilic than Cl- in SN2 type reactions (12, 13). As a result, vitamin B12 (CoI) is able to either reduce these molecules through electron transfer pathways or through a bond-forming nucleophilic attack on a carbon atom. Both are critical for the reductive dechlorination of chlorinated ethylenes: vitamin B12 ðCoI Þ þ PCE þ Hþ f vitamin B12 ðCoIII Þ þ TCE þ Cl -

ð3Þ

Cobalt experiences a two-electron oxidation in which one electron is used to reduce chlorine to chloride and the other to

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Vol. 88 No. 2 February 2011 pubs.acs.org/jchemeduc r 2010 American Chemical Society and Division of Chemical Education, Inc. 10.1021/ed1005344 Published on Web 10/25/2010

In the Laboratory

form the carbon-hydrogen bond in the products. When titanium(III) citrate is present in large excess, chlorinated ethylenes are degraded in a catalytic cycle with a small quantity of vitamin B12 until the bulk reducing agent is consumed. The exact mechanism for degradation of each chlorinated ethylene is not completely understood and has been an area of significant study including product analysis, kinetic analysis, and synthesis of possible catalytic-cycle intermediates (10, 11, 14-18). Other transition-metal complexes have been shown to catalytically degrade chlorinated ethylenes with dechlorination rates superior to vitamin B12 (19, 20). For example, the tetrakis(4-carboxyphenyl)porphyrin cobalt complex, (TCPP)Co (2), is a synthetic complex that mimics the chemistry of the biological molecule.

Figure 1. Reaction vessel showing the aqueous-gas phase partitioning of volatile organic compounds.

Overview of the Experiment Students are often exposed to chromatography, kinetics, and transition-metal complexes in separate courses, starting as early as general chemistry. This experiment challenges students to integrate, apply, and build upon skills obtained in previous courses and put them to use for understanding a chemical system with potential application to the remediation of a common groundwater contaminant. In addition, students gain experience with oxygen-sensitive reagents, catalysis, reaction monitoring, and kinetic data analysis. All of which can be completed during a standard 4-h laboratory period but can easily be extended over two laboratory periods. Experimental Section Materials Vitamin B12, titanium(III) chloride (20% w/w in 2 M HCl), methanol, and toluene were purchased from Acros. Perchloroethylene, trichloroethylene (unstabilized), and trisodium citrate were purchased from Sigma. Tris(hydroxymethyl)aminomethane (Tris), 40 mL EPA vials, and sodium hydroxide were purchased from Fisher. cis-Dichloroethylene was purchased from Pfaltz and Bauer. Tetrakis(4-carboxyphenyl)porphyrin cobalt was purchased from Porphyrin Systems. Additional materials include air-free storage tubes (Chem Glass), a 100 μL gas-tight, SampleLock syringe (Hamilton), a thermostatted water bath (Ika-Mag), and a nitrogen gas manifold (Ace Glass). Gas Chromatographic Conditions Gas chromatographic analysis reported here was performed using an SRI 8610C gas chromatograph (SRI, Torrance, CA) equipped with a split-splitless injector and flame ionization detector. Splitless injections with an inlet temperature of 155 °C were made on a MXT-1 column (30 m 0.5 mm, 5 μm film

r 2010 American Chemical Society and Division of Chemical Education, Inc.

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thickness, Restek). For PCE as the dechlorination substrate, the separation was carried out isothermally at 100 °C with a helium flow rate of 1 mL/min. For TCE as the substrate, an isothermal temperature program of 70 °C with a helium flow rate of 1 mL/min was used. The flame ionization detector temperature was 200 °C. Data Analysis During the course of reaction, the chromatographic peak areas for PCE, TCE, and toluene were recorded in addition to the time of sample collection. As the dechlorination reaction progressed, the headspace concentration of PCE decreased and became enriched in TCE (Figure 1). The toluene headspace vapor concentration remained constant because it does not undergo chemical change under the reaction conditions; thus, it functions as an internal standard. The toluene peak area was used to correct for variations in headspace gas volumes collected and injection technique variations, normalizing each injection's PCE and TCE peak area for any sampling differences. The PCE dechlorination step, eq 3, is an overall secondorder reaction with the rate law of rate = k[PCE][CoI catalyst] (10, 20). By assuming a constant cobalt catalyst concentration, the rate law simplifies to rate = kobs[PCE], where kobs = k[CoI catalyst] and the system can be studied with first-order kinetic models. Thus, the corrected PCE peak areas may be used without further transformation. The absolute [PCE] can be determined using an external calibration curve. If a calibration curve were to be employed, then vials should be prepared as for dechlorination experiments with different PCE concentrations but the catalyst should be omitted. The integrated rate law for the reaction is ! ½PCEt ð4Þ ¼ - kobs t ln ½PCE0

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In the Laboratory

Figure 2. Student-generated first-order kinetic data for the reductive dechlorination of PCE with vitamin B12 and (TCPP)Co; [CoI catalyst] = 1.8 μM, [TiIII] = 26 mM, pH 8, 35 °C, [PCE]i = 240 μM. The observed rate constants for vitamin B12 and (TCPP)Co were 3.88 10-5 s-1 and 1.58 10-4 s-1, respectively.

where [PCE]0 is concentration of substrate at the start of the reaction and [PCE]t is the concentration at the sampling time. Alternatively, because the chromatographic peak areas are linearly related to [PCE] and because first-order kinetics are obeyed, the actual [PCE] is not necessary. Thus, the concentrations in eq 4 can be replaced with PCE chromatographic peak areas for PCE (APCE) and toluene (Atoluene): 2 !3 APCE t 6 7 6 Atoluene 7 7 6 t !7 ¼ - kobs t ln6 ð5Þ 6 APCE 7 5 4 0 Atoluene 0 where the numerator is the PCE peak area at time, t, (AtPCE) corrected for injection variations by the internal standard area (Attoluene) and the denominator is the corrected PCE peak area at time zero. The slope of the best-fit line of either the log term in eq 4 or eq 5 versus time gives the observed rate constant (kobs) of the reaction (Figure 2). The rate of PCE degradation is dependent on the catalyst, catalyst concentration, pH, temperature, and substrate (PCE or TCE) (20). As a result, a class of students can be divided into individual groups to study the variable of their choice and determine the kobs for their group's set of experimental conditions. In a single 4-h laboratory, several variables may be studied, and students are able to compare their data on the effects of changing the catalyst, catalyst concentration, pH, and substrate (PCE or TCE) through comparison of the observed rate constants for the dechlorination reaction. Run-to-run precision for rate constants obtained under identical condition can be compared to published precision, if desired (20). Hazards Chlorinated ethylenes are suspected carcinogens and should be handled with appropriate care in the hood and should be disposed in suitable waste containers. Toluene is a flammable liquid, is harmful or fatal if swallowed, and causes irritation to skin, eyes, and respiratory tract. Tris(hydroxymethyl)aminomethane may 206


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Figure 3. Representative PCE dechlorination chromatograms for [(TCPP)Co] = 1.77 μM, [TiIII] = 26 mM, pH 8, 35 °C, [PCE]I = 240 μM. IS is toluene. The chromatograms are for headspace gas samples collected at: (a) 0 s, (b) 2400 s, (c) 4800 s, and (d) 7200 s.

cause irritation to skin, eyes, and respiratory tract. Methanol is extremely flammable and toxic by inhalation. Titanium(III) chloride is harmful if inhaled or swallowed and in contact with the skin. All chemicals should be handled with typical laboratory care to minimize exposure (lab coat, splash-proof goggles, and gloves). Because of the use of needles, students should take appropriate care to prevent accidental needle sticks. Results and Discussion Reductive dechlorination of PCE to TCE was monitored by analyzing headspace gas samples, and representative chromatograms taken over the course of the dechlorination reaction are provided in Figure 3. The PCE peak area decreases with a corresponding growth of the TCE peak area as the reaction progresses. The disappearance of PCE was fitted with first-order kinetic analysis for vitamin B12 and (TCPP)Co, where the experimental conditions were [CoI catalyst] = 1.8 μM, [TiIII] = 26 mM, pH 8, 35 °C, and [PCE]i = 240 μM in Figure 2. For the best fit lines, slope = -kobs for each reaction. As a result, the superior dechlorination activity of (TCPP)Co over vitamin B12 is evident and is consistent with literature reports (20). Reductive dechlorination observed rate constants for vitamin B12 and (TCPP)Co under a variety of reaction conditions are reported in Tables 1 and 2. Because kobs = k[CoI catalyst], the observed rate constant is expected to increase linearly with concentration, and this trend is observed for both vitamin B12 and (TCPP)Co. Also noteworthy is the increasing kobs as pH increases. Two factors are at play here. As pH increases, the reducing potential of titanium(III) citrate increases regenerating the active cobalt(I) catalyst more quickly (21). In addition, (TCPP)Co is sensitive to pH with the peripheral carboxyl moieties and through water/ hydroxyl groups coordinated to the cobalt center. The hydroxyl groups are likely to act as bridging ligands between titanium to cobalt providing a pathway for electron transfer (20, 22). The dechlorination kobs for PCE is greater than TCE for both catalysts, which is consistent with literature (20). This laboratory module has been used in the inorganic chemistry teaching laboratory because it integrates catalysis,


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In the Laboratory Table 1. Reductive Dechlorination Observed Rate Constants (kobs) for Chlorinated Ethylenes with Vitamin B12 Reaction Variable

kobs/(s-1 10-5)

[vitamin B12]/μM 0.74

1.21

1.11

2.50

Conclusions

1.47

3.01

1.84a

3.62a

7.38

10.00

The laboratory activities described here provide an integrative laboratory experience where students are required to use skills from several disciplines of chemistry for the determination of observed rate constants in the degradation of an environmental pollutant. Students gain valuable laboratory experience while working on a “real-world” problem that reflects ongoing work in the field.

Reaction pH 7.0

4.10

8.1

3.62

9.8

18.35

Substrate PCE

3.62

TCE

None observedb

a Standard conditions: [vit B12] = 1.84 μM, [PCE]i = 240 μM, [TiIIIcitrate] = 26 mM, pH 8, 35 °C. Rate constants are the average of two trials. b While TCE is dechlorinated by vit B12 under the conditions listed, no dechlorination was observed in 3 h of reaction time.

Table 2. Reductive Dechlorination Observed Rate Constants (kobs) for Chlorinated Ethylenes with (TCPP)Co Reaction Variable

kobs /(s-1 10-5)

[(TCPP)Co] (μM) 0.89

5.48

1.77a

14.50a

2.83

17.25

Reaction pH 7.0

8.88

8.1

14.50

9.8

29.85

Substrate PCE

14.50

TCE

5.22

a

Standard conditions: [(TCPP)Co] = 1.77 μM, [PCE]i = 240 μM, [TiIIIcitrate] = 26 mM, pH 8, 35 °C. Rate constants are the average of two trials.

transition-metal complexes, oxygen-sensitive compounds, firstorder kinetics, oxidation states, and environmental chemistry. For most students, working with oxygen-sensitive compounds presents new challenges and techniques as they remove oxygen from their reaction vials by sparging with N2 and providing an active N2 blanket while drawing the titanium(III) citrate from the stock solution container to initiate the reaction. As pH plays an important role in the experiment, the role and suitability of buffers across a pH range is discussed because the Tris buffer provides the greatest pH buffering near its pKa = 8.1. Because this laboratory module has a strong instrumentation component, students get hands-on experience with equipment common throughout the field. Variability of the sample injection volume is accounted for through the presence of the unreactive internal standard, toluene. Normalizing PCE peak area for injection volume provides a powerful correction factor


and improves the effectiveness of the data collection. As a result, students have the opportunity to analyze “real-world” kinetic data or apply the principles of first-order analysis. With multiple lab groups, an array of different experimental conditions (catalyst concentration, pH, and substrate) can be explored to find the optimal dechlorination conditions. Such a scenario reflects the increasingly collaborative nature of solving a difficult problem.

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Acknowledgment This work was funded in part through Tooma Undergraduate Research Fellowships and the generous support of Pepperdine University undergraduate research programs. We would like to thank Daphne Green and Paul Pestano for their laboratory support. Literature Cited 1. Doherty, R. E. Environ. Forensics 2000, 1 (2), 69–81. 2. Doherty, R. E. Environ. Forensics 2000, 1 (2), 83–93. 3. Squillace, P. J.; Scott, J. C.; Moran, M. J.; Nolan, B. T.; Kolpin, D. W. Environ. Sci. Technol. 2002, 36 (9), 1923–1930. 4. Westrick, J. J.; Mello, J. W.; Thomas, R. F. J. - Am. Water Works Assoc. 1984, 76 (5), 52–59. 5. Agency for Toxic Substances and Disease Registry (ATSDR). Toxicological Profile for Trichloroethylene; U.S. Department of Health and Human Services, Public Health Service: Atlanta, GA, 1997. 6. Agency for Toxic Substances and Disease Registry (ATSDR). Toxicological Profile for Tetrachloroethylene; U.S. Department of Health and Human Services, Public Health Service: Atlanta, GA, 1997. 7. National Institute for Occupational Safety and Health (NIOSH). Pocket Guide to Chemical Hazards: Trichloroethylene; Centers for Disease Control, National Institute for Occupational Safety and Health: Atlanta, GA, 2005. 8. NIOSH, Pocket Guide to Chemical Hazards: Tetrachloroethylene. In Centers for Disease Control and Prevention: Atlanta, GA, 2005. 9. Mohn, W. W.; Tiedje, J. M. Microbiol. Rev. 1992, 56 (3), 482–507. 10. Glod, G.; Angst, W.; Holliger, C.; Schwarzenbach, R. P. Environ. Sci. Technol. 1997, 31 (1), 253–260. 11. Glod, G.; Brodmann, U.; Angst, W.; Holliger, C.; Schwarzenbach, R. P. Environ. Sci. Technol. 1997, 31 (11), 3154–3160. 12. Schrauzer, G. N.; Deutsch, E.; Windgassen, R. J. J. Am. Chem. Soc. 1968, 90 (9), 2441–2442. 13. Schwarzenbach, R. P.; Angst, W.; Holliger, C.; Hug, S. J.; Klausen, J. Chimia 1997, 51 (12), 908–914. 14. McCauley, K. M.; Wilson, S. R.; van der Donk, W. A. J. Am. Chem. Soc. 2003, 125 (15), 4410–4411. 15. Rich, A. E. Mechanistic Studies on the Dehalogentaion of Chlorinated Ethylenes by Cobalamin Model Complexes. Master's Thesis, University of Minnesota, Minneapolis, MN, 2001.


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21. Zehnder, A. J. B.; Wuhrmann, K. Science 1976, 194 (4270), 1165–1166. 22. Pasternack, R. F.; Parr, G. R. Inorg. Chem. 1976, 15 (12), 3087– 3093.

Supporting Information Available Written materials for the students including prelab questions; spreadsheet for data analysis. notes for the instructor. This material is available via the Internet at http://pubs.acs.org.


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Degradation of Environmental Contaminants with Water-Soluble

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