Grantee Research Project Results

Final Report: Understanding Mechanisms of Green Oxidation Catalysis by Iron-TAML Peroxide Activators

EPA Grant Number: R832245
Title: Understanding Mechanisms of Green Oxidation Catalysis by Iron-TAML Peroxide Activators
Investigators: Collins, Terrence J. , Ryabov, Alexander D.
Institution: Carnegie Mellon University
EPA Project Officer: Hahn, Intaek
Project Period: September 1, 2004 through August 31, 2007
Project Amount: $242,300
RFA: Targeted Research Grant (2004) Recipients Lists
Research Category: Targeted Research

Objective:

The project was aimed at understanding the basic mechanistic behavior of the oxidation catalysis exhibited by iron(III)-Tetra-Amidato Macrocyclic Ligand (TAML) catalysts (1) that are important green chemistry catalysts. TAML^® catalysts utilize hydrogen peroxide as their most important source of oxidizing equivalents. The primary objective was to learn the underlying chemical reasons for their superior performance in a variety of environmentally important processes such as the bleaching of dyes, wood pulp, and colored effluents, the degradation of myriad persistent organic pollutants, and the decontamination of chemical and biological warfare agents. Iron-TAML catalysis reaction rates compare with what is typical for related enzymatic reactions and they are capable of at least 10³–10⁴ turnovers. We planned to explore the hypothesis that TAML activators function in a mechanistically parallel fashion to peroxidase and catalase enzymes. Once the chemical origin of the unique reactivity for synthetic peroxide activators had been identified, the information was used for further catalyst improvement. Understanding the first-order mechanistic features had already multiple positive impacts. Prior mechanistic studies had suggested ways for trying to advance the already unprecedented catalytic performance of Fe^III-TAMLs. One example involved learning how to replace hydrogen peroxide by dioxygen as the primary oxidizing agent. Another focused upon illuminating how to make superior activators. At the same time, the work was identifying novel tasks essential for further characterization the catalytic features of the activators. While this work was being pursuing this work, scholarship was also carried to characterize what is known about the health and environmental implications of pharmaceuticals in the environment. This led to an article referencing EPA support in the Green Chemistry special issue of Chemical Reviews entitled, Human Pharmaceuticals in the Aquatic Environment: A Challenge to Green Chemistry.^{[Publication 1]} And in the same period, a Scientific American article on TAML activators entitled, Little Green Molecules, was produced.^{[Publication 11]} Both these contributions were seen as important contributions to the broader understanding of major problems and developing solutions associated with the public health and environmental impacts of the chemical enterprise. In this context of our general studies on TAML activators and the ongoing scholarship involved with building the field of green chemistry, this specific project was undertaken along the following lines.

Specific Aims of the Project (as formulated)

New H₂O₂-activating catalysts could substantially green oxidation chemistry by providing economic alternatives to chlorine- and metal-based technologies. In the past, the rapid degradation of candidate catalysts by oxidation and/or hydrolysis has thwarted most development efforts. Iron-TAML activators are now the most significant exceptions. Fe^III-TAMLs activate, inter alia, hydrogen peroxide for numerous applications as described in the Background and Significance section of the proposal. We proposed to map the catalytic cycles by determining the mechanistic character of fundamental processes that are significant to the catalysis and by characterizing the reactive intermediate(s). The original specific goals of the proposed research were as follows.

First year:

• The speciation of the solid Fe^III-TAML complexes that represent the physical source of the catalysts in water will be determined by various spectroscopic techniques. The resistance of the catalysts to extreme conditions such as high/low pH will be measured to parameterize their stability over a range of possible operating conditions.
• A detailed mechanistic description will be obtained for each process that is a part of the catalytic cycle through the steady-state investigation of the kinetics of oxidation of model simple inorganic/organometallic 1e- donors such as ferrocene, ruthena(II)cycles, etc., and more complicated molecules such as phenols, organic dyes, etc. The processes that are currently thought to be critical on the basis of preliminary results are: peroxidase-like activity in which a substrate is oxidized by H₂O₂, catalase-like activity of Fe^III-TAMLs in which oxygen is produced from H₂O₂, catalase versus peroxidase activities of Fe^III-TAMLs, intramolecular oxidative degradation of Fe^III-TAMLs leading to catalyst inactivation, hydrolytic degradation of Fe^III-TAMLs leading to catalyst inactivation, the nature of the catalytically active reaction intermediate formed from Fe^III-TAMLs and H₂O₂—in our working hypothesis, it is an Fe-oxo complex.
• The mechanistic conclusions reached under the steady-state will be further investigated in single turnover experiments using a stopped-flow technique.
• Spectroscopic characterization (and possibly isolation and X-ray study) will be attempted for the purported reactive intermediate by electronic, EPR, Mössbauer, and possibly resonance Raman spectroscopy after its generation by different oxidants in water.

Second year:

• The knowledge of basic mechanisms of catalysis by Fe^III-TAMLs will be applied to the systematic studies of the catalyzed degradation of surrogates of chemical warfare agents where Fe^III-TAMLs display mixed oxidative and hydrolytic activity.
• Learn how to control general acid catalysis of the demetalation of Fe^III-TAML activators in neutral and acidic media.
• Explore micellar tuning of Fe^III-TAML activation of principal green oxidizing agents.
• Develop mechanistic understanding of the very high rate of activation of benzoyl peroxide by Fe^III-TAML activators.
• Quantify the reactivity of the transient oxidized form/s of Fe-TAML activators toward various electron donors (substrates) by steady state and transient kinetic measurements.
• Use hexacloroiridate(IV) as a probe for studying the proprieties of oxidized Fe-TAML activators.
• Move dioxygen chemistry of Fe^III-TAML activators into water: activation of O₂ in reverse micelles and photochemical catalytic cycles.
• Design and develop Fe^III-TAML oxidation processes mediated by electron carriers.
• Design and synthesize new “ideal” Fe^III-TAML catalysts suggested by the mechanistic studies.

Chart 1. Iron(III)-TAML activators and other key chemicals mentioned in this proposal.

Summary/Accomplishments (Outputs/Outcomes):

Objectives 1-1 and 1-3: Catalysts at work: High Oxidative Efficacy and Rapid Kinetics of Oxidation. The Catalase Activity: [To be submitted manuscript 1]
The exceptionally high peroxidase-like and catalase-like activities of iron(III)-TAML^® activators 1 have been established at pH 6-12.4 and 25-45 °C. While complexity precludes unambiguous data interpretation for the kinetics of Pinacyanol chloride oxidation, oxidation of “ruthenium dye”, the cyclometalated 2-phenylpyridine derivative 2 occurs via the stoichiometric equation

[Ru^II] + ½H₂O₂ + H⁺ [Ru^III] + H₂O (1)

The kinetics of reaction 1 using different TAML activators 1 has been investigated over a wide pH range by following a decrease in absorbance at 480 nm due to the oxidation of Ru^II to Ru^III. Enzyme-like kinetic traces without induction periods have had significant steady regions and displayed curvature at degrees of conversion around 85%, suggesting zero-order dependence on the concentration of 2. Absorbance changes in the absence of 1 were found to be negligible under these conditions. The steady-state rates have been measured at different concentrations of 1 (3.20×10^-8-4.80×10^-7 M), H₂O₂ (8.25×10^-5-1.11×10^-3 M), and 2 (2.58×10^-5-1.96×10^-4 M). As anticipated, the rates are directly proportional to the concentrations of 1 implying first order dependences on 1. First order dependence has been also observed in H₂O₂ concentration. The steady-state rate is independent of the 2 concentration. The zero-order in the Ru^II complex and the first order in H₂O₂ hold over the entire pH range (6-11.5). Thus, the experimental rate law given by eq 2 suggests that the rate-limiting step involves the interaction between 1 and H₂O₂ and followed by fast oxidation of the substrates.

rate = [1][H₂O₂] (2)

Free radical oxidation chemistry¹ is not involved in these Fe^III-TAML catalyzed oxidations to any detectable degree. The efficient hydroxyl radical scavenger, mannitol,^2,3 has no effect on the rate of reaction 1 catalyzed by 1 at pH 9.0. Similar results were obtained when urea was used. Based on these results, we conclude that this peroxide oxidation, catalyzed by 1, is not mediated by free radical processes.

Figure 1. pH profiles of the rate constants k2,obs for the oxidation of 2 by H2O2 catalyzed by 1a at different temperatures: [1c] 1.28×10-7 M, [2] 5.81×10-5 M, [H2O2] 3.3×10-4 M, 0.01 M phosphate.

Scheme 1. Proposed mechanism of catalyzed by 1 oxidation of “ruthenium dye” 2 by hydrogen peroxide.

The dependencies of on pH for 1c at 25-45 °C are shown in Figure 1. The rate constant increases dramatically at pH > 9 reaching a maximum around 10.5 and declines at pH > 11. The pH profiles in Figure 2 are rationalized quantitatively using the known pK_a’s of the catalysts⁴ and H₂O₂.⁵ Iron(III)-TAML activators are six-coordinated in water with two axial aqua ligands, the first pK_a1’s being in the range 9-10.⁴ The diaqua and aqua/hydroxo Fe^III complexes are likely reactive species that interact pairwise with either H₂O₂ or its conjugate base (pK_a2 ~ 11.2-11.6⁵) with the rate constants k₁-k₄ to give oxidized TAML, the nature of which is discussed below. Oxidized TAML reacts with the ruthenium(II) dye 2 very fast (Scheme 1) and therefore the reactions between iron(III)-TAML and hydrogen peroxide are treated as irreversible. Scheme 1 leads to the rate expression 3 with kinetically indistinguishable k₂ and k₃ pathways. The experimental data have been fitted to eq 3 neglecting either k₂ or k₃ pathways.

(3)

This allowed calculation of the rate constants k₂ and k₃, respectively, as well as k₁ and k₄. The best-fit values for the rate and equilibrium constants (Table 1) have been used for the emulation of the solid lines in Figure 1. The rate constants k₁ and k₄ are by a factor of ca. 100 and 10 lower compared to k₂, respectively, and this accounts for the sharp maxima in Figure 1. A trend in variation of the intrinsic rate constants k₁-k₄ is anticipated in view that the interaction between 1 and hydrogen peroxide is a redox reaction. The deprotonated species [FeL(OH)(H₂O)]^2- is more electron-rich than [FeL(H₂O)₂]- and therefore H₂O₂ oxidizes the former much faster than the latter (k₂>>k₁). In turn, deprotonated hydrogen peroxide HO₂- is more electron rich than H₂O₂ and therefore it reacts slower with [FeL(OH)(H₂O)]^2- (k₂>k₄) though a coulombic factor may also affect this step.

Table 1. Equilibrium and rate constants (in M^-1 s^-1), activation parameters for the Fe^III-TAML-catalyzed oxidation of 2 (eq 1) by hydrogen peroxide (0.01 M phosphate).

FeIII- TAML	T/Activation parameter	10-2×k1	10-4×k2	10-6×k3	10-3×k4	pKa1	pKa2
1a	25 °C	1.3±1.0	1.2±0.5	0.4±0.2	a)	9.7±1.0	10.8±2.0
1g	25 °C	4.0±1.8	1.8±0.7	0.8±0.1	1.5±1.0	9.5±0.4	10.9±0.4
	32 °C	8.5±11.6	3.3±0.2	1.3±0.1	4.0±2.4	9.3±0.1	10.4±0.1
	38 °C	9.0±8.8	3.5±0.2	1.2±0.1	3.8±1.7	9.6±0.7	10.4±0.1
	45 °C	12.5±0.7	3.6±0.1	1.8±0.1	5.1±2.3	9.2±0.1	10.9±0.1
	ΔH≠ kJ mol-1	46±12	28±13	30±10	49±19
	ΔS≠ J mol-1 K-1	-51±13	-77±25	-37±7	-32±11
a) could not be reliably estimated.

In the absence of electron donors, complexes 1 display a catalase-like activity (eq 4). Dioxygen evolution is visual at [H₂O₂] > 0.01 M. This catalytic feature has been studied kinetically by monitoring the initial rates of O₂ formation with a Clark electrode.

H₂O₂ H₂O + ½ O₂ (4)

The reaction stoichiometry is given in eq 4. Reaction 4 follows first-order kinetics in 1 ((0.6-9)×10^-6 M for 1a-δ) and H₂O₂ (2.3×10^-4-0.157 M) in the absence of an electron donor. Thus, rate expression 5 holds for the catalase-like activity, which is identical to eq 2 for the peroxidase-like activity.

rate = [Fe^III-TAML][H₂O₂] (5)

The rate constants are again pH-dependent and the pH profiles for reactions 2 (Figure 1) and 4, both catalyzed by 1g, are similar. Equation 3 has therefore also been applied for fitting the catalase data. The best-fit rate and equilibrium constants are summarized in Table 2. The rate constants in Tables 1 and 2 indicate that the "catalase-like" rate constants are somewhat lower. To relate the activation parameters, the temperature dependence of has been investigated at 25-45 °C and pH 10.1. At this pH, the k₂ pathway dominates—its contribution into the overall rate is more than 99%. Therefore, the effective enthalpy of activation obtained from the temperature dependence of should be associated with ΔH^≠ for the k₂ pathway. The corresponding ΔS^≠ should be calculated using k₂ from Table 2. Thus calculated ΔH^≠ and ΔS^≠ are shown in a footnote to Table 2. The similar rate laws (eqs 2 and 5), pH profiles, and the values of the observed second-order rate constants ( and ) suggest a common reaction intermediate in reactions 1 and 4.

Table 2. Rate (in M^-1 s^-1) and equilibrium constants, activation parameters for iron(III)-TAML catalyzed disproportionation of hydrogen peroxide (eq 4) at 25 °C and 0.1 M phosphate.

FeIII-TAML	10-2×k1	10-4×k2	10-3×k4	pKa1	pKa2
1a	1.4±1.2	0.51±0.03 a)	b)	11.0±0.2	12.1±0.4
1i	4.28±1.8	0.90±0.06	0.87±0.6	11.0±0.2	12.15±0.35
1g	1.3±0.7	0.64±0.04 c)	1.6±0.2	9.75±0.11	11.3±0.3
1c	6±3	1.68±0.08 δ)	3.2±0.2	10.5±0.2	11.6±0.2
a) ΔH≠ 23 ± 3 kJ mol-1; ΔS≠ -95 ± 20 J mol-1 K-1; b) could not be reliably estimated; c)ΔH≠ 14.6 ± 0.7 kJ mol-1; ΔS≠ -123 ± 13 J mol-1 K-1; d)ΔH≠ 8.1±3.3 kJ mol-1; ΔS≠ -137±63 J mol-1 K-1.

A plausible stoichiometric mechanism accounting both catalytic features of Fe^III-TAMLs is shown in Scheme 2. Taking all three steps into consideration the rate of the oxidative (peroxidatic) activity of 1 is given by eq 6.

(6)

Here ED is an appropriate electron donor. The catalase-like activity is described by eq 7.

(7)

The oxidation of ruthenium dye 2 is a zero-order reaction in 2. This implies that k_II[ED] >> {k_I + [H₂O₂](k_I+k_III)}. Equation 6 becomes very simple, i.e.

(8)

The first order dependence in hydrogen peroxide for the catalase-like activity holds when [H₂O₂](k_I+k_III) >> k_-I + k_II[ED]. Then, in the absence of added ED eq 7 becomes

(9)

Thus, both peroxidatic and catalase-like reactions have similar rate laws but the effective second order rate constant should be equal (when k_III>k_I) or lower (when otherwise) than provided the latter is zero-order in ED. The conditional rate constants k_I and k_III are pH dependent. The data in Figures 2 allow us to estimate the dependence of the rate constant k_III on pH using eqs 8 and 9. Thus the calculated plot of k_III versus pH in the range 8-11, where the accuracy in measuring and is the highest, is shown as a dashed line in Figure 2. The rate constant k_III is somewhat lower than k_I. This agrees with the fact that observation of product/s of the interaction between complexes 1 and H₂O₂ by spectral techniques is feasible. The shape of the pH profile for k_III in Figure 2 is explicable. Hydrogen peroxide becomes a better reducing agent upon deprotonation and therefore the rate constant k_III starts to rise at pH > 10. In fact, the reduction potential for the half-reaction O₂ + 2 H₂O + 2e → H₂O₂ + 2 OH- equals -0.146 V.⁶

Scheme 2. Mechanism of catalase and peroxidase activity of 1.

In conclusion, the catalytic activation of H₂O₂ by iron(III)-TAML activators is characterized by peroxidase-like and catalase-like activity. The rate determining step of the peroxidase-like activity is the interaction between 1 and H₂O₂. The step is strongly pH-dependent and the maximal activities are observed at pH around 10. The rate constants for the activation of H₂O₂ under the optimal conditions (>10⁴ M^-1 s^-1) approach a lower limit of the corresponding enzymatic peroxidase-catalyzed reactions. The activation of H₂O₂ results in the formation of the reactive intermediate formulated as the oxoiron(V) species which may react with substrate or conproportionate with more iron(III) to give the oxoiron(IV) intermediate (see next section of this report). The further oxidation of electron donors occurs in subsequent fast steps. In the absence of electron donors, the iron(III)-TAML activators display catalase-like activity with the production of dioxygen. The rate of this step is significantly slower than the oxidative bleaching of the dye Safranine O. In addition to underpinning the significant technological importance of the iron(III)-TAML activators,^7,8 these results signify that the activators are prominent functional models of the catalase-peroxidase enzymes. Iron(III)-TAML activators display both high catalase-like and peroxidase-like activities, but the latter dominates in the presence of many electron donor substrates other than H₂O₂ itself, thus eliminating unproductive consumption of hydrogen peroxide. This is a significant finding for industrially relevant oxidation processes.

In terms of publishing these extremely detailed studies, we have held back while we continue to carry out additional tests to continue to establish that this enticing mechanistic analysis does not suffer from any mechanistic oversight on our part. Thus, we have always been worried that the above analysis assumes that the Fe-TAML activators return always to a ferric resting state upon each revolution of the catalytic cycle. However, as described in the next section, the 1 activators in the presence of peroxide evolve to iron(IV)-containing solutions essentially instantly and in the presence of excess peroxide, continue to evolve oxygen by a catalase-like process. We have been concerned that the variable pH behavior could be explained, at least in part, by a partitioning of the total Fe-TAML in solution among various Fe(IV) species controlled by pH with each species having a differing rate of catalytic activity. In order to check on this possibility, it was essential to be able to identify and understand the catalytic role of any species that might be present in the steady state kinetics medium, and to this end, among other things, we conducted work at low temperature to characterize the speciation of the Fe-TAML under differing pHs in the presence of peroxide using Mössbauer spectroscopy. We found that there is a tightly pH dependent equilibrium between the iron(IV)-oxo monomer, [Fe(TAML)O]^2-, and bis-iron(IV)-oxo dimer, [{Fe(TAML)}₂O]^2-, with the former favored at higher pHs and the latter at lower pH’s. We are now studying the kinetics of the reactions of these species with 2. If the reactions of these iron(IV) species with 2 are fast wrt to the rate-determining process in our above scheme (a highly likely contingency which is implied by the results described above), then the conclusion that each revolution of the catalytic cycle returns the catalyst to the iron(III) resting state would be consistent and what we think is a final step toward publication of this very complex mechanistic story will have been made. In comparing the variation of the catalase-like behavior of the Fe-TAML catalysts with pH with the peroxidase-like behavior with 2, the same overall shape of the vs pH graph is observed, but the rates are slightly different and this may reflect that the Fe-TAML is not returning to the ferric state with peroxide. The manuscript for this work (available on request), will be held back until the one final test above is completed.

Objective 1-2: Catalysts at Work: The Nature of The Reactive Intermediate [Publication 2 and To Be Submitted Manuscript 2]
Iron(V)-oxo species have been proposed as key reactive intermediates in the catalysis of oxygen activating enzymes and synthetic catalysts. In Science this year, we reported the synthesis of [Fe(TAML)(O)]^- in nearly quantitative yield. Mass spectrometry, Mössbauer, electron paramagnetic resonance and X-ray absorption spectroscopies, reactivity studies, as well as density functional theory calculations, show that this long-lived (hours at -60°C) intermediate is an S = 1/2 iron(V)-oxo complex. Iron-TAML systems have proven to be efficient catalysts in the decomposition of numerous pollutants by hydrogen peroxide and the newly characterized species is a likely reactive intermediate in these reactions. This is the first case of a clearly characterized Fe(V)oxo complex, but thus far, we have only been able to isolate it in organic solvents where we can work at -60 to -80 °C.

The situation in water is quite different. We do not see evidence for the iron(V)-oxo, but we cannot work at low temperatures required for its isolation. Addition of H₂O₂ to aqueous solutions of iron(III)-TAML complexes 1 produces brownish-green colors. The spectral changes can be measured by uv-vis spectroscopy (Figure 2). Less than a stoichiometric amount of H₂O₂ causes a major increase in the absorbance in the range 350-550 nm. It requires several minutes to obtain invariable spectra at pH below 8.5, but the reaction becomes much faster at pH > 9. Two isosbestic points hold at 298 and 324 nm for 1e. At pH 11, the changes are fast and the final spectrum remains practically unchanged within 30 minutes. The inset in Figure 2 shows absorbance changes at different wavelengths as a function of added [H₂O₂], which allow for an estimate of the stoichiometry of the interaction. Matching data obtained at three wavelengths indicate that two iron(III) centers interact with one H₂O₂ molecule suggesting that the absorbing species are raised one oxidation equivalent above the resting state. Similar results were obtained with 1a.

Figure 2 (left). Spectral changes of 1e (2.5×10^-5 M) in the presence of 0, 2×10^-6, 4×10^-6, 6×10^-6, 8×10^-6, 12×10^-6, and 16×10^-6 M H₂O₂ at pH 11.4, 25 °C. Inset: Titration curves at 242 (), 280 (), and 420 () nm showing a 2:1 (Fe^III:H₂O₂) stoichiometry.

Figure 3 (right). Titration of the intermediate generated from 1e (X₁ = X₂ = Cl, R = Me) and H₂O₂ by [Fe(CN)₆]^4- showing a 1:1 stoichiometry at pH 11.35, 0.1 M KPF₆, 25 °C.

The same conclusion was reached using [Fe^II(CN)₆]^4- for the back reduction of the species derived from 1 and H₂O₂. The oxidized TAML was derived from 1e and a half equivalent of hydrogen peroxide (larger amount of H₂O₂ does not increase the absorbance). The data in Figure 3 shows that 1 equivalent of [Fe^II(CN)₆]^4- reduces the oxidized species. The optical density does not reach the initial level exactly due to absorbance of the oxidized titrant, i.e. [Fe^III(CN)₆]^3-, at 420 nm.

When an aqueous solution of tert-BuOOH is added to a solution of 1a at pH 14 (0.01 M phosphate) or 1 M KOH, a red color, stable at room temperature for hours, appears promptly. The same species can be obtained at any pH higher than 12. A titration of 1a by tert-BuOOH followed by the uv-vis spectroscopy revealed that ca. 0.5 equiv of the oxidant per Fe^III is needed to obtain the maximum change. Similar results have been obtained with H₂O₂ under identical conditions all suggesting the formation of an Fe^IV species.

The zero-field Mössbauer spectra of a frozen aqueous solution of the red compound (generated from 1×10^-3 M ⁵⁷Fe-1a and 0.5 eq t-BuOOH) at 4.2 K consists of one doublet with a quadrapole splitting, ΔE_Q = 4.2 mm s^-1 and isomer shift (versus Fe metal at 298 K) δ = -0.19 mm s^-1 (Figure 4). This highly negative isomer shift is indeed indicative of a high valent Fe^IV species and agrees with the titration studies. Furthermore, variable high field (1.5, 6.5, and 8 T) and high temperature (150 K) studies indicated the S = 1 species. Such intermediate spin Fe^IV species with a similar ligand and with the axial cyanides have been reported earlier from our group. Remarkably, the Mössbauer data accounts for ~100 and 80% of the total iron is this oxidized species at pH 14 and 12, respectively. Minor instability of the oxidized species has been detected at pH < 12.

Figure 4. The zero-field Mössbauer spectra of a frozen aqueous solution of the mixture of ⁵⁷Fe-1a (1×10^-3 M) and 0.5 eq t-BuOOH at 4.2 K. See text for details.

The uv-vis, Mössbauer, and EPR spectroscopic data obtained and reported previously⁹ suggest a rather detailed and complicated picture regarding the speciation of oxidized TAML species derived from 1 and various oxidants in aqueous solution (Scheme 3). If peroxides ROOH are primary two-electron oxidants, the oxidation of 1 gives presumably the oxoiron(V) intermediate 4. Such TAML-based species has never been detected in water but its existence in non-aqueous media has reliably been proved as part of this project and published earlier this year in Science.¹⁰ Similar to the sterically hindered anionic porphyrin chemistry,¹¹ the highly oxidized intermediate 4 is rapidly quenched by 1 into less oxidized iron(IV) species, the speciation of which is significantly pH dependent. At pH 14, all iron (ca. 95%) is monomeric oxoiron(IV) species 5', though theoretical DFT simulations (see below) do not exclude its aquated formulation as 5". The latter seems reasonable as the monomerization of the Fe^IV-O-Fe^IV dimer is induced by hydroxide ions.

The nature of the reactive intermediates as 5' and 5" ⁹ has been analyzed by DFT using Gaussian 03.¹² We have previously shown that the B3LYP functional and the 6-31G level of basis function predicts reliably electronic and structural properties of complexes 1.¹³ The same method was used for characterization of the structural and spectroscopic features of 5. The C_s symmetry is maintained in all cases to expedite the calculation time and full geometry optimization. The Mössbauer parameters in Table 3 are calculated using 6-311G with the established calibration curve. The parameters ΔE_Q and δ are known to be most accurately predicted by DFT. The results in Table 3 indicate that the formulation of 5 as a 5'/5" pair gives the best match with the experimental values of ΔE_Q and δ. An oxo-bridged dimer or Fe^III-radical cation (ligand-based oxidation)13 were ruled out at pH > 12 on the basis of either spin or Mössbauer parameter considerations.

Scheme 3. Speciation of oxidized TAML species derived from 1 in aqueous solution. Axial aqua ligands are not shown for clarity.

Our calculations also show that the S = 1 spin state is energetically more favorable than both high or low spin states. Time dependent (TD-DFT) analysis has been also applied to the geometry optimized structure of 5' (Figure 5) to ensure they indeed represent structures of the lowest energy. It is interesting to note that the theory predicts 5- and 6-coordinated iron(IV) in 5' and 5", respectively. The central metal in 5' does not hold axial water. This is obviously due to a strong trans effect of the axial oxo ligand of charge -2. It should however be added that the six coordinated iron with weakly bound water cannot be excluded.

Figure 5. DFT-predicted lowest energy structures of oxidized species 5'.

Table 3. Calculated by DFT anticipated parameters of the Mössbauer spectra of products generated from 1 and tert-BuOOH in basic (pH > 12) aqueous solution.

Model			S = 1	S = 2		Experimentally observed
ΔEQ / mm s-1	4.4	3.6	3.5	2.5	1.8	3.9
δ / mm s-1	-0.07	-0.15	-0.12	0.16	-0.07	-0.19

In contrast with organic solvents, the aqueous environment puts severe limitations on our ability to observe the Fe(V) state. All oxidizing agents tested so far increase the oxidation state of iron in 1 by one equivalent to produce iron(IV). Iron(IV) is generated by dioxygen in non-polar organic solvents,⁹ probably via an iron(V) species and indeed this very species is produced from 1 and organic peroxides in tert-butyronitrile.¹⁰ The speciation of these iron(IV) TAML derivatives is crucially pH dependent. The monomeric oxoiron(IV) complexes 5 dominate at pH > 12, whereas at pH 8-10 the major species is the μ-oxo dimer. Though it is just iron(IV), its reduction potential is similar to that of the [IrCl₆]^2-/3- couple, i.e. peroxides generate very powerful oxidizing agents from 1 in water (see below). Interestingly, the reduction potential is comparable to those of both to hexaaqua Fe^{II/III 6} and Compound II/resting state of horseradish peroxidase^14,15 couples though the oxidative chemistry of 1 is predominantly non-radical, i.e. biological like.

Objective 1-4: New Approaches For Evaluating The Catalyst Lifetimes Intermediate [Publication 12]
Kinetic measurements of the Fe^III-TAML-catalyzed oxidation of conventional dyes such as Pinacyanol chloride, Safranine-O and Orange II (Chart 1) by H₂O₂ have revealed that the oxidation does not follow monoexponential kinetics under pseudo-first-order conditions, i.e. under at least ten-fold excess of H₂O₂ over all dyes used. The data shown in Figure 6 illustrate the non-exponential kinetics. The 1a-catalyzed bleaching of Safranine-O was initiated by addition of an aliquot of H₂O₂. At low catalysts concentrations the bleaching does not go to completion even at [H₂O₂] >> [dye]. Addition of the second aliquot of H₂O₂ at time a does not restart catalytic bleaching (Figure 1) suggesting the catalyst has been deactivated. Accordingly, addition of a new aliquot of 1a resumes the reaction. The bleaching of Safranine-O can be completed by successive additions of the 1a (Figure 6 shows incomplete bleaching). The robustness of Safranine-O toward oxidation by Fe^III-TAMLs and H₂O₂ makes it convenient for studying Fe^III-TAML catalyst half-lives.

Figure 6 (left). Kinetics of 1a-catalyzed bleaching of Safranine-O (4.3×10^-5 M) by 0.012 M H₂O₂. Initial concentration of 1a (7.5×10^-8 M); aliquots of the same amount of 1a were added after complete inactivation of the catalyst giving rise to the stepped dependence. The dashed line shows that addition of 0.012 M H₂O₂ does not resume the catalytic bleaching. Inset shows linearization of the data obtained after each addition of 1a in terms of eq 13. Conditions: pH 11, 25 °C.

Figure 7 (right). Dependence of the amount of bleached dye expressed as A_o on the total amount of dye expressed as Ao for 1a-catalyzed (7.5×10^-8 M) bleaching of Safranine-O (initial concentration 4.3×10^-5 M) by H₂O₂ (0.012 M) in terms of eq 15, see text for details. Conditions: 25 °C, pH 11.0.

A methodology for quantifying the catalytic performance is based on Scheme 4. The intermolecular inactivation is assumed to be kinetically insignificant at this point, i.e. k_2i ≈ 0, because the majority of the dye bleaching experiments were run at very low catalyst concentrations (10^-6-10^-8 M). Therefore, second-order pathways should not play a substantial role. The medium-induced hydrolytic degradation is known to be kinetically insignificant. Complexes 1 are subjected to the H⁺-catalyzed demetalation for which the pseudo-first-order rate constant is given by k_obs = k_α[H⁺] + k_γ[H⁺]³.⁴ The demetalation becomes significant with decreasing pH at pH ~ 4 for 1a, a methyl-tailed TAML catalyst, but only at significantly lower pH for 1c, a fluorine-tailed TAML catalyst. The bleaching experiments reported in our paper were performed in the pH range 10-12 and therefore k_δ is negligible. There is an additional general-acid catalyzed demetalation of 1a at pH 5-7 which is also unimportant because its time scale is much longer than that of the bleaching. No general-acid-catalyzed demetalation is observed for the fluorine-tailed 1c.

Scheme 4. General reaction scheme showing the major catalytic steps and the steps leading to inactivation of the resting state and the active form of a catalyst involved in a two-substrate process.

The stoichiometric mechanism of catalysis in Scheme 4 is based on eq 6 when k_III is negligible. The rate constants k_I, k_-I, and k_II are effective, pH-dependent parameters. A straightforward evaluation of the rate constants k_i and k_II is feasible when k_-I is negligible and the interaction between 1 and H₂O₂ is significantly faster than the oxidation of substrate itself, i.e., when the relation k_I[H₂O₂] > k_II[Substrate] holds. This condition does not usually hold for low-molecular-weight catalysts of hydrogen peroxide such as iron porphyrins.16 The interaction of synthetic catalysts with H₂O₂ is usually the slowest step in the oxidation process. In general, this is also true for 1, although the second-order rate constants k_I are as high as 10⁴ M^-1 s^-1 at pH around 10 and the relation k_I[H₂O₂] > k_II[Substrate] can be secured by using difficult-to-oxidize dyes with low values of k_II. This opens a broader pH range for characterizing the operational stability of the Fe^III-TAML catalysts. Safranine-O belongs to the class of difficult-to-oxidize dyes by Fe^III-TAMLs. The relation k_I[H₂O₂] > k_II[Substrate] holds and the initial rate of its bleaching is directly proportional to [Safranine-O] and independent of [H₂O₂] (data not shown). Thus, eq 6 simplifies into eq 10.

k_II[Substrate][Fe^III] (10)

According to Scheme 4, the intramolecular inactivation of active forms of 1 is monoexponential. The corresponding rate constant is k_i. Therefore the related differential equation for dye bleaching is given by eq 11.

= k_II(D_t - x)([Fe^III]_t) (11)

Here D_t and [Fe^III]_t are total concentrations of the dye and an Fe^III-TAML catalyst, respectively, and x is the concentration of a bleached dye at time t. Integration of eq 11 under the boundary conditions x = x_∞ (x_∞ is the concentration of bleached dye obtained with single catalyst aliquot) when t = ∞ affords

(12)

or

(13)

Equation 13 involves easy to measure absorbances A_t and A_∞ at time t and ∞, respectively, instead of concentrations as in eq 12. Equation 13 suggests a simple procedure for analysis of data such as in Figure 6. If the double logarithm of the ratio A_t/A_∞ is plotted against time, the slope of the straight line will equal -k_i. The rate constant k_II can then be calculated from the intercept which equals ln(k_II[Fe^III]_t/k_i). A linearization of the kinetic traces of Figure 6 in terms of eq 13 is demonstrated in the Inset. All four steps in Figure 6 have been analyzed. The values of k_i and k_II are nearly identical for each trace in Figure 6. The value of k_i indicates that the half-life of 1a is 3.4 min under the operating conditions (25 °C, pH 11). The rate constant k_II equals approximately 10⁴ M^-1 s^-1.

Careful inspection of the data in Figure 6 indicates that the total amount of the bleached dye (x_∞) after adding a new aliquot of 1a is not constant and gradually decreases after each new addition. This phenomenon is understood in terms of the general kinetic model proposed. The solution of differential equation 11 using the boundary condition x = 0 at t = 0 leads to eq 14.

(14)

Here A₀ and A_t are optical densities at time t = 0 and t, respectively. Since A_t = A₀ – x, x → x_∞ when t → ∞. Substitution and rearrangement result in eq 15.

(15)

Eq 15 implies that x_∞ should be directly proportional to the starting concentration of dye or to the absorbance A₀. The straight line passing through the origin in Figure 7 confirms the proposed model. Moreover, the predicted slope of this line is 0.25±0.04 based on the known values of the rate constants k_II and k_i and the concentration of 1a (7.5×10^-8 M). The slope determined from the experimental data is 0.24±0.02 which agrees well with the calculated value.

Figure 8. A. Kinetics of oxidation of Safranine-O (4.3×10^-5 M) by H₂O₂ (0.012 M) by catalysts 1f, 1a, 1b, and 1c (7.8×10^-5 M). B. Linearization of the data in terms of eq 13. Conditions: 25 °C, pH 11.0.

Catalysts 1a-f display variable performance (Figure 8). The bleaching by 1c is both the deepest and the fastest. Equation 13 was applied to all kinetic traces in Figure 8A affording satisfactorily straight lines shown in Figure 8B. The rate constants k_i and k_II summarized in Table 4 were calculated from the slopes and intercepts, as above. Similar experiments were performed at different temperatures covering the 13-60 °C range. The inactivation (k_i) and bleaching (k_II) rate constants in Table 4 have been used for calculating the corresponding enthalpies (ΔH^≠) and entropies (ΔS^≠) of activation, using the Eyring equation.¹⁷

Table 4. Rate constants (k_i and k_II) calculated for different Fe^III-TAML catalysts at pH 11 using eq 13 from the data such as in Figure 8.

Catalyst	T / °C	103×ki / s-1	10-4×kII / M-1 s-1
1f	13	0.44±0.05	0.24±0.03
	25	1.5±0.1	0.7±0.1
	42	5.7±0.6	1.2±0.1
	60	33±3	3.6±0.4
1a	13	1.9±0.2	0.9±0.1
	25	3.4±0.3	1.1±0.1
	42	5.2±0.5	1.5±0.2
	60	12±1	3.1±0.3
1b	13	2.5±0.3	0.99±0.1
	25	6.2±0.4	1.9±0.2
	42	11±1	3.6±0.7
	60	18±2	5.6±0.6
1c	13	4.8±0.5	6.3±0.6
	25	13±1	10±2
	42	20±2	14±1
	60	45±2	23±3

The rate constants in Table 4 reveal interesting trends. (i) The catalytic activity of 1 in bleaching Safranine-O (k_II) increases more than ten-fold at 25 °C when the tail ethyl groups of 1f are replaced by fluorine atoms in 1c. The fluorine-tailed compound 1c is the most active. The same level of reactivity is typical of horseradish peroxidase Compound II toward anilines and phenols.¹⁸ This emphasizes that Fe^III-TAMLs are really efficient synthetic low-molecular weight peroxidase mimics. (ii) The inactivation rate constants k_i also increase on going from 1f to 1c and a similar 10-fold gap holds. Previously, we have reported qualitative data showing a better performance of the methyl-tailed catalyst 1a versus its ethyl-counterpart 1f in bleaching of Pinacyanol chloride.¹⁹ Despite the fact that different inactivation processes are operating here, the quantitative kinetic data reported here also conform to the previous qualitative observations concerning comparative catalyst performances.

Our mechanistic studies of the Fe^III-TAML catalysts convincingly show that the introduction of fluorine atoms onto the tail changes the reactivity in a dramatic way, allowing activation of H₂O₂ in acidic media (in addition to the basic range for these catalysts) while also producing significantly more aggressive oxidizing systems. The modification of the "head", in contrast, induces considerably more subtle changes in Fe^III-TAML catalyst reactivity through aromatic ring mediated electronic effects, the structural features of the catalysts being otherwise the same.⁴ We also examined how electron-donating or electron-withdrawing substituents at the aromatic ring affect the rate constants k_i and k_II. Eight structurally similar, but electronically different Fe^III-TAML catalysts 1, all with R = Me, have been investigated by applying eq 13. Practically useful, but rather curious results are found in the linear free energy relationship (LFER). There is a linear dependence between lnk_II and lnk_i with a slope of –1. Electron-withdrawing substituents (NO₂, NMe₃⁺, COOEt, Cl) increase the oxidizing power of the Fe^III-TAML catalysts with respect to Safranine-O (k_II), but retard the intramolecular inactivation (k_i). The most inactivation resistant catalyst is the nitro-substituted Fe^III-TAML, 1d. It is more stable than 1a by a factor of six. This observation suggests that one of the vulnerable fragments of Fe^III-TAML catalysts is the aromatic component. Presumably the active catalyst may destroy itself beginning via oxidative damage at this group. Electron-withdrawing substitutients should protect the ligand system from this type of damage. The slope of -1.0±0.3 is unusual because the slope of +1 might be anticipated with more reactive catalysts decomposing more rapidly.

Equation 13 can be rearranged to a form (eq 16) that describes the relative conversion of substrate, i.e. A_t/A₀, as a function of time. Equation 16 provides convenient means for simulating the catalyst performance using the current model.

(16)

Thus, theoretical predictions can be compared with the experimental data. Additional support for the mechanistic concepts developed here would consist of a good match between the experimental and predicted kinetic data calculated using the obtained values of k_i and k_II at known concentrations of Fe^III-TAML catalysts. Provided the corresponding concentration regime is chosen, the calculated and measured absorbance versus time plots should be similar. Such a comparison is shown in Figure 9, were the dynamics of 1a-catalyzed bleaching of Safranine-O by hydrogen peroxide performed at different 1a concentrations is plotted together with the calculated traces. There is excellent agreement between two experimental and calculated curves. The comparison proves that quantitative bleaching of the dye is in fact achieved just by increasing the 1a concentration. But even after the increase, the concentration is very low, specifically, 10^-6 M when the difficult-to-oxidize Safranine-O is bleached. The approach described and justified herein is principally applicable to any two-substrate catalytic system provided the activation of the catalyst is faster than a target chemical reaction.

Figure 9. Normalized experimental and simulated bleaching of Safranine-O (4.3×10^-5 M) by H₂O₂ (0.012 M) catalyzed by 1a at pH 11 and 25 °C. Experimental data are shown as and . The simulations, shown as solid lines, were made using the rate constants from Table 4.

In conclusion, these kinetic studies show that green catalytic oxidation technologies can be designed that are fully capable of supporting sustainable and economic chemical products and processes. Fe-TAML catalysts perform comparably to, if not better than, peroxidase enzymes, but they are much easier to make, study and control. Fe^III-TAML catalysts begin to meet the great promise of green chemistry (the design of chemical products and processes that reduce or eliminate the use and generation of hazardous substances) in ways that matter to environmental integrity.^8,20,21 The results of this work makes them even more attractive for use in technologies that send effluent streams to environmental media, because there is now a clear scenario of how the catalytic activity of the catalysts affects their half-lives so that process conditions can be set to ensure that no undegraded Fe^III-TAML catalyst remains at the point of exit of the effluent stream from the plant. A substantial gain in reactivity is brought about by fluorination of the tail component of Fe^III-TAMLs, but this is accompanied by a drop in the half-lives of these more reactive catalysts. Overall, the important process comparison is that the fluorinated catalysts work more rapidly. Suicide inactivation is most efficiently retarded by introducing electron-withdrawing groups at the aromatic ring in the head part of Fe^III-TAMLs. Thus, knowing the rate constants k_i and k_II provides a shortcut for developing green oxidation technologies where one can ensure the catalysts will die in the plant when the intended work is done.

Objective 1-5: Catalysts at Work: Mixed Oxidative and Hydrolytic Activity [Publication 3]
We have studied degradation of three OP pesticides, namely, fenitrothion, parathion and chlorpyriphos methyl, by 1-H₂O₂. Fe-TAML (1a, 40 μM or 1c, 10 μM)-H₂O₂ (2.0 M) treatment of fenitrothion (1 mM) in water with 10% ^tBuOH as co-solvent in buffered pH (pH 8.0, 0.1 M KH₂PO₄/KOH) led to the immediate formation of a yellow color, λ_max=396 nm in UV-Vis spectroscopy, characteristic of 3-methyl-4-nitrophenolate. This absorption band was seen diminishing soon after its formation, eventually nearly disappearing (>95% degradation, within 2 h. Independent treatment of 3-methyl-4-nitrophenolate with 1-H₂O₂ under the same conditions led to similar UV-Vis observations. The reaction kinetics followed by UV-Vis spectroscopy revealed that the rate of formation of 3-methyl-4-nitrophenolate from fenitrothion is considerably faster than its subsequent oxidative degradation. HPLC analysis confirmed the generation of 3-methyl-4-nitrophenolate as the major initial product (~90%) and its concomitant degradation. The formation of fenitrooxon as a minor product (~10%) was also observed. However, the presence of even small amounts of fenitrooxon was unacceptable due to its high mammalian toxicity (Oral LD₅₀ rat 3.3 mg/kg). Moreover, while fenitrothion and 3-methyl-4-nitrophenolate were easily destroyed by 1-H₂O₂ under the reaction conditions, it had no observable effect on fenitrooxon even after 2 h.

Phosphate esters undergo more facile base hydrolysis than their thiophosphate analogs. The strongly nucleophilic hydropero

Journal Articles on this Report : 14 Displayed | Download in RIS Format

Publications Views

Other project views:	All 21 publications	14 publications in selected types	All 14 journal articles

Publications

Type	Citation	Project	Document Sources
Journal Article	Banerjee D, Markley AL, Yano T, Ghosh A, Berget PB, Minkley Jr EG, Khetan SK, Collins TJ. "Green" oxidation catalysis for rapid deactivation of bacterial spores. Angewandte Chemie International Edition 2006;45(24):3974-3977.	R832245 (Final)	Abstract from PubMed Full-text: Carnegie Mellon University PDF Exit Abstract: Wiley Exit
Journal Article	Chahbane N, Popescu D-L, Mitchell DA, Chanda A, Lenoir D, Ryabov AD, Schramm K-W, Collins TJ. FeIII–TAML-catalyzed green oxidative degradation of the azo dye Orange II by H2O2 and organic peroxides: products, toxicity, kinetics, and mechanisms. Green Chemistry 2007;9(1):49-57.	R832245 (Final)	Abstract: RSC Abstract Exit
Journal Article	Chanda A, Popescu D-L, Tiago de Oliveira F, Bominaar EL, Ryabov AD, Munck E, Collins TJ. High-valent iron complexes with tetraamido macrocyclic ligands: structures, Mössbauer spectroscopy, and DFT calculations. Journal of Inorganic Biochemistry 2006;100(4):606-619.	R832245 (Final)	Abstract from PubMed Abstract: Science Direct Exit
Journal Article	Chanda A, Ryabov AD, Mondal S, Alexandrova L, Ghosh A, Hangun-Balkir Y, Horwitz CP, Collins TJ. Activity-stability parameterization of homogeneous green oxidation catalysts. Chemistry-A European Journal 2006;12(36):9336-9345.	R832245 (Final)	Abstract from PubMed Abstract: Wiley Exit
Journal Article	Chanda A, Khetan SK, Banerjee D, Ghosh A, Collins TJ. Total degradation of fenitrothion and other organophosphorus pesticides by catalytic oxidation employing Fe-TAML peroxide activators. Journal of the American Chemical Society 2006;128(37):12058-12059.	R832245 (Final)	Abstract from PubMed Abstract: ACS Exit
Journal Article	Chanda A, Shan X, Chakrabarti M, Ellis WC, Popescu DL, Tiago de Oliveira F, Wang D, Que Jr. L, Collins TJ, Munck E, Bominaar EL. (TAML)FeIV=O complex in aqueous solution: synthesis and spectroscopic and computational characterization. Inorganic Chemistry 2008;47(9):3669-3678.	R832245 (Final)	Abstract from PubMed Full-text: ACS Publications Full Text Exit Abstract: ACS Exit Other: ACS Publications PDF Exit
Journal Article	Collins TJ, Walter C. Little green molecules. Scientific American 2006;294(3):82-90.	R832245 (Final)	Abstract from PubMed Full-text: Glatfelter PDF Exit Abstract: Scientific American Exit
Journal Article	Ghosh A, Mitchell DA, Chanda A, Ryabov AD, Popescu DL, Upham EC, Collins GJ, Collins TJ. Catalase-peroxidase activity of iron(III)-TAML activators of hydrogen peroxide. Journal of the American Chemical Society 2008;130(45):15116-15126.	R832245 (Final)	Abstract from PubMed Abstract: ACS-Abstract Exit
Journal Article	Khetan SK, Collins TJ. Human pharmaceuticals in the aquatic environment:a challenge to Green Chemistry. Chemical Reviews 2007;107(6):2319-2364.	R832245 (Final)	Abstract from PubMed Full-text: ACS Publications Full Text Exit Abstract: ACS Exit Other: ACS Publications PDF Exit
Journal Article	Polshin V, Popescu D-L, Fischer A, Chanda A, Horner DC, Beach ES, Henry J, Qian Y-L, Horwitz CP, Lente G, Fabian I, Munck E, Bominaar EL, Ryabov AD, Collins TJ. Attaining control by design over the hydrolytic stability of Fe-TAML oxidation catalysts. Journal of the American Chemical Society 2008;130(13):4497-4506.	R832245 (Final)	Abstract from PubMed Full-text: ACS Publications Full Text Exit Abstract: ACS Exit Other: ACS Publications PDF Exit
Journal Article	Popescu DL, Vrabel M, Brausam A, Madsen P, Lente G, Fabian I, Ryabov AD, van Eldik R, Collins TJ. Thermodynamic, electrochemical, high-pressure kinetic, and mechanistic studies of the formation of oxo FeIV-TAML species in water. Inorganic Chemistry 2010;49(24):11439-11448.	R832245 (Final)	Abstract from PubMed Abstract: ACS-Abstract Exit
Journal Article	Popescu D-L, Chanda A, Stadler M, Tiago de Oliveira F, Ryabov AD, Munck E, Bominaar EL, Collins TJ. High-valent first-row transition-metal complexes of tetraamido (4N) and diamidodialkoxido or diamidophenolato (2N/2O) ligands: synthesis, structure, and magnetochemistry. Coordination Chemistry Reviews 2008;252(18-20):2050-2071.	R832245 (Final)	Abstract: ScienceDirect-Abstract Exit
Journal Article	Shappell NW, Vrabel MA, Madsen PJ, Harrington G, Billey LO, Hakk H, Larsen GL, Beach ES, Horwitz CP, Ro K, Hunt PG, Collins TJ. Destruction of estrogens using Fe-TAML/peroxide catalysis. Environmental Science & Technology 2008;42(4):1296-1300.	R832245 (Final)	Abstract from PubMed Full-text: USDA PDF Abstract: ES&T Exit
Journal Article	Tiago de Oliveira F, Chanda A, Banerjee D, Shan X, Mondal S, Que Jr. L, Bominaar EL, Munck E, Collins TJ. Chemical and spectroscopic evidence for an FeV-oxo complex. Science 2007;315(5813):835-838.	R832245 (Final)	Abstract from PubMed Abstract: Science Exit

Supplemental Keywords:

Media: water, drinking water. Risk Assessment: risk, human health, bioavailability. Chemicals, Toxics, Toxic Substances: chemicals, toxics, metals, heavy metals, solvents, oxidants, organics, effluent, discharge, intermediates. Ecosystem Protection: aquatic. Risk Management: alternatives, sustainable development, clean technologies, innovative technology, remediation, cleanup, oxidation. Public Policy: public good. Scientific Disciplines: environmental chemistry,, Scientific Discipline, Water, Engineering, Chemistry, & Physics, Analytical Chemistry, homeland security, peroxide activators, catalyst formulations, biological warfare agents, decontamination, catalytic oxidation, environmental chemistry, green chemistry, persistant organic pollutants

Relevant Websites:

http://www.chem.cmu.edu/groups/Collins/ Exit

Progress and Final Reports:

Original Abstract

2005

2006