Article pubs.acs.org/Organometallics
Oxidative Addition of Haloalkanes to Metal Centers: A Mechanistic Investigation Veeranna Yempally,† Salvador Moncho,† Sohail Muhammad,† Edward N. Brothers,† Bruce A. Arndtsen,*,‡ and Ashfaq A. Bengali*,† †
Department of Chemistry, Texas A&M University at Qatar, Doha, Qatar Department of Chemistry, McGill University, 801 Sherbrooke Street West, Montreal, Canada H3A 0B8
‡
S Supporting Information *
ABSTRACT: Photolysis of CpRe(CO)3 in the presence of dichloromethane results in the initial formation of the CpRe(CO)2(ClCH2Cl) complex followed by insertion of the metal into the C−Cl bond. The activation enthalpy is determined to be 20.4 kcal/mol, and with the assistance of DFT calculations, a radical mechanism is proposed for the oxidative addition reaction. Photolysis of Ni(CO)2(PPh3)2 with dihalomethanes also results in oxidative addition, but the intermediacy of a halogen-bound adduct has not been established.
P
oxidative addition motivated us to study the photochemistry of the tricarbonyl in the presence of dichloromethane to investigate the mechanism of possible metal insertion into the C−Cl bond. The analogous transformation was also probed with more electron rich nickel(0) complexes. The latter have found extensive use as catalysts in carbon−carbon bond forming reactions with organic halides. As described below, both of these systems undergo metal insertion into the C−Cl bond, although the results are more tentative with the nickel complex. In the case of the rhenium complexes, time-resolved IR data are consistent with initial formation of the Re− ClCH2Cl solvate followed by insertion of Re into a C−Cl bond to yield the cis-Re(Cl)(CH2Cl) complex (Scheme 1). The
olychlorinated solvents of industrial importance such as dichloromethane and chloroform are among the major volatile organic contaminants found in groundwater. As potential free chlorine release agents, their use can have an adverse effect on the environment.1 Efficient removal of these compounds requires a thorough understanding of their decomposition pathways and their reactivity toward metal catalysts which are used in catalytic hydrodechlorination2 and photodegradation3 of polychlorinated solvents. Investigating the interaction of organic halides with metal centers is also of considerable interest, since oxidative addition of the C−X bond is a key step in carbon−carbon bond formation reactions promoted by transition metals. This transformation has been extensively employed as a major synthetic tool in crosscoupling reactions.4 Despite the importance of this chemistry, many questions remain regarding the energetics of alkyl halide bond activation by transition metals and the nature of the intermediates involved. Such fundamental information has the potential to aid in the design of efficient catalysts for the photodegradation of chloroalkanes and cross-coupling reactions. Several rhenium complexes have been shown to thermally activate the C−Cl bond in dichloromethane.5 While photochemical C−X activation by photolysis of Cp*Re(CO)2(L) (L= N2, CO) in the presence of aryl halides has been well studied,6 the dynamics of the oxidative addition reaction with identification of the relevant intermediates has not been reported. It is well established that solution-phase photolysis of CpRe(CO)3 results in CO loss and the formation of the solvent complex CpRe(CO)2(solvent).7 This fact coupled with the observation that rhenium complexes often promote C−X © XXXX American Chemical Society
Scheme 1. Rhenium Insertion into the C−Cl Bond of Dichloromethane
experimental results supported by DFT calculations are consistent with a radical mechanism for C−Cl oxidative addition to the rhenium center. Unlike the Re systems, evidence for the prior formation of a Ni−XCH2X adduct was not obtained. Received: May 15, 2014
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RESULTS AND DISCUSSION (a). CpRe(CO)3 + CH2Cl2. Photolysis of a heptane solution of CpRe(CO)3 with 5 mM CH2Cl2 at 213 K results in the initial formation of the previously observed CpRe(CO)2(heptane) complex with CO stretching bands at 1946 and 1880 cm−1 (Table S1, Supporting Information). This solvate complex exhibits a single-exponential decay, and a dicarbonyl complex with CO bands at 1945 and 1874 cm−1 grows in at the same rate (Figure S1, Supporting Information). At 213 K this species is stable and demonstrates no further reactivity. Since CpMn(CO)2(ClCH2Cl) and other CpMn(CO)2(XR) (RX = Cl(CH2)2Cl, Cl(CH2)3Cl, Cl(CH2)6Cl) complexes have CO band positions similar to those observed here,8 this dicarbonyl species is identified as the chloro-bound complex CpRe(CO)2(ClCH2Cl) (1). DFT calculations are consistent with this assignment and predict a 24.3 kcal/mol Re−ClCH2Cl bond dissociation enthalpy (BDE). Two isomers of the chloro adduct were found with the uncoordinated Cl pointing toward (1-up) and away (1-down) from the Cp ring. The two rotamers have a low interconversion barrier of ΔG⧧298 K = 3.4 kcal/mol and similar calculated stabilities with the 1-up isomer favored by only 0.9 kcal/mol (Figure 1). The
evident in the case of the Cp* system discussed below. The experimental finding is therefore consistent with DFT modeling and points toward the potential presence of two isomers of the chloroadduct. Similar rotamers were also predicted in η2 coordination of benzene to a rhenium center by DFT calculations.9 Given the small calculated free energy difference between the two structures, the discussion below assumes that 1-up is the starting point for Re insertion into the C−Cl bond. While 1-up is stable at 213 K, it has a lifetime of less than 30 s at room temperature, converting rapidly to the dicarbonyl product 2, with CO bands at 2043 and 1967 cm−1 (Figure 2). The almost 100 cm−1 upfield shift in the CO band positions of 2 in comparison to those of 1 is indicative of a change in the oxidation state of the Re metal and is characteristic of a Re(III) center. For example, the CO bands in CpRe(CO)2(H)2 and CpRe(CO)2(H)(SnCl3) are observed at 2022, 1954 cm−1 and 2040, 1981 cm−1, respectively.10a,b Therefore, 2 is assigned to the CpRe(CO)2(Cl)(CH2Cl) complex.10c The formation of 2 is consistent with DFT modeling which predicts that oxidative addition of the C−Cl bond in dichloromethane to the Re center is favored over 1 with ΔG°298 K = −18.1 kcal/mol. Moreover, the similar relative intensities of the CO absorbances in this complex are consistent with a cis rather than trans geometry for 2. Calculations predict a relatively large barrier of ΔG⧧298 K = 32.6 kcal/mol for the cis → trans isomerization. with the trans isomer favored by ΔG°298 K = −2.6 kcal/mol. The oxidative addition reaction was studied at several temperatures, and an Eyring analysis yielded activation parameters of ΔH⧧ = 20.4 ± 0.9 kcal/mol and ΔS⧧ = 11 ± 3 eu (Figure S2, Supporting Information). Since the calculated Re−ClCH2Cl BDE is similar to the experimental activation enthalpy, it is possible that CH2Cl2 dissociates from the Re center in a rate-determining step before insertion. Prior dissociation of dichloromethane to afford the oxidative addition product was reported by Gladysz et al. in the case of the complex [Cp*Re(CO)(NO)(PPh3)(ClCH2Cl)][BF4].5a However, addition of 1-hexene as a trap for the CpRe(CO)2 transient which is expected to form upon CH2Cl2 dissociation does not affect the rate of insertion and, importantly, does not result in the formation of the previously observed CpRe(CO)2(1-hexene) complex.7b This finding indicates that CpRe(CO)2 is not generated during the oxidative addition reaction and that the rhenium−chloride bond in 1 is not broken once it forms. Supporting this conclusion is the finding that the decay rate constant for 1 is unaffected by a change in the concentration of CH2Cl2. As shown in Scheme 2, we considered three mechanisms for the oxidative addition process: (i) chloride ion abstraction, (ii) a concerted reaction with a three-centered transition state
Figure 1. Free energy profile (kcal/mol) at 298 K for the interconversion of the rotamers 1-up and 1-down.
spectroscopic characteristics of these two isomers are similar, with a calculated difference of only 10 cm−1 in the CO stretching band positions (Table S1). As shown in Figure 2, close inspection of the IR spectrum of 1 shows a slight but noticeable asymmetry and broadening of the lower wavenumber CO stretching absorbance. This asymmetry is more
Scheme 2. Possible Transition States for Oxidative Addition
Figure 2. Difference FTIR spectra obtained upon photolysis of a heptane solution of CpRe(CO)3 with 5 mM dichloromethane at 278 K. The inset shows the decay and growth of 1 and 2, respectively. B
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(b). Cp*Re(CO)3 + CH2Cl2. To ascertain the effect of increased metal electron density upon the rate of oxidative addition, the photolysis of Cp*Re(CO)3 with CH2Cl2 was studied. The reaction was significantly faster in this case. For example, unlike the Cp system, the reaction could not be monitored at ambient temperatures, since the initial adduct Cp*Re(CO)2(ClCH2Cl) (3) converted rapidly to cis-CpRe(CO)2(Cl)(CH2Cl) (4) under these conditions. However, as shown in Figure 4, the insertion reaction could be monitored
common to that found for C−H oxidative addition reactions, and (iii) chlorine radical abstraction to form CpRe(CO)2Cl• + • CH2Cl followed by radical recombination to yield 2. The measured activation parameters together with DFT modeling are not consistent with pathways i and ii, which have calculated barriers of ΔH⧧ = 164.6 and 37.3 kcal/mol, respectively, significantly greater than the experimentally determined value of 20.4 kcal/mol. For pathway iii the infinitely separated radical pair CpRe(CO)2Cl• and •CH2Cl was found to be 16.4 kcal/ mol higher in enthalpy than 1. We therefore favor a radical mechanism for C−Cl activation. Interestingly, in the reaction profile shown in Figure 3, the calculated transition state enthalpy of 13.9 kcal/mol is considerably lower than the experimental value.
Figure 3. Calculated enthalpy profile (kcal/mol) at 298 K for the oxidative addition of the C−Cl bond in CH2Cl2 to the Re center.
Figure 4. Difference FTIR spectra obtained upon photolysis of a heptane solution of Cp*Re(CO)3 with 5 mM CH2Cl2 at 243 K. The inset shows the decay and growth of complexes 3 and 4, respectively.
This observation indicates that the radicals in the transition state structure are closer to the infinite separation limit than the calculations suggest. The experimentally determined positive activation entropy also argues against a concerted oxidative addition process and is consistent with a separated radical pair in the transition state. However, experiments conducted in the presence of the free radical scavenger TEMPO showed no change in either the rate constant for the formation of 2 or its yield. Thus, it is likely that the radicals have not diffused out of the cage in the transition state and that therefore recombination is faster than the reaction of either of the radicals with TEMPO. Similar results were previously reported in oxidative addition reactions of alkyl halides to CpW(CO)3.11 A radical mechanism for dichloromethane oxidative addition to other metal centers has been reported as well.12 Photolysis of CpRe(CO)3 in the presence of 1,2-dichloroethane or 1,3-dichloropropane at room temperature results in the formation of the respective chloro adduct (Figure S3, Supporting Information). However, in contrast to the case for 1, these species are relatively stable with lifetimes of >50 s at 303 K. Importantly, the complexes decay without evidence for C−Cl oxidative addition. Presumably, the presence of two chlorine atoms on the carbon atom undergoing insertion in dichloromethane weakens the C−Cl bond relative to Cl(CH2)nCl (n = 1, 2) and promotes radical abstraction for oxidative addition. The lack of oxidative addition in this case is also consistent with the assignment of a radical mechanism for the reaction. DFT calculations predict that while for the concerted oxidative addition pathway ii the transition state enthalpies are similar for both ClCH2CH2Cl and CH2Cl2 (36.7 and 37.3 kcal/mol, respectively), it is higher by almost 7 kcal/ mol for the radical mechanism.
from 238 to 258 K. The CO bands of complex 3 are broad and asymmetrical, consistent with the presence of two rotamers, 3up and 3-down, discussed previously and similarly to the Cp analogue, DFT calculations predict the up isomer to be favored by 1.8 kcal/mol (ΔG°298 K). Eyring analysis yields activation parameters of ΔH⧧ = 19.1 ± 0.4 kcal/mol and ΔS⧧ = 20 ± 2 eu (Figure S4, Supporting Information). Extrapolation to 293 K suggests that the insertion rate is more than 600 times faster for the Cp* system in comparison to Cp at this temperature. DFT modeling also predicts that the enthalpic barrier for oxidative addition of the C−Cl bond (infinite separation of radicals) is 3 kcal/mol lower for the Cp* system. These results are consistent with an increase in the oxidation state of the Re center in the transition state, which is expected to be better stabilized by the electron-rich Cp* ligand, leading to a lower barrier and faster rate as observed. (c). Ni(CO)2(PPh3)2 + CH2X2 (X = Cl, Br, I). Thermally induced nickel-catalyzed coupling reactions of organic halides have been well explored.13 Both aryl and alkyl halides have been employed in these transformations. A key step in these reactions is the oxidative addition of the C−X bond to the nickel center. In the case of alkyl halides, there is considerable evidence for a radical mechanism of insertion.14 Nevertheless, direct observation of many of the intermediates in this chemistry has not been reported. We thought to generate the coordinatively unsaturated Ni(CO)(PPh3)2 intermediate photolytically from the parent Ni(CO)2(PPh3)2 complex in the presence of organic halides, with a view toward the identification of intermediates along the insertion pathway and for potential comparison with the rhenium system. Unfortunately, the parent dicarbonyl was photoactive only at 266 nm; as aryl halides strongly absorb at this wavelength, satisfactory IR spectra for the reactions could not be obtained. C
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calculated stabilities of the halide complexes shown in Figures S5−S7 (Supporting Information) also demonstrate that, in comparison to the reactants, the iodide product is 6.3 kcal/mol (ΔG°298 K) more stable than the chloride species. Because of the relative stability of the iodo complex and since photolysis of CpRe(CO)3 with 1,2-dichloroethane yielded only the chloro adduct, we photolyzed Ni(CO)2(PPh3)2 in the presence of 1,2diiodoethane to obtain evidence for similar coordination prior to oxidative addition. However, even at 248 K, the only species observed was a monocarbonyl complex with a CO stretching absorbance at 2028 cm−1, which by comparison with the CH2I2 system is tentatively assigned to the insertion product Ni(I)(CH2CH2I2)(CO)(PPh3)2. The available data indicate that, while the rate of C−X oxidative addition to a Ni center is faster, the resulting complex is less stable than in the case of the Re system.
We therefore focused on investigating nickel insertion into the C−X bonds of alkyl halides. As shown in Figure 5, photolysis of Ni(CO)2(PPh3)2 at 213 K in neat dichloromethane results in the formation of a
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CONCLUSION Photolysis of CpRe(CO)3 in the presence of dichloromethane results in the initial formation of the chloro adduct CpRe(CO)2(ClCH2Cl). At room temperature, this species rapidly converts to the complex CpRe(CO)2(Cl)(CH2Cl) formed upon oxidative addition of the C−Cl bond to the Re center. The activation enthalpy of 20.4 kcal/mol for the insertion process is consistent with a radical pathway for the reaction. Photolysis of Ni(CO)2(PPh3)2 with dihalomethanes yields highly reactive monocarbonyl species with CO stretching bands in the region expected for a Ni(II) species and are tentatively assigned as Ni(X)(CH2X)(CO)(PPh3)2 complexes.
Figure 5. Difference FTIR spectra obtained upon photolysis of Ni(CO)2(PPh3)2 in the presence of CH2X2 (X = Cl, Br, I).
transient monocarbonyl species with a CO stretching absorbance at 2013 cm−1. A weak peak at 2133 cm−1 is also observed, which is assigned to free CO generated upon photolysis of the Ni complex. The significantly higher wavenumber for this terminal carbonyl absorbance relative to the parent (Ni(0)) is evidence of a change in oxidation state of the metal center. By comparison with previously characterized complexes such as Ni(I)2(CO)(PPh3)2,15 the location of this CO band is in the region expected of a Ni(II) center and we therefore tentatively assign this species to the oxidative addition product Ni(Cl)(CH2Cl)(CO)(PPh3)2 (5). Even at 213 K, this complex is very reactive with a lifetime of less than 1 s. Unlike the Re species discussed above, no evidence for the formation of the σ complex Ni(CO)(PPh3)2(ClCH2Cl) prior to oxidative addition was obtained. The short lifetime of 5 and the poor quantum yield for its formation precluded further investigation into its reactivity. Consistent with the structural assignment, DFT calculations predict that 5 is favored over the reactants with ΔG°298 K = −5.1 kcal/mol. Initial chloro coordination is calculated to be slightly endoergic, which is consistent with the lack of observation of the σ complex. A pseudo-squarepyramidal geometry is calculated for 5 with the halide ligand in the axial position (Figure S5, Supporting Information). This structure is related to those postulated in nickel-catalyzed carbonylation reactions15 and suggests that oxidative addition of alkyl halides to the coordinatively unsaturated Ni(CO)(PPh3)2 complex is rapid and can lead to the formation of a five-coordinate nickel(II) monocarbonyl complex. It is important to note that we cannot comment on the mechanism of formation of this tentatively assigned complex. Previous studies have suggested the involvement of a radical mechanism for the oxidative addition of organic halides to the coordinatively unsaturated Ni(PEt3)3 complex.16 Similar results were obtained upon photolysis of Ni(CO)2(PPh3)2 with CH2Br2 and CH2I2, with monocarbonyl species observed with CO bands at 2016 and 2027 cm−1, respectively. Consistent with the trend in the metal−halogen bond strengths,17 the lifetime of the transient species increased in the order Cl < Br < I such that the iodide complex could be observed at room temperature for a few seconds. The
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EXPERIMENTAL SECTION
Time resolved IR spectra were obtained using a Bruker Vertex 80 FTIR equipped with step-scan and rapid-scan capabilities (2200−1800 cm−1). Sample photolysis was conducted using the fourth harmonic (266 nm) of a Nd:YAG laser (Quantel Brilliant B). To prevent multiple photolysis events, all spectra were obtained with a single shot of the laser. A temperature-controlled 0.5 mm path length IR cell with CaF2 windows (Harrick Scientific) was used to acquire IR spectra at or above ambient temperature. The temperature was monitored by a thermocouple located close to the photolysis solution and maintained by a water circulator to within ±0.1 °C. For low-temperature experiments, a 0.5 mm path length variable-temperature IR cell (Specac) was used. All spectra were obtained at 4 cm−1 resolution. Errors in the reported kinetic parameters were obtained from linear least-squares fits to the available data. Heptane and dichloromethane solvents were of anhydrous grade (Aldrich Sure Seal) and of >99% purity. The complexes CpRe(CO)3, Cp*Re(CO)3, and Ni(CO)2(PPh3)2 were obtained from Strem Chemicals and used as received. All calculations were performed in the development version of the Gaussian suite of programs using density functional theory.18 Geometries were optimized using the ωB97XD functional,19 which includes different fractions of exact exchange in the long and short ranges, as well as a dispersion correction. Rhenium complexes were described using the def2-TZVPP basis set and nickel complexes using the Def2-SVP basis set.20 Both basis sets describe the core electrons of the heavy atoms (Re and I) using an effective core potential. The computed gas-phase geometries were confirmed to be ground-state structures or transition states according to their number of imaginary frequencies. The energies reported in this paper are gas-phase enthalpies and free energies computed at 298.15 K and 1 atm and are expressed in kcal/mol. Figures of computed geometries included in this work were rendered using GaussView5.21 D
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(7) (a) Calladine, J. A.; Torres, O.; Anstey, M.; Ball, G. E.; Bergman, R. G.; Curley, J.; Duckett, S. B.; George, M. W.; Gilson, A. I.; Lawes, D. J.; Perutz, R. N.; Sun, X.-Z.; Vollhardt, K. P. C. Chem. Sci. 2010, 1, 622. (b) Bengali, A. A. J. Organomet. Chem. 2005, 690, 4989. (8) (a) Yang, P.-F.; Yang, G. K. J. Am. Chem. Soc. 1992, 114, 6937. (b) Bengali, A. A.; Fan, W. Y. Organometallics 2008, 27, 5488. (9) Clot, E.; Oelckers, B.; Klahn, A. H.; Eisenstein, O.; Perutz, R. N. Dalton Trans. 2003, 4065. (10) (a) Hoyano, J. K.; Graham, W. A. G. Organometallics 1982, 1, 783. (b) Dong, D. F.; Hoyano, J. K.; Graham, W. A. G. Can. J. Chem. 1981, 59, 1455. (c) Complex 2 may also reasonably be identified as the previously well characterized CpRe(CO)2Cl2 species. However, since the dichloride has νCO bands that are ∼20 cm−1 higher than those observed here, we favor identification of 2 as CpRe(CO)2Cl(CH2Cl): Einstein, F. W. B.; Klahn-Oliva, A. H.; Sutton, D.; Tyres, K. G. Organometallics 1986, 5, 53. (11) (a) Scott, S. L.; Espenson, J. H.; Zhu, Z. J. Am. Chem. Soc. 1993, 115, 1789. (b) Pelling, S.; Botha, C.; Moss, J. R. J. Chem. Soc., Dalton Trans. 1983, 1495. (12) (a) Pelling, S.; Botha, C.; Moss, J. R. J. Chem. Soc., Dalton Trans. 1983, 1495. (b) Olson, W. L.; Nagaki, D. A.; Dahl, L. F. Organometallics 1986, 5, 630. (c) Bartocci, C.; Maldotti, A.; Sostero, S.; Traverso, O. J. Organomet. Chem. 1983, 253, 253. (d) Algarra, A. G.; Braunstein, P.; Macgregor, S. A. Dalton Trans. 2013, 42, 4208. (13) (a) Semmelhack, M. F.; Helquist, P. M.; Jones, L. D. J. Am. Chem. Soc. 1971, 93, 5908. (b) Tsou, T.; Kochi, J. K. J. Am. Chem. Soc. 1979, 101, 6319. (c) Negishi, E.; King, A. O.; Okukado, N. J. Org. Chem. 1977, 42, 1821. (d) Colon, I.; Kesley, D. R. J. Org. Chem. 1986, 51, 2627. (e) Leadbeater, N. E.; Resouly, S. M. Tetrahedron Lett. 1999, 40, 4243. (14) (a) Phapale, V. B.; Buñuel, E.; Garcı ́a-Iglesias, M.; Cárdenas, D. J. Angew. Chem., Int. Ed. 2007, 46, 8790. (b) Zultanski, S. L.; Fu, G. C. J. Am. Chem. Soc. 2013, 135, 624. (c) Lin, X.; Phillips, D. L. J. Org. Chem. 2008, 73, 3680. (d) Luh, T.-Y.; Leung, M.-k; Wong, K.-T. Chem. Rev. 2000, 100, 3187. (15) Saint-Joly, C.; Mari, A.; Gleizes, A.; Dartiguenave, M.; Dartiguenave, Y.; Galy, J. Inorg. Chem. 1980, 19, 2403. (16) (a) Tsou, T. T.; Kochi, J. K. J. Am. Chem. Soc. 1979, 101, 6319. (b) Connor, J. A.; Riley, P. I. J. Chem. Soc., Chem. Commun. 1976, 634. (17) (a) Hartley, F. R. Nat. Phys. Sci. 1972, 236, 75. (b) Cottrell, T. L. The Strengths of Chemical Bonds, 2nd ed.; Butterworth: London, 1958. (c) Benson, S. W. J. Chem. Educ. 1965, 42, 502. (d) Kerr, J. A. Chem. Rev. 1966, 66, 465. (18) Frisch, M. J., et al. Gaussian Development Version, Revision H.32; Gaussian, Inc., Wallingford, CT, 2009. (19) Chai, J.-D.; Head-Gordon, M. Phys. Chem. Chem. Phys. 2008, 10, 6615. (20) Weigend, F.; Ahlrichs, R. Phys. Chem. Chem. Phys. 2005, 7, 3297. (21) Dennington, R.; Keith, T.; Millam, J. GaussView, Version 5; Semichem Inc., Shawnee Mission, KS, 2009.
ASSOCIATED CONTENT
S Supporting Information *
Figures, tables, and xyz files giving IR data, kinetic parameters, Eyring plots, and calculated structures. This material is available free of charge via the Internet at http://pubs.acs.org.
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AUTHOR INFORMATION
Corresponding Authors
*E-mail for B.A.A.:
[email protected]. *E-mail for A.A.B.:
[email protected]. Notes
The authors declare no competing financial interest.
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ACKNOWLEDGMENTS This publication was made possible by funding from the Qatar National Research Fund (member of Qatar Foundation). The experimental work was supported by NPRP award 5-156-1-037 and the theoretical studies by NPRP award 09-143-1-022. The statements made herein are solely the responsibility of the authors.
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REFERENCES
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