Metastable Electrons, High-Mobility Solvent Anions, and Charge

Kengo Itoh, Azusa Muraoka, Kazuo Watanabe, and Takashi Nagata , Masaru ... Influence of Electrode Characteristics on DC Point-to-Plane Breakdown in ...
0 downloads 0 Views 137KB Size
Download PDF
4520

J. Phys. Chem. B 2001, 105, 4520-4530

Metastable Electrons, High-Mobility Solvent Anions, and Charge Transfer Reactions in Supercritical Carbon Dioxide I. A. Shkrob* and M. C. Sauer, Jr. Chemistry DiVision, Argonne National Laboratory, Argonne, Illinois 60439 ReceiVed: January 25, 2001; In Final Form: March 16, 2001

Time-resolved dc photoconductivity has been used to observe quasifree electrons and solvent radical anions in supercritical CO2 (Fc ≈ 0.468 g/cm3; Tc ≈ 31 °C). The electrons have lifetime τe < 200 ps and mobility µe > 10 cm2/(V s). For F/Fc > 0.64, the product µeτe rapidly increases with the solvent density F, reaching 2.5 × 10-9 cm2/V at F/Fc ≈ 1.82. The mobility of the solvent anions exponentially increases with F, being 2-10 times higher than the mobilities of other ions. The activation energy of the solvent anion migration is 0.46 eV, whereas for solute ions, this energy is less than 20 meV. Electron detachment upon 1.0-3.5 eV photoexcitation of the solvent anion has been observed. The cross section of photodetachment linearly increases with photon energy above 1.76 eV end exhibits a shoulder at 2.82 eV; the spectrum is similar to photoelectron spectra of (CO2)n- clusters (n ) 6-9) in the gas phase. Both the electrons and solvent anions react with nonpolar electron acceptor solutes whose gas-phase electron affinity (EAg) is greater than 0.4 eV. At T ) 41 °C and F/Fc ≈ 1.77, the rates of the electron attachment and solvent anion scavenging correlate with each other and the solute EAg. Only solutes with EAg > 2 eV exhibit diffusion-controlled kinetics. For F/Fc > 0.85, the scavenging radii of these diffusion-controlled reactions decrease with F. For O2, the scavenging of the solvent anion is reversible (∆H° ≈ -0.44 eV).

1. Introduction Supercritical (sc) fluids allow a high degree of control over the solvent properties, in particular, a smooth transition between gaslike and liquidlike behavior. This tuneability accounts for the growing popularity of sc fluids in studies of reaction mechanisms in condensed media. Studies on solvation,1 electron transfer,2 exciplex and excimer formation,3 radical recombination,4 etc. have been reported. These studies, as well as molecular dynamics simulations,5 suggest that regions of low and high density coexist in sc fluids, and the latter may be viewed as rapidly shifting aggregates of molecular clusters.5,6 Even near critical density, the lifetime of these clusters could be tens of picoseconds.5 To the many classes of reactions that have already been studied,1-4 ion-molecule reactions are the most recent addition (e.g., ref 7). Such reactions are interesting per se since ions are more sensitive to their environment than neutral species. Anions and cations formed upon the ionization of the solVent itself are expected to be especially sensitive to the local solvent structure. Reactions of these solvent ions and their short-lived precursors, electrons and holes, are of much significance for radiation chemistry of technologically important sc fluids, such as water and CO2, considered as prospective reactor coolants. Carbon dioxide (with a critical temperature Tc of 31 °C and a critical density Fc of 0.468 g/cm3) is the most studied of these sc fluids. While the radiolysis of gaseous, liquid, and subcritical CO2 has been studied for many years, only recently have the solvent ions in densified CO2 been identified and their reactions examined. Using transient absorbance-pulse radiolysis, Dimitrijevic et al. observed solvent radical cation (C2O4•+ dimer) in sc-CO2 and studied its reactions with dimethylaniline, O2, * To whom the correspondence should be addressed.

CO, etc.7 In the same study, reactions of negative charge carriers with p-benzoquinone were observed. While 50-80% of the p-benzoquinone•- anions are formed within 100 ps after a short 20 MeV electron pulse, some of these anions are formed over tens of nanoseconds. In the low-density regime (F/Fc < 0.3), the reaction is diffusion-controlled, but in the high-density regime (F/Fc > 1), it is ≈10 times slower than charge-transfer reactions of C2O4•+. One way to rationalize these observations is to assume that the prompt formation of p-benzoquinone•- is due to rapid scavenging of a short-lived, mobile precursor of the solvent radical anion (e.g., quasifree electron), while the slow reaction is that of a less mobile, fully thermalized species (e.g., solvent radical anion). The short-lived precursor would also account for the observations of Yoshimura et al.8 who studied the yield of CO in radiolysis of sc-CO2 containing electron scavenger, SF6. It is believed that CO is formed by recombination of primary electron-hole pairs, via dissociation of excited solvent molecules. A rather high concentration, ≈30 mM of SF6, was needed to halve the yield of CO, suggesting that SF6 reacted with a short-lived species. Unfortunately, pulse radiolysis experiments leave uncertainty as to the nature of the radical anion since prethermalized electrons could attach to (CO2)n clusters both dissociatively and nondissociatively

e- + (CO2)n f (CO2)n•-

(1)

e- + (CO2)n f CO + O•-(CO2)n-1 f CO + CO3•-(CO2)n-2 (2) depending on the electron energy (for n ) 2, reaction 2 requires ≈3.6 eV more energetic electrons).9 In radiolytic spurs, a large

10.1021/jp010325h CCC: $20.00 © 2001 American Chemical Society Published on Web 04/17/2001

J. Phys. Chem. B, Vol. 105, No. 19, 2001 4521

Electrons and Solvent Radical Anions in sc-CO2 fraction of electrons have initial excess energies above 10 eV, and it is not clear which one of these negative-charge species is responsible for the slow reaction. In the gas phase, anions formed in reaction 1 have been extensively studied.10 A linear CO2 molecule has negative electron affinity (EA) of -0.6 ( 0.2 eV. The metastable C2V monomer anion, CO2•- (with OCO angle of 135°and autodetachment time of 60-100 µs) exhibits vertical detachment energy (VDE) of 1.33-1.4 eV.10a This energy increases to 2.6 eV10b (or 2.79 eV)10c in the stable dimer anion, C2O4•- (with D2d geometry and lifetime >2 ms), and further increases to 3.4 eV for n ) 6 clusters.10b In larger clusters, the VDE first drops to 2.49 eV (n ) 7), then monotonically increases to 3.14 eV (n ) 13), and then, for n g 14, suddenly increases to 4.55 eV.10c The onset of the photoelectron spectrum increases from ≈1.5 eV for n ) 2-7 clusters to 1.8-2 eV for n ) 8-13 clusters to 3 eV for n ) 15-16 clusters. The structure of these anions is poorly known. Johnson and co-workers argue that while the core of small (n e 6) and large (n g 14) clusters is a D2d dimer anion, the core of 6 < n 0.7 g/cm3 (Figure 34S(a)). In the intermediate density range, which did not change with temperature, the kinetics showed no discernible break, decaying almost linearly (Figure 34S(b)). When normalized, the TOF kinetics obtained in different fields E were the same when plotted as a function of the product (delay time × E), Figure 35S. This behavior suggests that the sloping kinetics are not caused by chemical transformation of the cation during the flight (the time scale of such a transformation would not depend on the electric field). Abnormal TOF kinetics were also observed in an experiment where the ions were generated uniformly between the electrodes (Figure 36S). Though we have not pursued this behavior, it seems that the only way to account for these observations is to postulate a distribution of cation mobilities (but not anion mobilities) for F ∼ Fc. A possible cause for this dispersion is polymerization of aromatic cations. Indeed, for a nonaromatic solute, triethylamine, the dispersion of the TOF kinetics is much lower, and we were able to determine the cation mobilities over the entire density range (Figure 37S). Once more, the mobility was a function of the solvent density only. Where comparison is possible, long-lived cations in benzene and triethylamine solutions have similar mobilities. A comparison of the diffusion coefficients for benzene22 and the corresponding cation (Figure 38S) indicates that the latter is ≈2 times lower. Presumably, most of these cations are dimer radical cations of benzene.23 Except for the near-critical density range, the mobilities of solute ions obtained in our photolytic experiments are similar to those observed in X-ray radiolysis of neat sc-CO2.21 We speculate that the species studied in these radiolytic experiments were impurity ions. 4. Discussion 4.1. Solvent Radical Anion. Anion-1 is the solvent radical anion. Indeed, (i) the same anion is produced in photoionization of different hydrocarbon solutes, (ii) it behaves as a radical anion in its reactions with electron-acceptor solutes, (iii) the UV excitation energy is too low to produce CO3•- via reaction 2 (the latter is also known to react with N2O in the gas phase,24 while anion-1 does not react with N2O), (iv) photoexcitation of anion-1 yields electrons, and (v) this anion has much higher mobility than regular solute ions. What kind of a solvent radical anion is it? First of all, in this anion, the electron is deeply trapped. The scavenging reactions become diffusion-controlled only for solutes with EAg > 2 eV. Though the electron photodetachment was observed even at 1.17 eV, at this energy the yield of electrons was very low, Iφ ≈ 3.2 M-1 cm-1 vs 870 M-1 cm-1 at 2.82 eV. The photodetachment threshold Eth ) EAliq + V0 is ≈1.76 eV, where EAliq is the liquid-phase electron affinity,

EAliq ) EAgas + Es

(5)

Es is the solvation energy (estimated using the Born equation), and V0 is the energy of the conduction band electron. The latter is unknown, but in other supercritical fluids, V0 is minimum

Electrons and Solvent Radical Anions in sc-CO2

J. Phys. Chem. B, Vol. 105, No. 19, 2001 4527

where the electron mobility µe is maximum25 (which is explained by the Basak-Cohen model of electron transport).26 Thus, we believe that in high-density sc-CO2 this energy V0 is a small fraction of an electron volt. In the vapor, the EA of (CO2)n clusters (n ≈ 6) is ≈0.9 eV. 15a The molar volume of CO2 is 42.7 cm3/mol,27 and the n ≈ 6 cluster has a radius ≈0.47 nm. In high-density sc-CO2,  ≈ 1.5,19 and the Born solvation energy of an n ≈ 6 cluster is 0.51 eV, which gives EAliq ≈ 1.4 eV. The solvent anion is reversibly scavenged by O2 (sections 3.3), yielding a CO4•- anion (O2•- covalently binds to CO2).28 The heat of this reaction is -0.44 eV (section 3.3), whereas the gasphase VDE for CO4•- is 1.22 eV.28 Thus, the solvent EAliq is 0.78 eV plus the solvation energy of CO4•-. The latter must be around 0.8 eV (Table 1), which gives EAliq ≈ 1.6 eV. All of these estimates suggest that the solvent EAliq is between 1.4 and 1.8 eV. The photodetachment spectrum of the solvent anion (Figure 9) resembles photoelectron spectra of small (CO2)n•- clusters with n ) 6-9.10b,c These n ) 6-9 anions exhibit an onset at ≈1.7 eV and a Gaussian peak centered at 2.5-2.9 eV. For CO2•- monomer in hydrocarbon solutions, the onset of photodetachment is ca. 1 eV and there is a well-resolved maximum at 3 eV.13 For anion-1, the cross section at 2.82 eV is 4 × 10-18 cm2, which is 10 times lower than the cross section of electron photodetachment for CO2•- in isooctane.13 Thus, in sc-CO2 the electron is much more deeply trapped; the solvent anion is certainly different from the monomer anion in hydrocarbon solutions. Qualitatively, our photodetachment spectrum is consistent with a bound-to-continuum transition postulated by Lukin and Yakovlev.13 Recently, Takahashi and Jonah7c found evidence for the solvent anion absorption in pulse radiolysis of sc-CO2 mixtures containing H2, CO, or N2O. None of these solutes reacts with the solvent anion. By contrast, the dimer cation (C2O4•+) reacts rapidly, as evidenced by the disappearance of the characteristic absorption band centered at 1.6 eV (Figure 39S, traces i-iii). In the gas phase, C2O4•+ is known to react with all of these solutes,24 for example

C2O4•+ + H2 f HCO2+ + CO2 + H•

(6)

The residual spectra are shown in Figure 39S, trace iv. The absorption band has an onset at 1.65 eV and reaches a plateau at 2 eV. At 80 ns delay, the product G I for this band is ∼370 M-1 cm-1 (100 eV)-1, where G is the radiolytic yield. This absorption band was not observed when O2 and SF6 were added to the reaction mixture. It is likely that it is from the solvent radical anion observed in our dc conductivity experiments. The spectrum (iv) shown in Figure 39S would account for the low cross section of electron photodetachment below 2 eV (due to poor light absorbance). One of the most intriguing of our findings is the rapid, thermally activated migration of the solvent anion. At least two models could account for this behavior: In the first model, the solvent anion is in dynamic equilibrium with the quasifree electron (eqf-). In such a case, the “lifetime” τe of this electron equals the settling time of the equilibrium reaction 7,

eqf- + (CO2)n h (CO2)n•-

(7)

and the apparent anion mobility µ1- ≈ µeQ, where Q ) {1 + Keq[(CO2)n]}-1 is the equilibrium fraction of eqf- and Keq is the equilibrium constant. Reaction 7 would explain the observed correlation between the reaction rates for electrons and solvent

anions (Table 1). In this model, the quantity µeτe/µ1- (≈µeτe/ µ1s plotted in Figure 27S) is the inverse rate of electron emission; this rate has an isochore activation energy between 0.35 and 0.45 eV. The main difficulty with this model is unfavorable energetics. For µe ∼ 10-100 cm2/(V s) and µ- ≈ 0.014 cm2/(V s), Keq[(CO2)n] ≈ 700-7000 and the free energy of electron trapping must be between -0.14 and -0.21 eV (for 〈n〉 ≈ 6). This energy is too low as compared with the heat of reaction 1, ≈-1.76 eV (estimated from the photodetachment threshold and thermophysical calculations above). Thus, one needs to postulate unrealistically large standard entropy ∆S° for electron trapping, at least -400 J mol-1 K-1. For gas-phase clustering, X-(CO)n-1 + CO2 ) X-(CO2)n (where X- is a halide anion and n e 5), ∆S° is -(80-90) J mol-1 K-1.29 In hydrocarbon solutions, ∆S° for reaction 1 is -51 to -59 J mol-1 K-1. 12 Having an overall ∆S° of -400 J mol-1 K-1 is, therefore, equivalent to producing a monomer anion and binding four CO2 molecules to it. Jacobsen and Freeman15a estimated that in dense CO2 vapor, the entropy for formation of n ≈ 6 clusters from individual CO2 molecules followed by electron attachment was about -274 J mol-1 K-1. Since the quasifree electrons in sc-CO2 must be attaching to preexisting (CO2)n clusters, this figure overestimates the entropy of reaction 1 many times; still it is much lower than needed to reconcile the data. It appears that the equilibrium model is not supported by our data. A reviewer pointed out another problem with the equilibrium model: reaction 7 would be accompanied by a large decrease in volume, as is the case for trapping reactions of quasifree electrons in liquid and supercritical hydrocarbons.14 Then, an increase in pressure would shift reaction 7 toward anion-1 and decrease µ1-, opposite to observations. This argument, however, could be incorrect, as the unproven assumption must be made that the decrease in the equilibrium fraction Q of the electrons with pressure is larger than the corresponding increase in the electron mobility, µe. In the second model, the solvent anion migrates by thermally activated degenerate electron transfer. Such rapidly hopping anions were previously observed in liquid C6F6.30 In this solvent, the electron rapidly attaches to C6F6; the solvent anion migrates with mobility of 0.02 cm2/(V s) (≈44 times faster than other ions) and activation energy of 0.11 ( 0.01 eV. The absorption band of C6F6•- in neat C6F6 is red-shifted by 0.5 eV as compared to the monomer anion. It was proposed that the charge is hopping via resonant electron transfer between the solvent and multimer (C6F6)n•- anions (n g 2). The analogy with the solvent anion in sc-CO2 is striking; even the mobilities of these two solvent anions are similar. Charge hopping qualitatively accounts for the increase in the anion mobility with the solvent density, due to reduction in the average distance (and the electron coupling integral) between (CO2)n clusters. As discussed in the Introduction, the VDE strongly varies for (CO2)n•- clusters of different sizes and there is core switching for the n ) 6-7 and n ) 13-14 clusters.10 The transfer of electron between the clusters could be caused either by their reorganization (due to collisions or monomer exchange) or by this core switching. The activation energy of 0.46 eV is equivalent to the loss or exchange of just two CO2 molecules. Assuming that the average jump length is 0.94 nm (the diameter of the 〈n〉 ≈ 6 cluster), the residence time of the electron trapped by a given molecular cluster would be ≈4 ps at 41 °C and ≈0.6 ps at 65 °C. These residence times are much shorter than mean lifetimes of solvent clusters (>40 ps) estimated from molecular dynamics simulations of Tucker et al.5

4528 J. Phys. Chem. B, Vol. 105, No. 19, 2001 Though the case for charge hopping seems convincing, only ultrafast studies of the electron dynamics can decide which model, equilibrium or hopping, is correct. Knowing this mechanism is prerequisite for understanding ion-molecule reactions in sc-CO2. Indeed, if there is an electron-anion equilibrium, the decay of the conductivity signal would depend on both scavenging constants, which have different density and temperature dependencies. In the following, we will assume that the electron is trapped irreversibly. As stated in section 3.2, charge-transfer reactions of solvent anions do not follow any obvious pattern. The reaction rates correlate with solute gas-phase EA's, albeit inexactly (Table 1). At face value, this correlation is what one would expect from the Marcus' theory of charge transfer;31 the latter has been used by Dimitrijevic et al. 7 to analyze the rate constants for solvent radical cation in sc-CO2. However, it is not clear why these gas-phase electron affinities would correlate with the enthalpies of the corresponding scavenging reactions: In the gas phase, the product of electron transfer from the (CO2)n•- anion (n g 2) to oxygen is CO4•- rather than O2•-.32 The overall reaction is exothermic by 0.54 eV; 32 direct measurement in sc-CO2 gives 0.44 eV (section 3.3). Reaction with SF6 yields SF5• and F-. Klots and Compton reported that low-energy electron attachment to CCl4 in a supersonic CO2 beam resulted in the formation of metastable CCl4•- and CCl4•-(CO2) anions.33 In our analysis (Table 1), it is assumed that CCl4•- does not dissociate on the time scale comparable to the lifetime of the reaction complex, but that could be incorrect. The fragmentation of SF6 and CCl4 yields small halide anions, which require high solvation energy, and radicals, for which this solvation energy is difficult to estimate. All of these “details” complicate thermophysical analysis. Table 1 gives the best estimates for EAliq (eq 5). Using corrected values clearly makes the correlation worse. In the Marcus' theory of charge transfer,31 rate constants for a series of homologous solutes correlate with the free energy ∆Gr of reaction, provided that the “solvent organization energy” is similar within the series. Thus, either the reaction entropy or this “reorganization energy” is different for different solutes, destroying the correlation. Alternatively, the Marcus theory is inapplicable: there is no correlation between ∆Gr and the reaction barrier. Since the rate constants correlate with EAg rather than the heat of the overall reaction, it appears that the scavenging is initiated by direct electron transfer to the solute. Solvation, fragmentation, and complexation of the resulting anion occur later and have little effect on the reaction rates (although considerable effect on the reaction heat). Phenomenologically, density dependencies of scavenging constants (Figure 7) may be accounted for by a systematic decrease in the reaction radii with density. For example, the reaction radius for CCl4 decreases from 1.2 nm at 0.38 g/cm3 to 0.4 nm at 0.85 g/cm3 (Figure 40S(a)). The cause of this behavior is unclear. If the solvent anion had a radius of ≈0.47 nm, the minimum radius of a diffusion-controlled reaction with CCl4 would be ≈0.78 nm (the molecular radius of CCl4 is ≈0.31 nm).31 Therefore, even the fastest reactions of solvent anions in high-density sc-CO2 are not diffusion-controlled. At subcritical densities, even the reactions that are slow at F > Fc become diffusion-controlled (Figure 13S(a)). For O2 and SF6, which have scavenging radii below 0.05 nm for densities greater than 0.4 g/cm3 (Figure 40S), the decrease in the radii is small or even reversed (Figure 40S(b)). With the exception of SF6, slow scavenging reactions have lower activation energies than

Shkrob and Sauer solvent anion diffusion. Such a situation implies that these reactions are kinetically rather than diffusion-controlled. 4.2. Quasifree Electron. In high-density sc-CO2 (F/Fc g 1.5), the precursor of the solvent radical anion has µe > 10 cm2/(V s), suggesting that it is a conduction band (quasifree) electron. Electron mobilities of 10-50 cm2/(V s) have been observed for many supercritical hydrocarbons, such as alkanes, cycloalkanes, and alkenes (see Table VII.2 in ref 24). It is surprising that this class of molecules provides a better reference system than diatomic molecules like O2 and N2 (both of which exhibit much lower electron mobilities). For F > Fc, the product µeτe rapidly increases with density (Figure 4a). Though our conductivity data do not indicate the density dependence of which parameter, µe or τe, causes the rapid growth of the product µeτe, almost certainly it is the electron mobility. Figure 4b shows the dependence of the density-normalized product (Fµe)τe as a function of reduced density, F/Fc. For F/Fc between 1.2 and 1.8, this product exponentially increases with the solvent density. Very similar dependencies of Fµe on F/Fc, in the same density range of 1 e F/Fc e 2, have been observed for excess electrons in supercritical methane, ethane, neopentane, and isobutane.25 In these fluids, the electron mobility at F ∼ Fc obeys the Cohen-Lekner equation, with Fµe ∝ Vs2T-3/2, where Vs is the speed of sound (the electron scattering is due to density fluctuations). [For scCO2, this model seems to be inappropriate, as Vs decreases with F/Fc ( 1 was found in supercritical ammonia;35 however, in this liquid, µe < 0.02 cm2/(V s). Pending ultrafast studies revealing the details of rapid electron dynamics, we may only speculate what causes the metastability of electron in high-density sc-CO2. High electron mobility with low activation energy means low efficiency of trapping; there must be relatively few shallow traps near the conduction band edge. In sc-CO2, deep electron traps are very deep indeed: attachment to the (CO2)n clusters is exothermic by 1.4-1.8 eV. Perhaps, this very exothermicity is what slows the electron trapping down, as known from Marcus theory (the famous “inverted region”).31 However, such an explanation is not entirely satisfactory: if the slow trapping were due to the unfavorable energetics, then an introduction of shallower electron traps would cause rapid electron decay. Experimentally, one can add 0.1-1 M N2O (with EAg of 0.22 eV24 and estimated EAliq of 1.14 eV) without any effect on the electron signal. We conclude that the electron is metastable not due to the scarcity of traps but due to extremely low trapping efficiency. This brings us to the rate constants for electron scavenging. As mentioned in section 3.2, both the prompt electron signal and the initial fraction of solvent anions obey the Stern-Volmer law. Using the WAS-like equation36 for the reduction in the prompt conductivity signal, eq 3, with γ ≈ 0.5-0.6, results in a poor fit, implying that the electron acceptor solute scavenges mainly free electrons (section 3.3S). The difference between coefficients Re′ extracted from the concentration dependencies of the prompt signal (eq 3) and Re extracted the fraction f0 of anions-1 (eq (5S)) is due to the interference from the geminate dynamics; it is seen from Table 1 that Re/Re′ ∼ 2-3, for all solutes. The coefficient Re ≈ ke τe, where ke is the electron scavenging constant and τe is the “natural” lifetime of the quasifree electron (section 2.1S). Since τe is the same for all solutes, a correlation between Re and k2 at T ) 41 °C and F/Fc

Electrons and Solvent Radical Anions in sc-CO2

J. Phys. Chem. B, Vol. 105, No. 19, 2001 4529

≈ 1.77 (Table 1) means that ke and k2 also correlate under these conditions. A similar correlation exists for reactions of “dry” and hydrated electrons in water.37 Though ke cannot be determined from our data, the radius Re of electron attachment can be determined from the ratio of Re and µeτe. For CCl4, this radius is 1.24 nm, typical of diffusioncontrolled electron-scavenging reactions in liquid hydrocarbons;38 for other solutes, this radius is much shorter: 0.02 nm for O2 and 0.12 nm for SF6. The EAliq for these two solutes is close to that for (CO2)n clusters (Table 1). Thus, there is, indeed, good reason to expect that the solvent clusters in sc-CO2 have a rather low cross section of electron attachment.

C. D. Jonah for sharing recent data on pulse radiolysis of scCO2, ref 7c. I.A.S. is grateful to Dr. R. A. Holroyd for prepublication versions of refs 14a and 21. This work was performed under the auspices of the Office of Basic Energy Sciences, Division of Chemical Science, US-DOE under contract No. W-31-109-ENG-38.

5. Conclusion

References and Notes

Time-resolved dc photoconductivity has been used to observe negative charge carriers in sc-CO2. It is shown that quasifree electrons are trapped in less than 200 ps, yielding long-lived (>10 µs) solvent radical anions. The mobility of these anions is several times higher than the mobility of stable solute ions and exponentially increases with the solvent density for F > 0.3 g/cm3. The activation energy for solvent anion migration is independent of the solvent density, being ca. 0.46 eV, whereas for the solute ions this activation energy is below 20 meV. Because of the short lifetime τe of the electron, the mobility µe cannot be determined. The product µeτe rapidly increases above Fc, reaching 2.5 × 10-9 cm2/V at 0.85 g/cm3; it depends weakly on temperature. The electron mobility µe is greater than 10 cm2/ (V s); probably, it is 102-103 cm2/(V s). One-photon excitation of the solvent anion at 1-3.5 eV causes electron detachment. The cross section of photodetachment linearly increases above 1.76 eV and exhibits a shoulder at 2.82 eV. A weak absorption band of the solvent anion with onset about 1.6 eV was observed in radiolysis of sc-CO2 by Takahashi and co-workers. The electron photodetachment spectrum resembles photoelectron spectra of n ) 6-9 anion clusters (CO2)n•- in the gas phase10 and does not resemble photodetachment spectra of CO2•- in liquid hydrocarbons.13 Charge transfer reactions of quasifree electrons and solvent anions with various nonpolar solutes have been studied. Only solutes with gas-phase EAg > 0.4 eV reacted with these species. Common electron scavengers, such as N2O, are unreactive toward the quasifree electron and the solvent anion. Oxygen (EAg ≈ 0.44 eV) reacts with the solvent radical anion reversibly, yielding a solvated CO4•- anion, with ∆G° ≈ - 0.44 eV. In high-density sc-CO2 (F > Fc), only electron transfer reactions with solutes that have EAg > 2 eV have nearly diffusioncontrolled rates; for p-benzoquinone, O2, and SF6 the reaction is 15-100 times slower. The reaction radii decrease with the solvent density; the rate constant of electron attachment and solvent anion scavenging correlate with each other and solute EAg. We argue that the solvent anion in high-density sc-CO2 is best viewed as an electron attached to a medium-size (CO2)n cluster. This anion does not appear to be in dynamic equilibrium with the quasifree electron, and the rapid anion migration must be due to thermally activated charge hopping. The metastability of electrons in sc-CO2 awaits explanation. As a concluding remark, sc-CO2 exhibits unusual negative charge dynamics. More research is needed, in particular, on the very short time scales.

(1) Kajimoto, O. Chem. ReV. 1999, 99, 353 (review). Biswas, R.; Lewis, J. E.; Maroncelli, M. Chem. Phys. Lett. 1999, 310, 485. Castanheira, E. M. S.; Martinho, J. M. G. Chem. Phys. Lett. 1991, 185, 319. Maiwald, M.; Schneider, G. M. Ber. Bunsen-Ges. Phys. Chem. 1998, 102, 960. Morita, A.; Kajimoto, O. J. Phys. Chem. 1990, 94, 6420. Kim, S.; Johnston, K. P. Ind. Eng. Chem. Res. 1987, 26, 1206; AIChE J. 1987, 33, 1603. (2) Kimura, Y.; Takebayashi, Y.; Hirota, N. Chem. Phys. Lett. 1996, 257, 429; J. Chem. Phys. 1998, 108, 1485. Schulte, R. D.; Kauffman, J. F. Appl. Spectrosc. 1995, 49, 31. Sun, Y.-P.; Fox. M. A. J. Am. Chem. Soc. 1993, 115, 747. Rollinson, A. M.; Drickamer, H. G. J. Chem. Phys. 1980, 73, 5981. (3) Rice, J. K.; Niemeyer, E. D.; Dunbar, R. E.; Bright, F. V. J. Am. Chem. Soc. 1995, 117, 5832. Inomata, H.; Hamatani, H.; Wada, N.; Yagi, Y.; Saito, S. J. Phys. Chem. 1993, 97, 6332. Sun, Y.-P. J. Am. Chem. Soc. 1993, 115, 3340. Brennecke, J. F.; Tomasko, D. L.; Peshkin, J.; Eckert, C. A. Ind. Eng. Chem. Res. 1990, 29, 1682; J. Phys. Chem. 1990, 94, 7692. Okada, T.; Kobayashi, Y.; Yamasa, H.; Mataga, N. Chem. Phys. Lett. 1986, 128, 583. (4) Brennecke, J. F.; Chateauneuf, J. E. Chem. ReV. 1999, 99, 433. Zhang, J.; Connery, K. A.; Brennecke, J. F.; Chateauneuf, J. E. J. Phys. Chem. 1996, 100, 12394. Randolh, T. W.; Carlier, C.; O’Brien, J. A. Int. J. Thermophys. 1996, 17, 471. Ganapathy, S.; O’Brien, J. A.; Randolf, T. W. J. Supercritic. Fluids 1996, 9, 51. Roberts, C. B.; Zhang, J.; J. F.; Chateauneuf; K. A.; Brennecke J. Am. Chem. Soc. 1995, 117, 6553; J. Phys. Chem. 1993, 97, 5618; J. Am. Chem. Soc. 1992, 114, 8455. Randolf, T. W.; Carlier, C. J. Phys. Chem. 1992, 96, 5146. Ganapathy, S.; Zawadzki, A. G.; Hynes, J. T. J. Phys. Chem. 1989, 93, 7031. (5) Tucker, S. C.; Maddox, M. W. J. Phys. Chem. B 1998, 102, 2437. G. Goodyear; Tucker, S. C. J. Chem. Phys. 1999, 111, 9673. (6) Nishikawa, K.; Takematsu, M. Chem. Phys. Lett. 1994, 226, 359; J. Phys. Chem. 1996, 100, 418 and references therein. Goodyear, G.; Madox, M. W.; Tucker, S. C. J. Chem. Phys. 2000, 112, 10327. Tucker, S. C. Chem. ReV. 1999, 99, 391 (review). (7) (a) Dimitrijevic, N. M.; Bartels, D. M.; Jonah, C. D.; Takahashi, K. Chem. Phys. Lett. 1999, 309, 61. (b) Dimitrijevic, N. M.; Takahashi, K.; Bartels, D. M.; Jonah, C. D.; Trifunac, A. D. J. Phys. Chem. A 2000, 104, 568. (c) Takahashi, K.; Jonah, C. D. Private communication. (8) Yoshimura, M.; Chosa, M.; Soma, Y.; Nishikawa, M. J. Chem. Phys. 1972, 75, 1626. (9) Klots, C. E.; Compton, R. N. J. Chem. Phys. 1978, 69, 1636. (10) (a) Bowen, K. H.; Eaton, J. G. In The Structure of Small Molecules and Ions; Naamna, R., Vagar, Z., Eds.; Plenum: New York, 1987; p 147. (b) DeLuca, M. J.; Niu, B.; Johnson, M. J. Chem. Phys. 1988, 88, 5857. (c) Tsukuda, T.; Johnson, M.; Nagata, T. Chem. Phys. Lett. 1997, 268, 429. (11) Zhou, M.; Andrews, L. J. Chem. Phys. 1999, 110, 2414 and 6820. Thompson, W. E.; Jacox, M. E. J. Chem. Phys. 1999, 111, 4487; J. Chem. Phys. 1989, 91, 1410. (12) R. A. Holroyd, Gangwer, T. E.; Allen, A. O. Chem. Phys. Lett. 1975, 31, 520. (13) Lukin, L. V.; Yakovlev, B. S. High Energy Chem. (Engl. Transl.) 1978, 11, 440. (14) (a) Holroyd, R. A.; Nishikawa, M.; Itoh, K. BehaVior of Charged Species in Nonpolar Supercritical Fluids. Electron Attachment Reactions and Ionic Transport; BNL reprint No. 66560, 1999. (b) Itoh, K.; Holroyd, R. J. Phys. Chem. 1990, 94, 8854. Nishikawa, M.; Itoh, K.; Holroyd, R. A. J. Phys. Chem. 1988, 92, 5262. (15) (a) Jacobsen, F. M.; Freeman, G. R. J. Chem. Phys. 1986, 84, 3396. (b) The kinetic analysis given in ref 15a is based upon the assumption that though the apparent electron mobility decreases with the gas density due to reversible trapping by n > 5 gas clusters, the density-normalized mobility of free electrons (scattered by CO2 molecules and small clusters) does not change. (16) Lehning, H. Phys. Lett. A 1968, 28, 103. Allen, N. L.; Prew, B. A. J. Phys. B 1970, 3, 1113. Warman, J. M.; Sowada, U.; Armstrong, D. A. Chem. Phys. Lett. 1981, 82, 458.

Acknowledgment. We thank Drs. K. Takahashi, J. Cline, N. Dimitrijevic, D. M. Bartels, and C. D. Jonah for their help and many useful discussions. We thank Dr. K. Takahashi and

Supporting Information Available: 1. Experimental setup and procedures. 2. Data analysis. 3. Geminate pair dynamics observed by dc conductivity. 4. Additional figures and figure captions. This material is available free of charge via the Internet at http://pubs.acs.org.

4530 J. Phys. Chem. B, Vol. 105, No. 19, 2001 (17) Cooper, H. W.; Goldfrank, J. C. Hydrocarbon Process. Petrol. Refiner 1967, 46, 141. (18) Yoon, P.; Thodos, G. AIChE J. 1970, 16, 300; 1962, 8, 59. (19) Moriyoshi, T.; Kita, T.; Uosaki, Y. Ber. Bunsen-Ges. Phys. Chem. 1993, 97, 589. (20) Hummel, A.; Schmidt, W. F. Radiat. Res. ReV. 1974, 5, 199. (21) Itoh, K.; Holroyd, R. A.; Nishikawa, M. J. Phys. Chem. A 2001, 105, 703. (22) (a) Swaid, I.; Schneider, G. M. Ber. Bunsen-Ges. Phys. Chem. 1979, 83, 969. (b) Bueno, J. L.; Sua´rez, J. J.; Dizy, J.; Medina, I. J. Chem. Eng. Data 1993, 38, 344. (23) Badger, B.; Brocklehurst, B. Trans. Faraday Soc. 1969, 65, 2582; 1970, 66, 2939. (24) Tabata, Y. et al., CRC Handbook of Radiation Chemistry; CRC Press: Boca Raton, FL, 1991; Ch. IV. (25) Itoh, K.; Nakagawa, K.; Nishikawa, M. Radiat. Phys. Chem. 1988, 32, 221; J. Chem. Phys. 1986, 84, 391. Nishikawa, M.; Holroyd, R. A.; Sowada, U. J. Chem. Phys. 1980, 72, 3081. (26) Basak, S.; Cohen, M. H. Phys. ReV. B. 1979, 20, 3404. See also: Vestal, M. L.; Mauclair, G. H. J. Chem. Phys. 1977, 67, 3758. (27) Grigoriev, I. S.; Meilikhov, E. Z. Handbook of Physical Quantities; CRC Press: Boca Raton, FL, 1997; p 399. (28) Pack, J. L.; Phelps, A. V. J. Chem. Phys. 1966, 45, 4316. Cosby, P. C.; Ling. J. H.; Peterson, J. R.; Moseley, J. T. J. Chem. Phys. 1976, 65, 5276.

Shkrob and Sauer (29) Keesee, R. G.; Castleman, A. W., Jr. J. Phys. Chem. Ref. Data 1986, 15, 1011 and J. Phys. Chem. 1984, 88, 2880. (30) van den Ende, C. A. M.; Nyikos, L.; Sowada, U.; Warman, J. M.; Hummel, A. J. Electrostatics 1982, 12, 97; Radiat. Phys. Chem. 1982, 19, 297; J. Phys. Chem. 1980, 84, 1155. (31) Marcus, R. A. J. Chem. Soc., Faraday Discuss. 1982, 74, 7; Annu. ReV. Phys. Chem. 1964, 15, 155. (32) Parkes, D. A. J. Chem. Soc., Faraday Trans. 1 1973, 69, 199. Moruzzi, J. L.; Phelps, A. V. J. Chem. Phys. 1966, 45, 4617. (33) Klots, C. E.; Compton, R. N. J. Chem. Phys. 1977, 67, 1779. (34) Hilsenrath, J.; et al. Tables of Thermodynamic and Transport Properties of Air, Argon, Carbon Dioxide, Carbon Monoxide, Hydrogen, Nitrogen, Oxygen, and Steam; Pergamon: New York, 1960; p 185. (35) Krebs, P. J. Phys. Chem. 1984, 88, 3702. Krebs, P.; Heintze, M. J. Chem. Phys. 1982, 76, 5484. (36) Warman, J. M.; Asmus, K. D.; Schuler, R. H. J. Phys. Chem. 1969, 73, 3, 931. (37) Jonah, C. D.; Miller, J. R.; Matheson, M. S. J. Phys. Chem. 1977, 81, 1618. (38) Warman, J. M. The Study of Fast Processes and Transient Species by Electron-Pulse Radiolysis; Baxendale, J. H., Busi, F., Eds.; Reidel: Amsterdam, 1982; p 433. (39) Boesch, S. E.; Grafton, A. K.; Wheeler, R. A. J. Phys. Chem. 1996, 100, 10085.