IR spectra of phenol+–(O2)n cation clusters (n = 1–4): Hydrogen bonding versus stacking interactions

Chemical Physics Letters 457 (2008) 298–302

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IR spectra of phenol+–(O2)n cation clusters (n = 1–4): Hydrogen bonding versus stacking interactions Alexander Patzer, Harald Knorke, Judith Langer, Otto Dopfer * Institut für Optik und Atomare Physik, Technische Universität Berlin, Hardenbergstrasse 36, 10623 Berlin, Germany

a r t i c l e

i n f o

Article history: Received 26 February 2008 In final form 28 March 2008 Available online 13 April 2008

a b s t r a c t IR photodissociation spectra of phenol–(O2)n cations, Ph+–(O2)n, are recorded in the O–H stretch range (mOH). Two isomers of Ph+–O2 are identified by their characteristic DmOH shifts upon complexation, namely the H-bonded global minimum featuring a nearly linear O–H–O hydrogen bond (DmOH = 86 cm1) and the less stable p-bonded local minimum (DmOH = +8 cm1). The diradical character of O2 has little influence on the Ph+–O2 interaction. IR spectra of larger Ph+–(O2)n clusters with n = 2–4 display a single mOH band with small incremental blue shifts from mOH of H-bonded Ph+–O2 (DmOH < 12 cm1), indicating that these clusters have a single H-bond and (n1) p-bonds. Quantum chemical calculations support the interpretation of the experimental data. Ó 2008 Elsevier B.V. All rights reserved.

1. Introduction Intermolecular forces involving aromatic molecules are vital for chemical and biological recognition phenomena [1]. Complexes of phenol with ligands, Ph–L, offer the interesting possibility to investigate two fundamental competing types of intermolecular interactions, namely hydrogen bonding and stacking (Fig. 1). The ligands L can interact either with the p electrons of the aromatic ring (p-bond, stacking) or with the acidic proton of the OH group (H-bond). Isolated Ph–Ln clusters have thus served as basic model systems to thoroughly investigate both theoretically and experimentally the interplay between both forces as a function of the solvent type (L), the degree of solvation (n), and the degree of electronic excitation or ionization of Ph. For example, neutral Ph–L complexes with nonpolar ligands, such as rare gas (Rg) atoms or CH4 prefer stacking [2–9], because dispersion forces between the ligand and the highly polarizable p electrons dominate the attraction. Ionization of these dimers causes, however, a switch of the preferred intermolecular interaction motif from p-bonding to H-bonding, as induction interactions between the ligand and the high positive partial charge of the OH proton override the dispersion forces [5,6,9–15]. On the other hand, Ph(+)–L dimers with quadrupolar and dipolar ligands (such as N2, CO2, C2H2, C6H6, CO, H2O) prefer H-bonds in both the neutral and the ionic ground electronic state due to dominant electrostatic forces [8–10,12,14,16–23]. The present work reports infrared photodissociation (IRPD) spectra of Ph+–(O2)n clusters (n = 1–4). This cluster system has been chosen for the following reasons. (1) To the best of our knowledge, no theoretical and experimental information is available

* Corresponding author. Fax: +49 30 31 423018. E-mail address: [email protected] (O. Dopfer). 0009-2614/$ - see front matter Ó 2008 Elsevier B.V. All rights reserved. doi:10.1016/j.cplett.2008.03.100

about the interaction of Ph(+) with O2. (2) The Ph+–O2 complex is the first Ph+–L dimer with an open-shell ligand ð3 R g Þ to be characterized and reveals the possible influence of the diradical character of L on the Ph+–L intermolecular potential. (3) The interaction and electronic excitation of aromatic molecules with O2 gives rise to well-known charge-transfer phenomena [24] leading to interesting photochemical processes [25,26]. The complex and broad chargetransfer excitation of complexes composed of aromatic molecules and O2 [24,27–29] may possibly be the reason for the lack of any size-selected spectroscopic studies of neutral and ionic Ph–(O2)n clusters, as broad transitions may prevent the application of mass-selective resonant photoionization techniques frequently applied to clusters with aromatic chromophores [4,16,17,30]. This obstacle is circumvented by the present approach, because Ph+–(O2)n clusters are generated in an electron impact ionization source [14,31] rather than by resonant photoionization of the neutral precursor. 2. Experimental and theoretical techniques IRPD spectra of Ph+–(O2)n clusters (n = 1-4) are recorded in a tandem quadrupole mass spectrometer (QMS1/2) coupled to an electron ionization (EI) source and an octopole ion trap [14,31]. Cold Ph+–(O2)n clusters are generated in a pulsed supersonic plasma expansion of phenol (T = 350 K) seeded in 6 bar O2. Cluster production is initiated by electron or chemical ionization of Ph, followed by three-body aggregation reactions [10]. A typical mass spectrum of the EI source reveals the presence of the ðO2 Þþ n and Ph+–(O2)n cluster series (Fig. 2a). The complexation efficiency (i.e., the Ph+–O2/Ph+ ratio) of the order of 1% confirms the generation of weakly-bonded Ph+–(O2)n clusters. The desired Ph+–(O2)n cluster is mass-selected by QMS1 and irradiated in the octopole

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Table 1 Selected properties of H-bonded Ph+–L dimers (L = Ar, O2, N2) calculated at the B3LYP/ 6-311G(2df,2pd) level compared to available experimental data (in brackets)

rOH/Å mOH/cm1 IOH/km mol1 De/cm1 D0/cm1 RHL/Å hOHL ms/cm1 a

Fig. 1. Structures of the H-bonded and p-bonded isomers of Ph+–O2. The planar H-bonded geometry (global minimum) corresponds to the equilibrium structure evaluated at the B3LPY/6-311G(2df,2pd) level. As the p-bonded isomer has not been studied computationally, the exact location and orientation of the O2 ligand is uncertain.

with the output of a single-mode optical parametric oscillator laser (mIR) leading to resonant photodissociation þ

þ

Ph  ðO2 Þn þ hmIR ! Ph  ðO2 Þm þ ðn  mÞO2

ð1Þ

+

The resulting Ph –(O2)m fragment ions are mass-selected with QMS2 and monitored as a function of mIR to obtain the IRPD spectrum of Ph+–(O2)n. The IR frequency is calibrated against atmospheric water absorptions to an accuracy of <1 cm1. For larger clusters several fragment channels (m) may occur. As an example, Fig. 2 shows mass spectra obtained by mass-selecting Ph+–(O2)4 with QMS1 and scanning QMS2 without (b) and with (c) resonant IR excitation (mIR = 3468 cm1). Spectrum (b) reveals weak signals of m = 0 and 1 fragment ions arising from metastable decay processes. Mass spectrum (c) displays additional signal in the m = 0 and m = 1 channels due to laser-induced dissociation. For Ph+–(O2)n with n 6 3, laser excitation exclusively leads to the evaporation of all ligands (m = 0), whereas for n = 4 the laser-induced fragments occurred in the m = 0 and 1 fragment channels with roughly equal intensity, providing useful experimental information about the dissociation energies of the ligands (vide infra). As IRPD spectra recorded in competing fragment channels are usually similar [10,32], only the dominant fragment channel is considered here.

+ Ph -(O ) 2n

(a)

m=0 m=1

(c) x50

(O ) 2n

0

50

100 m/u

150

+

200

(b) 80 100 120 140 m/u

Fig. 2. (a) Mass spectrum of the EI ion source for an expansion of Ph (T = 350 K) in 6 bar O2. (b) Mass spectrum obtained by mass-selecting Ph+–(O2)4 with QMS1 and scanning QMS2 with laser off. The Ph+–(O2)m fragment peaks with m = 0 and 1 arise from metastable decay. (c) Mass spectrum obtained by mass-selecting Ph+–(O2)4 with QMS1 and scanning QMS2 with the laser tuned to the mOH resonance of Ph+– (O2)4 at mIR = 3468 cm1. Additional m = 0 and 1 fragments are detected owing to laser-induced dissociation.

b c

Ph+

Ph+–Ar

Ph+–O2

Ph+–N2

0.9711 3534 (3534)a 262

0.9752 3447 (3464)b 771 533 455 2.3825 167.2 69

0.9772 3407 (3448) 979 1031 855 1.9832 174.6 67

0.9803 3351 (3365)b 1246 1745 1471 (1640 ± 10)c 1.9758 167.1 109 (113)b

Ref. [19]. Ref. [10]. Ref. [8].

Quantum chemical calculations are carried out for Ph+ and H-bonded Ph+–L dimers (L = Ar, O2, N2) at the unrestricted B3LYP/6-311G(2df,2pd) level (Table 1). Previous work has shown that this theoretical level describes the intermolecular interaction and resulting perturbation of the OH group in H-bonded Ph+–L dimers with satisfactory accuracy [15]. This statement is confirmed by the favourable comparison of theoretical with available experimental data on interaction energies and vibrational frequencies (Table 1). Stationary points on the potential corrected for basis set superposition error are obtained by relaxing all coordinates. Binding energies are corrected for harmonic zero-point energies. Harmonic vibrational frequencies are scaled by a factor of 0.9508 to bring mOH of Ph+ into agreement with its experimental value (3534 cm1 [19]). As the B3LYP level does not account for dispersion forces, the properties of the p-bonded Ph+–L dimers are not calculated. On the other hand, MP2 calculations suffer from spin contamination. However, as p-bonding has little influence on the OH group of Ph+, the properties of the O–H bond of p-bonded Ph+–L can be approximated by those of bare Ph+ [12,13]. 3. Results and discussion First the salient quantum chemical results are discussed (Table 1). Complexes of Ph+ (S = 1/2) and O2 (S = 1) may give rise to a doublet (S = 1/2) or a quartet (S = 3/2) electronic ground state. As the calculated interaction energy of H-bonded Ph+–O2 with S = 3/2 is slightly larger than that with S = 1/2 (by a few cm1), only dimers with S = 3/2 are considered in the present work. Apparently, the spin interactions have negligible influence on the interaction energy and the properties of the OH group. A single planar H-bonded Ph+–O2 isomer is identified (Fig. 1a), which is characterized by a nearly linear H-bond between the phenolic OH proton and the high electron density of the highest occupied molecular orbital of O2 (1pg), leading to a trans-configuration (De = 1031 cm1, D0 = 855 cm1, RHO = 1.98 Å, hOHO = 175°, hOOH = 130°, ms = 67 cm1). A corresponding cis-isomer could not be identified as a local minimum. The effects of complexation on the properties of the OH group are typical for H-bonding, with the major effects being an elongation of the O–H bond (DrOH = 0.0061 Å), a reduction of the O–H stretch frequency (DmOH = 127 cm1), and an increase of its IR intensity (DIOH = 274%). The magnitude of the perturbation of the OH group in the series L = Ar < O2 < N2 correlates with the strengths of the H-bonds (D0 = 455 < 855 < 1471 cm1). In general, the properties of Ph+–O2 are closer to those of Ph+-Ar than to those of Ph+–N2. H-bonding to the CH groups of Ph+ have also not been considered, as they are less stable than both H-bonding to the OH group and stacking [9]. Fig. 3 compares the IRPD spectra of Ph+–(O2)n with previous IRPD spectra of Ph+–Ar and Ph+–N2 [10]. The observed band centers are listed in Table 2, along with their widths and assignments. The

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Ph+

H

Ph+-N2

H

π

Ph+-Ar

Ph+-(O2)4 Ph+-(O2)3

Ph+-(O2)2

π

H 3300

3400

3500

Ph+-O2 3600

3700

νIR / cm–1 Fig. 3. IRPD spectra of Ph+–(O2)n with n = 1–4 are compared to IRPD spectra of Ph+Ar and Ph+–N2 [10]. All spectra have been recorded in the Ph+ fragment channel. Narrow dips are due to atmospheric water absorptions along the IR laser path. The asterisk indicates the mOH + ms transition of H-bonded Ph+–N2. The arrow marks the position of mOH of bare Ph+.

Table 2 Band centers, widths (fwhm, in parentheses) and assignments of the transitions observed in the IRPD spectra of Ph+–(O2)n n

m/cm1

Assignment

1 1 2 3 4

3542(8) 3448(14) 3460(10) 3465(6) 3468(10)

mOH(p) mOH(H) mOH(H/p) mOH(H/2p) mOH(H/3p)

IRPD spectrum of Ph+–O2 is dominated by an intense blue-shaded peak at 3448 cm1, which is readily assigned to mOH of the H-bonded isomer (trans-orientation) by comparison with the calculated value (3407 cm1). The observed red shift agrees qualitatively with the predicted one, DmOH = 86 versus 127 cm1, although the theoretical approach appears to slightly overestimate the frequency shift. The large width (14 cm1) and blue-shaded band contour are typical for proton donor stretch vibrations, supporting the given assignment [12,14,33]. There is no spectroscopic sign for the existence of a less stable H-bonded Ph+–O2 dimer with cis-orientation (in agreement with the calculations). The less intense and more symmetric band at 3542 cm1 exhibits a small blue shift (+8 cm1) from mOH of bare Ph+, consistent with an assignment to the p-bonded isomer of Ph+–O2 [12,14]. Using the ratios of the integrated band intensities (10) and the calculated IR oscillator strengths (3.7) of mOH of the H-bonded and p-bonded isomers of Ph+–O2, their relative abundance in the plasma expansion can be estimated to be of the order of 70 and 30% under the present experimental conditions, respectively. This result indicates that the H-bonded isomer is the global minimum, whereas the p-bonded dimer corresponds to a less stable local minimum. Previous studies demonstrated that the EI ion source predominantly produces the most stable isomer of a given cluster ion [10,14,34,35]. Significantly, the detection of p-bonded Ph+–O2 corresponds to the first observation a p-bonded Ph+–L dimer involving a quadrupolar or dipolar ligand.

The IRPD spectra of Ph+–(O2)n with n = 2–4 in Fig. 3 display a single transition with small and decreasing incremental blue shifts of 12, 5, and 3 cm1 with respect to mOH of H-bonded Ph+–O2 (Fig. 4a). These transitions are assigned to Ph+–(O2)n structures featuring a single H-bonded and (n1) p-bonded ligands. This cluster growth sequence is similar to that observed recently for Ph+–(N2)n and Ph+–Arn complexes (Fig. 4a) [10], as well as other related complexes of acidic aromatic ions with inert ligands, A+–Ln, including A = 1-naphthole, indole, aniline, imidazole, and cyclopropenyl and their protonated ions [14,34–41]. In general, the spectral positions of Ph+–(O2)n are much closer to those of Ph+–Arn than to those of Ph+–(N2)n (Fig. 4a). The small incremental blue shifts for n = 2–4 indicate that further solvation of the H-bonded Ph+–O2 dimer core leads to a slight stabilization of the intramolecular O–H bond, which is accompanied by a slight weakening of the adjacent intermolecular H-bond. Such noncooperative three-body effects are typical for interior ion solvation by neutral solvent molecules [14]. The lack of any absorption in the vicinity of mOH of bare Ph+ (3540 cm1) in the Ph+–(O2)2 spectrum demonstrates that the population of trimer isomers with two p-bonded ligands in the molecular beam expansion are below the detection limit, again confirming that H-bonding is favoured over the stacking motif. Moreover, the disappearance of the transition near 3540 cm1 in the Ph+–(O2)2 trimer spectrum validates that this band in the Ph+–O2 spectrum is indeed due to mOH of the p-bonded isomer and not arising from a combination band of mOH of H-bonded Ph+–O2 with an intermolecular mode. Such transitions have previously been detected in the IRPD spectra of Ph+–(N2)n (Fig. 3) [10]. The photofragmentation branching ratios can be utilized to roughly estimate ligand binding energies within the framework of a simple model outlined in detail elsewhere [14]. The core assumption of this model is that the absorbed photon energy (hmIR) is consumed to evaporate the most weakly bound ligands. As the branching ratio for m = 0 and m = 1 fragment ions of Ph+–(O2)4 is roughly unity (Fig. 2b), the absorbed photon energy of 3450 cm1 is estimated to be of the order of the sum of the dissociation energies of all four ligands, i.e., D0(H) + 3 D0(p). Assuming further that D0(H) is of the order of 1000 cm1 (Table 1), D0(p) can be approximated as 800 cm1, consistent with stacking interactions of related complexes with N2 [10,14,42]. Moreover, this value is compatible with an earlier measurement of the C6 Hþ 6  O2 binding energy, D0 = 1190 ± 105 cm1 obtained by photoionization spectroscopy [43]. The long-range part of the attractive potential between Ph+ and O2 is dominated by charge-induced dipole and charge-quadrupole interaction. For a positive charge and a negative quadrupole moment H (Table 3), the anisotropy of both contributions favours a linear approach of O2 over a T-shaped orientation [33,44,45]. However, the calculations predict a bent intermolecular orientation (hOOH = 130°) in the vicinity of the potential minimum of H-bonded Ph+–O2 arising from the interaction of the positive partial charge on the OH proton with the high electron density region of the partially filled highest occupied molecular orbital of O2 with 1pg symmetry. Such a structure is indeed also expected in the limit of the formation of a chemical bond as evidenced by the bent structure of O2H+ in O2H+–Rg dimers [46]. However, the weak intermolecular interaction of O2 with Ph+, as evidenced by the modest DmOH shift and the small calculated binding energy, does not indicate the formation of an incipient chemical bond. Such an inert behaviour is in fact predicted for complexes of diatomic molecules with 3 R g electronic states on the basis of ab initio calculations [45,47]. Similar to O2H+, OCOH+ is a nonlinear ion [48] and, as a consequence, Ph+–OCO also features a nonlinear H-bond (trans-configuration) via the interaction of the OH proton with one of the lone pairs of OCO [20]. Unfortunately, as theoretical calculations for p-bonded Ph+–O2 are not available, the accurate location and orientation of

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3550

3600 Ne

3550 Ph+-Arn

3450

He νOH / cm-1

νOH / cm-1

3500

Ph+-(O2)n Ph+-(N2)n

3400

3500 Ar O2

3450 3400

Kr

3350 3350

0

1

2

3

4

5

6

7

3300

CH4 N2

CO CO2

200

300

400

500

600

PAL / kJ mol-1

n

Fig. 4. (a) Plot of the mOH frequency of Ph+–Ln (L = Ar [10], O2, N2 [10]) as a function of n. (b) Plot of the mOH frequency of selected H-bonded Ph+–L dimers [10,15,20] as a function of the proton affinity of L. The line corresponds to a fit of the data to a linear polynomial without using the O2 data.

Table 3 Parallel polarizabilities (ak), quadrupole moments (H), and proton affinities (PA) of selected ligands L [45] L

Ar

O2

N2

CO2

ak/1025 cm3 H/1040 Cm2 PA/kJ mol1

16.42 0 198.8

23.5 1.33 421

23.8 5.00 493.8

40 14.01 540.5

the O2 ligand remains uncertain for this isomer, although a (nearly) linear approach toward the aromatic plane appears reasonable (Fig. 1) [45]. It is illustrative to compare the interaction strength in H-bonded Ph+–L dimers by considering the physical and chemical properties of L, such as their polarizabilities, a, quadrupole moments, H, and proton affinities, PA (Table 3). Inspection of Table 3 suggests that the interaction of Ph+ with L should increase in the order Ar < O2 < N2 < CO2. The bonding with O2 should be slightly stronger than with Ar on the basis of somewhat larger induction forces (a) and the small additional contribution arising form charge-quadrupole interaction (H). However, the difference is expected to be small, and indeed the Ph+–(O2)n spectra closely resembles those of the corresponding Ph+–Arn complexes (Figs. 3 and 4a), suggesting similar potential energy surfaces for both cluster systems. In particular, for both dimers, the H-bonded and p-bonded isomers are observed (Fig. 3). Changing to L = N2 significantly enhances the electrostatic forces (via larger H) and thus increases the difference between H-bonding and stacking forces. As a consequence, the spectrum of Ph+–N2 exhibits only features of the H-bonded isomer (Fig. 3). In the case of L = CO2, the larger a and H values lead to an even stronger H-bond in Ph+–CO2 [20]. Comparison of the calculated binding energy of H-bonded Ph+–O2 (De/0 = 12.3/10.2 kJ mol1) with experimental values for neutral O2 dimers of benzene (6.7 ± 1.3 kJ mol1), hexafluorobenzene (8.8 ± 1.7 kJ mol1), fluorobenzene (8.3 kJ mol1), and toluene (1.8 kJ mol1) [29,43], which probably all have a stacking geometry [49], illustrates the increase in interaction strength upon formation of the ionic O–H–O H-bond in aromatic molecule complexes with O2. Previous work has demonstrated that the red shift in H-bonded Ph+–L dimers, DmOH, is correlated with PA of L [10,20,21,33,50]. The mOH transition of Ph+–O2 derived from the IR spectrum, 3448 cm1, is slightly larger than the frequency of 3415 cm1 predicted from related complexes (Fig. 4b), using a PA value of 421 kJ mol1 for O2. This small discrepancy is attributed to the error limits of the

simple DmOH–PA correlation model, although a minor contribution arising from the diradical character of O2 cannot be completely ruled out. 4. Conclusions The IRPD spectra and quantum chemical calculations of Ph+–O2 provide the first impression of the interaction between the phenol cation and molecular oxygen. The IR spectrum demonstrates that the H-bonded isomer is more stable than the p-bonded one. Although no information is presently available for neutral Ph–O2, the close similarity between the Ph+–O2 and Ph+–Ar interaction suggests that stacking might be competitive with H-bonding in the neutral complex, if not more favourable. Such stacking contacts are predicted for O2 complexes with benzene [49]. The H-bond in Ph+–O2 is relatively weak (10 kJ mol1), in line with the low proton affinity and quadrupole moment of O2. There is no indication of incipient chemical bond formation arising from the diradical character of O2. IRPD spectra of larger Ph+–(O2)n clusters with n = 2–4 are consistent with a cluster growth sequence, in which p-bonded ligands are attached to a H-bonded dimer core. Acknowledgements This study was supported by the Deutsche Forschungsgemeinschaft (DO 729/2) and the Fonds der Chemischen Industrie. References [1] E.A. Meyer, R.K. Castellano, F. Diederich, Angew. Chem. Int. Ed. 42 (2003) 1210. [2] N. Gonohe, H. Abe, N. Mikami, M. Ito, J. Phys. Chem. 89 (1985) 3642. [3] A. Fujii, T. Sawamura, S. Tanabe, T. Ebata, N. Mikami, Chem. Phys. Lett. 225 (1994) 104. [4] X. Zhang, J.L. Knee, Faraday Discuss. 97 (1994) 299. [5] S. Ishiuchi et al., Angew. Chem. Int. Ed. 44 (2005) 6149. [6] S. Ishiuchi et al., J. Chem. Phys. 127 (2007) 114307. [7] S. Ishiuchi, Y. Tsuchida, O. Dopfer, K. Müller-Dethlefs, M. Fujii, J. Phys. Chem. A 111 (2007) 7569. [8] C.E.H. Dessent, K. Müller-Dethlefs, Chem. Rev. 100 (2000) 3999. [9] M.A. Vincent, I.H. Hillier, C.A. Morgado, N.A. Burton, X. Shan, J. Chem. Phys. 128 (2008) 044313. [10] N. Solcà, O. Dopfer, J. Phys. Chem. A 105 (2001) 5637. [11] N. Solcà, O. Dopfer, J. Mol. Struct. 563/564 (2001) 241. [12] N. Solcà, O. Dopfer, Chem. Phys. Lett. 325 (2000) 354. [13] N. Solcà, O. Dopfer, Chem. Phys. Lett. 369 (2003) 68. [14] O. Dopfer, Z. Phys. Chem. 219 (2005) 125. [15] A. Takeda, H.S. Andrei, M. Miyazaki, S. Ishiuchi, M. Sakai, M. Fujii, O. Dopfer, Chem. Phys. Lett. 443 (2007) 227. [16] K. Müller-Dethlefs, O. Dopfer, T.G. Wright, Chem. Rev. 94 (1994) 1845. [17] T. Ebata, A. Fujii, N. Mikami, Int. Rev. Phys. Chem. 17 (1998) 331.

302

A. Patzer et al. / Chemical Physics Letters 457 (2008) 298–302

[18] K. Kleinermanns, C. Janzen, D. Spangenberg, M. Gerhards, J. Phys. Chem. A 103 (1999) 5232. [19] A. Fujii, M. Miyazaki, T. Ebata, N. Mikami, J. Chem. Phys. 110 (1999) 11125. [20] A. Fujii, T. Ebata, N. Mikami, J. Phys. Chem. A 106 (2002) 10124. [21] A. Fujii, T. Ebata, N. Mikami, J. Phys. Chem. A 106 (2002) 8554. [22] O. Dopfer, G. Reiser, K. Müller-Dethlefs, E.W. Schlag, S.D. Colson, J. Chem. Phys. 101 (1994) 974. [23] O. Dopfer, K. Müller-Dethlefs, J. Chem. Phys. 101 (1994) 8508. [24] H. Tsubomura, R.S. Mulliken, J. Am. Chem. Soc. 82 (1960) 5966. [25] H. Frei, Science 313 (2006) 309. [26] M.J. Paterson, O. Christiansen, F. Jensen, P.R. Ogilby, Photochem. Photobiol. 82 (2006) 1136. [27] J.B. Birks, E. Pantos, T.D. Hamilton, Chem. Phys. Lett. 20 (1973) 544. [28] E.A. Gooding, K.R. Serak, P.R. Ogilby, J. Phys. Chem. 95 (1991) 7868. [29] J.J. Casero, J.A. Joens, J. Phys. Chem. A 103 (1999) 7136. [30] B. Brutschy, Chem. Rev. 100 (2000) 3891. [31] O. Dopfer, Int. Rev. Phys. Chem. 22 (2003) 437. [32] O. Dopfer, R.V. Olkhov, J.P. Maier, J. Chem. Phys. 111 (1999) 10754. [33] E.J. Bieske, O. Dopfer, Chem. Rev. 100 (2000) 3963.

[34] [35] [36] [37] [38] [39] [40] [41] [42] [43] [44] [45] [46] [47] [48] [49] [50]

N. Solcà, O. Dopfer, Eur. Phys. J. D 20 (2002) 469. N. Solcà, O. Dopfer, Phys. Chem. Chem. Phys. 6 (2004) 2732. N. Solcà, O. Dopfer, J. Phys. Chem. A 106 (2002) 7261. H.S. Andrei, N. Solcà, O. Dopfer, J. Phys. Chem. A 109 (2005) 3598. H.S. Andrei, N. Solcà, O. Dopfer, Phys. Chem. Chem. Phys. 6 (2004) 3801. O. Dopfer, D. Roth, J.P. Maier, Int. J. Mass Spectrom. 218 (2002) 281. N. Solcà, O. Dopfer, J. Chem. Phys. 120 (2004) 10470. F. Pasker, N. Solcà, O. Dopfer, J. Phys. Chem. A 110 (2006) 12793. N. Solcà, O. Dopfer, Chem. Eur. J. 9 (2003) 3154. J.R. Grover, G. Hagenow, E.A. Walters, J. Chem. Phys. 97 (1992) 628. R.V. Olkhov, S.A. Nizkorodov, O. Dopfer, J. Chem. Phys. 107 (1997) 8229. D. Roth, O. Dopfer, Phys. Chem. Chem. Phys. 4 (2002) 4855. S.A. Nizkorodov, D. Roth, R.V. Olkhov, J.P. Maier, O. Dopfer, Chem. Phys. Lett. 278 (1997) 26. G. Chalasinski, J. Klos, M. Cybulski, M. Szczesniak, Collect. Czech. Chem. Commun. 63 (1998) 1473. O. Dopfer, R.V. Olkhov, D. Roth, J.P. Maier, Chem. Phys. Lett. 296 (1998) 585. G. Granucci, M. Persico, Chem. Phys. Lett. 205 (1993) 331. R.V. Olkhov, O. Dopfer, Chem. Phys. Lett. 314 (1999) 215.