Journal of Fluorine Chemistry 145 (2013) 58–62

Intramolecular hydrogen bonds in fluorinated, methoxylated, or unsubstituted2-(anilino)-1,4-naphthoquinones. A theoretical study

Sarai Vega-Rodrıguez, Rogelio Jimenez-Catano *, Elisa Leyva, Silvia Elena Loredo-Carrillo

Facultad de Ciencias Quımicas, Universidad Autonoma de San Luis Potosı, Av. Manuel Nava No. 6, Zona Universitaria, C.P. 78210, San Luis Potosı, S.L.P., Mexico, Mexico

A R T I C L E I N F O

Article history:

Received 11 July 2012

Received in revised form 29 September 2012

Accepted 2 October 2012

Available online 17 October 2012

Keywords:

Anilino-1,4-naphthoquinone

DFT frequency

Intramolecular hydrogen bond

A B S T R A C T

Density functional calculations at the BP86/6-311G(d,p) level of theory were realized to analyze the

existence of intramolecular hydrogen bonds in a series of neutral 2-anilinonaphthoquinones. The ortho

position of the aniline ring (C2a) was unsubstituted in two of them, substituted by a fluorine atom in

another two compounds, and substituted by a methoxy group in an additional pair of compounds. The

characteristic features of hydrogen bond formation (elongation of O55C1, N–H, F–C2a and O–C2a bonds,

increase in the out-of-plane N–H bending frequency, decrease in the N–H stretching frequency) suggest

the formation of a regular (two-center) intramolecular hydrogen bond in 2-anilinonaphtoquinones,

involving nitrogen as donor and an oxygen in the naphtoquinone fragment as acceptor. The geometry

and frequency changes also suggest the formation of a second and weaker hydrogen bond, resembling a

bifurcated (three-center) intramolecular hydrogen bond, in fluorine- or methoxy-ortho-substituted 2-

anilinonaphtoquinones.

� 2012 Elsevier B.V. All rights reserved.

Contents lists available at SciVerse ScienceDirect

Journal of Fluorine Chemistry

jo ur n al h o mep ag e: www .e lsev ier . c om / loc ate / f luo r

1. Introduction

Naphthoquinones are biologically active organic compoundscontaining a quinone moiety in their structure which has beenassociated with a wide biological activity [1]. Naphthoquinone andits derivatives are used as anti-microbial, anti-inflammatory andanti-tumor medications [2,3]. The biological activity of quinoneshas been related to their redox properties and their capacity toaccept one or two electrons to form the corresponding radicalanion (Q��) and the dianion (Q2�) [4]. These intermediate speciesinteract with crucial cellular molecules such as oxygen, DNA, andproteins, modifying their biological activity [4,5]. Due to thebiological relevance of naphthoquinones, their electrochemicalbehavior has been widely studied [6–11]. From those studies it hasbeen demonstrated that the presence of intramolecular hydrogenbonds (O� � �H–N) is an important factor in the acceptance of thefirst electron on the electrochemical reduction of the quinonesystem [8]. In addition, quinones containing hydrogen bonds intheir structure are biologically more active than those lackingthem [9]. An example is provided by hydroxynaphthoquinones inwhich intramolecular hydrogen bonds (O� � �H–O) are formed [10].This type of hydrogen bond is also called regular or two-centerhydrogen bond, where the atom bonded covalently to hydrogen isacting as a proton donor and the other electronegative atom isacting as a proton acceptor [12].

* Corresponding author. Tel.: +52 444 826 2440x526; fax: +52 444 826 2371.

E-mail address: jimenezr@uaslp.mx (R. Jimenez-Catano).

0022-1139/$ – see front matter � 2012 Elsevier B.V. All rights reserved.

http://dx.doi.org/10.1016/j.jfluchem.2012.10.001

It has been proposed that some 2-(fluoroanilino)-1,4-naphtho-quinones form a three-center intramolecular hydrogen bond(O� � �H(N)� � �F) when an ortho-fluoro-substituent is present inthe aniline; this feature facilitates the reduction of the quinonesystem because the resonance hybrids stabilize the radical anion(Q��) [11]. This type of hydrogen bond is also called bifurcated orthree-center hydrogen bond; in this case nitrogen acts as a donorand fluorine and oxygen as acceptors [12]. Since the synthesis andcharacterization of some of the 2-(fluoroanilino)-1,4-naphthoqui-nones had not been reported by the time of that study, a structuralevaluation of these compounds, the main purpose of this work,would be essential to support the existence of intramolecularhydrogen bonds in such compounds.

Density functional theory (DFT) calculations have been used tostudy inter- and intramolecular hydrogen bonds in diversecompounds [13–21]. They show a good agreement amongcalculated molecular properties and experimental data in benzeneanalogs and other compounds containing hydrogen bonds [15–17,21]. In addition, studies where DFT methods have been used tocalculate the structure and vibration frequencies of severalquinones, along with interpretation of the IR spectra, show thatDFT methods can predict accurate vibration frequencies of the C55Cand C55O stretching modes of the quinone moiety [22,23].Furthermore, DFT methods have been used to evaluate thestability of systems with intermolecular hydrogen bonds, likequinone dimers and quinone-benzoic acid adducts [24,25]; theyare also used to calculate half wave potentials (E1/2) of a quinonefamily because of a good correlation between half wave potentialsand electron affinity [26].

S. Vega-Rodrıguez et al. / Journal of Fluorine Chemistry 145 (2013) 58–62 59

There are several features that reveal the presence of ahydrogen bond. In first place are geometry changes. If B1 and B2

are the atoms to which a donor (D) and an acceptor (A) arerespectively bonded, the distances H–D and A–B2 are expected toincrease if a hydrogen bond H� � �A is present; in contrast, the D–B1

distance is expected to decrease. Both effects are due to electrondensity redistributions [21]. In second place there are vibrationfrequency changes. A bond that stretches as a result of a hydrogenbond become slightly weaker and its stretching frequencydecrease; this is the case of the H–D and A–B2 stretchingfrequencies; in contrast, the D–B1 stretching frequency increase[19,21]. On the other hand, the H–D–B1 bending becomes harder inthe presence of a hydrogen bond, so the corresponding bendingfrequency moves to higher values.

For intermolecular hydrogen bonds, a quantitative assessmentof their strength can be obtained by calculating the energydifference between the hydrogen bonded complex and thefragments. However, this is not possible for intramolecularhydrogen bonds. For the latter, a variety of tools have beenemployed, such as Bader’s Quantum Theory of Atoms in Molecules[27], the electron localization function, and the Natural Bondstheory [28]. However, in this work we are interested in signals ofhydrogen bonds rather than evaluating the energy contribution ofthose bonds, so we use the geometry and vibration frequencychanges described above to provide theoretical evidence of theexistence of regular (two-center) and bifurcated (three-center)hydrogen bonds in several 2-(anilino)-1,4-naphthoquinones(Fig. 1, structures 3–8).

2. Computational details

We performed density functional theory (DFT) calculations toevaluate intramolecular hydrogen bond formation in 2-(anilino)-1,4-naphthoquinone and some of its fluorinated or methoxylatedderivatives. We studied, as a reference, 1,4-naphthoquinone, 1, and2-anilinonaphthalene, 2, which lack intramolecular hydrogenbonds. We also studied 2-anilino-1,4-naphthoquinone, 3; the

Fig. 1. Molecules theoretically investigated: 1,4-naphthoquinone, 1; 2-anilinon

naphthoquinone, 4; 2-(2,3-difluoro)-anilino-1,4-naphthoquinone, 5; 2-(3,5-difluoro)-a

(2,5-dimethoxy)-anilino-1,4-naphthoquinone, 8.

fluorinated derivatives 2-(2,5-difluoro)-anilino-1,4-naphthoqui-none, 4; 2-(2,3-difluoro)-anilino-1,4-naphthoquinone, 5; 2-(3,5-difluoro)-anilino-1,4-naphthoquinone, 6; the methoxylated deri-vatives 2-(2-methoxy)-anilino-1,4-naphthoquinone, 7; and 2-(2,5-dimethoxy)-anilino-1,4-naphthoquinone, 8, which show intramo-lecular hydrogen bonds as described below. Compounds 1–8 areshown in Fig. 1.

To choose among the several density functionals available andthe many sets of basis functions to accompany them, we musthave in mind that in this work we need a good estimation ofmolecular structure and vibration frequencies. One of the firstachievements of the electronic structure theory is an accurateprediction of molecular geometries, which can be reached even ata relatively simple level of theory; for example, HF/6-31G*,whereas a good calculation of vibration frequencies requires theinclusion of electron correlation at least as provided by MP2 orDFT methods with density gradients (GGA methods) [29].Although we are not focusing this work in energy calculations,a systematic comparison of calculated results versus experimen-tal data for a large sample of reference organic molecules wasreported by Delley [30]. Those results show that GGA functionalsand their hybrid counterparts perform within similar margins:13–19 kJ/mol mean average deviation respect to true energies.Then, in our experience using intermediate workstations, whenfinding similar results we choose the slightly faster calculationsperformed with DFT methods that exclude Hartree–Fock ex-change, such as the BP86 density functionals (Becke B88 exchangefunctional [31] and Perdew P86 correlation functional [32]). Forsimilar speedy considerations, while obtaining similar results, wechoose polarized Gaussian basis sets instead of correlationconsistent basis sets, such as the 6-311G(d,p) selected for thiswork [33].

The computational methodology consisted in geometry opti-mization and vibration frequency calculations for all molecules.The geometry of molecules 1–8 was optimized, without con-straints, using the BP86 density functionals along with a 6-311G(d,p) polarized basis set. Vibration frequency calculations

aphthalene, 2; 2-anilino-1,4-naphthoquinone, 3; 2-(2,5-difluoro)-anilino-1,4-

nilino-1,4-naphthoquinone, 6; 2-(2-methoxy)-anilino-1,4-naphthoquinone, 7; 2-

Table 1Interatomic distances (A) and angles (8) in molecules 1–8.a

Molecule 1 2 3 4 5 6 7 8

Distances

O55C1 1.233 1.236 1.235 1.235 1.236 1.235 1.235

N–H 1.017 1.029 1.030 1.030 1.029 1.030 1.030

N–C2 1.401 1.366 1.369 1.368 1.369 1.365 1.365

F–C2ab 1.365 1.357

F–C3a 1.352 1.356

F–C5a 1.359 1.356

O–C2a 1.372 1.376

O–C5a 1.373

(N)H� � �O(C1)c 2.026 2.033 2.027 2.021 2.053 2.057

(N)H� � �F(C2a) 2.281 2.293

(N)H� � �O(C2a) 2.168 2.164

N� � �O(C1) 2.604 2.600 2.599 2.600 2.609 2.610

N� � �F(C2a) 2.681 2.689

N� � �O(C2a) 2.611 2.614

Angles

H–N–C2 115.3 110.7 111.4 111.2 110.7 111.9 112.0

H–N–C1a 115.3 117.1 116.5 116.7 117.3 115.5 115.3

N–H� � �O(C1) 112.9 112.0 112.3 112.9 111.2 115.1

N–H� � �F(C2a) 101.3 101.1

N–H� � �O(C2a) 103.7 104.2

a Calculated with BP86/6-311G(d,p).b Cna indicates the carbon atom on the ring of aniline where the substituent is bonded.c Atoms in parentheses are not included in the definition of distances or angles.

S. Vega-Rodrıguez et al. / Journal of Fluorine Chemistry 145 (2013) 58–6260

were carried out at the same level of theory. All calculations wereperformed with the Gaussian03 suite of programs [34]. Formolecules 1 and 3–7, calculated vibration frequencies werecompared with experimental values obtained from infraredspectroscopy (IR).

3. Results and discussion

3.1. Geometrical parameters

All molecules were optimized in C1 symmetry. Table 1 showsthe main interatomic distances and angles related to the possibleformation of intramolecular hydrogen bonds in molecules 3–8,additionally, a schematic representation of these molecules isshown in Fig. 2. The O55C1, N–H, F–C2a and O–C2a distances areexpected to increase if an atom from each pair is involved in ahydrogen bond; in contrast, the N–C2 distance is expected todecrease [12]. In molecules 3–8 the O55C1 and N–H bond lengthsare longer than O55C1 in 1 or N–H in 2 which lack intramolecularhydrogen bonds. Similarly, in 4 and 5 the F–C2a bond lengths(fluorine on carbon 2 of the aniline ring) are slightly longer than theF–C5a bond length in 4 or the F–C3a bond length in 5 where fluorine

Fig. 2. Schematic definition of interatomic d

cannot form hydrogen bonds; in contrast, the F–C3a and F–C5a bondlengths in 6 are identical and they cannot form intramolecularhydrogen bonds. The same phenomenon is observed in 8, wherethe O–C2a bond length (oxygen on carbon 2 of the aniline ring) isslightly longer than the O–C5a bond length where oxygen cannotform a hydrogen bond. The shortening of the N–C2 distanceproduced by a hydrogen bond is shown for 3–8, compared to thedistance observed for 2.

The H–N–C2 and H–N–C1a angles for 2–8 are also shown inTable 1. In 2, without intramolecular hydrogen bonds, those angleshave the same values. However, in 3 the presence of oxygen in C1

causes the H–N–C2 angle to decrease about 58 compared to 2,whereas the H–N–C1a angle increases about 28; this may beattributed to a hydrogen bond. In addition, the H� � �O distance(2.026 A) and the donor–H–acceptor angle (112.98) are very closeto similar values in hydrogen bonded five-member rings [35];similarly, the donor-acceptor (N� � �O) distance (2.604 A) fallswithin ranges that qualify the bond as strong if it were betweenidentical atoms but, being different, would be less strong [36]. Onthe other hand, 4 has a fluorine atom on C2a which pulls thehydrogen atom on nitrogen opening slightly the H–N–C2 angle,compared to 3, while closing slightly the H–N–C1a angle; this could

istances and angles reported in Table 1.

Table 2Dihedral angles (8) in molecules 2–8.a

Molecule 2 3 4 5 6 7 8

Distances

C1–C2–C1a–C2a �68.5 �36.6 �30.3 �30.4 �36.8 �29.0 28.2

C1–C2–N–H �21.5 �4.8 �6.0 �5.9 �5.1 6.5 6.5

H–N–C1a–C2ab �25.8 �21.9 �16.0 �16.4 �21.9 14.8 14.4

a Calculated with BP86/6-311G(d,p).b Cna indicates the carbon atom on the ring of aniline where the substituent is

bonded.

Table 3Comparison between calculated and experimental vibration frequencies (cm�1).a

Vibration mode Molecule Calculated

frequencies

Experimental

frequencies

N–H out-of-plane bending 2 367

3 715 690

4 738 719

5 723 698

6 692 665

7 724 �700b

8 743

C55O stretch 1 1651, 1663 1658

3 1623, 1655 1580, 1600

4 1627, 1659 1591, 1610

5 1626, 1659 1608w

6 1628, 1657 1573, 1600

7 1621, 1657 1594, 1615

8 1623, 1658

N–H stretch 2 3522

3 3355 3310

4 3347 3316

5 3345 3314

6 3354 3193

7 3347 3302

8 3349

a Only calculated frequencies for 2 and 8.b Mixed with the absorption at 719 cm�1 (out-of-plane C–H bending).

S. Vega-Rodrıguez et al. / Journal of Fluorine Chemistry 145 (2013) 58–62 61

be attributed to a N–H� � �F hydrogen bond although weaker thanthe N–H� � �O bond; the N� � �F distance (2.681 A) longer than theN� � �O distance (2.600 A) points to the same direction. Thecombined effect of oxygen and fluorine may be interpreted as abifurcated hydrogen bond, although it would be asymmetric alsoin strength. 5 has also fluorine on C2a and shows about the same H–N–C angles as 4. 6, however, does not have a fluorine atom closeenough to form a bifurcated hydrogen bond, so its H–N–C2 and H–N–C1a angles are about the same as in 3. Compounds 7 and 8 have amethoxy group on C2a where oxygen may join the oxygen atom onC1 to form a bifurcated hydrogen bond; however, the H–N–C2 angleis still smaller than the H–N–C1a angle in each, showing a slightlyweaker N–H� � �O(C2a) hydrogen bond compared to the N–H� � �O(C1)bond; the distances (N)H� � �O(C1) (2.053, 2.057 A) smaller than(N)H� � �O(C2a) (2.168, 2.164 A) confirm the asymmetry in strength.

A possible explanation about the N–H� � �O(C1) bond strengthslightly bigger than the N–H� � �F(C2a) or N–H� � �O(C2a) bondstrength could be the effect of ‘‘resonance-assisted hydrogenbonding’’ (RAHB) [36], that apparently appears when the donorand acceptor atoms are connected by a sequence of conjugateddouble bonds or, equivalently, an extended and stable p molecularorbital. Since extended p orbitals require planar surfaces, theseappear in two main regions in molecules 2–8: the naphtoquinoneand benzene fragments. The angle between those planes may bedescribed by the dihedral angle C1–C2–C1a–C2a that is reported inthe first row of Table 2. In 2, without hydrogen bonds, the anglebetween the naphtoquinone and benzene planes is 688. Then, weexpect, and actually find, the most stable p molecular orbitalbelonging to the naphtoquinone fragment and another pmolecular orbital, slightly higher in energy, corresponding to thebenzene ring. The lone pair of electrons in nitrogen may interactstrongly with one of those extended p molecular orbitals if it islocated in a p orbital that mixes well in the p system and, as aconsequence, the hydrogen on nitrogen would be in the sameplane of the atoms forming the extended p molecular orbital. Thedihedral angles C1–C2–N–H and H–N–C1a–C2a, reported in Table 2,indicate the closeness of the hydrogen on nitrogen to thenaphtoquinone and benzene planes, respectively. In 2, the C1–C2–N–H and H–N–C1a–C2a angles are 218 and 258 respectively, sothe hydrogen is located in an intermediate position between theplanes indicating almost no preference of the lone pair on nitrogentoward one of the plane fragments. In 3 and 6, however, the anglebetween the naphtoquinone and benzene planes decreases to 378and the hydrogen is very close to the naphtoquinone plane asanother indication of the presence of a hydrogen bond. In 4, 5, 7and 8 the presence of a second, although weaker, hydrogen bondcloses the angle between the naphtoquinone and benzene planesto about 308. It is important to note that the hydrogen on nitrogenmoves closer to the naphtoquinone plane as a result of thehydrogen bond and not that the hydrogen bond is a consequence ofthe extended p system involving nitrogen. Then, in thesemolecules the RAHB is not a factor that strengthens that bond.

Trying to understand why the second hydrogen bond is weakerthan the first one, we think the explanation resides in smallbut unavoidable geometry differences: the O–C1 distances

(1.235–1.236 A) are always smaller than the F–C2a (1.357,1.365 A) or the O–C2a (1.372, 1.376 A) distances and the rigidityof the 2-anilinonaphtoquinones preserves those qualitativedifferences all the way to the (N)H� � �O(C1) distances (2.021–2.057 A) which are smaller than the (N)H� � �F(C2a) (2.281, 2.293 A)or (N)H� � �O(C2a) (2.164, 2.168 A) distances. Then, we are close tothe researchers who say the stability of intramolecular hydrogenbonds is due to interatomic distances fixed by the s skeletoninstead of a RAHB effect [37,38]; in our study, the rigidity of thesigma skeleton influences the strength of the intramolecularhydrogen bonds.

3.2. Vibration frequencies

All calculated vibration frequencies were real numbers so theoptimized geometries correspond to energy minima. The mainfeatures of the experimental IR spectra compare well withcalculated frequencies, as can be observed for molecules 1 and3–7 in Table 3.

In our systems containing a N–H bond, some characteristicfeatures of hydrogen bond formation are an increase in the N–Hbending (out of the C2–N–C1a plane) frequency and a decrease inthe N–H and C55O stretching frequencies [19,21]. These changescan be observed in some of the most representative frequencies formolecules 3–8:

(a) The out-of-plane N–H bending frequency between 360 and750 cm�1. The calculated frequencies for 3–8 are significantlyshifted to higher values than the N–H bending frequencycalculated for 2 where an intramolecular hydrogen bondcannot be present. Additionally, for molecules 4, 5, 7 and 8 thecalculated frequencies present a larger shift, which can be anindicator of the rigidity imposed by a second hydrogen bondcompared to 3 and 6; however, for experimental values thesedifferences are not clear enough, presumably because theexperimental spectra was realized in the solid state where theout-of-plane N–H bending would be more difficult than in thegas state.

(b) The C55O bond stretching with strong intensity bands from1570 to 1660 cm�1. In 1, with equivalent carbonyls and

S. Vega-Rodrıguez et al. / Journal of Fluorine Chemistry 145 (2013) 58–6262

without intramolecular hydrogen bonds, the calculated C55Obond stretching modes involve both carbonyls (1651 symmet-ric, 1663 cm�1 asymmetric). For 3–8, however, carbonyls havedifferent environment and the calculated frequency of thecarbonyl closer to the N–H bond is the higher value shown inTable 3. Those higher values are smaller than the correspond-ing frequency for 1 in both, calculated and experimental values,suggesting the formation of intramolecular hydrogen bonds in3–8 [39].

(c) The N–H bond stretching signals from 3190 to 3530 cm�1. In 2,with no intramolecular hydrogen bonds, the calculated N–Hbond stretching frequency (3522 cm�1) shift to lower frequen-cies in 3–8 (3345–3355 cm�1); this agrees with the elongationof the N–H bond observed in 3–8 suggesting the formation ofintramolecular hydrogen bonds in these.

4. Conclusion

The analysis of geometrical parameters and vibration frequen-cies in molecules 1–8, at the BP86/6-311G(d,p) level oftheory, strongly suggest the formation of two-center (N–H� � �O)intramolecular hydrogen bonds in molecules 3 and 6; andbifurcated, although asymmetric in strength, intramolecularhydrogen bonds in molecules 4, 5, 7 and 8. For key changes invibration frequencies that accompany the formation of hydrogenbonds, the experimental frequency changes resemble the calcu-lated changes.

5. Experimental IR

IR spectra were recorded on a Thermofisher Nicolet iS10 FTIRspectrophotometer using diamond attenuated total reflectancesystem (ATR). The compounds 3–7 were synthesized using apreviously reported methodology [11].

Acknowledgments

We thank the Consejo Nacional de Ciencia y Tecnologia ofMexico for financial support (Grant 155678) and one of us (SVR) forher PhD scholarship (no. 33999).

Appendix A. Supplementary data

Supplementary data associated with this article can be found, in the

online version, at http://dx.doi.org/10.1016/j.jfluchem.2012.10.001.

References

[1] R.A. Morton, Biochemistry of Quinones, Academic Press, New York, 1965.[2] L.I. Lopez, E. Leyva, R.F. Garcıa de la Cruz, Revista Mexicana de Ciencias Farm-

aceuticas 42 (2011) 6–17.[3] K.C.G. De Moura, F.S. Emery, C. Neves-Pinto, M.F.R. Pinto, A.P. Dantas, K. Salomao,

S.L. de Castro, A.V. Pinto, Journal of the Brazilian Chemical Society 12 (2001) 325–338.

[4] G.E.W. Wolstenholm, C.M. O’Conner, Quinones in Electron Transport, Churchill,London, 1961.

[5] L.F.C. Medina, V. Stefani, A. Brandelli, Letters in Applied Microbiology 42 (2006)381–385.

[6] P.W. Crawford, E. Carlos, J.C. Ellegood, C.C. Cheng, Q. Dong, D.F. Liu, Y.L. Luo,Electrochimica Acta 41 (1996) 2399–2403.

[7] J. Tonholo, L.R. Freitas, F.C. de Abreu, D.C. Azevedo, C.L. Zani, A.B. de Oliveira, M.O.F.Goulart, Journal of the Brazilian Chemical Society 9 (1998) 163–169.

[8] N. Macıas-Ruvalcaba, G. Cuevas, I. Gonzales, M. Aguilar-Martınez, Journal ofOrganic Chemistry 67 (2002) 3673–3681.

[9] J. Johnson Inbaraj, R. Gandhidasan, R. Murugesan, Free Radical Biology andMedicine 26 (1999) 1072–1078.

[10] M.O.F. Goulart, C.L. Zani, J. Tonholo, L.R. Freitas, F.C. de Abreau, A.B. Oliveira, D.S.Raslan, S. Starling, E. Chiari, Bioorganic and Medicinal Chemistry Letters 7 (1997)2043–2048.

[11] E. Leyva, L.I. Lopez, S.E. Loredo-Carrillo, M. Rodrıguez-Kessler, A. Montes-Rojas,Journal of Fluorine Chemistry 132 (2011) 94–101.

[12] R.D. Parra, J. Ohlssen, Journal of Physical Chemistry A 112 (2008) 3492–3498.[13] J. Ireta, J. Neugebauer, M. Scheffler, Journal of Physical Chemistry A 108 (2004)

5692–5698.[14] A. Nowroozi, H. Roohi, H. Hajiabadi, H. Raissi, E. Khalilinia, M. Najafi Birgan,

Computational and Theoretical Chemistry 963 (2011) 517–524.[15] G. Chung, O. Kwon, Y. Kwon, Journal of Physical Chemistry A 101 (1997) 4628–

4632.[16] H.G. Korth, M.I. de Heer, P. Mulder, Journal of Physical Chemistry A 106 (2002)

8779–8789.[17] A. Simperler, W. Mikenda, Monatshefte fur Chemie 128 (1997) 969–980.[18] R.D. Parra, M. Furukawa, B. Gong, X.C. Zeng, Journal of Chemical Physics 115

(2001) 6030–6035.[19] R.D. Parra, B. Gong, X.C. Zeng, Journal of Chemical Physics 115 (2001) 6036–6041.[20] R.D. Parra, S. Bulusu, X.C. Zeng, Journal of Chemical Physics 118 (2003) 3499–

3509.[21] A. Kovacs, A. Szabo, I. Hargittai, Accounts of Chemical Research 35 (2002) 887–

894.[22] M. Nonella, Photosynthesis Research 55 (1998) 253–259.[23] D.V. Berdyshev, V.P. Glazunov, A.Ya. Yakubovskaya, T.Yu. Kochergina, V.F. Anu-

friev, Journal of Applied Spectroscopy 73 (2006) 798–806.[24] J. van de Bovenkamp, J.M. Matxain, F.B. van Duijneveldt, T. Steiner, Journal of

Physical Chemistry A 103 (1999) 2784–2792.[25] J. Garza, R. Vargas, M. Gomez, I. Gonzalez, F.J. Gonzalez, Journal of Physical

Chemistry A 107 (2003) 11161–11168.[26] C. Frontana, A. Vazquez-Mayagoitia, J. Garza, R. Vargas, I. Gonzalez, Journal of

Physical Chemistry A 110 (2006) 9411–9419.[27] R.F.W. Bader, Atoms in Molecules: A Quantum Theory, Oxford University Press,

Oxford, 1990.[28] F. Fuster, S.J. Grabowsky, Journal of Physical Chemistry A 115 (2011) 10078–

10086, and references therein.[29] M. Head-Gordon, Journal of Physical Chemistry 100 (1996) 13213–13225.[30] B. Delley, Journal of Physical Chemistry A 110 (2006) 13632–13639.[31] D. Becke, Physical Review A: Atomic, Molecular, and Optical Physics 38 (1988)

3098–3100.[32] J.P. Perdew, Physical Review B: Condensed Matter and Materials Physics 33

(1986) 8822–8824.[33] R. Krishnan, M.J. Frisch, J.A. Pople, Journal of Chemical Physics 72 (1980) 4244–

4245.[34] M.J. Frisch, G.W. Trucks, H.B. Schlegel, G.E. Scuseria, M.A. Robb, J.R. Cheeseman,

J.A. Montgomery Jr., T. Vreven, K.N. Kudin, J.C. Burant, J.M. Millam, S.S. Iyengar, J.Tomasi, V. Barone, B. Mennucci, M. Cossi, G. Scalmani, N. Rega, G.A. Petersson, H.Nakatsuji, M. Hada, M. Ehara, K. Toyota, R. Fukuda, J. Hasegawa, M. Ishida, T.Nakajima, Y. Honda, O. Kitao, H. Nakai, M. Klene, X. Li, J.E. Knox, H.P. Hratchian,J.B. Cross, C. Adamo, J. Jaramillo, R. Gomperts, R.E. Stratmann, O. Yazyev, A.J.Austin, R. Cammi, C. Pomelli, J.W. Ochterski, P.Y. Ayala, K. Morokuma, G.A. Voth,P. Salvador, J.J. Dannenberg, V.G. Zakrzewski, S. Dapprich, A.D. Daniels, M.C.Strain, O. Farkas, D.K. Malick, A.D. Rabuck, K. Raghavachari, J.B. Foresman, J.V.Ortiz, Q. Cui, A.G. Baboul, S. Clifford, J. Cioslowski, B.B. Stefanov, G. Liu, A.Liashenko, P. Piskorz, I. Komaromi, R.L. Martin, D.J. Fox, T. Keith, M.A. Al-Laham,C.Y. Peng, A. Nanayakkara, M. Challacombe, P.M.W. Gill, B. Johnson, W. Chen,M.W. Wong, C. Gonzalez, J.A. Pople, Gaussian 03, Revision B.04, Gaussian, Inc.,Pittsburgh, PA, 2003 .

[35] V.F. Sidorkin, E.P. Doronina, N.N. Chipanina, T.N. Aksamentova, B.A. Shainyan,Journal of Physical Chemistry A 112 (2008) 6227–6234.

[36] P. Gilli, V. Bertolasi, V. Ferretti, G. Gilli, Journal of the American Chemical Society116 (1994) 909–915.

[37] I. Alkorta, J. Elguero, O. Mo, M. Yanez, J.E. Del Bene, Molecular Physics 102 (2004)2563–2574.

[38] P. Sanz, O. Mo, M. Yanez, J. Elguero, ChemPhysChem 8 (2007) 1950–1958.[39] D.L. Pavia, G.M. Lampman, G.S. Kriz, Introduction to Spectroscopy, third ed., J.

Vondeling and E. Barrosse Publishers, Harcourt Inc., Orlando, 2001.