Cleavage of hydrogen by activation at a single non-metal centre - towards new hydrogen storage materials.

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Cite this: Phys. Chem. Chem. Phys., 2015, 17, 13539

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Cleavage of hydrogen by activation at a single non-metal centre – towards new hydrogen storage materials Sławomir J. Grabowskiab Molecular surfaces of non-metal species are often characterized by both positive and negative regions of electrostatic potential (EP) at a non-metal centre. This centre may activate molecular hydrogen which further leads to the addition reaction. The positive EP regions at the non-metal centres correspond to s-holes; the latter sites are enhanced by electronegative substituents. This is why the following simple moieties; PFH2, SFH, AsFH2, SeFH, BrF3, PF(CH3)2 and AsF(CH3)2, were chosen here to analyze the H2 activation and its

Received 14th January 2015, Accepted 20th April 2015

subsequent splitting at the P, As, S, Se and Br centres. Also the reverse H–H bond reforming process is anal-

DOI: 10.1039/c5cp00219b

processes. The sulphur centre in the SFH moiety is analyzed in detail since the potential barrier height for the

yzed. MP2/aug-cc-pVTZ calculations were performed for systems corresponding to different stages of these addition reaction for this species is the lowest of the moieties analyzed here. The results of calculations show

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that the SFH + H2 - SFH3 reaction in the gas phase is endothermic but it is exothermic in polar solvents.

1. Introduction The role of molecular hydrogen in chemical reactions and biological processes is often a subject of analysis. One can mention the addition of H2 to organic compounds1,2 or reactions of molecular hydrogen in microorganisms.3 On the other hand hydrogen as an alternative to traditional sources of energy is also a subject of intensive investigations.4,5 In spite of its routine use in industry the cost of hydrogen storage and transportation is still too high. This is why topics connected with the properties of H2 are also discussed in numerous analyses.6 There are various experimental7,8 and theoretical6 studies where the activation of H2 is analyzed. Reactions of dihydrogen with transition metal complexes are well known and often investigated with related aspects of hydrogenation catalysis.7,8 Special attention has been paid in recent studies to small molecules, among them molecular hydrogen, activated by frustrated Lewis pairs.9,10 The Lewis base and Lewis acid centres existing in the same or in two different molecular species usually interact to form an adduct; in the case of a frustrated Lewis pair they can not combine because of steric hindrance. For example, the reactions of trityl borates with Lewis bases, such as amines, pyridines and phosphines were examined.11 It was found that classical Lewis acid–Lewis base adducts are a

Faculty of Chemistry, University of the Basque Country and Donostia, International Physics Center (DIPC), P.K. 1072, 20080 Donostia, Spain b IKERBASQUE, Basque Foundation for Science, 48011 Bilbao, Spain. E-mail: [email protected]

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formed in most cases; however the reactions of sterically encumbered PR3 phosphines (R = i-Pr, Cy, t-Bu) proceed in a different way, not directly through the Lewis acid–Lewis base centres contact.11 It is worth mentioning that such frustrated Lewis pair combinations, like phosphine–borane, may lead to the activation of dihydrogen. The mechanisms of the molecular hydrogen activation and its splitting are the subject of numerous studies. Two ways to split dihydrogen are commonly known; the homolytic cleavage and the heterolytic cleavage.12,13 In both cases the dihydrogen activation proceeds at the transition metal centres, this process for non-metal centres being rather rare. In the first case of the dihydrogen activation, the binding of the H2 molecule to the metal centre is observed which is stabilized by the orbital– orbital interaction between the filled d-orbital of the metal and the antibonding s* orbital of molecular hydrogen. The increase of antibonding orbital occupation results in weakening and elongation of the H–H s bond. In extreme cases the breaking of this bond is observed followed by the formation of two metal– hydrogen s-bonds. In the case of heterolytic cleavage, the H2 molecule is polarized due to the interaction with the positive metal site, further leading to the dihydrogen splitting, the binding of hydridic hydrogen with the metal centre and migration of the protic hydrogen to the other reaction site. Other mechanisms for specific systems were analyzed; for example the electrophilic activation mechanism was discussed for the phosphine–borane species14 where the electrophilic boron centre combined with the nucleophilic phosphorus site leads to the polarization of dihydrogen followed by its splitting.

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The splitting of ammonia and dihydrogen by a single carbon centre activation was analyzed13 where the singlet carbenes mimic transition metal centres since they possess a lone electron pair (Lewis base site) as well as an accessible vacant orbital (Lewis acid site). Recently two reactivity models have been proposed for the heterolytic cleavage of dihydrogen by frustrated Lewis pairs.15 One can see that the dihydrogen activation followed by its splitting may also be classified as a single centre or two centre process. In the former case there is the H2 activation at the transition metal or at the singlet carbene mentioned earlier here, while the activations by the frustrated Lewis pairs are mostly classified as the latter process. The aim of this study is to analyze if non-metal centres may be utilized in the processes of the activation of dihydrogen. Such centres often possess a dual character since they may act simultaneously as a Lewis base and as a Lewis acid, similarly to the singlet carbene centre mentioned earlier here.13 For example, the elements of Groups IV–VII are commonly known as electronegative centres playing the role of the Lewis bases in numerous reactions and interactions. However it was explained that they possess the electron deficient regions (s-holes) on the extensions of the covalent bonds to those centres.16 Such electron deficient regions are often characterized by the positive electrostatic potential (EP) thus they may act also as Lewis acids. The halogen bond17 being the interaction between the Lewis base centre and the halogen playing the role of the Lewis acid is classified as such an interaction (s-hole bond), since the halogen atom is often characterized by the positive EP in the elongation of the covalent bond to the halogen.16 On the other hand the halogen atoms may play the role of the Lewis base centres due to the lone electron pairs forming the belt of the negative EP around the halogen atom. The Lewis acid properties of other non-metal atoms may be explained in a similar way as the properties of halogen atoms; chalcogen,18 pnicogen19 and tetrel bonds20,21 are observed in a case of VI, V and IV Group elements, respectively. Thus there are cases where the same non-metal centre may act as the Lewis base or as the Lewis acid in the process of the activation of dihydrogen. The molecular hydrogen also possesses such dual character since it is characterized by the regions of positive and negative electrostatic potential attributed to the edges of the molecule and to the H–H s-bond, respectively.22 The reactions of simple non-metal species with H2 are analyzed here. The SFH moiety is analyzed in detail for the reasons described later here; the sulphur centre in this molecule is characterized as having negative and positive EP regions since it possesses two lone electron pairs as well as the depletion of the electron charge density at the S-centre on the extension of F–S and H–S bonds. The influence of solvents on the activation and splitting processes of H2 is also analyzed as well as the reverse process – the reforming of the H–H bond.

2. Computational methods The calculations were carried out with the Gaussian09 set of codes23 using the second-order Møller–Plesset perturbation

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theory (MP2),24 and the aug-cc-pVTZ basis set.25 Frequency calculations were performed at the same computational level to confirm that the obtained structures correspond to energetic minima or to transition states at the reaction paths analyzed. The self-consistent reaction field (SCRF) method using the polarizable continuum model (PCM) and the integral equation formalism variant (IEFPCM) is applied to include solvent effects.26 The following solvents are taken into account in this study: cyclohexane, benzene, ethanol and water. For the initial complexes of Z-species (PFH2, SFH, AsFH2, SeFH, BrF3, PF(CH3)2 and AsF(CH3)2; Z = P, S, As, Se and Br) with molecular hydrogen, the binding energy, Ebin, is calculated as the difference between the energy of the complex and the sum of energies of monomers in their energetic minima. The Quantum Theory of ‘Atoms in Molecules’ (QTAIM)27 was also applied to analyze critical points (BCPs) in terms of the electron density (rBCP), its Laplacian (r2rBCP) and the total electron energy density at BCP (HBCP). The regions of the concentration and depletion of the electron density are also analyzed by the QTAIM approach. The QTAIM calculations were performed with the use of the AIMAll program.28 The Natural Bond Orbital (NBO) method29,30 was applied to analyze electron charge shifts being the result of complexation, orbital occupancies as well as orbital–orbital interactions. The latter interactions are calculated as the second-order perturbation theory energies.29,30 The NBO calculations were carried out at same level as the previous geometry optimizations were performed (MP2/aug-cc-pVTZ). For the Natural Bond Orbital (NBO) calculations the NBO 6.0 program31 was used.

3. Results and discussions The r-holes as centres of dihydrogen activation The activation of dihydrogen at the Z-centres (Z = P, S, As, Se or Br) of the PFH2, SFH, AsFH2, SeFH, PF(CH3)2, AsF(CH3)2 and BrF3 species is analyzed. The fluorine moieties were chosen here since the electron withdrawing F-substituent enhances the Z-centre s-hole which results in an increase of its positive electrostatic potential (EP). The s-hole concept was firstly applied to explain the Lewis acid properties of halogen atoms (designated usually as X) in the C–X covalent bonds.16 This concept is based on simple analyses of the valence electronic structures and hybridizations in the species considered. For example, for the CY3X molecule (X – halogen, Y – another substituent) the electronic structure of the halogen may be represented as ns2npx2npy2npz1.32 The half-filled halogen npz orbital located on the C–X line interacts with a carbon orbital; this leads to the localization of the npz electron in the C–X bond region and the depletion of the electron charge density in the region corresponding to the outer, non-involved in such interaction, lobe of the npz orbital. This depletion region situated at the X-centre in the elongation of C–X bond is the s-hole and if it is sufficient enough then it is characterized by positive EP. This is why the linear C–X B halogen bonds (B designates the Lewis base centre) are mainly formed.17 The s-hole in the CY3X


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molecule is enhanced by the electron withdrawing Y-substituents, also a greater electron charge density depletion is observed for the more polarizable halogen (X).16 That is why the positive EP increases in the following order F o Cl o Br o I; in the same order the increase of the strength of the C–X B halogen bond is observed for the same B Lewis base centre. Two remaining halogen p-orbitals, perpendicular to the C–X bond, cause the ‘‘belt’’ of the negative EP around this bond. That is why the X-centre may interact with the Lewis acid species in the direction perpendicular to C–X bond, or nearly so. It was confirmed that covalently bonded atoms of Groups IV–VI can also interact through the positive regions of EP (s-holes) with the Lewis base centers.33 For example, for the SFH species analyzed here the sulphur centre is characterized by six valence electrons; four of them belong to two lone electron pairs being responsible for the existence of the negative regions of EP, and the remaining two electrons are involved in the formation of S–H and S–F bonds which should lead to the depletion of the electron density at the S-centre on the extensions of these bonds. The latter regions (s-holes) may be characterized by positive EP, the region situated in the elongation of the F–S bond should be ‘‘more positive’’ than the region in the elongation of the H–S bond since there is a greater electron density shift to the fluorine than to the hydrogen. That is why halogen (Group VII), chalcogen (Group VI) and pnicogen (Group V) centres characterized by s-holes as well as by lone electron pairs often have dual Lewis acid–Lewis base character, thus they may interact with nucleophiles and electrophiles, respectively. A different situation is observed for the tetrel (Group IV) centres (especially for sp3 hybridized ones) where the regions of negative EP do not result from the existence of lone electron pairs. The tetrel centres were analyzed recently in the simple ZH4, ZFH3 and ZF4 species (Z = C, Si and Ge).21 For example, for the CFH3 molecule the maximum positive EP is localized at the C-centre on the extension of F–C bond (s-hole), while the minimum EP (negative) is situated at the fluorine. However for the CH4 and SiH4 species the maximum positive EPs are located at the H-atoms while the minimum negative EPs at the four face regions of the tetrel centre. The Lewis acid and Lewis base properties of methane were analyzed earlier.34 It was found for the latter species that the C–H bonds may play a role of the proton donors while the C-centre may act as the proton acceptor in hydrogen bonds.34 It was mentioned earlier here that the dual Lewis acid–base character is also observed for dihydrogen.22 Fig. 1 presents the electrostatic potential map for the SFH molecule as an example. One can see that the sulphur atom involves the region of the positive EP (s-hole) in the elongation of the F–S bond; this is the maximum EP amounting +0.066 au. The s-hole is also observed on the extension of the H–S bond but the positive electrostatic potential is much lower here than in the former case since it is equal to +0.022 au. A minimum EP of 0.023 au is observed for the fluorine centre. One can also observe the local EP minima of 0.015 au corresponding to the lone electron pairs of the S-atom; two lone pairs are situated above and below the plane of the molecule (Fig. 1, the orange


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Fig. 1 Computed electrostatic potential on the 0.001 au electron density molecular surface of SFH (left); blue colour corresponds to the maximum positive electrostatic potential and red to the negative minimum; the molecular graph of SFH (right) with isolines of the laplacian of electron density; positive values are depicted in solid lines and negative values in broken lines (the laplacian isodensity lines in the plane of the molecule); the negative laplacian region above the sulphur atom corresponds to lone electron pairs; MP2/aug-cc-pVTZ results presented here and in other figures and tables.

colour around the S-centre corresponds to these local minima). Fig. 1 also presents the isolines of the laplacian of electron density for SFH molecule. According to the QTAIM approach the negative values of laplacian show the regions of concentration of the electron density; this is often used to localize lone electron pairs. On the other hand a negative value of laplacian at a bond critical point (BCP) of a bond path between two interacting centres indicates a covalent character of such interaction.27 For the SFH moiety analyzed here, the negative values of laplacian at the BCPs of the S–F and S–H bond paths indicate the covalent character of these interactions. This is obvious since these bond paths correspond to the S–F and S–H covalent bonds. Besides there are regions of negative laplacian close to the sulphur atom which correspond to the lone electron pairs situated symmetrically below and above the plane of the molecule (see Fig. 1). As it was mentioned above, Fig. 1 shows the positive EP regions (s-holes) indicated by blue and green colors in the elongation of F–S and H–S bonds, respectively (see EP map in the figure). This is confirmed partly by the laplacian map presented in the same figure, since approximately in the elongation of the F–S bond the area of the positive laplacian is observed. The region of the negative laplacian corresponding to lone electron pairs becomes thinner on the extension of the H–S bond thus the electron density is outweighed here by the influence of the nucleus; the latter results in the positive EP. Table 1 presents the EP maxima and the local minima corresponding to the F–Z s-holes and to the lone electron pairs regions, respectively, for the Z-species analyzed here. The global minima of EP for these Z species are attributed to the fluorine substituents. The existence of the EP maxima and minima for those species may be explained in a similar way as it was performed above for the SFH moiety. This is interesting if the regions of positive EP described above, corresponding to the s-holes, are strong enough Lewis acid centers to activate the H2 molecule and to induce its splitting. Since for the initial stage of the dihydrogen splitting the complexes linked by the Z H2 interaction are created, these interaction characteristics are also included in Table 1. Fig. 2 presents molecular graphs of the SeFH H2 and


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Table 1 The characteristics of the H2 complexes with the species specified in the first left column; Ebin – the binding energy and ENBO energies in kcal mol1; E1NBO corresponds to the s(H2) - s*(Z–F) overlapping while E2NBO to n(Z) - s*(H2); ET is the charge of the H2 molecule in the complex; EPmax (at s-holes) and EPmin (local minima at lone electron pairs) values at the 0.001 au molecular surface are presented for monomers specified in the first left column and not involved in interactions with H2

Species

EPmax

EPmin

ET

Ebin

E1NBO

E2NBO

PFH2 SFH AsFH2 SeFH BrF3 PF(CH3)2 AsF(CH3)2

0.0602 0.0657 0.0683 0.0750 0.0810 0.0369 0.0487

0.0194 0.0145 0.0107 0.0141 0.0278 0.0330 0.0218

0.001 0.003 0.003 0.006 0.001 0.006 0.005

1.2 1.4 1.6 2.1 1.3 0.7 0.9

1.4 2.3 2.3 4.6 0.5 0.1 0.2

0.9 1.2 0.8 2.2 0.3 1.1 1.1

Fig. 2 Molecular graphs of SeFH H2 and PFH2 H2 complexes; big circles correspond to attractors, small ones to bond critical points, the solid and broken lines designate bond paths.

PFH2 H2 complexes where one can see the bond paths linking the Z-centre (Se or P) with the bond critical point of the H2 molecule. In other words these interactions may be classified as Z s ones where s-electrons of dihydrogen play the role of the Lewis base. For the SFH H2, AsFH2 H2 and BrF3 H2 complexes, similar Z s interactions are observed. The interactions with dihydrogen acting as the Lewis base have been analyzed before in several studies. For example, the A–H s hydrogen bonds were analyzed35,36 and compared with the dihydrogen bonds;37 also other types of interactions with the s-electrons of molecular hydrogen playing the role of the Lewis base were investigated.38 This kind of interaction was analyzed recently for numerous complexes with the pnicogen centre acting as the Lewis acid,22 its interaction with the molecular hydrogen leading to the weak binding since the binding energy, Ebin, does not exceed 2 kcal mol1. The strongest interaction was observed for the AsFH2–H2 complex where Ebin is equal to 1.6 kcal mol1, however the processes connected with the splitting of H2 were not analyzed there. The QTAIM observations described above are supported by the Natural Bond Orbitals (NBO) results since for most of complexes presented in Table 1 the s(H2) - s(Z–F)* overlap is the most important orbital–orbital interaction. It means that the Z-species play the role of the Lewis acid in the complexes formed. However there is also the n(Z) - s(H2)* interaction which for two complexes, the PF(CH3)2 H2 and AsF(CH3)2 H2, is dominant, meaning that for the latter complexes the Lewis acid properties of dihydrogen are observed. The molecular graphs for the latter species show the Z H bond paths, i.e. the Z-centre is not connected by the bond path with the BCP of dihydrogen but with the H-atom attractor (see Fig. 3 where one of these complexes is presented – PF(CH3)2 H2).


Fig. 3 Molecular graph of PF(CH3)2 H2 complex; big circles correspond to attractors, small ones to bond critical points, the solid and broken lines designate bond paths.

Table 1 includes the charges of H2 molecules (ET-values calculated within the NBO approach) involved in the interactions analyzed. For two last complexes these values are negative, meaning that there is an electron charge shift from the Z-species, PF(CH3)2 or AsF(CH3)2, to the hydrogen molecule. Hence it is confirmed that the Lewis acid properties of the molecular hydrogen are stronger than the Lewis base ones in those complexes. For the remaining complexes the ET-values are positive meaning that the Lewis base properties for H2 are more distinct here. In general, one may state that for the complexes listed in Table 1 the molecular hydrogen acts as the Lewis base and as the Lewis acid. For some of the complexes the dihydrogen reveals more its acidity and for others more basicity. It is worth mentioning that the properties of dihydrogen coincide with those of the Z-centre – if the stronger acidic character is observed for the dihydrogen, thus simultaneously the stronger basic character occurs for the Z-centre and vice versa. It seems that the basicity of dihydrogen is more important than its acidity since there is the correlation between the s(H2) - s(Z–F)* interaction energy and the binding energy (Fig. 4), while such correlation is not observed for the n(Z) s(H2)* interaction energy. The latter correlation (Fig. 4) shows that the processes connected with the electron charge density shifts (polarization) are important here. A correlation between the maximum positive EP localized at the Z-centre s-hole and the binding energy is observed (Fig. 5, the quadratic polynomial regression is denoted) which also shows the electrostatic nature of the Z H2 interaction. In both latter relationships the BrF3 H2 complex was omitted since it is out of the trends observed, i.e. it is characterized by the greatest EP at Z-centre (Br atom) and a relatively small value of Ebin (1.3 kcal mol1). The latter complex is out of the correlations described above since the bromine of the BrF3 moiety possesses different characteristics than those observed for the other Z-species. For the BrF3 molecule there is an extremely large shift of the electron charge to the F-substituents which results in the large positive charge of the Br-atom (+1.65 and +1.70 au for NBO and QTAIM approaches, respectively) and the great polarization of the Br–F bonds. Besides BrF3 H2 is the only complex linked


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Table 2 The potential barrier height, DE2/1, (the difference between energies of the transition state and reactants) and the difference between energies of products and reactants, D E3/1, are listed; corresponding enthalpy (H) and free Gibbs energy (G) values are included (all results in kcal mol1). The numbers 1, 2 and 3 assigned here for energy barriers and differences correspond to those assigned for structures presented in Fig. 6

Species

DE2/1

DH2/1

DG2/1

DE3/1

DH3/1

DG3/1

PFH2 SFH AsFH2 SeFH BrF3 PF(CH3)2 AsF(CH3)2

40.2 23.6 53.1 25.8 58.7 46.0 55.9

40.6 23.6 52.8 25.2 57.6 47.7 56.0

43.7 26.3 55.6 27.2 61.0 52.1 60.9

1.2 3.2 15.2 7.9 31.7 1.7 14.5

4.6 6.5 17.6 10.0 34.5 5.3 17.2

8.5 9.1 21.3 12.2 38.3 9.5 21.2

Fig. 4 The linear correlation between ENBO energy (in kcal mol1) corresponding to the s(H2) - s*(Z–F) overlap and the binding energy (kcal mol1).

Fig. 5 The relationship between the maximum electrostatic potential at the Z-centre (in au) and the binding energy (kcal mol1).

through the specific halogen bond where the multivalent Br-centre exists; the specific properties of halogen bonds with multivalent halogen centres were described recently.39 The other complexes presented here are linked through the chalcogen and pnicogen bonds.

The dihydrogen splitting at r-holes Table 2 presents results corresponding to the H2 activation at the Z-centre followed by breaking of the H–H bond and forming two new Z–H bonds. The reactions in the gas phase are analyzed. These processes, for the selected SFH + H2 reaction, are presented in Fig. 6. One can see that for the local minimum, E1, there is the interaction of the H2 molecule with SFH through the sulphur s-hole, then the breaking of the H–H bond and formation of two S–H bonds is observed (E2 transition state and E3 local minimum). The valency of sulphur increases from two (for E1) to four (for E3). The energy (E), enthalpy (H) and free Gibbs energy (G) for reaction stages designated by numbers 1, 2 and 3 and corresponding to the SFH + H2 local minimum, to the transition state and to the SFH3 local minimum (Fig. 6), respectively, are also presented for the other Z-species (Table 2 presents the potential barrier


Fig. 6 The reaction of the splitting of hydrogen for SFH + H2 reactants, energies of species at vertical axis are related to the global minimum of the H2S HF system (right side of the figure); the case of the gas phase.

heights and differences between energetic minima related to these values). The lowest potential barrier height (DE2/1) for the reaction of hydrogen splitting, amounting to B24 kcal mol1, is observed for the SFH species (Table 2). DE2/1 designates the difference between energies corresponding to reaction stages 2 and 1. Similar designations are applied later here for the other energy, enthalpy and free Gibbs energy differences. However the latter reaction for the sulphur centre is endothermic since DE3/1 amounts to B3 kcal mol1. For the PF(CH3)2 species even the lower value of DE3/1 is observed (B2 kcal mol1) but the potential barrier height, DE2/1, is much greater at B46 kcal mol1. The corresponding H and G values (Table 2) are in line with energies since the lowest potential barrier height (DH2/1 and DG2/1) and the relatively small DH3/1 and DG3/1 values are observed for the SFH + H2 system. This is why the dihydrogen activation at the sulphur centre is analyzed in further detail here.


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Fig. 6 shows that the SFH3 species being the product of the SFH + H2 reaction may be further decomposed, i.e. the decrease of the valency of S-atom with the breaking of the F–S bond is observed for the E4 transition state which further leads to the E5 global minimum of the H2S HF hydrogen bonded system. Table 3 shows that the potential barrier height for the latter reaction in the gas phase amounts B7 kcal mol1 (DE4/3) and the process is exothermic (DE5/3 E 53 kcal mol1). It is important that for the reverse process to reform the H–H bond the relatively low potential barrier height (DE2/3) of B20 kcal mol1 corresponding to the SFH3 - SFH + H2 reaction is observed. It was mentioned earlier here that different mechanisms for the hydrogen splitting were proposed. It has been pointed out in earlier studies that there are three ways by which this process occurs; homolytic cleavage (observed mostly for metals), heterolytic cleavage and electrophilic activation.12 It was discussed in the previous section that the H2 molecule interacting with the Z-centre acts simultaneously as the Lewis acid and as the Lewis base since two kinds of orbital–orbital interactions are detected for Z H2 contacts; however the Lewis base properties are more evident for the H2 molecule. This is why the mechanisms of reactions analyzed here are partly similar to the heterolytic cleavage where at the first stage of reaction the H2 molecule is linked with the electrophile via its s-bond electrons. Fig. 7 presents molecular graphs with the laplacian of electron density isolines for the SFH + H2 reaction stages corresponding to the local energetic minima (Fig. 6); the negative laplacian values indicate the regions of the concentration of the electron density. The laplacian of the electron density distributions presented in Fig. 7 may be interpreted in the following way. For the initial

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structure, E1, there are two lone electron pairs of the SFH molecule. For the E3 structure (SFH3 molecule) one lone electron pair is observed; the sulphur centre is tetravalent thus it may be classified as hypervalent (10 electrons in the sulphur valence shell). Note that for three S–H bonds of the SFH3 molecule the laplacian of electron density at the corresponding bond critical points (BCPs) is negative which is attributed to covalency, for the S–F bond the laplacian is positive but the total electron energy density at the B–F BCP is negative which is also treated as a satisfactory condition of covalency.40 The final H2S HF structure is linked through the F–H S hydrogen bond and it corresponds to the global energetic minimum; the octet rule is obeyed here for the sulphur atom which obviously is characterized by two lone electron pairs and two S–H s-bonds. The Natural Bond Orbitals (NBO) analysis29,30 confirms partly the latter observations, two lone electron pairs for the S-atom in SFH and H2S molecules and one electron pair for the SFH3 molecule. However for the latter structure there are three s-orbitals corresponding to S–H bonds with the occupancy very close to 2 and the S–F bond is not detected within the NBO approach, but the F anion interacting with the H3S+ cation. In other words E3 is the SFH3 species characterized by the hypervalent S-centre according to the AIM approach and it is the H3S+ F ion pair according to the NBO method. This means that the condition of the covalent character of the bond if the total electron energy density at the corresponding BCP is negative should be treated with caution. The more that the F-charge in the SFH3 moiety is close to 1 au for both approaches, the NBO charge is equal to 0.789 au while the QTAIM integrated charge amounts to 0.760 au. However both approaches show that there is the H3S+ F ion pair for the E4 transition state structure. The influence of solvent on the H2 splitting reactions

Fig. 7 The molecular graphs of E1, E3 and E5 structures (see Fig. 6) with isolines of the laplacian of electron density; positive values are depicted in solid lines, negative in broken ones, the plane projections presented, E1 – plane parallel to H2 molecule and containing S–H bond of SFH molecule; E3 – plane containing –SH2 fragment; E5 – plane perpendicular to H2S molecule and containing the F–H S link.


The splitting process at the sulphur centre in the gas phase is characterized by the relatively high potential barrier height (B24 kcal mol1 – Table 2). This is why solvent effects are also analyzed here with the use of the self-consistent reaction field (SCRF) theory.26 The values of energy (DE3/1) collected in Table 3 show that the SFH + H2 - SFH3 reaction is exothermic for all solvents considered while it is endothermic in the gas phase. This means that in the solvent the E3 energy for the SFH3 molecule is lower than the E1 energy of the SFH H2 complex. For this dihydrogen splitting reaction the potential barrier height (DE2/1) decreases in solvent if it is compared with the gas phase. Table 3 shows that the greater decreases are observed for the polar water and ethanol solvents (dielectric constants are equal to 80.1 and 24.5, respectively) than for the non-polar benzene and cyclohexane (dielectric constants are equal to 2.3 and 2.0, respectively). The SFH3 molecule is lower in energy than the SFH H2 reactants for polar solvents, by B10 and B9 kcal mol1 for water and ethanol, respectively (DE3/1 values in Table 3), while for benzene and cyclohexane only by about 1 kcal mol1. The corresponding enthalpies (H) and free Gibbs (G) energies for polar water and ethanol solvents confirm that the splitting reactions are exothermic but in a case


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Table 3 The potential barrier heights, DE2/1 and DE4/3 (the differences between energies of the transition state and reactants) and the difference between energies of products and reactants, DE3/1 and DE5/3, are listed; corresponding enthalpy (H) and free Gibbs energy (G) values are included (results in kcal mol1). The numbers 1, 2, 3, 4 and 5 assigned here for energy barriers and differences correspond to those assigned for structures presented in Fig. 6 and 8

Solvent

DE2/1

DH2/1

DG2/1

DE4/3

DH4/3

DG4/3

DE3/1

DH3/1

DG3/1

DE5/3

DH5/3

DG5/3

Cyclohexane Benzene Ethanol Water Gas phase

21.3 21.0 17.5 17.1 23.6

21.4 21.0 17.4 17.1 23.6

24.0 23.7 20.1 19.8 26.3

4.4 4.1 1.6 1.4 6.5

2.7 2.5 0.3 0.2 4.6

2.6 2.4 0.4 0.3 4.5

0.7 1.4 8.8 9.6 3.2

2.4 1.8 5.8 6.7 6.5

5.0 4.3 3.8 4.7 9.1

50.7 50.2 44.1 43.4 53.3

51.2 50.8 44.5 43.8 53.9

53.3 52.8 45.9 45.1 56.4

of non-polar benzene and cyclohexane they are endothermic. The potential barrier heights (DH2/1 and DG2/1) are still high for all solvents; 17–20 kcal mol1 for polar solvents and even more for non-polar ones. However the results show that for all solvents the hydrogenation process is more probable than for the gas phase and that for the reverse reforming of the H–H bond, SFH3 - SFH + H2, the potential barrier heights for the gas phase and for the solvents are characterized by similar values of B20 kcal mol1 (DE2/3, DH2/3 and DG2/3). The latter observation is very important for the designing of hydrogen storage. However there is a disadvantage for the SFH3 species for the polar solvents. There is the low potential barrier height, DE4/3, for the further process leading to the hydrogen bonded H2S HF system corresponding to the global minimum (Table 3). The DE4/3 energy for the latter process is equal to B1 and B2 kcal mol1 for water and ethanol, respectively. This means that in polar solvents this process proceeds very easily; note that the reforming of the H–H bond from the system being in the global minimum (E5) is much less probable than such reforming from the SFH3 molecule (from E3 structure). The

above-mentioned barriers related to enthalpies and free Gibbs energies are even lower and close to zero for polar solvents (see DH4/3 and DG4/3 for water and ethanol, Table 3). The scheme of the processes of the dihydrogen activation, its splitting and finally the formation of the H2S HF hydrogen bonded complex corresponding to the global energetic minimum is presented in Fig. 8 for the water solvent; very similar schemes may be constructed from results collected in Table 3 for the other solvents.

Conclusions The processes of molecular hydrogen activation and of its splitting were analyzed mostly for the metal centres and recently also for the boron centre. However it was shown in this study that such a process is also possible for non-metal halogen, chalcogen and pnicogen centres (Z-centres) characterized by the regions of the positive (s-holes) and negative (attributed to the lone electron pairs) electrostatic potential. This is why the complex mechanism of the dihydrogen activation is observed here since also the latter species is characterized by the dual character. It may act as the Lewis acid as well as the Lewis base, due to the existence of the regions of positive and negative EP, similarly to the Z-centres. The dihydrogen activation and its splitting were analyzed here in the gas phase for the halogen, chalcogen and pnicogen centres (Z-centres). It was found that the potential barrier height for the Z + H2 - ZH2 process exceeds 20 kcal mol1 and that all those reactions are endothermic in the gas phase. This is why the SFH + H2 - SFH3 reaction characterized by the lowest potential barrier height was also analyzed in different solvents. It was found that for polar solvents (water and ethanol) and nonpolar ones (benzene and cyclohexane) that the potential barrier height for the SFH + H2 - SFH3 reaction decreases if it is compared with the gas phase. The energy, enthalpy and free Gibbs energy values indicate that this reaction is exothermic in polar solvents. However the SFH3 moiety corresponds to the local energetic minimum and its further SFH3 - H2S + HF decomposition is possible, leading to the global energy minimum.

Acknowledgements Fig. 8 The reaction of the splitting of hydrogen for SFH + H2 reactants, energies of species at vertical axis are related to the global minimum of the H2S HF system (right side of the figure); water solvent.


Financial support came from Eusko Jaurlaritza (GIC IT-588-13) and the Spanish Office for Scientific Research (CTQ2012-38496-C0504). Technical and human support was provided by Informatikako Zerbitzu Orokora – Servicio General de Informatica de la Universidad


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del Pais Vasco (SGI/IZO-SGIker UPV/EHU), Ministerio de Ciencia e ń (MICINN), Gobierno Vasco Eusko Jaurlanitza (GV/EJ), Innovacio European Social Fund (ESF) is gratefully acknowledged.


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