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Covalency of hydrogen bonds in solids revisited† Cite this: Chem. Commun., 2014, 50, 11547
Volker L. Deringer,a Ulli Englert*a and Richard Dronskowski*ab
Published on 08 August 2014. Downloaded on 05/10/2014 17:42:47.
Received 20th June 2014, Accepted 7th August 2014 DOI: 10.1039/c4cc04716h www.rsc.org/chemcomm
The covalent nature of short hydrogen bonds has been under debate for long. Here we show that the crystal orbital Hamilton population (COHP) bonding indicator gives new, complementary evidence of covalent hydrogen acceptor interactions in the molecular solid state.
‘‘What is the covalency of hydrogen bonding?’’ asked Grabowski in a recent review article.1 The ubiquity of hydrogen bonds (HBs) both in crystal chemistry and molecular biology2 has stimulated much thought about the issue.1 Short HBs were long assumed to be covalent to some degree,3 and particular milestones were set by Gilli et al. with the introduction of the resonance-assisted hydrogen bond (RAHB).4 Nonetheless, there remain several interesting questions. For example, if short HBs are covalent and others are not, should one draw a line between both regimes, or is there a gradual transition?5 Furthermore, the H acceptor distance is an important criterion to identify HBs in solids, but the mere existence of short contacts does not suffice to explain the stability of crystalline networks.6,7 Reliable knowledge about the covalency of HBs is very much sought after, and in particular, one would like to have tools for analysing hydrogen-bonded crystal structures. There have been previous efforts to gauge the strength of HBs in solids—for example, from DFT energies,8 and there are also useful estimators such as the PIXEL method.9 These tools assign a numerical energy value to a given intermolecular contact or crystalline direction, and ‘‘bond energy’’ is a tangible quantity without doubt, even if no rigorous observable directly corresponds to it. On the other hand, topological analysis of the electronic charge density has proven to be valuable, and the latter may be derived from theory5 or from high-resolution diffraction experiments.10,11 It would now seem to be rewarding to combine these conceptual approaches: to determine the nature and strength of HBs in a solid a
Institute of Inorganic Chemistry, RWTH Aachen University, Landoltweg 1, 52056 Aachen, Germany. E-mail: [email protected], [email protected]; Fax: +49 241 8092288, +49 241 8092642 b ¨lich–Aachen Research Alliance (JARA-HPC), RWTH Aachen University, Ju 52056 Aachen, Germany † Electronic supplementary information (ESI) available: Details of theoretical computations. See DOI: 10.1039/c4cc04716h
This journal is © The Royal Society of Chemistry 2014
directly from its underlying electronic structure. In this work, we seek to explore HBs by invoking the familiar model of covalent bonding between atomic orbitals. To this end, we performed periodic DFT computations for molecular solids with pronounced HBs, and then dissected the electronic structure of the crystals into orbital-wise contributions to reveal the bonding nature between neighbouring atoms. The method we used has been dubbed ‘‘crystal orbital Hamilton population’’ (COHP) analysis and is described in ref. 12. In solid-state and materials chemistry and physics, COHP is an established tool which has been applied to a plethora of compounds, e.g., to tailor-made ferromagnets or datastorage alloys.13 Using COHP analysis to investigate organic crystals and HBs may seem unconventional in this light, but there is no fundamental reason for not doing so.14 We began our study with an intriguing compound that contains many different HBs in its crystal structure: from very short (O–H O type) contacts to those on the verge of being bonded at all. The compound, N,N-dimethylbiguanidinium bis(hydrogensquarate) 1, was previously characterised using high-resolution X-ray and neutron diffraction.11 We have now performed periodic electronic-structure computations15 on the neutron-derived crystal structure, in which we kept the structure frozen but re-optimised the positions of hydrogen atoms.16 From the plane-wave based DFT output, we then computed COHP curves in the framework of a recently developed projection scheme,17 and the result is hence denoted as ‘‘projected COHP’’. This new method makes it possible to treat even complex molecular structures as presented here. First, we examined the characteristic anion dimer in crystalline 1, as sketched in Fig. 1a. Two O–H O bonds are formed, one (highlighted in red) notably shorter than the other one (blue); this asymmetry is caused by the crystalline network.11 The electronic DOS (Fig. 1b), which is the solid state analogue of a molecularorbital scheme, reveals contributions of the acceptor oxygen over the entire valence region, from 12 eV upwards, and the local DOS of the hydrogen atom in the shortest HB is largest at 12 eV, too. Due to this very orbital overlap, a covalent bonding interaction between both atoms arises, as reflected in pCOHP values plotted to the right of the vertical axis. We reiterate that this technique reveals
Chem. Commun., 2014, 50, 11547--11549 | 11547
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Fig. 2
Integrated –pCOHP at eF for all relevant H O contacts in crystalline 1.
Fig. 1 (a) Structural drawing of 1;11 the hydrogensquarate anions dimerise in the solid state as shown here. The strongest (red) and second (blue) O–H O bonds have been highlighted, and the labelling (H1/H5) corresponds to the original article11 for convenience. (b) Computed densities of states (DOS, for hydrogen and acceptor atoms of the strongest HB in 1); projected crystal orbital Hamilton populations (pCOHP) for both above contacts, and their integrals (in eV).
covalent interactions (in terms of pairwise contributions to the electronic band-structure energy), but not electrostatic and dispersive terms which would govern weaker HBs.2 The higher-lying energy bands (from 8 eV on) also contain H acceptor interactions, but these are nonbonding in good approximation because stabilising and destabilising pCOHP regions cancel each other. When integrated up to the highest filled orbital (i.e., up to the Fermi energy eF), the (p)COHP provides a measure of the covalent bond strength,12 just like energy-integrating the DOS would give the number of electrons in the system. We have hence plotted the integrated pCOHP (ICOHP) for the two contacts on the right of Fig. 1b. Defining stabilising interactions to have positive numerical values, we obtain 149 and 104 kJ mol 1, respectively. This indicates significant covalent bonding, on the order of what is found for resonance-assisted hydrogen bonds (RAHBs), a well-studied family of strong HB contacts.4 We stress once more that the ICOHP does not equal the bond energy but provides a reasonable estimate.12 We then proceeded to investigate the entire range of HB interactions in 1. For each, the ICOHP was computed, and the results are plotted in Fig. 2 against the respective H O distance. The covalent contributions decrease for the longer HBs, and a clear trend is seen: several N–H O interactions in 1 have lower ICOHP values than the O–H O ones but still exhibit some covalent bonding. The unconventional C–H O HBs, on the other hand, have values close to zero, in line with the expectation that they are weak and noncovalent in nature.2 The difficulty of identifying single bonds at such distances (and thereby possibly overlooking important, but less localised electrostatic and dispersive attraction) has been discussed before.6 For validation, we compared the ICOHP of each bond to the respective charge density at the bond-critical point, rbcp, which counts among the traditionally established criteria for HB covalency, together with the Laplacian (cf. ref. 11 and references
11548 | Chem. Commun., 2014, 50, 11547--11549
Fig. 3 Scatter plot of integrated pCOHPs vs. experimentally determined electron densities at the respective bond-critical points. The latter data are taken from ref. 11 and 19a–d. For the very short HB in benzoylacetone,19a the hydrogen atoms were either frozen in their neutron-derived structure or re-optimised.
therein) or energy densities.18 Generally, covalent HBs show significant rbcp values in the range of 0.1–1 e Å 3, whereas weaker HBs, dominated by electrostatic forces, do not exhibit rbcp values far from zero. To rely on a larger dataset, we extended our study to a number of high-quality charge-density and neutron diffraction studies from the previous literature.19 Fig. 3 clearly evidences that both methods of estimating HB covalency do exhibit a linear correlation—this is not to be expected per se, because ICOHP and rbcp stem from fundamentally different approaches. It is also appreciable that the dataset for 1 (blue symbols in Fig. 3) is consistent with the additional data points (red); the latter pertain to different structures and measurements as detailed in the ESI.† This provides confidence for future use of the method. For very strong HBs (rbcp 4 0.6 e Å 3), the results are visibly influenced by the decision whether hydrogen positions are re-optimised with DFT or locked to the neutron-derived values. We illustrate this for the O–H O contact in benzoylacetone, for which both options are visible in Fig. 3 (circled data points): the optimised covalent O–H bond length is shorter compared to neutron data (cf. ref. 16), and hence the computed H O distance is larger. While the limitations of proton localisation become apparent in such extreme cases, we argue that there
This journal is © The Royal Society of Chemistry 2014
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We thank the Studienstiftung des deutschen Volkes for a ¨lich–Aachen Research Alliance scholarship to V.L.D. and the Ju for JARA-HPC computer time.
Notes and references
Published on 08 August 2014. Downloaded on 05/10/2014 17:42:47.
Fig. 4 (a) Fragment from the crystal structure of 1-(2-hydroxy-5nitrophenyl)ethanone, drawn with data from ref. 19b. (b) The fragment from the crystal structure of benzene,20 a prototypical molecular solid with ‘‘T-stacking’’; see text.
remains significant correlation between ICOHP and rbcp data regardless of how the H positions are treated. The new method appears to provide an advantage when looking at the intramolecular H O contacts which exhibit short distances but would not be classified as HBs. Fig. 4a shows a structural fragment from one of the previous studies we built upon:19b the H5 atom forms a C–H O bond with a neighbouring O2 atom; furthermore, it is close to the O3 atom of the nitro group, but the latter contact would not be seen as a hydrogen bond as its oCHO angle is close to 90 degrees. If the electron density in the H O bond-critical point is taken as the only output of the experimental charge-density analysis (the authors of the original work did not do that!), it might suggest a higher degree of covalency (rbcp = 0.08(1) e Å 3) in the intramolecular contact than in the intermolecular H bond (0.04(1) e Å 3). This clearly underlines that even a highly useful analysis technique must be used with (chemical) caution. The pCOHP data in Fig. 4a clearly differentiate between the two types of contacts, yielding a slightly stabilising interaction for the H bond but none for the intramolecular H5 O3 contact. We stress that this ‘‘nonbonding’’ intramolecular contact has been omitted from Fig. 3, for obvious reasons. Finally, an alert reviewer pointed out to us the C–H p interactions which, by their name, could suggest the involvement of orbital-pair interactions. Test computations for the crystal structure of benzene (Fig. 4b),20 however, gave no indications of covalent twocentre bonding (with ICOHP values very slightly below 0 eV), which is in accord with the quite long H C distances (optimised: 42.85 Å; cf. the trends in Fig. 2). For these families of interactions, different theoretical tools will be needed, and the argumentation of Dunitz and Gavezzotti6 holds. In conclusion, we have demonstrated an alternative, orbitalbased way of probing hydrogen-bond covalency in solids, which complements previous careful charge-density studies. We observe a gradual transition from strong covalent contributions (4100 kJ mol 1) in the O–H O bonded dimer anion of 1 to a moderate covalent character in the N–H O case, and finally to clearly noncovalent interactions as would be expected for C–H O hydrogen bonds close to the van-der-Waals limit. This approach seems to be useful to characterise known and new molecular crystal structures whenever a detailed look at their chemical-bonding nature is warranted.
This journal is © The Royal Society of Chemistry 2014
1 S. J. Grabowski, Chem. Rev., 2011, 111, 2597–2625, see also references therein. 2 (a) G. C. Pimentel and A. L. McClellan, The Hydrogen Bond, W. H. Freeman, San Francisco, London, 1960; (b) G. A. Jeffrey, An Introduction to Hydrogen Bonding, Oxford University Press, Oxford, 1997; (c) W. W. Cleland, P. A. Frey and J. A. Gerlt, J. Biol. Chem., 1998, 273, 25529–25532. 3 Indeed, early considerations based on molecular-orbital theory have already been made in the 1970’s: P. A. Kollman and L. C. Allen, J. Am. Chem. Soc., 1970, 92, 6101–6107. 4 (a) G. Gilli, F. Bellucci, V. Ferretti and V. Bertolasi, J. Am. Chem. Soc., 1989, 111, 1023–1028; (b) A key contribution is also in G. Gilli and P. Gilli, J. Mol. Struct., 2000, 552, 1–15. ´ski, J. Phys. 5 (a) S. J. Grabowski, W. A. Sokalski, E. Dyguda and J. Leszczyn Chem. B, 2006, 110, 6444–6446; (b) R. Parthasarathi, V. Subramanian and N. Sathyamurthy, J. Phys. Chem. A, 2006, 110, 3349–3351; (c) I. Mata, I. Alkorta, E. Molins and E. Espinosa, Chem. – Eur. J., 2010, 16, 2442–2452. 6 J. D. Dunitz and A. Gavezzotti, Angew. Chem., Int. Ed., 2005, 44, 1766–1787. ¨ller, R. Dronskowski and 7 V. L. Deringer, F. Pan, J. George, P. Mu U. Englert, CrystEngComm, 2014, 16, 135–138. 8 (a) C. A. Morrison and M. M. Siddick, Angew. Chem., Int. Ed., 2004, 43, 4780–4782; (b) V. Hoepfner, V. L. Deringer and R. Dronskowski, J. Phys. Chem. A, 2012, 116, 4551–4559. 9 (a) A. Gavezzotti, Mol. Phys., 2008, 106, 1473–1485; (b) J. D. Dunitz and A. Gavezzotti, Cryst. Growth Des., 2012, 12, 5873–5877. 10 J. Netzel and S. van Smaalen, Acta Crystallogr., Sect. B: Struct. Sci., 2009, 65, 624–638. 11 M. Serb, R. Wang, M. Meven and U. Englert, Acta Crystallogr., Sect. B: Struct. Sci., 2011, 67, 552–559. ¨chl, J. Phys. Chem., 1993, 97, 12 R. Dronskowski and P. E. Blo 8617–8624; for a more comprehensive introduction, we may refer the reader to R. Dronskowski, Computational Chemistry of Solid State Materials, Wiley-VCH, Weinheim, 2005. 13 For these examples, see: (a) G. A. Landrum and R. Dronskowski, Angew. ¨sebrink, Chem., Int. Ed., 2000, 39, 1560–1585; (b) M. Wuttig, D. Lu D. Wamwangi, W. Wełnic, M. Gilleßen and R. Dronskowski, Nat. Mater., 2007, 6, 122–128. 14 Indeed, the orbital picture has been invoked before by Gilli et al.4 and the COHP has been used in model computations on idealised HB chains in M. Springborg, Int. J. Quantum Chem., 2000, 77, 843–858. 15 Computations were done using gradient-correcteda DFT and projector augmented wavesb as implemented in VASP.c,d Full computational details are in the ESI.† (a) J. P. Perdew, K. Burke and M. Ernzerhof, ¨chl, Phys. Rev. B: Phys. Rev. Lett., 1996, 77, 3865–3868; (b) P. E. Blo Condens. Matter Mater. Phys., 1994, 50, 17953–17979; (c) G. Kresse and ¨ller, Phys. Rev. B: Condens. Matter Mater. Phys., 1996, 54, J. Furthmu 11169–11186; (d) G. Kresse and D. Joubert, Phys. Rev. B: Condens. Matter Mater. Phys., 1999, 59, 1758–1775. 16 V. L. Deringer, V. Hoepfner and R. Dronskowski, Cryst. Growth Des., 2012, 12, 1012–1021. ´eff and R. Dronskowski, J. Phys. 17 (a) V. L. Deringer, A. L. Tchougre Chem. A, 2011, 115, 5461–5466; (b) S. Maintz, V. L. Deringer, ´eff and R. Dronskowski, J. Comput. Chem., 2013, 34, A. L. Tchougre 2557–2567. 18 Y. A. Abramov, Acta Crystallogr., Sect. A, 1997, 53, 264–272. 19 (a) G. K. H. Madsen, B. B. Iversen, F. K. Larsen, M. Kapon, G. M. Reisner and F. H. Herbstein, J. Am. Chem. Soc., 1998, 120, 10040–10045; (b) D. E. Hibbs, J. Overgaard and R. O. Piltz, Org. Biomol. Chem., 2003, 1, 1191–1198; (c) P. M. B. Piccoli, T. F. Koetzle, A. J. Schultz, E. A. Zhurova, J. Stare, A. A. Pinkerton, J. Eckert and D. Hadzi, J. Phys. Chem. A, 2008, 112, 6667–6677; (d) M. Schmidtmann, L. J. Farrugia, D. S. Middlemiss, M. J. Gutmann, G. J. McIntyre and C. C. Wilson, J. Phys. Chem. A, 2009, 113, 13985–13997. 20 G. E. Bacon, N. A. Curry and S. A. Wilson, Proc. R. Soc. London, Ser. A, 1964, 279, 98–110.
Chem. Commun., 2014, 50, 11547--11549 | 11549
COMMUNICATION
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Covalency of hydrogen bonds in solids revisited† Cite this: Chem. Commun., 2014, 50, 11547
Volker L. Deringer,a Ulli Englert*a and Richard Dronskowski*ab
Published on 08 August 2014. Downloaded on 05/10/2014 17:42:47.
Received 20th June 2014, Accepted 7th August 2014 DOI: 10.1039/c4cc04716h www.rsc.org/chemcomm
The covalent nature of short hydrogen bonds has been under debate for long. Here we show that the crystal orbital Hamilton population (COHP) bonding indicator gives new, complementary evidence of covalent hydrogen acceptor interactions in the molecular solid state.
‘‘What is the covalency of hydrogen bonding?’’ asked Grabowski in a recent review article.1 The ubiquity of hydrogen bonds (HBs) both in crystal chemistry and molecular biology2 has stimulated much thought about the issue.1 Short HBs were long assumed to be covalent to some degree,3 and particular milestones were set by Gilli et al. with the introduction of the resonance-assisted hydrogen bond (RAHB).4 Nonetheless, there remain several interesting questions. For example, if short HBs are covalent and others are not, should one draw a line between both regimes, or is there a gradual transition?5 Furthermore, the H acceptor distance is an important criterion to identify HBs in solids, but the mere existence of short contacts does not suffice to explain the stability of crystalline networks.6,7 Reliable knowledge about the covalency of HBs is very much sought after, and in particular, one would like to have tools for analysing hydrogen-bonded crystal structures. There have been previous efforts to gauge the strength of HBs in solids—for example, from DFT energies,8 and there are also useful estimators such as the PIXEL method.9 These tools assign a numerical energy value to a given intermolecular contact or crystalline direction, and ‘‘bond energy’’ is a tangible quantity without doubt, even if no rigorous observable directly corresponds to it. On the other hand, topological analysis of the electronic charge density has proven to be valuable, and the latter may be derived from theory5 or from high-resolution diffraction experiments.10,11 It would now seem to be rewarding to combine these conceptual approaches: to determine the nature and strength of HBs in a solid a
Institute of Inorganic Chemistry, RWTH Aachen University, Landoltweg 1, 52056 Aachen, Germany. E-mail: [email protected], [email protected]; Fax: +49 241 8092288, +49 241 8092642 b ¨lich–Aachen Research Alliance (JARA-HPC), RWTH Aachen University, Ju 52056 Aachen, Germany † Electronic supplementary information (ESI) available: Details of theoretical computations. See DOI: 10.1039/c4cc04716h
This journal is © The Royal Society of Chemistry 2014
directly from its underlying electronic structure. In this work, we seek to explore HBs by invoking the familiar model of covalent bonding between atomic orbitals. To this end, we performed periodic DFT computations for molecular solids with pronounced HBs, and then dissected the electronic structure of the crystals into orbital-wise contributions to reveal the bonding nature between neighbouring atoms. The method we used has been dubbed ‘‘crystal orbital Hamilton population’’ (COHP) analysis and is described in ref. 12. In solid-state and materials chemistry and physics, COHP is an established tool which has been applied to a plethora of compounds, e.g., to tailor-made ferromagnets or datastorage alloys.13 Using COHP analysis to investigate organic crystals and HBs may seem unconventional in this light, but there is no fundamental reason for not doing so.14 We began our study with an intriguing compound that contains many different HBs in its crystal structure: from very short (O–H O type) contacts to those on the verge of being bonded at all. The compound, N,N-dimethylbiguanidinium bis(hydrogensquarate) 1, was previously characterised using high-resolution X-ray and neutron diffraction.11 We have now performed periodic electronic-structure computations15 on the neutron-derived crystal structure, in which we kept the structure frozen but re-optimised the positions of hydrogen atoms.16 From the plane-wave based DFT output, we then computed COHP curves in the framework of a recently developed projection scheme,17 and the result is hence denoted as ‘‘projected COHP’’. This new method makes it possible to treat even complex molecular structures as presented here. First, we examined the characteristic anion dimer in crystalline 1, as sketched in Fig. 1a. Two O–H O bonds are formed, one (highlighted in red) notably shorter than the other one (blue); this asymmetry is caused by the crystalline network.11 The electronic DOS (Fig. 1b), which is the solid state analogue of a molecularorbital scheme, reveals contributions of the acceptor oxygen over the entire valence region, from 12 eV upwards, and the local DOS of the hydrogen atom in the shortest HB is largest at 12 eV, too. Due to this very orbital overlap, a covalent bonding interaction between both atoms arises, as reflected in pCOHP values plotted to the right of the vertical axis. We reiterate that this technique reveals
Chem. Commun., 2014, 50, 11547--11549 | 11547
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Published on 08 August 2014. Downloaded on 05/10/2014 17:42:47.
Fig. 2
Integrated –pCOHP at eF for all relevant H O contacts in crystalline 1.
Fig. 1 (a) Structural drawing of 1;11 the hydrogensquarate anions dimerise in the solid state as shown here. The strongest (red) and second (blue) O–H O bonds have been highlighted, and the labelling (H1/H5) corresponds to the original article11 for convenience. (b) Computed densities of states (DOS, for hydrogen and acceptor atoms of the strongest HB in 1); projected crystal orbital Hamilton populations (pCOHP) for both above contacts, and their integrals (in eV).
covalent interactions (in terms of pairwise contributions to the electronic band-structure energy), but not electrostatic and dispersive terms which would govern weaker HBs.2 The higher-lying energy bands (from 8 eV on) also contain H acceptor interactions, but these are nonbonding in good approximation because stabilising and destabilising pCOHP regions cancel each other. When integrated up to the highest filled orbital (i.e., up to the Fermi energy eF), the (p)COHP provides a measure of the covalent bond strength,12 just like energy-integrating the DOS would give the number of electrons in the system. We have hence plotted the integrated pCOHP (ICOHP) for the two contacts on the right of Fig. 1b. Defining stabilising interactions to have positive numerical values, we obtain 149 and 104 kJ mol 1, respectively. This indicates significant covalent bonding, on the order of what is found for resonance-assisted hydrogen bonds (RAHBs), a well-studied family of strong HB contacts.4 We stress once more that the ICOHP does not equal the bond energy but provides a reasonable estimate.12 We then proceeded to investigate the entire range of HB interactions in 1. For each, the ICOHP was computed, and the results are plotted in Fig. 2 against the respective H O distance. The covalent contributions decrease for the longer HBs, and a clear trend is seen: several N–H O interactions in 1 have lower ICOHP values than the O–H O ones but still exhibit some covalent bonding. The unconventional C–H O HBs, on the other hand, have values close to zero, in line with the expectation that they are weak and noncovalent in nature.2 The difficulty of identifying single bonds at such distances (and thereby possibly overlooking important, but less localised electrostatic and dispersive attraction) has been discussed before.6 For validation, we compared the ICOHP of each bond to the respective charge density at the bond-critical point, rbcp, which counts among the traditionally established criteria for HB covalency, together with the Laplacian (cf. ref. 11 and references
11548 | Chem. Commun., 2014, 50, 11547--11549
Fig. 3 Scatter plot of integrated pCOHPs vs. experimentally determined electron densities at the respective bond-critical points. The latter data are taken from ref. 11 and 19a–d. For the very short HB in benzoylacetone,19a the hydrogen atoms were either frozen in their neutron-derived structure or re-optimised.
therein) or energy densities.18 Generally, covalent HBs show significant rbcp values in the range of 0.1–1 e Å 3, whereas weaker HBs, dominated by electrostatic forces, do not exhibit rbcp values far from zero. To rely on a larger dataset, we extended our study to a number of high-quality charge-density and neutron diffraction studies from the previous literature.19 Fig. 3 clearly evidences that both methods of estimating HB covalency do exhibit a linear correlation—this is not to be expected per se, because ICOHP and rbcp stem from fundamentally different approaches. It is also appreciable that the dataset for 1 (blue symbols in Fig. 3) is consistent with the additional data points (red); the latter pertain to different structures and measurements as detailed in the ESI.† This provides confidence for future use of the method. For very strong HBs (rbcp 4 0.6 e Å 3), the results are visibly influenced by the decision whether hydrogen positions are re-optimised with DFT or locked to the neutron-derived values. We illustrate this for the O–H O contact in benzoylacetone, for which both options are visible in Fig. 3 (circled data points): the optimised covalent O–H bond length is shorter compared to neutron data (cf. ref. 16), and hence the computed H O distance is larger. While the limitations of proton localisation become apparent in such extreme cases, we argue that there
This journal is © The Royal Society of Chemistry 2014
View Article Online
ChemComm
Communication
We thank the Studienstiftung des deutschen Volkes for a ¨lich–Aachen Research Alliance scholarship to V.L.D. and the Ju for JARA-HPC computer time.
Notes and references
Published on 08 August 2014. Downloaded on 05/10/2014 17:42:47.
Fig. 4 (a) Fragment from the crystal structure of 1-(2-hydroxy-5nitrophenyl)ethanone, drawn with data from ref. 19b. (b) The fragment from the crystal structure of benzene,20 a prototypical molecular solid with ‘‘T-stacking’’; see text.
remains significant correlation between ICOHP and rbcp data regardless of how the H positions are treated. The new method appears to provide an advantage when looking at the intramolecular H O contacts which exhibit short distances but would not be classified as HBs. Fig. 4a shows a structural fragment from one of the previous studies we built upon:19b the H5 atom forms a C–H O bond with a neighbouring O2 atom; furthermore, it is close to the O3 atom of the nitro group, but the latter contact would not be seen as a hydrogen bond as its oCHO angle is close to 90 degrees. If the electron density in the H O bond-critical point is taken as the only output of the experimental charge-density analysis (the authors of the original work did not do that!), it might suggest a higher degree of covalency (rbcp = 0.08(1) e Å 3) in the intramolecular contact than in the intermolecular H bond (0.04(1) e Å 3). This clearly underlines that even a highly useful analysis technique must be used with (chemical) caution. The pCOHP data in Fig. 4a clearly differentiate between the two types of contacts, yielding a slightly stabilising interaction for the H bond but none for the intramolecular H5 O3 contact. We stress that this ‘‘nonbonding’’ intramolecular contact has been omitted from Fig. 3, for obvious reasons. Finally, an alert reviewer pointed out to us the C–H p interactions which, by their name, could suggest the involvement of orbital-pair interactions. Test computations for the crystal structure of benzene (Fig. 4b),20 however, gave no indications of covalent twocentre bonding (with ICOHP values very slightly below 0 eV), which is in accord with the quite long H C distances (optimised: 42.85 Å; cf. the trends in Fig. 2). For these families of interactions, different theoretical tools will be needed, and the argumentation of Dunitz and Gavezzotti6 holds. In conclusion, we have demonstrated an alternative, orbitalbased way of probing hydrogen-bond covalency in solids, which complements previous careful charge-density studies. We observe a gradual transition from strong covalent contributions (4100 kJ mol 1) in the O–H O bonded dimer anion of 1 to a moderate covalent character in the N–H O case, and finally to clearly noncovalent interactions as would be expected for C–H O hydrogen bonds close to the van-der-Waals limit. This approach seems to be useful to characterise known and new molecular crystal structures whenever a detailed look at their chemical-bonding nature is warranted.
This journal is © The Royal Society of Chemistry 2014
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