Are beryllium-containing biphenyl derivatives efficient anion sponges?

  • Oriana Brea
  • Otilia Mó
  • Manuel Yáñez
  • M. Merced Montero-Campillo
  • Ibon Alkorta
  • José Elguero
Original Paper
Part of the following topical collections:
  1. P. Politzer 80th Birthday Festschrift


The structures and stabilities of 2,2′-diBeX-1,1′-biphenyl (X = H, F, Cl, CN) derivatives and their affinities for F, Cl, and CN were theoretically investigated using a B3LYP/6–311 + G(3df,2p)//B3LYP/6–31 + G(d,p) model. The results obtained show that the 2,2′-diBeX-1,1′-biphenyl derivatives (X = H, F, Cl, CN) exhibit very high F, Cl, and CN affinities, albeit lower than those reported before for their 1,8-diBeX-naphthalene analogs, in spite of the fact that the biphenyl derivatives are more flexible than their naphthalene counterparts. Nevertheless, some of the biphenyl derivatives investigated are predicted to have anion affinities larger than those measured for SbF5, which is considered one of the strongest anion capturers. Therefore, although weaker than their naphthalene analogs, the 2,2′-diBeX-1,1′-biphenyl derivatives can still be considered powerful anion sponges. This study supports the idea that compounds containing –BeX groups in chelating positions behave as anion sponges due to the electron-deficient nature and consequently high intrinsic Lewis acidity of these groups.

Graphical Abstract

Compounds containing –BeX groups in chelating positions, such as 2,2′-diBeX-1,1′-biphenyl (X = H, F, Cl, CN) derivatives, behave as anion sponges due to the electron-deficient nature of these groups


Anion sponges Be-containing biphenyl derivatives Density functional theory 


Noncovalent interactions play a crucial role in chemistry; indeed, the last decades of the twentieth century as well as the first decades of the present century have seen a significant increase in the number of noncovalent interactions described in the literature [1, 2, 3, 4, 5, 6, 7, 8, 9, 10], and many of them have been fully characterized. The contributions of Peter Politzer and coworkers in this field of research are numerous and of great relevance [11]. His pioneering work established the basis for obtaining quantitative information on the strength of these interactions from appropriate electronic densities, and it showed how the electrostatic and polarization energies associated with noncovalent interactions can be obtained from them. Also, new and interesting noncovalent interactions such as halogen bonds have been characterized [12, 13, 14]. Perhaps one of the most useful concepts to be introduced into this domain was that of the σ-hole [15, 16, 17], which allows the nature of some of these interactions, such as the halogen bonds mentioned above, to be rationalized [18]. However, as stated above, the most fundamental finding was that noncovalent interactions control an enormous number of processes, including selective interactions between molecules and anions [19, 20, 21, 22, 23, 24, 25, 26, 27, 28, 29, 30, 31, 32, 33, 34, 35, 36, 37, 38, 39, 40, 41, 42, 43, 44, 45, 46, 47, 48, 49, 50, 51, 52, 53, 54, 55, 56, 57, 58]. As a matter of fact, the selective extraction of anions [59]—a crucial process in many fields such as medicinal and biological chemistry, environmental science, and even basic chemistry—has received a great deal of attention, mainly from an experimental viewpoint [60], but also, more recently, from a theoretical perspective [32, 61, 62, 63, 64]. In particular, works aiming at the specific design of functionalized compounds for the capture and immobilization of toxic anions have been very popular over the last few years. Indeed, it was very recently shown that compounds containing electron-deficient elements, such as beryllium, behave as highly efficient anion sponges [63, 64], exhibiting anion binding energies that are estimated to be up to 30% higher than those of compounds such as SbF5 or AsF5, which are among the systems that had previously been considered to possess the highest anion affinities [19, 65]. This is the case for 1,8-diBeX-naphthalene (X = H, F, Cl, CN, CF3) derivatives [63] as well as for 4,5-bis(BeX)-fluorene (X = H, F, Cl, CN, NC, OCH3) derivatives [64].

However, an intrinsic characteristic of these two families of compounds is the rigidity of the molecular skeleton, which may imply a certain constraint when binding certain anions. This constraint is readily apparent for the 1,8-diBeX-naphthalene derivatives, where the distance between the two Be atoms interacting with the anions is limited by the σ-skeleton of the naphthalene moiety. Hence, we thought that it could be interesting to explore whether more flexible molecular frameworks allow enhanced interactions with anions. We considered 2,2′-diBeX-1,1′-biphenyl derivatives to be ideal candidates for such a study (see Scheme 1) because the free rotation of the two aromatic rings around the C1–C1′ bridge allows the Be···Be distance to be adjusted to achieve the optimal interaction with the incoming anion. In this paper, we explore the behavior of these derivatives (X = H, F, Cl, CN) when they interact with a set of anions Y = F, Cl, CN.
Scheme 1.

The 2,2′-diBeX-1,1′-biphenyl derivatives included in this study

Computational details

The geometries of the 2,2′-diBeX-1,1′-biphenyl (X = H, F, Cl, CN) neutral derivatives and those of the complexes of these derivatives with a set of Y anions (Y = F, Cl,CN) were optimized using the B3LYP hybrid functional [66, 67], which includes the three-parameter functional developed by Becke [66] and the correlation functional of Lee, Yang, and Parr [67], together with a 6–31 + G(d,p) basis set expansion. The final energies of these optimized structures were determined in single-point B3LYP/6–311 + G(3df,2p) calculations. This theoretical model was found to provide anion affinities that were only slightly larger than the values obtained through the use of high-level ab initio methods such as G4(MP2) [68], which typically provides thermodynamic magnitudes to an accuracy of ±4 kJ mol−1. The harmonic vibrational frequencies for both neutral and anion complexes were calculated at the B3LYP/6–31 + G(d,p) level of theory in order to check that the stationary points found were local minima of the potential energy surface, and to determine the vibrational corrections to the corresponding enthalpies.

To analyze the electronic characteristics of the anionic complexes under investigation, we used two different but complementary approaches: the quantum theory of atoms in molecules (QTAIM) [69] and the natural bond orbital (NBO) method [70]. The first approach is based on an analysis of the topology of the electron density of the system, which usually exhibits different critical points, such as local maxima, associated with the position of the nuclei, and saddle points (usually called bond critical points, BCPs), the electron density of which reflects the strength of the bond. The sign of the so-called energy density at each BCP permits us to assess the covalent character of the bonding. Simultaneously, the existence of rings is evidenced by the presence of ring critical points (RCPs). The NBO approach allows us to detect and quantify the existence of charge donation and backdonation between the interacting moieties that stabilize the complex through the use of localized occupied and empty orbitals. All these calculations were carried out using the AIMAll v1.0 [71] and NBO6G [72] programs.

Results and discussion

Structures and relative stabilities of the neutral compounds

We initially focused on analyzing the structures of the neutral systems. As we mentioned in the “Introduction,” the systems under investigation are characterized by the flexibility permitted by rotation around the C1–C1′ bond. Also, the neutral systems present different conformations depending on whether the BeX groups are arranged trans or cis. Our B3LYP results show that the trans conformers always exhibit a CC1C1′C dihedral angle of around 130° or slightly larger, as depicted in Fig. 1. These trans conformers are also systematically less stable than the corresponding cis conformers (see the relative energies in Fig. 1), and were therefore discarded from our analysis of the interactions of these derivatives with the anions under consideration.
Fig. 1

Trans conformers of 2,2′-diBeX-1,1′-biphenyl (X = H, F, Cl, CN) derivatives. The CC1C1′C dihedral angle (°) between the planes containing the two aromatic rings and the energy gap with respect to the most stable low-lying cis conformer are also indicated. Values were obtained at the B3LYP/6–311 + G(3df,2p) level. The H atoms are shown in white, Be atoms in yellow, C atoms in grey, N atoms in dark-blue, F atoms in cyan, and Cl atoms in green

Two different cis conformers were found to be local minima of the potential energy surface for each X substituent, as shown in Fig. 2. The first one (cis1, left column of Fig. 2) corresponds to a C 2v structure in which the BeX substituents are symmetrically located above and below the plane defined by the coplanar aromatic rings. The second local minimum (cis2, central column of Fig. 2) has only C 1 symmetry, and is characterized by the formation of a Be–X–Be bridge in which one of the Be atoms accepts electron density from the lone pairs of X in the neighboring BeX group.
Fig. 2

Molecular graphs of the stable cis conformers of 2,2′-diBeX-1,1′-biphenyl (X = F, Cl, CN) derivatives are displayed in the first (F), second (Cl) and third (CN) rows, respectively. The molecular graph for the X = F trans conformer is also shown in the first row. The left column corresponds to C 2v structures, whereas the central column shows nonsymmetric minima. Green and red dots indicate BCPs and RCPs, respectively. The electron densities (e/bohr3) of the most relevant BCPs are given, as well as the relative energies of these conformers in kJ mol−1

It is interesting to note that the enhanced stability of the cis structures is very likely to be related to the fact that the two Be atoms in all of these structures appear to be tricoordinated rather than dicoordinated as in the trans isomers. In the C 2v cis1 structures, both Be atoms are bound to one X substituent and two C atoms, whereas in the nonsymmetric cis2 structures, only one of the Be atoms presents this coordination pattern; the other Be is bound to one C and two X substituents. This increased coordination of the Be atoms in the cis conformers seems to be the reason for the enhanced stability of these forms with respect to their trans analogs. Indeed, although the two C–Be bonds in the trans conformer are stronger than the four C–Be bonds in the cis1 isomer according to the corresponding electron densities (see the BeF derivatives in Fig. 2), the overall stabilizing effect is larger for the cis1 conformer, as there are four bonding interactions instead of two. In addition, an extra bonding interaction between these two C atoms further contributes to the stabilization of the cis1 conformer. This description is coherent with that obtained through NBO analysis. In this approach, the stabilization associated with the individual C–Be linkages is greater for the trans than for the cis1 forms (531 vs. 368 kJ mol−1), but according to the total number of interactions, the overall stabilizing effect is larger for cis1. The situation for the cis2 conformer is similar to that described for cis1 as far as the C–Be bonds are concerned. Again, these bonds are stronger in the trans conformer, but we replace two bonding interactions with three upon converting to the cis2 form, yielding a greater stabilizing effect overall. A new bonding interaction appears between the bridging F atom and one of the C atoms, contributing further to the stabilization of this cis2 form. Once more, the bonding picture obtained through the NBO analysis is coherent with that obtained using the QTAIM approach. In this case, the C–Be interactions involve similar energies for both conformers (ca. 1020 kJ mol−1 for the cis2 conformer vs. 1050 kJ mol−1 for the trans one), whereas the Be–F interactions are 110 kJ mol−1 stronger for the cis2 structure.

It should also be noted that the relative stabilities of the two cis isomers change depending on the nature of the X substituent. With –BeF and –BeH, the cis2 nonsymmetric minimum is the most stable, whereas –BeCl and –BeCN derivatives have the cis1 C 2v structure as their global minimum. This finding probably reflects the fact that, because the F and H atoms are relatively small, a F···C (H···C) stabilizing interaction occurs for the BeF (BeH) derivatives (see Fig. 2), which is not observed for the –BeCl and –BeCN derivatives.

Structures and relative stabilities of the anionic complexes

In order to obtain the anion affinities of our set of biphenyl derivatives (X = H, F, Cl, CN), we considered different anion binding patterns (Y = F, Cl, CN) for the two most stable isomers cis1 and cis2.

Taking the cis1 neutral isomer as a starting point, symmetry considerations suggest that two anionic conformers are possible, depending on whether the new structure preserves the original C 2v symmetry or not. If not, the structure found when Y and X are equal is identical to the one obtained when the anion attaches to the dicoordinated Be atom of the cis2 isomer. In other words, for the case where X = Y both cis1 and cis2 lead to the same anion. Therefore, for the sake of simplicity, we will first discuss this case. The structures obtained are shown in Fig. 3. Attachment that preserves the C 2v symmetry of the cis1 isomer leads to the anions shown in the right column of Fig. 3, whereas attachment to cis1 that does not preserve the C 2v symmetry or attachment to isomer cis2 yields the structures in the left column of the same figure. Note that for the F and Cl derivatives, the corresponding anionic species have C 2 symmetry, whereas the anionic CN derivative belongs to the C 1 symmetry group.
Fig. 3

Structures and relative stabilities (in kJ mol−1) of the anionic species formed upon the attachment of Y to the neutral 2,2′-diBeX-1,1′-biphenyl derivatives when X = Y (X = F, Cl, CN; Y = F, Cl, CN). The symmetry of the newly formed anionic complex is also indicated. The H atoms are shown in white, Be atoms in yellow, C atoms in grey, N atoms in dark-blue, F atoms in cyan, and Cl atoms in green

The first conspicuous fact is that, in all cases, the most stable anion is the least symmetric species, although the energy gap decreases significantly upon switching from the F- to the Cl-containing derivative, and it becomes almost zero for CN-containing analogs. As an example, let us discuss the F-containing systems (molecular graphs for them are shown in Fig. 4). QTAIM analysis shows that whereas the Be–F interactions are stronger in the C 2 than in the C 2v structure, there is an additional interaction between the central F atom and the two carbon atoms in the latter, which likely compensates for the difference in Be–F interaction strength. Similarly, C–Be interactions are stronger in the C 2 than in the C 2v structure, but in the latter, each C is bound to the two Be atoms. A more quantitative investigation can be performed using NBO analysis, which shows that the F–Be interaction per bond is slightly stronger for the C 2 structure (259 vs. 247 kJ mol−1) but that the difference is larger for the C–Be interactions (431 vs. 376 kJ mol−1). A second factor that contributes to the enhanced stability of the C 2 structure is steric repulsion between the two aromatic rings, which are necessarily coplanar in the C 2v structure but not in the C 2 structure. Our estimates indicate that the rotation around the C1–C1′ bond connecting both rings alleviates this steric repulsion by about 10 kJ mol−1.
Fig. 4

Molecular graphs of the C 2 and C 2v complexes between 2,2′-diBeF-1,1′-biphenyl and F. Conventions used are the same as those in Fig. 2

Lower C 2v anion stability is also observed for the complexes in which X ≠ Y, with the only exception being the Cl complex of the 2,2′-diBeCN-1,1′-biphenyl derivative (see Fig. 5), as we will discuss later. As shown in that figure, the nature of the global minimum depends on the X substituent. For –BeF derivatives, the global minimum corresponds to the asymmetric structure arising from the attachment of the anion to the cis2 isomer of the neutral derivative. Conversely, for the –BeCl derivatives, the global minimum is the C 2 structure that can be considered to result from the attachment of the anion to the C 2v isomer of the neutral derivative. This is also the case for the –BeCN derivatives when they interact with F but not with Cl, in which case the global minimum is predicted to be the C 2v adduct, which is almost degenerate with the attachment to the cis2 isomer.
Fig. 5

Structures and relative stabilities (in kJ mol−1) of the anions formed by the attachment of Y = F, Cl, CN to the cis1 and cis2 forms of the 2,2′-diBeX-1,1′-biphenyl (X = F, Cl, CN) derivatives when X ≠ Y. The H atoms are shown in white, Be atoms in yellow, C atoms in grey, N atoms in dark-blue, F atoms in cyan, and Cl atoms in green

As we will see, these results imply that the relative stabilities of the anions depend on a subtle balance of different factors. To illustrate this point, let us consider the complexes between 2,2′-diBeF-1,1′-biphenyl and Cl. As indicated in the first row of Fig. 5, the A3 anion is the global minimum, with the A2 anion being about 50 kJ mol−1 higher in energy. Upon switching from A2 to A3, the number of Be–F bonds increases from two to three, whereas the number of Be–Cl bonds decreases from two to one. No other changes in bond type occur, but the environment of the C–Be bonds undoubtedly changes. While both C–Be bonds are identical in A2, but not in A3, as evidenced by the electron densities at the BCPs in Fig. 6.
Fig. 6

Molecular graphs of the A2 and A3 anions formed upon the interaction of the 2,2′-diBeF-1,1′-biphenyl derivative with Cl. Conventions used are the same as those in Fig. 2. Recall that the A2 structure is 50.7 kJ mol−1 higher in energy than the A3 isomer

The C–Be bonds are also predicted to be slightly stronger in the global minimum, A3. However, the reduction in the number of the Be–Cl bonds overrides this, even though the Be–Cl bond is stronger in A3 than in A2, the electron density is not twice as strong. Accordingly, a destabilization of the system would be expected. The increase in the number of Be–F bonds upon switching from A2 to A3 stabilizes the system, because even though the two F–Be bonds involving the central F atom in A3 are weaker than each of the Be–F bonds in A2, the formation of two bonds results in an overall stabilizing effect. If this latter effect is the dominant one, we would expect the A3 structure to be more stable than A2. Unfortunately, the QTAIM approach does not provide information about the energetics of the aforementioned electron density redistributions, but this information can be deduced from the NBO analysis, which shows that the greater number of Be–F interactions in the A3 isomer stabilizes it by about 69 kJ mol−1, while the decreased number of Be–Cl interactions destabilizes it by around 52 kJ mol−1 with respect to A2. The NBO results also show that the C–Be interactions are slightly reinforced (by 20 kJ mol−1). Finally, because the Cl atom is larger than F, it forces a wider separation between the Be atoms (2.72 Å in A3 vs. 3.11 Å in A2) and therefore a greater distortion of the molecular skeleton, which has a corresponding energetic cost, estimated to be around 65 kJ mol−1. The small geometrical distortion of the aromatic system when the bridging atom is F is very likely the main factor that causes isomer A3 to be the most stable one for –BeF derivatives when they interact with Cl and CN. Instead, isomer A2 is the most stable one for –BeCl derivatives interacting with F and CN, as well as for –BeCN derivatives interacting with F, although, in the latter case, the A3 form is found to be slightly more stable when the interaction takes place with CN (see Fig. 4).

Anion affinities

We have summarized in Table 1 the calculated anion affinities of the 2,2′-diBeX-1,1′-biphenyl derivatives under consideration, defined as the enthalpy difference between the most stable anionic adduct and the most stable neutral form.
Table 1

Comparison between the calculated anion affinities (in kJ mol−1) of the 2,2′-diBeX-1,1′-biphenyl derivatives (Biph) and those of the 1,8-BeX-naphthalene derivatives (Napht)


X = H

X = F

X = Cl

X = CN





































These values are compared in the table with those calculated to the same level of accuracy for the corresponding 1,8-BeX-naphthalene derivatives. However, some considerations should be taken into account when comparing these two families. The values for the 1,8-BeX-naphthalene derivatives with X = Cl, CN were taken from [63]. For X = H, F, the values reported in [63] correspond to the symmetric structure (a) shown in Scheme 2, whereas the most stable neutral conformation for X = H, F is structure (b), in which the substituent bridges both Be atoms. Hence, for the sake of coherence, the anion affinities for these two compounds were calculated in this work at the same level of theory as used for the biphenyl analogs, using the more stable isomer (b) as a reference for the neutral compound.
Scheme 2

Possible local minima for the 1,8-BeX-naphthalene derivatives

The first conspicuous fact is that the anion affinities for the 2,2′-diBeX-1,1′-biphenyl derivatives are systematically lower than those for the corresponding naphthalene analogs. This result is rather unexpected, because we would expect the rigidity of naphthalene to enhance the anion affinity (AA) rather than decreasing it.

It is also interesting to note that there is a reasonably good correlation between the biphenyl and the naphthalene anion affinities (see Fig. 7), as summarized below:
$$ \mathrm{AA}\left(\mathrm{biphenyls}\right)=\left(0.85\pm 0.02\right)\mathrm{AA}\left(\mathrm{naphthalenes}\right),\kern.8em {r}^2=\mathrm{0.995.} $$
Fig. 7

Correlation between the F, Cl, and CN anion affinities of 2,2′-diBeX-1,1′-biphenyl derivatives (Biph) and those of 1,8-BeX-naphthalene derivatives (Napht); values in kJ mol−1

This correlation presents a residual RMS of 27.0 kJ mol−1. This value indicates that (with an error of about 27 kJ mol−1) biphenyls behave as anion sponges with anion affinities that are 85% of those of the naphthalene analogs. Indeed, the F anion affinity for the 2,2′-diBeCN-1,1′-biphenyl derivative is estimated to be 18 kJ mol−1 higher than that of SbF5, which is considered one of the strongest F capturers in the literature [65].

The high anion affinities of the biphenyl derivatives investigated can also be explained by considering the molecular electrostatic potentials (MESPs). As shown in Fig. 8, the 2,2′ substitution at the biphenyl molecule by BeX leads to the appearance of an electron-deficient area between the Be atoms. Also, in agreement with the higher anion affinities obtained for the complexes with X = CN than with the other substituents, it is apparent that this positive area becomes increasingly positive upon switching from X = F to Cl to CN.
Fig. 8

Top view of the MESPs for the biphenyl and cis conformers of 2,2′-diBeX-1,1′-biphenyl (X = F, Cl, CN). The MESPs were calculated at the same level of theory as the optimized geometries, which are represented by electron density isosurfaces of 0.001 a.u. (roughly corresponding to the van der Waals surface); the values shown are in kJ mol−1


Our high-level DFT calculations show that 2,2′-diBeX-1,1′-biphenyl derivatives (X = H, F, Cl, CN), just like their 1,8-diBeX-naphthalene analogs, exhibit very high anion affinities when interacting with F, Cl, and CN. However, rather unexpectedly, the biphenyl derivatives—although more flexible—are weaker Lewis acids than their naphthalene counterparts. In spite of this, it must be emphasized that some of the derivatives investigated are predicted to have larger anion affinities than those measured for SbF5, which is generally considered to be one of the strongest anion capturers. Therefore, although they have weaker anion affinities than the naphthalene analogs, the 2,2′-diBeX-1,1′-biphenyl derivatives can still be considered anion sponges. Thus, the main conclusion of this study is that compounds with BeX groups in positions which favor the chelation of anions behave as anion sponges due to the electron-deficient nature and therefore high intrinsic Lewis acidities of the BeX groups.



This work was supported by the projects CTQ2015-63997-C2 and CTQ2013-43698-P of the Ministerio de Economía y Competitividad of Spain, by the project FOTOCARBON-CM S2013/MIT-2841 of the Comunidad Autónoma de Madrid, and by the COST Action CM1204. Computational time at the Centro de Computación Científica (CCC) of Universidad Autónoma de Madrid is also gratefully acknowledged.


  1. 1.
    Muller-Dethlefs K, Hobza P (2000) Noncovalent interactions: a challenge for experiment and theory. Chem Rev 100(1):143–167CrossRefGoogle Scholar
  2. 2.
    Karshikoff A (2006) Non-covalent interactions in proteins. World Scientific, SingaporeGoogle Scholar
  3. 3.
    Riley KE, Pitonak M, Jurecka P, Hobza P (2010) Stabilization and structure calculations for noncovalent interactions in extended molecular systems based on wave function and density functional theories. Chem Rev 110(9):5023–5063CrossRefGoogle Scholar
  4. 4.
    Hobza P, Müller-Dethlefs K (2010) Non-covalent interactions: theory and experiment. Royal Society of Chemistry, CambridgeGoogle Scholar
  5. 5.
    Adeli M, Soleyman R, Beiranvand Z, Madani F (2013) Carbon nanotubes in cancer therapy: a more precise look at the role of carbon nanotube–polymer interactions. Chem Soc Rev 42(12):5231–5256Google Scholar
  6. 6.
    Yan QF, Luo ZY, Cai K, Ma YG, Zhao DH (2014) Chemical designs of functional photoactive molecular assemblies. Chem Soc Rev 43(12):4199–4221CrossRefGoogle Scholar
  7. 7.
    Mahadevi AS, Sastry GN (2016) Cooperativity in noncovalent interactions. Chem Rev 116(5):2775–2825CrossRefGoogle Scholar
  8. 8.
    Rodgers MT, Armentrout PB (2016) Cationic noncovalent interactions: energetics and periodic trends. Chem Rev 116(9):5642–5687Google Scholar
  9. 9.
    Parrill AL, Lipkowitz KB, DiLabio GA, Otero-de-la-Roza A (2016) Noncovalent interactions in density functional theory (Reviews in Computational Chemistry, vol 29). Wiley, HobokenGoogle Scholar
  10. 10.
    Hermann J, DiStasio RA, Tkatchenko A (2017) First-principles models for van der Waals interactions in molecules and materials: concepts, theory, and applications. Chem Rev 117(6):4714–4758CrossRefGoogle Scholar
  11. 11.
    Ma YG, Politzer P (2004) Electronic density approaches to the energetics of noncovalent interactions. Int J Mol Sci 5(4–7):130–140CrossRefGoogle Scholar
  12. 12.
    Politzer P, Lane P, Concha MC, Ma YG, Murray JS (2007) An overview of halogen bonding. J Mol Model 13(2):305–311CrossRefGoogle Scholar
  13. 13.
    Politzer P, Murray JS, Concha MC (2007) Halogen bonding and the design of new materials: organic bromides, chlorides and perhaps even fluorides as donors. J Mol Model 13(6–7):643–650CrossRefGoogle Scholar
  14. 14.
    Politzer P, Murray JS, Clark T (2010) Halogen bonding: an electrostatically-driven highly directional noncovalent interaction. Phys Chem Chem Phys 12(28):7748–7757CrossRefGoogle Scholar
  15. 15.
    Clark T, Hennemann M, Murray JS, Politzer P (2007) Halogen bonding: the sigma-hole. J Mol Model 13(2):291–296CrossRefGoogle Scholar
  16. 16.
    Politzer P, Murray JS, Lane P (2007) Sigma-hole bonding and hydrogen bonding: competitive interactions. Int J Quant Chem 107(15):3046–3052CrossRefGoogle Scholar
  17. 17.
    Murray JS, Lane P, Politzer P (2009) Expansion of the sigma-hole concept. J Mol Model 15(6):723–729CrossRefGoogle Scholar
  18. 18.
    Kolar MH, Hobza P (2016) Computer modeling of halogen bonds and other sigma-hole interactions. Chem Rev 116(9):5155–5187Google Scholar
  19. 19.
    Haartz JC, McDaniel DH (1973) Fluoride-ion affinity of some Lewis acids. J Am Chem Soc 95(26):8562–8565Google Scholar
  20. 20.
    Pflugrath JW, Quiocho FA (1985) Sulfate sequestered in the sulfate-binding protein of Salmonella typhimurium is bound solely by hydrogen bonds. Nature 314(6008):257–260Google Scholar
  21. 21.
    Luecke H, Quiocho FA (1990) High specificity of a phosphate-transport protein determined by hydrogen-bonds. Nature 347(6291):402–406CrossRefGoogle Scholar
  22. 22.
    Stephan H, Gloe K, Schiessl P, Schmidtchen FP (1995) Lipophilic ditopic guanidinium receptors—selective extractants for tetrahedral oxoanions. Supramol Chem 5(4):273–280Google Scholar
  23. 23.
    Mangani S, Ferraroni M (1997) Natural anion receptors: anion recognition by proteins. In: Bianchi A, Bowman-James K, Garcia-Espafia E (eds) Supramolecular chemistry of anions. Wiley–VCH, Weinheim, pp 63–78Google Scholar
  24. 24.
    Ihm H, Yun S, Kim HG, Kim JK, Kim KS (2002) Tripodal nitro-imidazolium receptor for anion binding driven by (C-H)(+)-X- hydrogen bonds. Org Lett 4(17):2897–2900CrossRefGoogle Scholar
  25. 25.
    Kang SO, Jeon S, Nam KC (2002) Anion recognition by urea derivatives of anthraquinone: dihydrogen phosphate ion selective neutral receptors. Supramol Chem 14(5):405–410CrossRefGoogle Scholar
  26. 26.
    Guo W, Wang J, He JQ, Li ZC, Cheng JP (2004) Polymethylene-bridged cystine-glycine-containing cyclopeptides as hydrogen-bonding electroneutral anion receptors: design, synthesis, and halide ion recognition. Supramol Chem 16(3):171–174CrossRefGoogle Scholar
  27. 27.
    Zhang Y, Li MX, Lu MY, Yang RH, Liu F, Li KA (2005) Anion chelation-induced porphyrin protonation and its application for chloride anion sensing. J Phys Chem A 109:7442–7448Google Scholar
  28. 28.
    Melaimi M, Sole S, Chiu CW, Wang HD, Gabbai P (2006) Structural and electrochemical investigations of the high fluoride affinity of sterically hindered 1,8-bis(boryl)naphthalenes. Inorg Chem 45(20):8136–8143CrossRefGoogle Scholar
  29. 29.
    Blondeau P, Segura M, Perez-Fernandez R, de Mendoza J (2007) Molecular recognition of oxoanions based on guanidinium receptors. Chem Soc Rev 36(2):198–210CrossRefGoogle Scholar
  30. 30.
    Caltagirone C, Gale PA, Hiscock JR, Brooks SJ, Hursthouse MB, Light ME (2008) 1,3-Diindolylureas: high affinity dihydrogen phosphate receptors. Chem Commun 26:3007–3009CrossRefGoogle Scholar
  31. 31.
    Caltagirone C, Mulas A, Isaia F, Lippolis V, Gale PA, Light ME (2009) Metal-induced pre-organisation for anion recognition in a neutral platinum-containing receptor. Chem Commun 41:6279–6281CrossRefGoogle Scholar
  32. 32.
    Chen ZH, Amine K (2009) Computational estimates of fluoride affinity of boron-based anion receptors. J Electrochem Soc 156(8):A672–A676CrossRefGoogle Scholar
  33. 33.
    Hudnall TW, Chiu CW, Gabbai FP (2009) Fluoride ion recognition by chelating and cationic boranes. Acc Chem Res 42(2):388–397Google Scholar
  34. 34.
    Kubik S (2009) Amino acid containing anion receptors. Chem Soc Rev 38(2):585–605CrossRefGoogle Scholar
  35. 35.
    Zhang ZG, Schreiner PR (2009) (Thio)urea organocatalysis—what can be learnt from anion recognition? Chem Soc Rev 38(4):1187–1198Google Scholar
  36. 36.
    Amendola V, Fabbrizzi L, Mosca L (2010) Anion recognition by hydrogen bonding: urea-based receptors. Chem Soc Rev 39(10):3889–3915CrossRefGoogle Scholar
  37. 37.
    Hong SJ, Yoo J, Yoon DW, Yoon J, Kim JS, Lee CH (2010) Superior anion-binding properties of a cryptand-like oligopyrrolic macrocycle. Chem Asian J 5(4):768–772CrossRefGoogle Scholar
  38. 38.
    Li AF, Wang JH, Wang F, Jiang YB (2010) Anion complexation and sensing using modified urea and thiourea-based receptors. Chem Soc Rev 39(10):3729–3745CrossRefGoogle Scholar
  39. 39.
    Steed JW (2010) Anion-tuned supramolecular gels: a natural evolution from urea supramolecular chemistry. Chem Soc Rev 39(10):3686–3699CrossRefGoogle Scholar
  40. 40.
    Zhao HY, Gabbai FP (2010) A bidentate Lewis acid with a telluronium ion as an anion-binding site. Nat Chem 2(11):984–990CrossRefGoogle Scholar
  41. 41.
    Bergamaschi G, Boiocchi M, Monzani E, Amendola V (2011) Pyridinium/urea-based anion receptor: methine formation in the presence of basic anions. Org Biomol Chem 9(24):8276–8283CrossRefGoogle Scholar
  42. 42.
    Kraft A, Beck J, Krossing I (2011) Facile access to the pnictocenium ions Cp*ECl (+) (E = P, As) and (Cp*)(2)P (+): chloride ion affinity of Al(ORF)(3). Chem Eur J 17(46):12975–12980Google Scholar
  43. 43.
    Wenzel M, Light ME, Davis AP, Gale PA (2011) Thiourea isosteres as anion receptors and transmembrane transporters. Chem Commun 47(27):7641–7643CrossRefGoogle Scholar
  44. 44.
    Caltagirone C, Bazzicalupi C, Bencini A, Isaia F, Garau A, Lippolis V (2012) Anion recognition properties of pyridine-2,6-dicarboxamide and isophthalamide derivatives containing L-tryptophan moieties. Supramol Chem 24(2):95–100Google Scholar
  45. 45.
    Baggi G, Boiocchi M, Ciarrocchi C, Fabbrizzi L (2013) Enhancing the anion affinity of urea-based receptors with a Ru(terpy)(2)(2+) chromophore. Inorg Chem 52(9):5273–5283Google Scholar
  46. 46.
    Zhao HY, Leamer LA, Gabbai FP (2013) Anion capture and sensing with cationic boranes: on the synergy of Coulombic effects and onium ion-centred Lewis acidity. Dalton Trans 42(23):8164–8178CrossRefGoogle Scholar
  47. 47.
    Datta S, Halder M (2014) Effect of encapsulation in the anion receptor pocket of sub-domain IIA of human serum albumin on the modulation of pK(a) of warfarin and structurally similar acidic guests: a possible implication on biological activity. J Photochem Photobiol B Biol 130:76–85CrossRefGoogle Scholar
  48. 48.
    Elmes RBP, Yuen KKY, Jolliffe KA (2014) Sulfate-selective recognition by using neutral dipeptide anion receptors in aqueous solution. Chem Eur J 20(24):7373–7380CrossRefGoogle Scholar
  49. 49.
    Sekutor M, Mlinaric-Majerski K (2014) Adamantyl aminoguanidines as receptors for oxo-anions. Tetrahedron Lett 55(49):6665–6670CrossRefGoogle Scholar
  50. 50.
    Elmes RBP, Jolliffe KA (2015) Anion recognition by cyclic peptides. Chem Commun 51(24):4951–4968CrossRefGoogle Scholar
  51. 51.
    Pandian TS, Kang J (2015) Participation of aliphatic C-H hydrogen bonding in anion recognition. Tetrahedron Lett 56(28):4191–4194CrossRefGoogle Scholar
  52. 52.
    Tepper R, Schulze B, Jager M, Friebe C, Scharf DH, Gorls H, Schubert US (2015) Anion receptors based on halogen bonding with halo-1,2,3-triazoliums. J Org Chem 80(6):3139–3150Google Scholar
  53. 53.
    Amendola V, Bergamaschi G, Boiocchi M, Fusco N, La Rocca MV, Linati L, Lo Presti E, Mella M, Metrangolo P, Miljkovic A (2016) Novel hydrogen- and halogen-bonding anion receptors based on 3-iodopyridinium units. RSC Adv 6(72):67540–67549CrossRefGoogle Scholar
  54. 54.
    Amendola V, Bergamaschi G, Boiocchi M, Legnani L, Lo Presti E, Miljkovic A, Monzani E, Pancotti F (2016) Chloride-binding in organic–water mixtures: the powerful synergy of C-H donor groups within a bowl-shaped cavity. Chem Commun 52(72):10910–10913Google Scholar
  55. 55.
    Edwards SJ, Marques I, Dias CM, Tromans RA, Lees NR, Felix V, Valkenier H, Davis AP (2016) Tilting and tumbling in transmembrane anion carriers: activity tuning through n-alkyl substitution. Chem Eur J 22(6):2004–2011Google Scholar
  56. 56.
    Nehra A, Bandaru S, Yarramala DS, Rao CP (2016) Differential recognition of anions with selectivity towards F− by a calix 6 arene-thiourea conjugate investigated by spectroscopy, microscopy, and computational modeling by DFT. Chem Eur J 22(26):8903–8914Google Scholar
  57. 57.
    Qi J, Jinghan H, Chen JJ, Sun Y, Li JB (2016) Cyanide detection using azo-acylhydrazone in aqueous media with high sensitivity and selectivity. Curr Anal Chem 12(2):119–123Google Scholar
  58. 58.
    Wu HW, Chen YY, Rao CH, Liu CX (2016) Anion receptors based on CH donor group. Progr Chem 28(10):1501–1514Google Scholar
  59. 59.
    Molina P, Zapata F, Caballero A (2017) Anion recognition strategies based on combined noncovalent interactions. Chem Rev 117(15):9907–9972CrossRefGoogle Scholar
  60. 60.
    Kubik S (2010) Anion recognition in water. Chem Soc Rev 39(10):3648–3663CrossRefGoogle Scholar
  61. 61.
    Jenkins HDB, Krossing I, Passmore J, Raabe I (2004) A computational study of SbnF(5n) (n=1-4)—implications for the fluoride ion affinity of nSbF(5). J Fluor Chem 125(11):1585–1592CrossRefGoogle Scholar
  62. 62.
    Li SG, Dixon DA (2006) Molecular and electronic structures, Bronsted basicities, and Lewis acidities of group VIB transition metal oxide clusters. J Phys Chem A 110(19):6231–6244CrossRefGoogle Scholar
  63. 63.
    Brea O, Corral I, Mó O, Yáñez M, Alkorta I, Elguero J (2016) Beryllium-based anion sponges. Close relatives of proton sponges. Chem Eur J 22:18322–18325CrossRefGoogle Scholar
  64. 64.
    Montero-Campillo MM, Corral I, Mó O, Yáñez M, Alkorta I, Elguero J (2017) Beryllium-based fluorenes as efficient anion sponges. Phys Chem Chem Phys 19(34):23052–23059CrossRefGoogle Scholar
  65. 65.
    Christe KO, Dixon DA, McLemore D, Wilson WW, Sheehy JA, Boatz JA (2000) On a quantitative scale for Lewis acidity and recent progress in polynitrogen chemistry. J Fluor Chem 101(2):151–153CrossRefGoogle Scholar
  66. 66.
    Becke AD (1993) A new mixing of Hartree–Fock and local-density-functional theories. J Chem Phys 98(2):1372–1377Google Scholar
  67. 67.
    Lee C, Yang W, Parr RG (1988) Development of the Colle–Salvetti correlation-energy formula into a functional of the electron density. Phys Rev B 37(2):785–789Google Scholar
  68. 68.
    Curtiss LA, Redfern PC, Raghavachari K (2007) Gaussian-4 theory. J Chem Phys 126(8):12CrossRefGoogle Scholar
  69. 69.
    Bader RFW (1990) Atoms in molecules. A quantum theory. Clarendon, OxfordGoogle Scholar
  70. 70.
    Reed AE, Curtiss LA, Weinhold F (1988) Intermolecular interactions from a natural bond orbital, donor–acceptor viewpoint. Chem Rev 88(6):899–926Google Scholar
  71. 71.
    Biegler-König F, Schonbohm J, Bayles D (2001) Software news and updates—AIM2000—a program to analyze and visualize atoms in molecules. J Comput Chem 22(5):545–559Google Scholar
  72. 72.
    Glendening ED, Badenhoop JK, Reed AE, Carpenter JE, Bohmann JA, Morales CM, Weinhold F (2004) NBO6.G. Theoretical Chemistry Institute, University of Wisconsin, Madison.

Copyright information

© Springer-Verlag GmbH Germany, part of Springer Nature 2017

Authors and Affiliations

  1. 1.Stockholm University, Department of Organic Chemistry, Arrhenius LaboratoryStockholmSweden
  2. 2.Departamento de Química, Facultad de Ciencias, Módulo 13, and Institute of Advanced Chemical Sciences (IadChem)Universidad Autónoma de MadridMadridSpain
  3. 3.Instituto de Química Médica, CSICMadridSpain

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