Strength and Character of Halogen Bonds in Protein–Ligand

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Strength and Character of Halogen Bonds in Protein
Ligand Complexes Published as part of the Crystal Growth & Design virtual special issue on Halogen Bonding in Crystal Engineering: Fundamentals and Applications Kevin E. Riley*,† and Pavel Hobza†,‡ †

Institute of Organic Chemistry and Biochemistry, Academy of Sciences of the Czech Republic and Center for Biomolecules and Complex Molecular Systems, 166 10 Prague, Czech Republic ‡ Regional centre of Advanced Technologies and Materials, Department of Physical Chemistry, Palacky University, 771 46 Olomouc, Czech Republic

bS Supporting Information ABSTRACT: In this study we investigate the strength and character of eight halogen bonding interactions from six protein
ligand complexes. The halogen bonding complexes investigated here were selected because of their favorable halogen bond characteristics. Interaction energies of model systems derived from protein
ligand complexes are computed at the MP2/aug-cc-pVDZ level of theory, and the relative contributions of electrostatics and dispersion are estimated by computing ΔE(HF)/ΔE(MP2) ratios. The relationship between these ratios and DFT-SAPT Eelec/Edisp results is calibrated using smaller model systems in order to gain a qualitative understanding of the relative roles that electrostatics and dispersion play in these halogen bonds. Electrostatic potentials for the halogen bonding ligands are also generated in order to study the relationship between halogen bond strengths and halogen σ-hole size (and charge). It is found that the strength and character of the protein
ligand halogen bonds investigated here are strongly dependent on geometric factors and σ-hole characteristics. Many of the halogen bonds studied here, especially those with favorable geometric and electrostatic properties, are found to be of sufficient magnitude to make significant contributions to protein
ligand binding.

’ INTRODUCTION The phenomenon of halogen bonding has been recognized as being important in many areas of chemistry, biochemistry, and material science.1
5 Within the realm of biochemistry, halogen bonds have been observed within protein
ligand complexes,5
14 RNA
ligand complexes,15 and modified nucleic acid structures.16,17 Halogens are present in many ligands within protein
ligand complexes, and halogen bonds have been identified in complexes of native biological (hormones; see Figure 1),10 synthetic (pharmaceuticals),11 and environmental ligands11 with various protein receptors. By substituting various halogens with other atoms (or chemical groups), Hardegger, Diedrich, and co-workers have shown that the presence of a halogen bond has a strong effect on protein
ligand binding for inhibitors of human Cathepsin L.13 It is also shown in this study that the identity of the halogen bonding halogen and the halogen’s chemical environment strongly affect binding. Still, relatively little is known about the mechanisms by which halogen bonds contribute to the overall stabilities of protein
ligand complexes, and it is extremely important that studies seeking to elucidate the energetic and geometric roles that halogen bonds play in these complexes continue to be carried out. r 2011 American Chemical Society

Halogen bonds (X-bonds) are interactions that occur between an organic halide (-Cl, -Br, -I) and a bound electronegative atom (Lewis base), such as oxygen, nitrogen, or sulfur; the former of these is referred to as the X-bond donor while the latter is called the X-bond acceptor.4,18 Computational studies on halogen bonds indicate that these interactions depend strongly on both electrostatic and dispersion effects.19
21 Organic halogens, which generally carry an overall negative charge, have a very anisotropic charge distribution, with an equatorial ring of negative charge and a region of positive charge along the extension of the C
X bond, termed the σ-hole (see electrostatic potential figures below).4,18,22 It is the strong electrostatic interaction between the halogen σ-hole and the Lewis base that characterizes halogen bonds and is responsible for the strong degree of directionality exhibited by these interactions. The strongest X-bonds will tend to have C
X 3 3 3 Y (Y = O, N, S) angles approaching 180°, as this will result in the most direct interaction between the σ-hole and the Lewis base.21 The ideal halogen bond distance for most complexes Received: July 12, 2011 Revised: August 26, 2011 Published: September 14, 2011 4272

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Figure 1. Thyroid hormone in the thyroid receptor binding pocket. Among the two binding pocket interactions are two halogen bonds involving iodine (represented by dashed red lines).

is about 3.0 Å and generally decreases with increasing halogen size.19,23 The size and charge of the halogen σ-hole, along with the strength and character of the corresponding halogen bond, is highly variable and depends on two key factors, namely the identity of the halogen and the chemical environment in which the halogen is found.3,4,18,19,21 Larger halogens tend to form larger (more positive) σ-holes, leading to stronger interactions that are more electrostatic in nature.19 Fluorine, the smallest halogen, only forms σ-holes in very special circumstances and, for all practical purposes, is not regarded as a halogen bonding halogen. The size of a σ-hole is strongly affected by the electronegativity of the atoms (or chemical groups) that are located near the C
X bond. For example, the size of the σ-hole in F3CBr is much larger than that of H3CBr.19 The electron withdrawing effects of electronegative atoms tend to lower the overall negative charge on a halogen and, therefore, lead to more positively charged σ-holes. A large number of pharmaceutical ligands have halogens incorporated into their structures. There are several reasons for this, the most important of which are that halogens tend to increase the membrane permeability of small molecules while also increasing their catabolic stability.11 Protein data bank survey studies done in the past several years have identified many halogen bonds that occur within protein
ligand complexes.6,24 One of the principle findings in these surveys is that most of the halogen bonds in these complexes involved halogens bound to aromatic groups. Halogen bonds involving aromatically bound halogens are generally stronger than those of aliphatic halogens because aromatic moieties have electron withdrawing properties that lead to larger σ-holes. Among the halogen bonds identified in PDB surveys as having near-ideal geometric characteristics (C
X 3 3 3 Y angles in the range 165
180°), most involve aromatic moieties containing electron withdrawing groups, including structures with multiple halogens on one aromatic ring. It has also been found in PDB surveys that the halogen bond acceptors in a large majority of protein
ligand systems are backbone carbonyl oxygens.

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Generally speaking, there are four main factors that will have a strong effect on the strength and character of a halogen bonding interaction; these are the X-bond distance (d(X 3 3 3 O)), the X-bond angle (a(C
X 3 3 3 O)), the charge on the halogen σ-hole, and the magnitude of the dispersion contribution to the total interaction energy. The latter of these is largely determined by the halogen bond distance and the identity of the halogen. Halogen bonds involving iodine tend to have the largest contributions from dispersion because iodine is the largest halogen (considered here) and has the largest polarizability. However, as noted above, iodines will generally have larger σholes than bromines or chlorines and, thus, will also have larger contributions from electrostatic interactions. In this work we investigate the strengths and characters of eight halogen bonds found within six protein
ligand systems containing either bromine or iodine. In all cases, these halogens are bound to aromatic groups and four of the ligands (accounting for five interactions) have tetrahalo structures in which four halogen atoms are bound to one aromatic ring. Backbone carbonyl oxygens are the halogen bond acceptors in each of the complexes. Five of the investigated interactions represent near-ideal geometric halogen bonding configurations, with C
X 3 3 3 O angles greater than 165° and X 3 3 3 O distances smaller than 3.2 Å. The remaining three interactions have C
X 3 3 3 O angles of approximately 160°. It is not our intention in this work to determine binding energies and interaction characteristics for a large variety of halogen bonding interactions in biomolecules; this work should not be regarded as a survey of halogen bonds in protein
ligand interactions. In fact, most of the halogen bonds that have been selected for study here display particularly favorable characteristics: aromatically bound halogens, electron withdrawing substituents (as in the tetrahalo ligands), near-linear C
X 3 3 3 O angles, and X 3 3 3 O distances very close to 3.0 Å. Thus, many of the complexes studied here represent “best case scenarios”, in terms of the interaction energies that can be expected of halogen bonds in protein
ligand complexes. Ideally it would be possible to carry out DFT-SAPT analyses of these interactions in order to compute the relative contributions of dispersion, electrostatics, and induction to total attraction in each of the complexes. Unfortunately, computational demands make this type of analysis extremely difficult for such large systems. Another complication associated with the use of DFT-SAPT for large atoms is that the method cannot be used with relativistic corrections and can often yield inaccurate results when it is used along with pseudopotentials, which inherently account for relativistic effects. With no reliable way to account for the relativistic effect, the treatment of very large atoms, such as iodine, will generally be deficient. A classical way to gain some understanding of the relative roles played by dispersion and electrostatics in a noncovalent interaction is to compare the Hartree
Fock (HF) interaction energy with that of a correlated method, such as MP2 or CCSD(T). These types of analyses, usually comparing HF and MP2 results, have long been used to estimate the dispersion characteristics of an interaction.25 With a reasonably large basis set, Hartree
Fock describes electrostatic interactions quite well but does not account for correlation effects and, thus, fails to describe dispersion interactions.26 Correlated methods account for both intramolecular and intermolecular correlation effects, the former of which is relevant in the description of a molecule’s dipole moment and the latter of which chiefly describes the dispersion 4273

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Figure 3. Electrostatic potentials for iodine-containing ligands and model complexes used for interaction energy calculations. Red represents the most positive charges (those above a value of +12 kcal/mol), and blue represents negative charge values. Top, 1BSX-1; middle, 1BSZ-2; bottom, 3KXN.

Figure 2. Electrostatic potentials for bromine-containing ligands and model complexes used for interaction energy calculations. Red represents the most positive charges (those above a value of +12 kcal/mol), and blue represents negative charge values. From top to bottom: 1, 2OXY, 2, 2OXD; 3, 1ZOH-1; 4, 1ZOH-2; 5, 1GXZ.

interaction within the complex.27 Generally, the effects of intramolecular correlation are small compared to those of intermolecular correlation; thus, the main difference between a Hartree
Fock treatment and a correlated one should be in the description of the dispersion within the complex (to varying accuracy, depending on the particular method). Within this framework, the difference between the correlated and the HF interaction energies should largely be determined by the dispersion contribution to the interaction.26,28

’ COMPUTATIONAL METHODS In this study we compute MP2/aug-cc-pVDZ and HF/aug-ccpVDZ interaction energies for eight halogen bonding interactions, involving bromine and iodine, from six protein
ligand complexes. The aug-cc-pVDZ-PP basis set, which contains pseudopotentials and implicitly describes relativistic effects, is used to describe bromine and iodine. This is important for large atoms such as these halogens (especially iodine), as relativistic effects can play a large role in their electronic behavior. All binding energies are computed using counterpoise corrections to account for basis set superposition error (BSSE).29 The ratios of the HF and MP2 interaction energies (ΔE(HF)/ΔE(MP2)) are used to estimate the relative contributions from electrostatic and dispersion forces. These calculations give insight into the strengths of the halogen bonds in these complexes as well as their character.

In most cases, the model systems used for the calculations consist of a ligand complexed with the protein backbone carbonyl group bound to the neighboring backbone nitrogen, with hydrogens added to complete the valence shells (Figures 2 and 3). It should be pointed out that using a larger model of the halogen bond acceptor molecule might more accurately reflect the charge characteristics of the carbonyl oxygen and, thus, provide more accurate interaction energies; however, the effect of using smaller models should be rather small compared to the accuracy of the MP2/aug-cc-pVDZ method. The small model acceptor molecules were chosen to minimize the effects of secondary (non-Xbond) interactions that can occur in the complexes. In the case of the thyroid hormone, which consists of two iodine-containing aromatic rings and a carboxylic acid group, model systems are constructed from each of the aromatic rings, along with the corresponding halogen bond acceptor (Figure 4). It was necessary to carry out this division of the ligand in order to isolate the two X-bond interactions. The structures of the halogen bonding complexes used in this study are derived from crystal structures of protein
ligand complexes. Brief descriptions of these protein
ligand complexes are given in Table 1. The structure of each of the protein
ligand complexes was optimized using the PM6-DX dispersion augmented semiempirical method, which includes empirical corrections for halogen bonding interactions.30,31 Optimizations were carried out using the FIRE algorithm.32 Here each interaction is labeled according to its PDB ID code, in the cases of 1ZOH and 1BSX, which exhibit two halogen bonding interactions, the strongest halogen bond is labeled with a
1 and the weakest with a
2 (ie 1ZOH
1 and 1ZOH
2). PM6-DX optimizations were carried out using the MOPAC semiempirical code.33 HF, MP2, and DFT-SAPT computations were done using MOLPRO 2010.34 As noted above, the ΔE(HF)/ΔE(MP2) ratios can serve as indicators of the relative contributions of electrostatics and 4274

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dispersion to the stabilization of a complex. These values can be made more meaningful if they are compared to DFT-SAPT results obtained for smaller model systems. DFT-SAPT can be used to decompose the interaction energy into physically meaningful contributions, namely those from electrostatics, exchange, dispersion, and induction.35,36 The DFT-SAPT interaction energy is given as Eint ¼ Eelec þ Eexch þ Edisp þ Eind þ δHF

giving a general impression of the relative contributions of electrostatic and dispersion forces in these complexes. Table 2 gives the HF, MP2, and DFT-SAPT interaction energies for the three model complexes, along with DFT-SAPT decomposition terms, the Eelec/Edisp ratio, and the ΔE(HF)/ ΔE(MP2) ratio. A detailed analysis of the DFT-SAPT terms shown here will be reserved for a future article, which will also include DFT-SAPT analyses of other halogen bonding systems. The main points of interest for the purposes of this discussion are the DFT-SAPT electrostatic and dispersion terms, which are the most dominant attractive terms, and the Eelec/Edisp ratios. It can be seen that the electrostatic component, as a percentage of the total attractive interaction in these complexes, increases with successive fluorine substitutions, accounting for 40% of the attraction in bromobenzene 3 3 3 acetone and 51% of the

ð1Þ

The quantity δHF represents higher-order effects, which are usually mainly inductive.37 The Edisp and Eind terms also contain contributions from exchange-dispersion and exchange-induction effects, respectively. In this work DFT-SAPT computations have been carried out using the LPBE0AC potential along with the aug-cc-pVTZ basis set. The density fitting procedure was employed to significantly reduce the computational cost of these calculations.36 Here, in order to “calibrate” the HF/MP2 comparison technique, we have carried out DFT-SAPT computations, as well as HF/MP2 comparisons on three model complexes, bromobenzene 3 3 3 acetone, 2,6-diflourobromobenzene 3 3 3 acetone, and pentafluorobromobenzene 3 3 3 acetone (Figure 5). The successive substitution of aromatic fluorines in these systems results in the bromine σ-hole becoming larger and more positive, leading to interactions that are stronger and more electrostatic in nature. The most straightforward way to relate the quantities obtained using these two techniques is to compare the DFT-SAPT Eelec/ Edisp ratio to ΔE(HF)/ΔE(MP2). Of course, as the ΔE(HF)/ ΔE(MP2) values have been calibrated to only three complexes, the estimated Eelec/Edisp values should be seen as being purely qualitative in nature. Nonetheless, the values can still be useful in

Figure 5. Model complexes used for calibration of the ΔE(HF)/ΔE(MP2) ratios to DFT-SAPT Eelec/Edisp results. From left to right: bromobenzene 3 3 3 acetone, 3,5-difluorobromobenzene 3 3 3 acetone, and pentafluorobromobenzene 3 3 3 acetone.

Table 2. Comparison of DFT-SAPT/aug-cc-pVTZ, HF/augcc-pVDZ, and MP2/aug-cc-pVDZ Results, as well as DFTSAPT Eelec/Edisp and ΔE(HF)/ΔE(MP2) Ratios for Model Complexesa F subsituents

Figure 4. Separation of thyroid hormone (1BSX) into two models used in interaction energy computations.

a

0

2

5
6.06 (51%)

Eelec


2.91 (40%)


4.02 (46%)

Eexch

5.15

6.17

8.03

Eind


0.60 (8%)


0.76 (9%)


1.08 (9%)

Edisp


3.33 (46%)


3.46 (39%)


3.99 (33%)

dHF


0.41 (6%)


0.54 (6%)


0.81 (7%)

Eint


2.09


2.62


3.90

HF MP2

0.29
2.23


0.50
2.77


1.70
4.08

ΔE(HF)/ΔE(MP2)


0.13

0.18

0.42

SAPT Eelec/Edisp

0.88

1.16

1.52

All values in kcal/mol.

Table 1. Description of Protein
Ligand Interactions Studied in This Work PDB-ID

X=

description

ligand properties

2OXY

Br

CK2 protein kinase inhibitor

tetrabromo arrangement with one X-bond

2OXD

Br

CK2 protein kinase inhibitor

tetrabromo arrangement with one X-bond

1ZOH 1GXZ

Br Br

CK2 protein kinase inhibitor ADP-ribosyl transferase ligand

tetrabromo arrangement with two X-bonds 5-bromonicotinamide

1BSX

I

native thyroid receptor ligand

bicyclic with one X-bond per ring

3KXN

I

CK2 protein kinase inhibitor

tetraiodo arrangement with one X-bond

4275

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Figure 6. Calibration of ΔE(HF)/ΔE(MP2) ratios to DFT-SAPT Eelec/Edisp ratios for model complexes.

attraction in pentafluorobromobenzene 3 3 3 acetone. The relative contribution from dispersion interactions follows the opposite trend and decreases as hydrogens are substituted with fluorines, going from 46% (bromobenzene 3 3 3 acetone) to 33% (pentafluorobromobenzene 3 3 3 acetone). Both of these trends are reflected in the Eelec/Edisp ratios as well as the ΔE(HF)/ΔE(MP2) ratios, which both increase with successive fluorine substitutions. It should be noted that, in the case of bromobenzene 3 3 3 acetone, the HF binding energy is positive (resulting in a negative ΔE(HF)/ΔE(MP2) ratio), indicating a repulsive HF interaction. Within the range of values considered here, the relationship between the Eelec/Edisp and ΔE(HF)/ΔE(MP2) ratios is approximately linear, and in a plot of Eelec/Edisp as a function of ΔE(HF)/ΔE(MP2), the y-intercept is approximately one (Figure 6). The linear relationship is given as: Eelec =Edisp ¼ 1:17ΔEðHFÞ=ΔEðMP2Þ þ 1:01

ð2Þ

Thus, a negative ΔE(HF)/ΔE(MP2) value indicates an interaction in which Edisp is responsible for more than half of the attraction. Again following the linear relationship, a ΔE(HF)/ ΔE(MP2) value of +0.42 indicates an electrostatic interaction that is greater than the dispersion contribution by a factor of 1.5 (i.e. Eelec/Edisp = 1.5). Equation 2 will be used to estimate the Eelec/Edisp for the eight protein
ligand halogen bonds considered in this study. To gain insight into the nature and directionality of the halogen bonding interactions being considered here, we have generated electrostatic potentials for each of the halogen bond donating molecules. These electrostatic potentials have been computed on the molecular “surfaces”, taking this to be an outer contour of the electronic density, generally the 0.001 au (electrons/bohr3) surface, as proposed by Bader et al.38 The most positive and most negative values of the potentials (local maxima and minima) are referred to as VS,max and VS,min.39 There may be several of each on a given molecular surface, and the VS,max value on a halogen bonding halogen directly indicates the potential for that halogen to participate in a strong electrostatic interaction. Here electrostatic potentials are computed at the B3LYP/6-31+G* level of theory using the WFA program.39 In the description of electrostatic potentials, it is not necessary to use methods that treat correlation effects (as in the case of interaction energies). Here we compute electrostatic potentials at a level of theory similar to that used in many other works.18,22,40

Figure 7. Electrostatic potential for 2OXY ligand and corresponding schematic of VS,max values for each of the four σ-holes in the molecule.

’ RESULTS AND DISCUSSION As has been noted above, a halogen’s chemical environment has a strong effect on the size and charge of the halogen σ-hole. Figure 7 shows the electrostatic potential for the 2OXY ligand along with the VS,max values for each of the four bromines it contains. Here it can be seen that the σ-hole on the bromine closest to the aromatic N
H group (VS,max = 20.5 kcal/mol) is much larger and more positive than the neighboring bromine’s σhole (VS,max = 14.3 kcal/mol). This trend of decreasing σ-hole size continues in the molecule, with the next bromines having VS, max values of 11.9 and 10.5 kcal/mol, respectively. Similar trends are observed for the other tetrabromine ligands examined in this work. Here we will not discuss the differences in σ-hole sizes in these types of molecules in great detail, and we will reserve such analyses for future studies. Figure 2 shows the electrostatic potentials for each of the bromine-containing halogen bond donor molecules considered here as well as the interactions of the ligands with the halogen bond acceptor models. It can be seen that the largest σ-hole occurs for the 1ZOH-1 bromine (VS,max = 17.3 kcal/mol) and the smallest σ-hole occurs for the 1ZOH-2 bromine (VS,max = 9.2 kcal/mol). Focusing on the first five entries in Table 3 (X-bonds involving bromine), it can be seen that there are three interactions with near-ideal halogen bond distances and angles; these are 2OXY (3.04 Å, 177.2°), 2OXD (3.01 Å, 177.9°), and 1ZOH-1 (3.05 Å, 175.1°). Here there are also two interactions with less than ideal halogen bond distances and/or angles, namely 1ZOH-2 (3.49 Å, 159.4°) and 1GXZ (3.06 Å, 157.2°). The MP2 interaction energies for these complexes largely reflect the geometric characteristics of halogen bonds, with the near-ideal geometry complexes having interaction energies much larger (
3.52 to
4.28 kcal/mol) than those of the far from ideal geometries (
1.97 to
2.30 kcal/mol). Among these structures, the interaction energy increases (in absolute value) with σ-hole charge, as indicated by the VS,max values, with the strongest interaction occurring for 1ZOH-1 (VS,max = 17.3 kcal/mol, ΔE(MP2) =
4.28 kcal/mol). The ΔE(HF)/ΔE(MP2) and estimated Eelec/Edisp values for the bromine-containing complexes also largely conform to the expected behavior, with the most electrostatic interactions corresponding to the bromines with the most positively charged σ-holes. Among the three near-ideal geometry complexes, it can be seen that the ΔE(HF)/ΔE(MP2) and estimated Eelec/Edisp values consistently increase with increasing VS,max values, with the most electrostatic interaction having an Eelec/Edisp value of 1.58, which indicates that this halogen bond is strongly dominated by 4276

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Table 3. Characteristics of Protein
Ligand Halogen Bonds Studied Hereb X-bond

X=

d(X 3 3 3 O)

a(C-X 3 3 3 O)

VS,max

ΔE(MP2)

ΔE(HF)

ΔE(HF)/ΔE(MP2)

Eelec/Edispa

2OXY

Br

3.04

177.2

11.9


3.52


1.04

0.30

1.36

2OXD

Br

3.01

177.9

15.9


3.97


1.64

0.41

1.49

1ZOH-1

Br

3.05

175.1

17.3


4.28


2.08

0.48

1.58

1ZOH-2

Br

3.49

159.4

9.2


1.97


0.40

0.20

1.25

1GXZ

Br

3.06

157.2

13.1


2.30

0.22


0.10

0.90

1BSX-1

I

3.01

173.8

21.3


4.47


1.73

0.39

1.46

1BSX-2

I

3.13

161.8

25.9


4.36


1.66

0.38

1.46

3KXN

I

3.06

169.0

29.4


5.40


2.86

0.53

1.63

Qualitative Eelec/Edisp values estimated from ΔE(HF)/ΔE(MP2) values. b Distances in Å, angles in degrees. VS,max values, MP2 interaction energies, and HF interaction energies are in kcal/mol. a

electrostatic forces. This trend is not followed by the far from ideal geometry complexes, for which the 1ZOH-2 interaction is predicted to be more electrostatic in nature than the 1GXZ interaction. The halogen bond angles for 1ZOH-2 and 1GXZ are very similar (159.4° and 157.2°, respectively), while 1ZOH-2 has a significantly larger halogen bond distance (3.49 Å and 3.06 Å, respectively) and a smaller σ-hole (9.2 and 13.1 kcal/mol, respectively). Nonetheless, the ΔE(HF)/ΔE(MP2) value for 1ZOH-2 (0.20) is larger than that of 1GXZ (
0.10), indicating that electrostatics are slightly dominant in 1ZOH-2 and slightly less important than dispersion in 1GXZ. The apparent reason for this contradiction is a secondary interaction in 1ZOH-2 that occurs between an X-bond acceptor NH hydrogen and a region of negative charge located between the two nearest bromines. This is the only complex considered here for which such a secondary interaction involving the X-bond acceptor NH2 group is possible, indicating that this secondary interaction is likely responsible for this contradictory result. Figure 3 shows the electrostatic potentials, as well as the model interaction complexes, for the three systems containing iodine. Here it can clearly be seen that, in all three cases, iodine has a larger σ-hole than any of the bromine-containing systems. This is an expected result, as it is known that iodine generally forms larger σ-holes than bromine. The computed VS,max results, given in Table 3, also indicate σ-holes that are more positively charged than those of the halogen bonding bromines, with values in the range from 21.3 to 29.4 kcal/mol. Here only the 1BSX-1 interaction has a halogen bond angle greater than 170°, with a value of 173.8°, which is slightly smaller than the halogen bond angles for the near-ideal bromine cases. The smallest halogen bond angle of 161.8° occurs for the 1BSX-2 complex, which also exhibits the longest halogen bond distance of 3.13 Å. Interestingly, the MP2 interaction energies for both 1BSX-1 and 1BSX-2 are very similar (
4.47 and
4.36 kcal/mol, respectively). This can most likely be attributed to compensatory effects, as the 1BSX-1 X-bond exhibits better geometric parameters while the 1BSX-2 σ-hole is larger and more positively charged. The size of the σ-hole probably plays an important role here, as a larger σ-hole will allow for strong electrostatic interactions across a broader range of C
X 3 3 3 O angles. The strongest interaction occurs for the 3KXN system, with an MP2 interaction energy of
5.40 kcal/mol, which is not surprising considering the large VS,max value of the iodine (29.4 kcal/mol) and the reasonably large C
I 3 3 3 O angle (169.0°). The MP2 interaction energies for all three iodine-containing complexes are larger (in absolute value) than all interactions for bromine-containing complexes.

1BSX-1 and 1BSX-2 have very similar ΔE(HF)/ΔE(MP2) values (0.39 and 0.38, respectively), which correspond to the same SAPT Eelec/Edisp value (1.46). As in the case of the MP2 binding energies, the fact that these two interactions have similar characteristics can likely be explained by a compensatory effect. With a C
I 3 3 3 O angle of 173.8, the 1BSX-1 interaction represents a better alignment of the iodine σ-hole with the oxygen acceptor, while the larger, more positive, σ-hole of 1BSX2 will tend to form stronger interactions that are more electrostatic in nature. It should be noted that the Eelec/Edisp values for these interactions are (slightly) smaller than those of the two most electrostatic interactions involving bromine (2OXD and 1ZOH-1). This may seem to be a contradictory result, as all of the iodine σ-holes are larger than those of the bromines considered here. However, it should be kept in mind that the magnitude of attractive dispersion interactions should also be greater for interactions involving iodine than for those involving bromine. The result of the increase in the relative contribution of dispersion for an iodine-containing system is a lower Eelec/Edisp ratio than would be expected for a similar interaction involving bromine. Not surprisingly, the largest ΔE(HF)/ΔE(MP2) value of 0.53 occurs for the 3KXN complex; this corresponds to an estimated SAPT Eelec/Edisp value of 1.63, indicating a halogen bond that is strongly dominated by electrostatic forces. The most important conclusion to be drawn from the results presented here is that halogen bonds are strong interactions that play an important role in the binding of small molecule ligands to proteins. The MP2/aug-cc-pVDZ interaction energies for the eight halogen bonds examined in this study are in the range from
1.97 kcal/mol to
5.40 kcal/mol, with the strongest interactions generally occurring for halogen bonds with large σ-holes and C
X 3 3 3 O angles in the range from 165 to 180°. As expected, the strongest halogen bonds are those that involve iodine as the halogen bond donor, as iodine can form larger σ-holes than bromine and, because of its larger size, also forms stronger dispersion contacts. Among the eight X-bond interactions studied here, there are five that have interaction energies greater than approximately
4.0 kcal/mol (in absolute value), a value that is comparable with the interaction energies of common hydrogen bonds; for example, the interaction energy of the water dimer is approximately
5 kcal/mol. ΔE(HF)/ΔE(MP2) and corresponding estimated SAPT Eelec/ Edisp values obtained for these interactions confirm the halogen bonding behavior that has been observed on model system calculations, namely that the electrostatic/dispersion nature of a halogen bond is highly dependent on the size of the σ-hole and the C
X 3 3 3 O angle. Here only one interaction was found 4277

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Crystal Growth & Design to depend more strongly on dispersion than electrostatics (1GXZ; secondary interactions in 1ZOH
2 probably mask a dominant dispersion contribution). Contributions from electrostatics are estimated to clearly be greater than those of dispersion for six of the complexes studied here, with only one of these (1BSX
2) having a C
X 3 3 3 O angle smaller than 165°. The large, very positively charged, σ-hole found in the 1BSX
2 complex is likely responsible for the strong electrostatic attraction.

’ ASSOCIATED CONTENT

bS

Supporting Information. Cartesian coordinates (.xyz files) for all of the model complexes considered here. This material is available free of charge via the Internet at http:// pubs.acs.org.

’ AUTHOR INFORMATION Corresponding Author

*E-mail: [email protected]. Fax: +420 220 410 320.

’ ACKNOWLEDGMENT This work was a part of Research Project No. Z40550506 of the Institute of Organic Chemistry and Biochemistry, Academy of Sciences of the Czech Republic, and was supported by Grants No. LC512 and MSM6198959216 from the Ministry of Education, Youth and Sports of the Czech Republic. It was also supported by the operational program Research and Development for Innovations of European Social Fund (CZ.1.05/2.1.00/ 03.0058). The support of Praemium Academiae, Academy of Sciences of the Czech Republic, awarded to P.H. in 2007 is also acknowledged. ’ REFERENCES (1) Metrangolo, P.; Neukirch, H.; Pilati, T.; Resnati, G. Acc. Chem. Res. 2005, 38, 386. (2) Metrangolo, P.; Resnati, G. Halogen Bonding: Fundamentals and Applications; Springer: Berlin, 2008. (3) Murray, J. S.; Riley, K. E.; Politzer, P.; Clark, T. Aust. J. Chem. 2010, 63, 1598. (4) Politzer, P.; Lane, P.; Concha, M. C.; Ma, Y. G.; Murray, J. S. J. Mol. Model. 2007, 13, 305. (5) Lu, Y. X.; Shi, T.; Wang, Y.; Yang, H.; Yan, X.; Luo, X.; Jiang, H.; Zhu, W. J. Med. Chem. 2009, 52, 2854. (6) Auffinger, P.; Hays, F. A.; Westhof, E.; Ho, P. S. Proc. Natl. Acad. Sci. U.S.A. 2004, 101, 16789. (7) Battistutta, R. Cell. Mol. Life Sci. 2009, 66, 1868. (8) Liu, L. J.; Baase, W. A.; Matthews, B. W. J. Mol. Biol. 2009, 385, 595. (9) Matter, H.; Nazare, M.; Gussregen, S.; Will, D. W.; Schreuder, H.; Bauer, A.; Urmann, M.; Ritter, K.; Wagner, M.; Wehner, V. Angew. Chem., Int. Ed. 2009, 48, 2911. (10) Valadares, N. F.; Salum, L. B.; Polikarpov, I.; Andricopulo, A. D.; Garratt, R. C. J. Chem. Inf. Model. 2009, 49, 2606. (11) Parisini, E.; Metrangolo, P.; Pilati, T.; Resnati, G.; Terraneo, G. Chem. Soc. Rev. 2011, 40, 2267. (12) Voth, A. R.; Ho, P. S. Curr. Top. Med. Chem. 2007, 7, 1336. (13) Hardegger, L. A.; Kuhn, B.; Spinnler, B.; Anselm, L.; Ecabert, R.; Stihle, M.; Gsell, B.; Thoma, R.; Diez, J.; Benz, J.; Plancher, J. M.; Hartmann, G.; Banner, D. W.; Haap, W.; Diederich, F. Angew. Chem., Int. Ed. 2011, 50, 314. (14) Voth, A. R.; Khuu, P.; Oishi, K.; Ho, P. S. Nature Chem. 2009, 1, 74.

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