Interactions Involving Polar Groups

Polar groups decrease lipophilicity by characteristic increments. When two or more such groups are present, the solute is often found to be more lipophilic than calculated by the simple addition of increments, implying that two or more polar groups may interact to prevent each other from expressing its full polarity. This phenomenon is very well known and amply documented in the literature [7-9], There is thus little need for us to go into too many details in describing such intramolecular effects, and we shall simply restrict ourselves to discriminating between the various mechanisms by which polar groups may interact intramolecularly to increase lipophilicity. However, because such mechanisms act seldom alone, their discrimination is not always straightforward.

4.3.2.1 Proximity Effects Between Two Neutral Polar Groups

The fragmental system of Rekker afforded the first incremental method of calculation derived from and applicable to aliphatic moieties and molecules. In this system, a limited number of correction factors are necessary to take into account intramolecular interactions such as electronic conjugation and proximity effects. These occur between polar groups defined as being electronegative functionalities. In the 1979 version of the system [7], a correction factor of 4- 0.84 must be added to the sum of fragmental values when such groups are separated by one sp3-carbon. The correction factor decreases to + 0.56 when the two groups are separated by two carbons, and becomes statistically nonsignificant for three carbons.

The point of relevance in Rekker's fragmental system is that the correction for proximity is independent of the chemical nature of the polar groups, which may be halogens, H-bond acceptors, or H-bond donors/acceptors. This implies that in Rekker's fragmental system internal H-bond formation is not regarded as playing an explicit role. A simple molecular explanation for such proximity effects is to consider each polar group as surrounded by a hydration sphere, proximity bringing these spheres to overlap and loose influence. The Molecular Lipophilicity Potential (MLP), which is to be discussed in chapter 12, allows a simpler picture to emerge. In this approach, the polarity on the solvent-accessible molecular surface is computed for its intensity and extension. The former is seen to increase, and the latter to decrease, when two polar groups become proximal. However, the deep reason for this proximity effect remains to be understood.

While empirically very useful, Rekker's assumption is an oversimplification for two major reasons. First, there is a problem of intensity of effects since it is difficult to understand how different polar groups could contribute identically to proximity effects. The second reason is related to the first, namely that the formation of internal H-bonds should be given explicit consideration to obtain more precise log P estimates.

It is a feature of the fragmental system of Hansch and Leo [8] that it acknowledges the formation of internal H-bonds and classifies polar groups into two classes, namely H-bond-forming (designated as H) and non-H-bond-forming (i.e., halogens, designated as S). As a result, three types of proximity effects are recognized, namely those resulting from S/S, H/S and H/H interactions.

Multiple halogenation on the same or adjacent carbon atoms, a typical example of S/S interaction, results in higher log P values than predicted by simple additivity. This is assumed to be due to the localized dipoles being partly shielded from water by the neighbouring halogens. In the fragmental system of Hansch and Leo (version 1979), the correction factors for N gemical halogens are + 0.30 N, + 0.53 N and + 0.72 N for N = 2, 3 or 4, respectively [8]. In the case of vicinal halogens, the correction factors are -I- 0.28 (N - 1). These numbers are average values that neglect other structural factors such as conformational behavior and electronic distribution. Nevertheless, they are clearly significant and testify of the significant of polar proximity effects not involving H-bonds.

In the fragmental system of Hansch and Leo, distinct correction factors are used to account for proximity effects between H-bond-forming groups. These factors, which are always positive, depend on the nature of the polar groups and decrease with increasing number of intervening carbon atoms (1, 2, or 3). Like in Rekker's system, these factors do not incorporate internal H-bonds, which are factorized separately as discussed below. Hence, the system of Hansch and Leo also recognizes the decrease in hydrophilicity due to the mere proximity of polar groups.

4.3.2.2 Internal H-Bonds

The most important intramolecular electrostatic interactions affecting lipophilicity are ionic bonds (see section 4.3.2.3), H-bonds discussed here, and perhaps also dipoledipole interactions. As stated above, the system of Hansch and Leo allows the average influence of H-bonds on log P to be estimated [8]. Thus, intramolecular H-bonds involving oxygen or nitrogen receive a correction factor of + 1.0 and +0.60, respectively. This is a rather marked effect, but the difficulty when calculating a log P value is to decide whether such a bond exists or not in a given solute. Indeed, a correction factor for H-bond can be introduced a priori (subject to experimental verification) based on knowledge or expectation, or a posteriori if a discrepancy is found between measured and calculated log P.

Water being a dipole with strong H-bond donor and acceptor properties, it will interact electrostatically with polar groups in a solute and prevent it from forming internal electrostatic bonds. Water-saturated octanol (approximately four molecules of octanol for one of water) is a relatively polar solvent with a H-bond acceptor basicity as good as that of water, and a H-bond donor acidity markedly smaller than that of water. As a result, the tendency of solutes to form internal H-bonds is usually comparable in octanol and in water. Thus, octanol is by far not the best solvent to observe the formation of internal H-bonds, and more generally hydrophilic folding (see section 4.3.2.4). In contrast, apolar solvents such as alkanes or poorly polar solvents such as dibutyl ether (which dissolves about 0.1 % water) will strongly favor internal H-bonds and more generally hydrophilic collapse.

Thus, a thermodynamic study of the partitioning of isomeric and homologous pyri-dylalkanamides (see Fig. 2) afforded some insight into the underlying mechanisms [14]. In dibutyl ether/water, a single mechanism prevailed for the partitioning of all solutes, and the solutes able to form internal H-bonds (i.e., the 2-pyridyl derivatives)

dbe

oct

h20

h2o

Figure 3. Mechanistic difference in the partitioning of 2-pyridylalkanamides in the systems dibu-tyl ether/water (DBE/H20) and n-octanol/watcr (0CT/H20). In DBE, the solutes appear to exist excusively in an internally H-bonded conformation that masks in part the hydrophobic segment of the side-chain. In octanol, the solutes are in equilibrium between a small population of internally H-bonded conformers and a predominant population of extended conformers [14].

Figure 3. Mechanistic difference in the partitioning of 2-pyridylalkanamides in the systems dibu-tyl ether/water (DBE/H20) and n-octanol/watcr (0CT/H20). In DBE, the solutes appear to exist excusively in an internally H-bonded conformation that masks in part the hydrophobic segment of the side-chain. In octanol, the solutes are in equilibrium between a small population of internally H-bonded conformers and a predominant population of extended conformers [14].

were more lipophilic than expected. This indicated that the 2-pyridylalkanamides existed exclusively as internally H-bonded conformers in dibutyl ether. In contrast, all evidence from octanol/water partition coefficients indicated that in this solvent the 2-pyridylalkanamides existed in a conformational equilibrium with only a small population of internally H-bonded conformers (Fig. 3) [14].

coo o o

na 2

highly localized intermediate delocalized

Figure 4. A schematic representation of zwitterions ranging from the highly localized (e.g. (3- and y-amino acids) to the delocalized (e.g. guanidinium/enolates).

4.3.2.3 The Case of Zwitterions

Zwitterions represent a particular and as yet insufficiently explored type of solutes. A priori, their intramolecular and intermolecular interactions differ from those of other solutes, although one must be aware that the differences between highly polarized nonionic solutes and zwitterions may be more quantitative than qualitative, especially for delocalized zwitterions. Indeed, we have come to feel that there is a need to distinguish between two major types of zwitterions, namely the more usual ammonium-carboxylates, and the more delocalized guanidinium or amidinium enolates. Figure 4 presents a continuum between (3- and y-amino acids and some highly delocalized structures as for example found in oxicams and other nonsteroidal antiinflammatory agents

In Rekker's fragmental system, the aliphatic and aromatic carboxylate fragments (COO") receive an incremental value of — 5.00 and — 4.13, respectively [7]. The case of the ammonium group is less clear, but a value of — 3.73 has been proposed for the aliphatic-NH3+ fragment [16]. By simply adding fragments, one would arrive at a predicted log P value for zwitterionic glycine ( OOC-CH2-NH3+) of — 8.21, when the actual value is — 3.00 [17]. The difference between the predicted and experimental values clearly demonstrates that the two charged groups in a-amino acids interact strongly, and gives a fair estimate of the importance of this interaction. The nature of this interaction is certainly a dual one, involving partial neutralization via derealization across the sp3-carbon (through-bond interaction), plus an internal ionic bond (through-space interaction) which further contributes to partial neutralization.

Figure 5. The structure of zwitterionic piroxicam as existing in the pH range 2-5.

In zwitterions such as piroxicam in the pH range 2-5 (Fig. 5), the two charges are formally separated by 5 atoms, but molecularorbital calculations reveal that due to marked derealization the effective distance between the centroid of positive and negative regions becomes less than5 i. This appears to result in a marked partitioning of the zwitterion into octanol [15].

4.3.2.4 Hydrophilic Collapse

Hydrophilic collapse is defined as a conformational change by which a solute maximizes the number and strength of internal electrostatic bonds (mainly H-bonds) and thus partly masks some of its polar groups from the solvent. The drive for hydrophilic collapse comes from a nonpolar solvent, the solute hiding its polar groups away from this nonpolar solvent in order to become less polar and resemble that solvent. As such, hydrophilic collapse is the opposite of hydrophobic collapse discussed in section 4.3.3.3, and alone or together with the latter may account for a chameleonic behavior (see section 4.4.4).

An example of hydrophilic collapse is offered by cyclosporin A (CsA), an immunosuppressive cyclic undecapeptide widely used in clinical organ transplantation. In water, CsA exists as a mixture of conformers characterized by H-bonding groups (in Abu2, Val5, Ala7 and Ala8) pointing away from the ring, i.e., towards the solvent [18]. This is in fact the active conformation of CsA as bound to cyclophilin. In apolar solvents, the conformational state of CsA is very different, being characterized by four major internal H-bonds (Abu2-to-Val5, Val5-to-Abu2, Ala7-to-MeValn, and Ala8-to-MeLeu6). Thus, CsA in apolar solvents turns a number of its polar groups towards the interior of the ring. The driving force for the creation of this polar interior is the formation of the intramolecular H-bonds.

Table 2. Partition coefficients (log P) and H-bonding capacity [A(log Poctanoi-heptane)] of model -solutes and cyclosporin A [18]

Solute

log Pocanol

log Pheptane

Alog P

Not forming internal H-bonds

phenol

1.46

-0.82

2.22

p-nitrophenol

1.77

-2.11

3.88

benzamide

0.65

-2.45

3.10

/j-fluorobenzamide

0.96

-2.34

3.25

acetanilide

1.16

-1.54

2.70

p-fluoroacctanilide

1.47

-1.57

3.04

cyclo(Phe-Phe)

1.59

<-3.0

>4.5

cyclo(Trp-Tyr)

1.05

<-3.0

>4.1

Forming internal H-bonds

o-nitrophenol

1.68

1.04

0.64

o-fluorobenzamide

0.64

-1.47

2.11

o-fluoroacetanilide

0.96

-0.69

1.65.

CyclosporinA

2.92

1.40

1.52

The conformational behavior of CsA is reflected in its partitioning. As shown in Table 2, solutes unable to form internal H-bonds have Alog P values (i.e., log f0ctanoi minus log /\eptane, a measure of the H-bonding capacity) of > 2, whereas for those able to form internal H-bonds the value is < 2 [18]. These numbers reflect the fact, already discussed in section 4.3.2.2, that hydrophilic collapse and particularly the formation of internal H-bonds are favored in apolar solvents (e.g., alkanes) significantly more than in solvents of low polarity (e.g., octanol).

4.3.2.5 Proximity Effects Between Polar and Nonpolar Groups

In sections 4.3.2.1 and 4.3.2.3, we have seen how the proximity of two polar groups in a solute decreases its expected hydrophilicity, i.e., produces a higher than expected lipophilicity. The interpénétration of hydration spheres, or a decrease of the polar molecular surface, are two complementary models to explain the many observations of this type and visualize their mechanism.

Interestingly, the same pictorial models allow another important phenomenon to be understood, namely the decrease in their hydrophobic increment experienced of apolar moieties in the proximity of highly polar groups. A highly illustrative example is provided by amino acids, where the carboxylate and ammonium groups decrease the hydrophobicity of neighboring CH2 and CH3 units in a distance-dependent manner [17]. When examining the log D of a-amino acids (Fig. 6) determined at isoelectric pH (i.e., the log P of the zwitterions), it was found that the hydrophobic increment of the CH2 groups did not increase as predicted. Indeed, the increments in the series R = H (glycine), R = CH3 (alanine), R = CH2CH3 (a-aminobutyric acid), R = CH2CH2CH3 (norvaline) and R = CH2CH2CH2CH3 (norleucine) were 0.23, 0.24, 0.42, and 0.57. Thus, only the fourth CH2 unit could express a full hydrophobicity (0.57), suggesting that the first three CH2 units are partly masked from the solvent by the polar groups [17].

A comparable observation was made with zwitterionic co-amino acids (Fig. 6), with the first six homologs (n = 1 to 6) having log P values between - 3.0 and - 3.1. Only the seventh CH2 group did contribute to an increased lipophilicity (log P for n = 7: - 2.55) [17]. Here, the results indicate that the first three CH2 groups attached to the carboxylate or the ammonium group are masked from the solvent.

Figure 6. Amino acids used to demonstrate the "masking" of hydrophobic groups by polar groups.

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