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184 BASIC CONCEPTS IN MEDICINAL CHEMISTRY
FIGURE 6-13.Hydrogen bonding interactions of isoproterenol and a biological protein
target.
The importance of collinearity in forming hydrogen bonds can be further exemplified by examining DNA base pairing. Shown in Figure 6-14 is the base pairing of thymidine with adenosine and
guanosine with cytidine. Notice the collinear nature of the hydrogen bonds between these nucleotide bases. Additionally, note that the oxygen, nitrogen, and hydrogen atoms of the nucleotide bases
can interact with functional groups present within the structures of drug molecules via ion–dipole
interactions (as the dipole), dipole–dipole interactions, or hydrogen bonds. The sugar-phosphate
backbone of DNA can also participate in a variety of drug binding interactions. The phosphate groups
are primarily ionized and can participate in ionic bonds and ion–dipole bonds with functional groups
FIGURE 6-14.The importance of hydrogen bonding in nucleotide base pair formation.
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present on drug molecules whereas the oxygen atoms in the sugars can participate in ion–dipole
bonds (as the dipole) and hydrogen bonds.
Binding interactions analogous to hydrogen bonds can also enhance the strength of an ionic
bond. The resulting attraction has been referred to as a reinforced ionic bond and has the potential to
enhance the binding energy to the original ionic bond by approximately 5 kcal/mol. Shown below is an
example of a reinforced ionic bond. In this example, the carboxylic acid of furosemide initially forms
an ionic bond with the primary amine of lysine. Once this occurs, a hydrogen atom from the primary
amine can form a second interaction with the carbonyl oxygen. This second interaction is similar to a
hydrogen bond and does enhance the overall interaction; however, due to the resonance delocalization of the negative charge and the resulting changing interactions, it is not a true hydrogen bond.
Finally, intramolecular hydrogen bonds can play a significant role in the overall conformation
of a drug molecule or a biological macromolecule. Two examples are shown in Figure 6-15. The
first example is seen with metoclopramide, an antiemetic agent that is indicated for the prophylactic treatment of chemotherapy-induced nausea and vomiting as well as for the treatment of
FIGURE 6-15.Examples of intramolecular hydrogen bonds.
186 BASIC CONCEPTS IN MEDICINAL CHEMISTRY
gastroesophageal reflux disease. The intramolecular hydrogen bond between the amide hydrogen
atom and the oxygen atom of the methoxy group locks the conformation of this portion of the
molecule and allows for the formation of a virtual ring system. The second example is seen with
a β-sheet (also known as a β-pleated sheet), which is a common structure found in various types
of proteins. Multiple intramolecular hydrogen bonds between carbonyl oxygen atoms and amide
hydrogen atoms stabilize these structures.
Summary of the Advantages of Hydrogen Bonds
• Hydrogen bonds are generally stronger than other types of dipole–dipole interactions.
• Hydrogen bonds can enhance the strength of an ionic bond.
• Hydrogen bonds are important for enhancing the water solubility of a drug molecule.
• Intramolecular hydrogen bonds can be important for determining the preferred con-
formation for a drug molecule or a biological macromolecule.
Other functional groups are capable of forming hydrogen bonds; however, their prevalence in
drug structures is not common. To be complete, these additional functional groups are shown in
Figure 6-16. Please note the following key points. Once an amide nitrogen atom is substituted with
FIGURE 6-16.Additional functional groups capable of forming hydrogen bonds.
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two additional groups, it no longer has a hydrogen atom and cannot function as a hydrogen bond
donor. The same is true with sulfonamides and sulfonylureas. Sulfones, imines, and azides also lack
a hydrogen atom and can only function as hydrogen bond donors. As described in Chapter 3, the
acidity and ionization of sulfonamides can vary widely depending on adjacent functional groups.
Therefore, sulfonamide functional groups that are primarily unionized can participate in hydrogen
bonds as either the acceptor or the donor. Similarly, due to their respective pKa ranges and adjacent
functional groups, it is possible for a sulfonylurea or a hydrazine to be primarily unionized and participate in a hydrogen bond. As discussed in Chapter 3, the pKa values for the two phosphonates of a
bisphosphonate are very different. The pKa value for the first hydrogen atom is approximately 1.5 to
2.5 whereas the pKa value for second hydrogen atom is higher, ranging from 6.5 to 7.5. Thus, while
one phosphonate is primarily ionized, the second phosphonate may be able to participate in hydrogen bonds if it is significantly unionized.
Prior to discussing binding interactions with functional groups containing hydrocarbon, let’s
evaluate three drugs and apply what has been discussed thus far. Shown below are the structures
of benztropine, ezetimibe, and clindamycin. Their respective acidic and basic functional groups have
been highlighted and the pK
values have been provided. As discussed above, prior to determining
a
the types of binding interactions that these acidic and basic functional groups can undergo, it is
essential to first consider their predominant ionization states.
Let us first evaluate benztropine. Its tertiary amine is greater than 99% ionized at a physiologic
pH of 7.4 and can participate in an ionic bond, a reinforced ionic bond, or an ion–dipole interaction
(as the ion) with its biological target. The ether oxygen atom within the structure of benztropine
can participate in an ion–dipole interaction (as the dipole), a dipole–dipole interaction, or a hydrogen bond (as the hydrogen bond acceptor). In evaluating the structure of ezetimibe, its aromatic
hydroxyl group (or phenol) is approximately 99% unionized at a cellular pH of 7.4 and most likely
participates in an ion-dipole interaction (as the dipole), a dipole–dipole interaction, or a hydrogen
bond (as either the donor or acceptor) with its biological target. The secondary hydroxyl group and
the amide can form similar types of interactions although the amide can only act as a hydrogen
bond acceptor. The two fluorine atoms can also participate in hydrogen bonds as an acceptor.
Clindamycin provides an example in which the pKa of the functional group is close to the physiologic pH of 7.4. Its tertiary amine has a reported pKa of 7.5. In a physiologic pH of 7.4, this basic
functional group is 56% ionized and 44% unionized; thus, the probability that it will form an ionic
bond, a reinforced ionic bond, or an ion–dipole interaction with its biological target is similar to the
probability that it will form a hydrogen bond (as an acceptor), an ion–dipole interaction (as the
188 BASIC CONCEPTS IN MEDICINAL CHEMISTRY
dipole), or a strict dipole–dipole interaction that doesn’t involve a hydrogen bond. The hydroxyl
groups and the amide can participate in ion–dipole interactions (as the dipole), dipole–dipole interactions, and hydrogen bonds (as either the donor or acceptor). The sulfur atom can participate in a
hydrogen bond as an acceptor. As a key and closing point here: you should be able to provide similar
evaluations for any given drug molecule.
van der Waals Interactions
The term van der Waals interactions, or van der Waals forces, actually refers to three different types
of interactions or forces: Keesom forces, Debye forces, and London dispersion forces. As previously
mentioned, Keesom forces is another name for a dipole–dipole interaction. Many texts, including this
one, discuss dipole–dipole interactions, or Keesom forces, separate from Debye and London dispersion forces. In terms of the following discussion, the term van der Waals interactions pertains only
to Debye and London dispersion forces. These two types of interactions involve induced dipoles and
are thus different than the interactions between two permanent dipoles. Debye forces are those that
occur between a permanent dipole (e.g., a hydroxyl group) and an induced dipole whereas London
dispersion forces are those that occur between two induced dipoles. In terms of drug binding interactions, London dispersion forces are much more common than Debye forces. As such, the following
discussion focuses primarily on London dispersion forces.
Unlike polar groups, such as a hydroxyl or an amide functional group, hydrocarbon rings and
chains do not have measurable dipoles. This, however, does not mean that the electrons between
a carbon–carbon bond are static. On the contrary, there is a rapid fluctuation of partial positive and
partial negative charges between the two carbon atoms, which is illustrated below. At any given
moment, the electron movement within the carbon–carbon bond could result in the carbon atom on
the left having a partial positive charge and the carbon atom on the right having a partial negative
charge. In the next instant, the opposite could be true. Both carbon atoms have the same electronegativity, so they equally share the electrons back and forth.
While hydrocarbon chains and rings lack a measurable dipole, a dipole is induced whenever
these functional groups are aligned with either a polar functional group, such as a hydroxyl group,
or another nonpolar hydrocarbon chain or ring. Shown in Figure 6-17A is an example of a dipoleinduced dipole interaction between the n-butyl side chain of losartan, an angiotensin II receptor
blocking agent used to treat hypertension, and a hypothetical serine hydroxyl group present at the
angiotensin II receptor. As previously discussed, a permanent dipole exists within a hydroxyl group,
with the oxygen atom possessing a partially negative charge and the hydrogen atom possessing a
partially positive charge. When this functional group is close enough to the n-butyl alkyl chain, it
induces a dipole in the carbon–carbon bond resulting in the formation of a dipole-induced dipole
bond. Again, this is just a hypothetical example. Figure 6-17B shows a more likely interaction of
this n-butyl chain with a leucine residue present on the angiotensin II receptor. As the two hydrophobic chains approach one another, a mutual polarization of the carbon–carbon bonds occurs. This
induces the formation of complementary dipoles, resulting in numerous induced dipole-induced
dipole interactions.
Individual van der Waals interactions are weak, with bond strengths ranging from 0.5 to 1.0
kcal/mol. A value of 0.7 kcal/mol has been used as an approximate bond strength between two carbon atoms. The ability to form van der Waals interactions depends highly on the distance between
the participating functional groups. The ability for a dipole-induced dipole interaction to occur is
proportional to 1/r
4
, where r represents the distance between the interacting functional groups.
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FIGURE 6-17.Examples of (A) dipole-induced dipole interactions and (B) induced dipole-
induced dipole interactions.
Similarly, the ability for an induced dipole-induced dipole interaction to occur is proportional to 1/r6.
If only a single carbon–carbon interaction is examined, it is easy to dismiss this single van der Waals
interaction as relatively insignificant to the overall binding ability or affinity of the drug molecule for
its biological target. However, when one considers that alkyl chains and aromatic rings are capable
of multiple van der Waals interactions, the overall contribution can become significant. Consider
the n-butyl alkyl chain of losartan shown in Figure 6-17B. If all four carbon atoms are able to interact with the side chain of a leucine residue (or other hydrophobic amino acid), then this functional
group has the potential to provide an overall bond strength of 2.8 kcal/mol. Another example of
this can be seen with the unsubstituted aromatic ring of meperidine shown in Figure 6-18. If this
aromatic ring were able to closely interact with the aromatic ring of a phenylalanine residue, there
would be a total of six individual interactions and a potential bond strength of 4.2 kcal/mol.
FIGURE 6-18.Example of van der Waals interactions possible for the aromatic ring within
the structure of meperidine.
The above examples with losartan and morphine are meant to illustrate how van der Waals
interactions can provide meaningful binding interactions. Additional examples can be seen with salmeterol and tamoxifen (Figure 6-19). These drugs contain multiple aromatic rings, aliphatic chains,
and/or alicyclic rings. As a result, the total contribution of van der Waals interactions to the binding
of these drugs to their biological targets can be quite substantial.
190 BASIC CONCEPTS IN MEDICINAL CHEMISTRY
FIGURE 6-19.Salmeterol and tamoxifen.
An additional consideration for induced dipole-induced dipole interactions is that these interactions are most favorable when they occur between similar types of functional groups. In other
words, an aliphatic chain present on a drug molecule is able to form stronger van der Waals interactions with other aliphatic chains as compared with an aromatic ring. The reason here is that similar
types of functional groups are better able to make close contact with one another. This is especially
important for van der Waals interactions due to their distance requirements. The previous examples
of losartan and meperidine illustrate this concept of similar groups interacting with each other.
Hydrophobic Effects (aka Hydrophobic Interactions)
The term hydrophobic effects refers to the tendency of lipid-soluble, or hydrophobic, molecules
to interact with one another in an aqueous, or hydrophilic, environment. In terms of drug binding
interactions, a hydrophobic effect (or a hydrophobic interaction) is not directly related to bond formation but rather to the gain in entropy that occurs when two hydrophobic functional groups are
attracted to one another. As described above, the actual binding interaction between two hydrophobic groups is known as a van der Waals interaction; however, when these types of bonds occur,
water is displaced and there is a gain of entropy that results in the release of energy. Water molecules have the ability to become highly ordered. As shown below, each individual water molecule
can form hydrogen bonds with four other water molecules.
Within the body, drug molecules as well as their biological targets reside in an aqueous environment. Even biological targets that are embedded into cell membranes have some portion of the
macromolecule exposed to the aqueous physiologic environment. As shown in Figure 6-20A, water
molecules can become highly ordered whenever they encounter nonpolar functional groups. For
these two nonpolar groups to participate in van der Waals interactions, they must be very close
to one another. Whenever this occurs, there is a subsequent “squeezing out” or release of water
molecules from their original, highly ordered state (Figure 6-20B). The overall result of this action
is an increase in entropy and a release of energy. This process is often referred to as a hydrophobic
effect, or a hydrophobic interaction; however, please note that this designation actually describes
the exclusion of water rather than the attraction between two functional groups.
The release of energy through this process enhances the bond strength of the initial van der
Waals interaction. The hydrophobic effect, or interaction, between two carbon atoms has the potential to add up to 0.7 kcal/mol in binding energy and theoretically double the bond strength of a van
der Waals interaction. Thus, the maximum bond strength of a single carbon–carbon interaction is
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FIGURE 6-20.An illustration of hydrophobic effects. (A) Highly ordered water molecules
initially surround the lipid/water interface of hydrophobic functional groups. (B) When the
hydrophobic groups form a van der Waals interaction, they displace the water molecules
and increase the entropy of the system.
1.4 kcal/mol. In applying this concept to meperidine (Figure 6-18), it is found that the total binding
energy of meperidine’s phenyl ring with the phenyl ring of phenylalanine can approach 8.4 kcal/mol.
This finding reinforces the previous statement that the overall contribution of van der Waals interactions to drug binding can be substantial.
Caution is warranted in this situation. In theory, each carbon atom present in an aromatic ring,
aliphatic chain, or aliphatic ring of a drug molecule has the potential to contribute a maximum of
1.4 kcal/mol toward the overall binding strength associated with the interaction of the drug to its
biological target. In reality, this potential maximum value cannot always be attained due to steric fac-
tors and/or the location of the hydrocarbon groups. To illustrate this fact, let us evaluate meperidine
a little closer. The phenyl ring of meperidine shown in Figure 6-17 is a good example of a functional
group that lacks any significant steric hindrance and has an optimal location within the drug molecule. It is a terminal substituent of the piperidine ring as opposed to being located in the middle of
the molecule. As such, it has more conformational flexibility, has a greater chance to align itself with
a complementary region on its biological target, and can significantly contribute to meperidine’s
binding. Shown in Figure 6-21 is another view of morphine that highlights four additional carbon
atoms that could potentially contribute to meperidine’s binding. Carbon A and its adjacent methylene carbon are part of a side chain ester and have significant conformational freedom. Similar to the
phenyl ring, they should be able to align with complementary functional groups and be involved in
van der Waals interactions and hydrophobic effects. Carbon B is also located at an external position
and is sterically unhindered. A potential problem with this carbon is that it is directly adjacent to
a basic nitrogen atom. In all likelihood, this amine will be ionized and the methyl group will reside
192 BASIC CONCEPTS IN MEDICINAL CHEMISTRY
FIGURE 6-21.Additional carbons atoms present within the structure of morphine that
could undergo van der Waals and hydrophobic interactions.
in a polar region of the binding site. Because polar regions of a binding site are less likely to contain
complementary nonpolar functional groups, the probability that carbon B will participate in a van
der Waals interaction and a hydrophobic effect is diminished. Carbon C has a similar problem as it is
also adjacent to the basic amine. Further, this methylene is part of a piperidine ring. Due to the conformation of the ring as well its substituents, steric factors could prevent this group from attaining
the required distance to participate in van der Waals interactions. Carbon D is also located within
the piperidine ring and has bonds to four different carbon atoms. Although it may very well reside in
a nonpolar region of a binding site, the steric factors would eliminate the possibility of this carbon
atom getting close enough to form a van der Waals interaction.
In summary, each aromatic, aliphatic, and alicyclic carbon atom can theoretically contribute up
to 1.4 kcal/mol of binding energy by virtue of van der Waals interactions and hydrophobic effects.
Due to steric hindrance and the location of the carbon atom within the drug molecule, some carbon
atoms are not able to contribute to the overall binding or are only able to contribute a portion of the
maximum binding energy. Functional groups that are most likely to contribute the maximum binding energy are those that are the least sterically hindered, have the most conformational flexibility,
and have the ability to be located in a nonpolar region of the biological target’s binding site.
Additional Aromatic Interactions
Due to the overall electron distribution and density in an aromatic ring, these functional groups
can provide additional binding interactions beyond the van der Waals and hydrophobic interactions
discussed above. The first type of interaction involves the electron clouds of adjacent aromatic rings
and is known as a π-π or aryl-aryl stacking interaction. This interaction is similar to a van der Waals
interaction but also involves polar movements, or uneven charge distributions, of interacting rings.
As shown below, parallel stacking interactions can be either face-to-face (aka sandwiched) or displaced.
Additionally, it is possible for an edge-to-face, or T-shaped, interaction to occur.
A large number of drug molecules contain aromatic rings and ring systems that can potentially
form π-π stacking interactions with their biological targets. For drugs that bind to a protein receptor, π-π stacking interactions can occur with the side chains of phenylalanine, tyrosine, tryptophan,
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and histidine residues. Additionally, many antineoplastic drugs provide their therapeutic effects by
intercalating with DNA. In this process, the aromatic portion of the drug molecule is inserted in
between stacked nucleotide bases and can form several types of interactions. Figure 6-22A shows
an example of a π-π stacking interaction between the unsubstituted phenyl ring of tolterodine and
the side chain of a tyrosine residue within the structure of the cholinergic muscarinic receptor, the
biological target for tolterodine. The structure of mitoxantrone, a drug that binds to DNA through
intercalation, is shown in Figure 6-22B.
FIGURE 6-22.Examples of (A) a o-o stacking interaction and (B) a drug that interacts with
DNA through intercalation.
The second type of interaction is a cation–π interaction that occurs due to the attraction of a
positively charged group to the electrons in the π cloud of an aromatic ring. Within the structure
of a protein or enzyme receptor, the ionized side chains of arginine, lysine, and histidine can function as the cation, while the side chains of phenylalanine, tyrosine, tryptophan, and histidine can
provide the π electrons. Two examples of this type of interaction are shown in Figure 6-23 using
the antidepressant atomoxetine. Atomoxetine produces its antidepressant effects by blocking the
presynaptic norepinephrine transporter. The secondary amine is primarily ionized at physiologic pH
and could undergo a cation–π interaction with a phenylalanine residue located within the structure
of the norepinephrine transporter, while either of the aromatic phenyl rings could interact with a
lysine residue within the transporter.
FIGURE 6-23.Examples of cation–o interactions.
The third type of interaction occurs between electron rich and electron deficient aromatic rings
and is known as a charge transfer interaction. These types of interactions are, in essence, dipole–
dipole interactions; however, due to the overall number of atoms and electrons involved in these
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