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174 BASIC CONCEPTS IN MEDICINAL CHEMISTRY
The ability of any single noncovalent bond to form is inversely proportional to the distance
between the functional group present on the drug molecule and a complementary functional group
present on the biological target. In other words, the respective functional groups must be close
enough to one another for recognition to occur and the noncovalent interaction to occur. Each type
of noncovalent bond has a unique distance requirement. Ionic bonds can form over the greatest distance whereas certain types of van der Waals interactions can only form if the respective functional
groups are extremely close. This is also true for the covalent bonds that were previously discussed.
For covalent bonds to form, the reactive functional groups must be close enough to one another for
the reaction to occur. As an example, let’s review the reaction in Figure 6-6. Please note that the serine hydroxyl group must be close enough to isofluorophate to be phosphorylated. The same is true
for the alkylations and acylations shown in Figures 6-2 through 6-5. The distance requirement for a
covalent bond is usually the same as that for an ion–dipole or a dipole–dipole interaction. Specific
distance requirements are discussed with each type of noncovalent bond.
The ionization state of a functional group plays a key role in determining what types of noncovalent bonds it can and cannot form. Because most drug binding interactions occur at a physiologic pH of 7.4, this is the primary focus for this chapter. Therefore, prior to evaluating the types
of binding interactions that a functional group can participate in, it is first necessary to determine if
the functional group is primarily ionized or primarily unionized at a pH of 7.4. Functional groups that
are primarily ionized at physiologic pH can undergo binding interactions in which the functional
group has a full positive or negative charge. It is extremely unlikely that they would participate in
binding interactions in which the functional group needs to be unionized or carry a partial charge.
In contrast, functional groups that are primarily unionized at physiologic pH can participate in binding interactions in which the functional group provides a dipole, partial charge, or induced partial
charge. It is extremely unlikely that they would participate in binding interactions that require the
functional group to have a full positive or negative charge. As an example of this, let’s look at two
functional groups: a carboxylic acid and a primary aliphatic hydroxyl group. As discussed in Chapter
3, carboxylic acids are acidic functional groups and have a pK
range of 2.5 to 5. Applying the con-
a
cepts discussed in Chapter 4, we can conclude that this functional group will be primarily ionized at
physiologic pH and will participate in drug binding interactions in which it provides a full negative
charge. It does not participate in drug binding interactions that require it to be unionized and have a
partial charge. In contrast, a primary aliphatic hydroxyl group is neither acidic nor basic and is unionized at physiologic pH. As discussed in this chapter, it can participate in drug binding interactions in
which it provides a dipole bond with partial charge separation. It is not able to participate in drug
binding interactions that require it to be ionized.
At the end of Chapter 3, Table 3-1 provides the approximate pKa ranges for each acidic and basic
functional group. Using these ranges, the approximate ionization state for each of these functional
groups at a physiologic pH of 7.4 can be determined. This information is provided in Table 6-2. In
examining this table, the acidic and basic functional groups can be divided into three categories:
(1) those that will almost always be primarily ionized at a physiological pH of 7.4, (2) those that will
almost always be unionized at a physiological pH of 7.4, and (3) those whose pKa ranges span both
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TABLE 6-2.Approximate Ionization States for Common Acidic and
Basic Functional Groups at a Physiologic pH of 7.4
Functional Group Acidic or Basic pKa Range and Primary Ionization State at pH of 7.4
Carboxylic acid Acidic 2.5-5 (primarily ionized)
β-Dicarbonyl groups (includes
imides)
Sulfonamides Acidic 4.5-11 (varies depending on pK
Sulfonylureas Acidic 5-6 (primarily ionized)
Tetrazoles Acidic 4.5-6 (primarily ionized)
Phenols Acidic 9-10 (primarily unionized)
Thiols Acidic 10-11 (primarily unionized)
Sulfates Acidic 1-2 (primarily ionized)
Phosphates and phosphonates Acidic 1.5-2.5 (first phosphate; primarily ionized) 6.5-7.5
Aliphatic amines and alicyclic
amines (aka saturated
heterocycles)
Aromatic amines (aka anilines) Basic 2-5 (primarily unionized)
Imine Basic 3-5 (primarily unionized)
Hydrazine Basic 7.5-8.5 (primarily ionized but could vary)
Amidine Basic 10-11 (primarily ionized)
Guanidine Basic 12-13 (primarily ionized)
Nitrogen containing aromatic
heterocycles
Acidic 4.5-8.5 (varies depending on pK
(second phosphate; primarily ionized but could vary)
Basic 9-11 (primarily ionized)
Basic 1-6 (primarily unionized)
value)
a
value)
a
sides of 7.4 and require individual evaluation. The term almost is included in the previous statements because the steric and electronic factors discussed in Chapter 3 can result in rare outliers.
Steric factors and the stereochemistry of drug molecules are also important for the interactions
between a drug and its biological target. These topics are discussed in more detail in Chapter 7; however, based on the distance requirements for drug-target binding, drug molecules that cannot make
close contact with their respective targets would be expected to exhibit decreased binding. If steric
factors preclude a functional group on a drug molecule from getting close enough to a functional
group on its biological target, the bond will be either weak or nonexistent. Additionally, based on
the distance requirements for drug-target binding, one should expect a sequential process for multiple bond formation. Because ionic bonds can form over the longest distance, they generally occur
first, followed by ion–dipole interactions, and then dipole–dipole interactions, dipole-induced dipole
interactions, and finally induced dipole-induced dipole interactions.
Ionic Bonds
Ionic bonds occur between ionized acidic and basic functional groups or between ionized acidic and
quaternary ammonium functional groups. In other words, an ionic bond occurs between two atoms
that bear opposite full charges (i.e., a “+” and a “−” charge). The ability of these bonds to form
depends on the pKa of the functional groups and their ionization at a specific physiologic pH. Ionic
bonds have a bond strength of approximately 5 to 10 kcal/mol and are the strongest of all noncovalent interactions. The ability to form an ionic bond is proportional to 1/r, where r represents the
176 BASIC CONCEPTS IN MEDICINAL CHEMISTRY
distance between the ionized functional groups. An example of an ionic bond between nateglinide
and a positively charged guanidine group present on its biological target is shown below.
Nateglinide binds to SUR1, a specific receptor on the pancreatic β cell, and stimulates the
release of insulin in a glucose dependent fashion. Because ionic bonds can form over the greatest
distance, the electronic attraction between the positively charged guanidine group on SUR1 and the
negatively charged carboxylic acid on nateglinide serves here as the initial recognition between the
drug and its biological target. The importance of this initial recognition can be seen with the methyl
ester analog of nateglinide. There is no ionic attraction between this analog and SUR1, so this drug
is unable to recognize its biological target and completely lacks any anti-inflammatory or analgesic
activity.
Since drugs primarily bind to proteins and enzymes, the acid/base nature of the amino acids
comprising these macromolecules is essential to their ability to form ionic bonds. Figure 6-8 highlights the naturally occurring basic functional groups that are the most likely to be present on an
enzyme or protein.
FIGURE 6-8.Amino acids containing basic functional groups.
The side chains of both arginine and lysine contain strong, basic functional groups. The pKa value
of the guanidine group of arginine is approximately 12.5, while the pKa of the primary amine of lysine
is approximately 10.5. Both of these functional groups are primarily ionized in all physiologic environments. The imidazole ring of histidine is significantly less basic with a pKa of 6.0. As such, it is less
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than 10% ionized at physiologic pH. Despite this fact, it is capable of forming ionic bonds, especially
if it is in close proximity to other functional groups. An example of this is seen in Figure 6-9.
FIGURE 6-9.Enhanced ionization of histidine by an adjacent aspartic acid.
The interaction between the side chain of histidine and an adjacent aspartic acid enhances the
dipole between the imidazole nitrogen atom and the hydrogen atom. The hydrogen atom becomes
more partially positive and the ring becomes more partially negative, which allows the electrons on
the other imidazole nitrogen atom to become more available and hence more basic and more capable of forming ionic bonds. This is an example of what is known as a local environment within the
structure of a protein or enzyme. As seen in this example, local environments can provide situations
where functional groups may be ionized to a greater extent than that predicted by comparing their
values to a pH of 7.4.
pK
a
The N-terminal primary amine is also capable of forming an ionic bond with an acidic functional
group; however, this amine is often acetylated or otherwise modified for the purpose of forming a
neutral, unionized functional group. This common posttranslational modification masks the amino
end of the protein and prevents degradation by aminopeptidases, exopeptidases that sequentially
remove amino acids from the N-terminus of the protein or enzyme.
Figure 6-10 highlights the acidic functional groups present on enzymes and proteins. Both
aspartic acid and glutamic acid have side chains that contain a carboxylic acid. The pKa values for
these two acidic functional groups are 3.6 and 4.2, respectively, and therefore are primarily ionized
at physiologic pH. The difference in the pKa values of these two carboxylic acids is due to differences
in adjacent functional groups. Within the structure of glutamic acid, there is an additional adjacent methylene (CH2) group that is not present in aspartic acid. Because hydrocarbons are electron
donating, this additional methylene group decreases the acidity of the side chain carboxylic acid in
glutamic acid, compared with the analogous carboxylic acid in aspartic acid. The C-terminal carboxylic acid has the lowest pKa (range = 1.8 to 2.4); however, similar to the N-terminal amine, it is often
FIGURE 6-10.Amino acids containing acidic functional groups.
178 BASIC CONCEPTS IN MEDICINAL CHEMISTRY
esterified or otherwise modified to form a neutral, unionized functional group that protects it from
carboxypeptidase degradation. Thus, it is often unavailable to form ionic bonds.
Similar to what was discussed with histidine, adjacent functional groups can create local environments that can either enhance or detract from the ionization of aspartic acid or glutamic acid.
Figure 6-11A illustrates a glutamic acid moiety present in a polar binding pocket. The functional
groups on the adjacent lysine and serine amino acids help to enhance the ionization of glutamic
acid. This enhancement increases the tendency of the carboxylic acid to become ionized and hence
decrease its pKa. In contrast, Figure 6-11B illustrates a glutamic acid moiety present in an otherwise
nonpolar binding region. The surrounding amino acids, indicated in the figure by valine and phenylalanine, are not able to either stabilize or share in the ionic charge. This decreases the tendency of
the carboxylic acid to become ionized and increases its pKa. Please note that the same effects would
occur with an aspartic moiety.
FIGURE 6-11.The effects of local environments on the ionization of glutamic acid.
Another Way to Look at Local Environments
Within any given biological target, adjacent functional groups can interact with one another
and thus influence the ability to form specific types of bonds. This was illustrated in
Figures 6-9 and 6-11. To help understand this concept, let’s step outside the realm of chemistry and look at an ordinary social interaction. Consider a student who just received his first
“A” grade in a very challenging course. In the friendly environment of his fraternity’s lunch
table, he is more than willing to share this good news with his friends, especially those who
have helped him learn and study the course material. Later in the day, this same student is
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traveling home on a bus and is seated beside a total stranger who is somewhat engrossed in
the local newspaper. In this environment, the student is much less likely to share his good
news about his exam grade. The key point in this analogy is that the only thing that has
changed is the environment. The student is the same, the good news is the same, but the
environment or the people surrounding him is different. In the examples using histidine and
glutamic acid, the functional groups remained the same; however, the environments, or the
surrounding functional groups, altered their ionization and ability to form ionic bonds.
Dipole Interactions
Dipole interactions occur between functional groups in which there is a partial charge separation
between the atoms involved in the individual bonds of the functional group. This partial separation
of the charge is due to the varying electronegativity values of the atoms comprising the functional
groups. These electronegativity values were introduced in Chapter 2, in which the concept of intrinsic induction was discussed. They are reproduced here in Table 6-3, but we examine them with a
slightly different emphasis.
TABLE 6-3.Electronegativity Values for
Atoms Commonly Seen in Drug Molecules
Atom Electronegativity Value
F 3.98
O 3.44
Cl 3.16
N 3.04
Br 2.96
I 2.66
S 2.58
C 2.55
H 2.20
P 2.19
Listed below are the key points to learn in terms of applying electronegativity values to dipole
interactions. Items 1 and 2 are the most significant due to the relative prevalence of oxygen, nitrogen,
carbon, and hydrogen atoms in drug molecules.
1.
Oxygen and nitrogen are more electronegative than carbon and hydrogen; thus, whenever
an oxygen or nitrogen atom is bound to either a carbon or hydrogen atom, the oxygen or
nitrogen atom has a partial negative charge and the carbon or hydrogen atom has a partial
positive charge.
2. The electronegativity values of carbon and hydrogen atoms are similar to one another. As
such, there is no significant dipole in a C—H bond.
180 BASIC CONCEPTS IN MEDICINAL CHEMISTRY
3. The electronegativity values of sulfur and hydrogen atoms are also similar to one another;
however, due to the larger size of the sulfur atom, a dipole does exist in a S—H bond, with
the sulfur atom having a partial negative charge and the hydrogen atom having a partial
positive charge. This dipole is not as great as that seen for an O—H bond.
4. Similar to a carbon atom, a phosphorus atom has a much lower electronegativity than an
oxygen atom; thus, a phosphorous atom has a partial positive charge when bound to an
oxygen atom, and the oxygen atom has a partial negative charge.
5. All halogen atoms (i.e., F, Cl, Br, and I) are much more electronegative than carbon, so all
C—halogen bonds contain a partially positive carbon atom and a partially negative halogen
atom. With the exception of interactions with a C—F bond, interactions with C—halogen
bonds are not that common in terms of dipole interactions between a drug and its biological target.
Ion–Dipole and Dipole–Dipole Interactions
Dipole interactions can be subdivided into ion–dipole interactions, dipole–dipole interactions, and
hydrogen bonds. Dipole–dipole interactions are also known as Keesom forces, and hydrogen bonds
are a specialized type of dipole–dipole interaction. The bond strength of a dipole interaction varies
from 1 to 7 kcal/mol, depending on the specific interaction. Dipole interactions require functional
groups to be closer in proximity than that for an ionic bond. This is reflected in the fact that the ability
to form an ion–dipole interaction is proportional to 1/r2 whereas the ability to form a dipole–dipole
interaction is proportional to 1/r3. Similar to ionic bonds, r represents the distance between the
interacting functional groups. The mathematical relationships (e.g., 1/r, 1/r2) provided in this chapter
may vary slightly among references; however, the overall trends remain the same. As the electronic
interactions become smaller, the distance required to form the bond becomes more stringent.
An ion–dipole interaction occurs between an ionized acidic or basic functional group (full charge)
and a functional group with a partially charged dipole. An example of an ion–dipole interaction is
shown below. Chlorpheniramine is an antihistamine indicated for the treatment of allergic conjunctivitis, allergic rhinitis, hay fever, and symptoms of the common cold. It contains a basic tertiary
amine that is primarily ionized at physiologic pH. It is capable of forming an ion–dipole interaction
with the hydroxyl group of threonine. In this interaction, the full positive charge on the tertiary
amine is attracted to the partially negative charge of oxygen atom of the hydroxyl group. In identifying ion–dipole interactions, it is important that you identify which functional group is participating
as the ion and which functional group is participating as the dipole.
A disclaimer is appropriate at this point. Because the emphasis of this chapter is to identify the
types of bonds that can be formed between drug molecules and their biological targets, the example here and those that follow are simply examples of what is possible. Unless otherwise stated,
these examples are not meant to portray any specific known binding interactions. In other words,
the example here is simply showing that chlorpheniramine could form an ion-dipole bond with a
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threonine residue on a biological target; it is not stating that this is the actual interaction that occurs
between chlorpheniramine and the histamine H1 receptor. For the sake of illustrating what is possible, various amino acids and functional groups are used as examples.
A dipole–dipole interaction occurs between functional groups that have complementary, or
opposite, partial charges. The interaction of the hydroxyl oxygen of ezetimibe with the carbonyl
carbon of an asparagine residue present on a biological target of ezetimibe provides an example of
a dipole–dipole interaction.
Application of Concepts
An important skill to master is the ability to examine the structure of a drug molecule and identify the possible types of interactions that it can form with its biological targets. Thus far, we
have discussed the following noncovalent bonds: ionic bonds, ion–dipole bonds, and dipole–
dipole bonds. Using only these three types of bonds, let’s evaluate atropine. The tertiary
amine (Group A) is primarily ionized at physiological pH and thus is able to participate in
either ionic bonds or ion–dipole bonds (as the ion). The hydroxyl group (Group B) can participate in either ion–dipole bonds (as the dipole) or dipole–dipole bonds similar to what was
seen with chlorpheniramine and ezetimibe. The three atoms of the ester functional group can
all participate in either ion–dipole bonds (as the dipole) or dipole–dipole bonds. The carbon
atom (Group C) is partially electropositive whereas the carbonyl and ester oxygen atoms
(Groups D and E) are partially electronegative.
Hydrogen Bonds
A hydrogen bond is a specialized type of dipole–dipole interaction and in general is stronger than the
dipole–dipole bonds shown above. A discussion of hydrogen bonds and their role in the water solubility of a drug molecule can be found in Chapter 2. As initially discussed in that chapter, a hydrogen
bond is one that occurs whenever a hydrogen atom serves as a bridge between two electronegative
182 BASIC CONCEPTS IN MEDICINAL CHEMISTRY
atoms. In this type of bond, the hydrogen atom is covalently bound to one atom and noncovalently
bound to the other. A general representation of a hydrogen bond is shown below.
In a hydrogen bond, the atom that is covalently bound to the hydrogen atom is known as the
hydrogen bond donor, and the atom that is noncovalently bound to the hydrogen atom is known as
the hydrogen bond acceptor. Thus, in the representation above, atom X is the hydrogen bond donor
and Y is the hydrogen bond acceptor. Oxygen, nitrogen, and sulfur can serve as both hydrogen bond
donors and acceptors, with oxygen and nitrogen being the most common and providing the strongest hydrogen bonds. Fluorine atoms can serve as hydrogen bond acceptors. A representative summary of the most common functional groups that are capable of forming hydrogen bonds is shown
in Figure 6-12. A key point in identifying hydrogen bonds is to designate which functional group is act-
ing as the donor and which functional group is acting as the acceptor.
An example of hydrogen bonding is shown below with morphine and a serine residue that could
be present within a biological target. In this example, the phenolic group of morphine is able to
form two individual hydrogen bonds with the hydroxyl group of serine. In the leftmost interaction,
the hydrogen atom of the hydroxyl group of morphine is the hydrogen bond donor, and the oxygen
atom of the hydroxyl group of serine is the hydrogen bond acceptor. In the rightmost interaction,
the roles of these two hydroxyl groups are reversed, with the oxygen atom of morphine serving as
the hydrogen bond acceptor and the hydrogen atom of serine serving as the donor.
The bond strength of a hydrogen bond ranges from 3 to 7 kcal/mol, and two major factors affect
the overall bond strength. First, the bond strength depends on the alignment of the atoms involved
in the hydrogen bond. Using the R1—X—H---Y—R2 designation, the strongest hydrogen bonds occur
when X, H, and Y are collinear. Second, the bond strength depends on which atom serves as the
donor and which serves as the acceptor. In comparing oxygen and nitrogen, the two most common
electronegative atoms involved in a hydrogen bond, it is found that oxygen is a better hydrogen
bond donor than nitrogen, and nitrogen is a better hydrogen bond acceptor than oxygen. Using the
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FIGURE 6-12.Common functional groups capable of hydrogen bonds.
above XHY sequence, in which XH is the hydrogen bond donor and Y is the hydrogen bond acceptor,
the approximate rank order of hydrogen bond strength is OHN > OHO > NHN ≥ NHO. Figure 6-13
provides an example that includes both of these factors. In this example, the catechol hydroxyl
groups of isoproterenol can form hydrogen bonds with a nitrogen atom present on the backbone of
the protein (hydrogen bond 1) as well as a hydroxyl group present on a tyrosine residue of that same
protein (hydrogen bond 2). Based on the above information, hydrogen bond 2 would be expected to
be stronger based on the fact that the OHO atoms are collinear, whereas the NHO atoms are at a
120° angle. Additionally, an OHO hydrogen bond is stronger than an NHO bond.
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