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. • Important Types of Protein–Ligand Interactions
happen. There is no increased distribution of energy over
more degrees of freedom, and the system does not enter
amore disordered state. However, this is aprerequisite
for spontaneous processes. To quantify this nding, en-
tropy,S, was introduced. It is ameasure of the order of
asystem. It estimates over how many degrees of freedom
agiven amount of energy is distributed in asystem and
how much it will increase during aspontaneous process.
In the case of protein–ligand complexes, a degree of
freedom can be, for example, acertain vibration of the
system or arotation of individual groups against each
other. Ahighly ordered system, where energy is trapped
in only afew degrees of freedom, has alow entropy content. Increasing disorder increases entropy.
These aspects are particularly important for multicomponent systems, such as protein–ligand complexes,
which also include the surrounding solvent molecules.
To fully describe such asystem, we need to consider not
only the energy terms involved in the formation of noncovalent interactions between the two binding partners,
but also how the energy is distributed over the multiple
degrees of freedom of the formed protein–ligand complex. Therefore, we need athermodynamic property that
takes into account not only the energy contributions, but
also how the energy is distributed over the system. In
this way, the formed protein–ligand complex migrates to
amore disordered state, otherwise its formation would
not occur spontaneously.
The Gibbs free energy (∆G) is the appropriate property to describe the formation of such complexes. Another important thermodynamic relationship is that the
Gibbs free energy consists of two additive terms, an en-
thalpic term ∆H and an entropic term −T∆S. The latter
is temperature weighted. It takes into account not only
the energy balance of the process, but also the changes
in entropy (Eq.4.4).
(4.4)
The entropic component is temperature weighted and,
thus, receives the dimension of an energy. It makes abig
difference whether the entropy in asystem is changed
at low temperature, where all particles are in alargely
ordered state, or at high temperature, where the disorder
is already very high. Because of the negative sign, an
increase in entropy causes adecrease in ∆G and therefore an increase in binding afnity. When discussing the
thermodynamic aspects of binding proles, we always
have to consider ∆H and −T∆S in addition to ∆G. If the
enthalpic and entropic contributions are both negative,
this will improve afnity and the corresponding process
is called exergonic. If one of them is positive, this will be
detrimental to ligand binding. However, the sum of both
quantities must be negative for the binding reaction to
reach equilibrium. Otherwise, the complex formation will
not occur and the process is called endergonic.
In biology, many processes require an energy source
because they would be endergonic and, thus, not possible if considered alone. However, an overall exergonic
balance is required for these processes to occur spontaneously for thermodynamic reasons. Examples are muscle contraction, transport of substances against concentration gradients, or the synthesis of many biomolecules.
Such processes must be coupled to areaction that provides the required Gibbs free energy and gives the overall
process an exergonic inventory. Of central importance
in this context is the hydrolysis of adenosine triphosphate
(ATP) to adenosine diphosphate (ADP) and adenosine
monophosphate (AMP). The triphosphate unit is high
in energy and its hydrolysis releases alarge amount of
Gibbs free energy. Therefore, ATP is continuously produced by the organism and then hydrolyzed, usually in
aspecic ATPase, to allow the biological processes of
interest to occur spontaneously.
Finally, in Eq.4.4 the term “enthalpy” is used. Why
term “energy”? In chemistry and biology, processes take
place as so-called open systems under atmospheric pressure. Since the volume of the environment is enormous,
it can be assumed that the external pressure remains
unchanged even in processes in which gas is produced.
Therefore, these processes are considered to be under
constant pressure conditions. Nevertheless, a gas produced during areaction must rst nd its place among
the surrounding particles in the air. Therefore, some energy must be spent and reduces the maximum possible
energy amount that can be transferred by the system
(so-called internal energy, ∆U). Accordingly, the energy
reduced by this pressure–volume work is called enthalpy
(∆H). It is therefore the energy converted during aprocess corrected for the pressure–volume work. In biological processes, the emission of agas rarely plays arole.
Therefore, internal energy and enthalpy are usually equal
here (Note: in German usage, ∆G, which also refers to
constant pressure conditions, is called a“free enthalpy.”
In English, the term “Gibbs free energy” is used instead).
4.4 Important Types of Protein–Ligand
Interactions
Organic molecules can bind to proteins by forming noncovalent interactions between the ligand and the protein.
In some cases, covalent bonds are also formed. For example, achemically modied product of omeprazole reacts
with its protein target and forms acovalent bond to the
thiol group of acysteine residue (Sects.9.5 and30.9). In
this section, we will limit ourselves to ligands that bind
to the protein by forming noncovalent interactions. For
the following discussion, it is helpful to classify protein–
ligand interactions into different categories. The different
types of interactions are summarized in . Fig.4.3.
4
Chapter • Protein–Ligand Interactions as the Basis for Drug Action
. Fig. 4.3 Frequently occurring protein–ligand interactions. Im-
portant polar interactions are hydrogen bonds and ionic interactions.
Metalloproteases contain, for example, zinc ions as acofactor, the interaction of which with aligand often yields important contributions
to the binding afnity. Nonpolar parts of the protein and ligand contribute hydrophobic interactions. Because of the particular electron
distribution in aromatic rings, the interaction between unsaturated
ring systems is particularly large
Hydrogen bonds (H-bonds) between proteins and
ligands are frequently observed. The proton-bearing
partner in abiological system is usually an NH or OH
group, termed the hydrogen-bond donor. The opposite
group is an electronegative atom with apartial negative
charge and is termed the hydrogen-bond acceptor. Examples of hydrogen-bond acceptors are oxygen and nitrogen
atoms. Hydrogen bonds are predominantly electrostatic
interactions. They achieve their extraordinary strength
because the hydrogen atom of the donor group is bonded
to astrongly electronegative atom, causing the electron
density of the hydrogen atom to shift to the neighboring atom. The sphere of inuence of the hydrogen atom
becomes effectively smaller. This allows the acceptor to
come closer to the proton than the sum of the van der
Waals radii would allow. The electrostatic attraction between the partners thus becomes greater. The atoms of
ahydrogen bond, for example of N–H⋯O=C, assume
an almost linear arrangement to each other. The distance
N⋯O is between 2.5 and 3.2 Å. The angle N⋯H⋯O is
nearly always larger than 150°. Looking from the opposite side, amuch larger variation between 100 and 180°
is observed for the angle C=O⋯H.
It is often found that charged groups of the ligand
bind to oppositely charged groups on the protein and
they frequently overlap with ahydrogen bond. Then
they are called charge-assisted hydrogen bonds. If both
bear aformal charge, they are also named ionic inter-
actions (also known as salt bridges) and they are par-
ticularly strong when the two groups are separated by
only 2.7–3.0 Å. We will see that in many protein–ligand
complexes the association is largely determined by such
charge-assisted or ionic interactions. Afew proteins contain metal ions as cofactors, such as Zn2+ in metalloproteases (Chap.25). It is often the attractive interactions
between the metal ion and the opposite charge on the
ligand functional groups that contribute signicantly
to the afnity in these structures. In addition, there are
afew groups that are particularly well suited to form
complexes with transition metals. These include R–SH
thiols, R–CONHOH hydroxamic acids, carbonic acid
groups, and many nitrogen-containing heterocycles.
Whether the charge can increase the afnity contribution of hydrogen bonds depends strongly on the pro-
tonation state of the involved functional groups. Drugs are
usually weak acids or bases, which means they contain
so-called titratable groups (Sect.19.4). Whether these
groups, e.g., acarbonic acid, an acidic sulfonamide, or
anitrogen-containing heterocycle, can release or accept
aproton and transform into acharged state depends
strongly on the local pH conditions and the polarity and
the distribution of electrons in amolecule. The same
can apply for the functional groups of the acidic or
basic amino acid residues. These groups can then form
charge-assisted hydrogen bonds that contribute more to
the binding afnity (Sect.4.8).
The pKa value is used to estimate whether agroup
is in the protonated or deprotonated state. It indicates
the pH at which the two forms in equilibrium are present in equal amounts. The situation can become more
complicated because the pKa value can be shifted by the
local environment. The pKa value is in fact dened for
an aqueous medium, but we also use it for solvent mixtures and even transfer it to the environment in aprotein binding pocket. In ahydrophobic environment,
adopting acharged state is less favorable for acidic and
basic groups; this means ashift to aless acidic or basic
character is the result. If an already protonated, positively charged group in the ligand faces an amino acid
of the protein with the same charge, its protonated state
becomes even more difcult to accomplish. Formally,
the group will therefore exhibit less basic behavior. The
opposite is the case when putatively positively charged
basic groups bind in anegatively charged protein envi-
0
ion
. • Important Types of Protein–Ligand Interactions
. Fig. 4.4
protonation states. In the case of trypsin, the interaction is mediated by
an ordered water molecule; in the case of thrombin, the water molecule
is distributed over three positions. The differences in the protonation
states can be determined by isothermal titration calorimetry from sever-
Ligand 4.1 binds to thrombin and trypsin with different
ronment. Here, the charged state is formed even more
easily, corresponding to a stronger basic character.
The same principle is true for acidic groups only with
opposite signs. Here, apositively charged protein environment makes an acidic group appear more acidic,
whereas anegatively charged environment will make an
acidic group appear less acidic. In this way, the protein
environment can induce asignicant pKa shift of the
titratable groups of the ligand. Uncharged H-bonds can
become charge-assisted contacts that contribute signicantly more to binding afnity (Sect.21.9). Electrostatic
calculations can be used to estimate the pKa shift during
complex formation (Sect.15.4).
Khang Ngo in Marburg examined the following
example that illustrates these differences (. Fig.4.4).
Trypsin and thrombin are structurally very similar serine
al buffers, which exhibit different heat of ionization (
tion to the molecular scaffolds, the electron density around the ligand
is shown (green chicken wire mesh), which determines the placement
of the molecules in crystallography (Sect.13.5)
H
)
. In addi-
proteases (Sect.23.4). They have almost identical binding
pockets into which ligands bind (the so-called S1 pocket,
Sects.14.5 and23.3). An aspartate residue at position 189
(Asp 189) in this pocket is important for ligand binding.
This occurs by formation of ahydrogen bond between
the ligand and Asp 189 which is mediated via awater
molecule. It is interesting that ligand4.1 binds to thrombin with its pyridine ring in the deprotonated state, and
to trypsin in protonated form. In both cases, ahydrogen
bond is formed with the aspartate mediated via awater
molecule. In thrombin, the water molecule is spatially
scattered over three sites. In trypsin, the water molecule
occupies only one position, indicating amore xed geometry. Differences in protonation states can be measured by isothermal titration calorimetry (Sect.7.7), since
different heat signals are obtained when titrations are
4
Chapter • Protein–Ligand Interactions as the Basis for Drug Action
. Fig. 4.5 Typical lipophilic groups in ligands are aliphatic and ar-
omatic hydrocarbons, halogen substituents, as well as nonpolar heterocycles such as furan and thiophene. Halogens, such as chlorine or
bromine, have apositively polarized tip (so-called σ-hole) on the side
facing away from the aromatic ring in extension of the C–Cl bond,
with which they can enter into an attractive electrostatic interaction
with the π-electron system of aneighboring aromatic system
the inuence of direct attractive forces between the lipophilic groups is small. Hydrophobic interactions usually
involve the displacement, or more precisely, the release
of water molecules from the lipophilic environment of the
binding pocket. In addition, the ligand with its lipophilic
substituents leaves the bulk water phase in the vicinity of
the protein. The solvent “cave” in which the ligand was
hosted in water collapses. At rst glance, this release of
water molecules from hydrophobic surfaces into the bulk
solvent increases the disorder of the system and, hence,
has afavorable entropic contribution to the change in
the free energy of binding. The role of water molecules
is discussed in Sect.4.6. Another important interaction
should be mentioned here. Quaternary amines bind particularly well in binding pockets formed by the aromatic
side chains of aprotein. This contact is largely based on
the polarization interaction between the positive charge
and the electronic π-system of the aromatic rings.
performed in buffers with different heats of ionization. If
binding is accompanied by achange in protonation state,
the titrations allow quantication of the molar amount
of protons transferred during the binding reaction. For
thrombin, there is hardly any buffer dependence, while
for trypsin there is astrong effect (. Fig.4.4). In thrombin, asodium ion is found next to the aspartate residue.
This ion is absent in trypsin. Most likely, the positively
charged sodium ion attenuates the charge on the adjacent
acidic aspartate residue in thrombin to such an extent
that the induced pKa shift on the pyridine moiety is insufcient to reach aprotonated state. In the geometrically analogous trypsin, however, the polarization effect
is strong enough for protonation to occur.
Hydrophobic interactions are formed by the close
proximity of nonpolar amino acid side chains of the protein to lipophilic groups on the ligand. Lipophilic groups
include aliphatic or aromatic hydrocarbon groups, as well
as halogen substituents (e.g., achlorine) and many heterocycles such as thiophene and furan (. Fig.4.5). All
areas that cannot form H-bonds or other polar interactions count as lipophilic parts of the surface of aprotein
and ligand. Unlike hydrogen bonds, hydrophobic interactions are not directional. The relative orientation of
the lipophilic groups to one another does not matter. An
exception is the interactions formed between aromatic
rings, for which there is apreferred relative orientation.
Halogens, such as chlorine or bromine, have apositively
polarized tip on the side opposite to their covalent attachment. If they face with this tip, an aromatic ring,
e.g., of aPhe of Tyr residue, an attractive electrostatic
interaction will be formed (. Fig.4.5).
It has been shown that for ligands with large lipophilic groups, hydrophobic interactions often make asignicant contribution to the binding afnity. However,
4.5 The Strength of Protein–Ligand
Interactions
When evaluating the strength of protein–ligand interactions, it is useful to rst consider the noncovalent
interactions between isolated small molecules. Information about these interactions is available from quantum
mechanical calculations (Sect.15.5) and spectroscopic
studies. In this way, molecular pairs can be studied experimentally in the gas phase. The association energies
obtained for the molecules give an indication of the
strength of the direct interactions. Of course, the inuence of effects resulting from the release of the solvent
water (desolvation) is missing in such experiments. Some
of these data are summarized in . Table4.2.
The results show that electrostatic interactions are
the dominant energetic factor. The interaction between
acation and an anion in vacuum is more than −400 kJ/
mol. This is equivalent to the strength of a covalent
bond! This amount is enormous compared to the typical protein–ligand interactions in water, which were
discussed in Sect.4.4. The binding of an ion pair in
the gas phase is, therefore, much larger than the typical
strength of aprotein–ligand interaction in water. Two
water molecules bind with −22 kJ/mol. This interaction
is also predominantly electrostatic in nature, as the large
dipole moment of awater molecule is responsible for
the strong binding. Interactions between small, nonpolar molecules are much weaker. Two methane molecules
bind with about −2 kJ/mol. This is less than 10% of
the H2O⋯H2O interaction. Correspondingly, methane
boils at 90 K, while water is aliquid at room temperature. The direct interactions between polar groups are,
therefore, orders of magnitude stronger than those between nonpolar groups.
. • Blame It All on Water!
4.6 Blame It All on Water!
The data presented in the previous section might suggest
that protein–ligand interactions are mainly determined
by H-bonds and ionic interactions. It is all the more surprising that the acetate ion CH3COO− does not form
adimer with the guanidinium ion H2NC(═+)NH2 in
water. Analogously, amides do not associate at all in water, although hydrogen bonds between two amide groups
almost always occur inside protein structures. How can
this be? The answer is that water is to blame for all of it!
All biochemical reactions take place in water. Indeed,
aqueous solvation is an absolute requirement! How is aligand solvated prior to protein binding? Water molecules
form ashell around the ligand and the protein. This shell
must be shed as the ligand enters the protein binding
pocket. It involves breaking some of the hydrogen bonds
formed by the polar functional groups of the ligand to
the water molecules. At the same time, the “cavity” in
which the ligand was trapped in the bulk water phase before binding to the protein will collapse. In consideration
of this, we should not forget what the structure of pure
water looks like. The structure of ice can serve as arst
model. In fact, liquid water can transiently adopt the ice
structure, although in the temperature interval 0–100 °C
there are considerable modications between the local
volumes and their sizes where this occurs. In crystalline
water, each water molecule is tetrahedrally surrounded by
four other water molecules, resulting in aspatial structure
similar to that of diamond. The individual water molecules are connected to one another by hydrogen bonds.
However, the distribution of these H-bonds is not strictly
ordered and static. Instead, ahighly dynamic network is
formed. The individual water molecules can spatially rotate, so that, on average, each water molecule participates
in four equivalent H-bonds, twice as adonor and twice as
an acceptor. Such awater structure is, thus, highly disordered, which will have ahuge impact on the entropy contribution when structural changes occur. If this dynamic
network is perturbed, for example by the introduction
of ahydrophobic group of adissolved drug molecule,
the water molecules in the immediate vicinity of the hydrophobic group will no longer be able to establish this
dynamic hydrogen bonding network. Hence, they can no
further move and rotate randomly with the other water
molecules as they can in the bulk water phase. They must
choose one of the neighboring water molecules as their
sole interaction partner. This automatically reduces the
degree of disorder and leads to an unfavorable change in
entropy. If aligand is released from the water phase on
its way into the protein pocket, these previously frozen
degrees of freedom will be activated again, providing an
entropically favorable contribution.
When the ligand enters asupposedly uncomplexed
binding pocket of aprotein, it almost always encounters
acluster of water molecules that must be displaced during
the binding process. However, the local density of water
molecules can vary greatly, from apocket that is actually
empty to one that is lled with water molecules similar
in density to the bulk water phase. Some of the water
molecules form hydrogen bonds with the protein and are
in aspatially xed, energetically favorable arrangement.
Other water molecules are in contact with lipophilic regions of the protein surface and cannot form aperfect
hydrogen bonding network. At most, they form alocal
cluster of interconnected water molecules. The exchange
of these structurally very different water molecules will
be associated with very different desolvation costs. An
extreme example is the 0 pocket of the metalloprotease thermolysin (Sect.25.3). In this pocket, the enzyme
binds its substrates via the hydrophobic side chains of
the amino acids leucine or isoleucine. If the butyl moiety
of these side chains is replaced by those of the smaller
amino acids such as glycine, alanine, or valine, the afnity for substrates with these amino acids decreases dramatically. The transition from the glycine derivative to
the leucine derivative represents a41,000-fold increase in
afnity! Any medicinal chemist would dream of achieving this increase in afnity with only four carbon atoms
of an aliphatic substituent. Experimentally, Stefan Krim
mer was able to show that the 0 pocket of uncomplexed
thermolysin is practically free of water. Thus, binding of
the substrate does not require desolvation of the pocket.
The result is an extremely potent, almost exclusively
enthalpically preferred binding. The enzyme needs this
trick to efciently discriminate between substrates with
small and large aliphatic residues of ahomologous series.
This example shows how useful it can be for any drug designer to study the substrates between which their target
proteins can discriminate.
The displacement of xed water molecules usually
results in an entropic binding contribution, since the released hydrogen bonds of the displaced water molecules
are replaced by comparably strong hydrogen bonds to the
incoming ligand (. Fig.4.6). The enthalpic inventory is,
thus, to arst approximation balanced. However, the released water molecules have many more degrees of freedom in the bulk water phase, which implies an entropic
advantage. This consideration is even more valid when
we focus on hydrophobic regions of aligand. We have
already learned that release from the original cavity in
the water phase is entropically favored. Now we add the
entropic contribution for the displacement of xed water
molecules from the pocket (. Fig.4.6). Based on these
considerations, it was suggested that the hydrophobic
binding contribution is essentially entropic in nature (the
so-called “hydrophobic effect”). In the meantime, this
picture no longer seems to hold. Hydrophobic binding
can also be enthalpic. This is again due to the properties
of the water molecules released from the protein during
-
4
Chapter • Protein–Ligand Interactions as the Basis for Drug Action
ligand binding (. Fig.4.7). If one encounters alargely
empty water-free pocket, or apocket lled with only
afew isolated water molecules that are already highly disordered in the protein-bound state prior to their displacement, there is hardly any entropic gain to be achieved
upon their release. In such cases, the hydrophobic effect
may have apreferential enthalpic signature.
However, water can be critical to the energy balance
at acompletely different point in the binding process.
As mentioned above, aligand is solvated in water before
binding to the protein. In the water phase, it can adopt
several conformations. Barbara Wienen-Schmidt and
Tobias Hüfner have studied the binding of a series of
ligands to protein kinaseA. Surprisingly, the ligand in
the series with the most rotatable bonds shows the most
entropically favorable binding signature. Intuitively, one
would expect the opposite, since aconformationally exible ligand that is locked into asingle conformation in
the protein binding pocket should pay ahigh entropic
cost for the loss of these degrees of freedom on protein
binding. To investigate this phenomenon, the rst aspect
to check was whether the protein undergoes changes that
are associated with aloss of entropy. However, no evidence for this could be found. It was then necessary to
rule out the possibility that achange in the solvation
water inventory of the binding process could explain
the entropic advantage. Again, this could not explain
the observed entropic effect. Finally, it was investigated
whether the more exible ligand exhibits any peculiarities in aqueous solution prior to binding. It was observed that the conformation of the unbound ligand
entraps awater molecule so tightly that it can hardly
exchange with the surrounding bulk water. Evidence for
this behavior was provided by NMR studies (Sect.13.7)
and molecular dynamics simulations (Sect.15.8). Upon
binding of this ligand to the protein, the tightly bound
water molecule had to be released, which was accompanied by asignicant increase in entropy. In the end, this
contribution was so large that it became determinant for
the entire binding prole.
Water can still make an important contribution at
the nal step of the protein–ligand complex formation.
When anew protein–ligand complex forms, the ligand
adopts its binding pose in the protein pocket. In most
cases, substituents of the bound ligand protrude into
the surrounding water forming part of the surface of
the newly generated complex. Both binding partners
create anew, commonly exposed surface. Water molecules cluster around this surface and generate anew
solvate shell, which can display ahigh order. The network of these water molecules can be used to enhance
ligand binding through asuitable choice of side chain.
In particular, when fused ring structures of water molecules are formed, this improves protein binding. For
example, Stefan Krimmer and Jonathan Cramer were
able to increase the binding afnity through tailored de-
. Fig. 4.6 Inuence of water molecules on the strength of protein–
ligand interactions. Top row The formation of an H-bond between
protein and ligand requires the displacement of water molecules,
which themselves form H-bonds to the protein or ligand. The H-bond
inventory, meaning the number of H-bonds present before and after
the binding, is balanced in this case. Bottom row When ahydrophobic
contact is formed, water molecules are released from an environment
that is unfavorable for them. In the bulk water phase, they can form
H-bonds among one another
sign by afactor of50 for the thermolysin ligands shown
in . Fig.4.8.
4.7 Thermodynamic Contributions
to the Formation of Protein–Ligand
Complexes
Having considered the individual factors involved in
ligand binding, let us now turn to the successive steps
of complex formation. As mentioned above, the binding afnity is crucial for assessing the strength of aprotein–ligand complex. As outlined in Sect.4.3, it can be
described in terms of the binding constant (Eq.4.1) or
the change in Gibbs free energy of binding (Eq.4.3). This
is based on equilibrium thermodynamics. However, biological systems are so-called open systems for which
a“steady state” can be assumed as arst estimate. Of
course, application of equilibrium thermodynamics is
only acrude approximation. However, describing open
systems would further complicate the already very complex relationships. It is crucial for our considerations that
besides purely energetic contributions, an entropic component must always be taken into account.
. Fig.4.9 summarizes the steps involved in the for-
mation of acomplex between aligand and its protein.
At room or body temperature, protein and ligand can
ab cd
. • Thermodynamic Contributions to the Formation of Protein–Ligand Complexes
. Fig. 4.7 Example of the displacement of two water molecules from
the S1 pocket of thrombin upon binding of ligands4.2, 4.3, and4.4.
Ligand4.2 binds and two water molecules are found in the pocket(a).
One water molecule (orange) connects4.2 to Asp 189 via ahydrogen
bond and is in axed position. The second water molecule (green) is
dynamic and is located above Tyr 228 (compare(d), results of aneutron scattering study on the related trypsin). Comparing the binding
of4.2(a) with that of4.3(b), the xed orange water molecule is dis-
placed. In the next step to4.4(c), the mobile green water molecule also
move in all spatial directions. In simple terms, during
complex formation, two independent particles become
one common particle. This reduction in degrees of freedom is expressed as aloss of afnity, ∆G, of approximately 16–20 kJ/mol. This contribution must rst be
produced by the newly formed interactions in the complex. From the considerations in Sect.4.6, it is clear that
this system, which at rst glance appears to be binary,
must be extended by alarge number of water molecules.
The water molecules of the solvation shell are also mobile, diffusing back and forth. Some of them are locally
xed for extended periods of time, especially if they are
bound to the ligand or protein by multiple H-bonds.
Such water molecules can often be identied during
X-ray structure determination of aprotein. Other water molecules are freely mobile and, therefore, much
more difcult to detect in an X-ray structure determination. As described in Sect.13.5, hydrogen atoms cannot usually be spatially resolved in X-ray structures of
protein–ligand complexes. However, when neutron diffraction is used for structural analysis, H-atoms can be
characterized with ahigh degree of condence. Astudy
by Johannes Schiebel on ligand complexes of trypsin
showed that water molecules in the binding pocket exhibit different degrees of order depending on the bound
ligand (. Fig.4.10). The dynamic properties of these
water molecules affect the enthalpy/entropy prole of
the ligand binding.
Conformationally exible ligands can adopt multiple
conformations in aqueous solution, which, as described
above, can also efciently entrap water molecules. In the
binding pocket of the protein, the conformational freedom
leaves the pocket. The change in thermodynamic properties between
the three states going from4.2(a) to4.3(b), and to4.4(c) is summarized by the relative differences ∆∆G, ∆∆H, and T∆∆S. Displacement
of the orange water molecule leads to astrong entropic contribution
(negative value, indicated here and in the following gures as relative
difference (−T∆∆S
disordered green water molecule also leaves the S1 pocket(c). However,
in this case its displacement now proves to be an entropically unfavorable process (positive value for −T∆∆S
)). When4.4 is bound, compared with4.3, the
4.2⇨4.3
)
4.3⇨4.4
of the ligand is signicantly limited. The conformation ad-
opted at the protein site may differ from that in solution.
In addition, the ligand must be desolvated. Desolvation of
acharged ligand requires signicantly more energy than
desolvation of an uncharged ligand. The price to be paid
is an unfavorable enthalpic binding contribution.
As described above, the incorporation of the ligand
into the protein requires the displacement of water mole-
cules. The associated thermodynamic signature depends
strongly on the properties of the water molecules prior
to their displacement. It should be noted, however, that
ligand insertion can change the properties of all water
molecules remaining in the pocket (. Fig.4.10). In some
cases, the water molecules participate in ligand binding
and mediate contacts between the binding partners. It is
also important to remember that proteins can undergo
conformational changes during binding, which also contribute to the energy balance. Sometimes binding opens
pockets that were not present in the uncomplexed structure. Two extreme cases can be distinguished. Either the
protein adopts several conformations on average over
time, including the geometry of the ligand-bound state.
The ligand selects and stabilizes this geometry during
binding (conformational selection). However, there are
also situations in which the desired conformation is so
energetically unfavorable that it practically does not occur in the ensemble of possible conformations of the uncomplexed protein. Only when the ligand docks onto the
protein surface does the pocket temporarily unlock and
the ligand occupies the opening region. This induced t
represents the other extreme for binding processes that
occur under structural adaptation of the protein.
4
Chapter • Protein–Ligand Interactions as the Basis for Drug Action
. Fig. 4.8 Binding of a series of ligands with scaffold4.5 to the
metalloprotease thermolysin. The ligands place their side chainR in
the bowl-shaped 0 pocket (indicated by the gray surface), which is
open to the surrounding solvent, and form anew surface of the complex together with the protein. It is solvated by water molecules. In
the gure, the surface water molecules (red spheres) in contact with
the protein and the hydrophobic side chain are shown for three singled-out complexes. Depending on their geometry, an increasingly
4.8 What Is the Contribution of
aHydrogen Bond to the Strength of
Protein–Ligand Interactions?
Now that we have qualitatively gone through the individual steps of protein–ligand complex formation, we turn to
the question of how large the contribution of aparticular
interaction, e.g., ahydrogen bond, actually is to the bind-
ing afnity. This question can be answered experimentally
by comparing two protein–ligand complexes that differ in
their nal composition by only one hydrogen bond. Such
acomparison can be made, for example, by using protein
mutants in which an amino acid that forms ahydrogen
well-adapted network of water molecules wraps around the ligand like
ahood. In particular, the formation of ring structures can improve the
strength of binding to the protein. Within the series, the binding afnity (expressed as arelative change ∆∆G) of the ligands for thermolysin
increases by afactor of 50, enhancing the inhibitory potency of the
ligands. (Figure is taken from Krimmer etal., J.Med. Chem., 59 (23),
10530–10548 (2016), courtesy of the American Chemical Society)
bond with the ligand is replaced by another amino acid
that is unable to form this interaction.
In the enzyme aldose reductase (Sect. 27.4), the
amino acid Leu 300 has been replaced by proline, thereby
losing its ability to form an H-bond with the inhibitor
darestat4.6 (. Fig.4.11a). This was conrmed crystal-
lographically. The loss of the H-bond is accompanied by
adrop in afnity of 7.8 kJ/mol, and this loss is largely determined by enthalpy. In another experiment with the enzyme, Tyr48 was replaced by Phe, again with the loss of
an H-bond, here to the charged carboxylate group of the
ligand IDD594(4.7). Thus, acharge-assisted H-bond is
lost in this case. The afnity loss is somewhat larger with
. • What Is the Contribution of aHydrogen Bond to the Strength of Protein–Ligand Interactions?
. Fig. 4.9 Illustration of thermodynamic contributions to the
change in Gibbs free energy ∆G. Before binding, the ligand can move
freely in water. It has acertain translational and rotational entropy. In
addition, the ligand is usually exible and adopts different conformations in solution. Protein and ligand are solvated, forming H-bonds
to water molecules (blue spheres). Some water molecules are in loose
contact with the protein or ligand, while others are forming strong
H-bonds. Translational and rotational degrees of freedom are lost
during binding. The associated decrease in entropy is unfavorable for
binding. In addition, the protein and ligand must shed some of their
hydrate shell, also aprocess unfavorable for binding. The binding of
8.5 kJ/mol, but this is partitioned into alarge unfavorable enthalpic contribution and an opposite entropically
favorable component. The strongly divergent proles
show that for charged groups additional contributions
from stronger desolvation, electrostatic interactions and
changes in the solvate structure have to be considered.
The effects are even larger for asalt bridge, e.g., when an
amide function is added to the thrombin inhibitor4.8
(Sect.23.4) that has an unsubstituted phenyl ring to produce4.9. The binding mode is retained and the added
group forms asalt bridge with Asp 189 in the enzyme’s
S1 pocket. The change in Gibbs free energy for binding
of −14.4 kJ/mol demonstrates asubstantially increased
afnity. This improvement is due to astrong negative
exothermic enthalpy contribution, but is partly compensated by an unfavorable entropic contribution. These and
other examples show that the ∆G contribution of ahydrogen bond can vary between about −5 and −15 kJ/mol,
largely depending on the given local charge conditions.
The partitioning into enthalpy and entropy can extend
over an even larger energy range, indicating that both can
compensate their contributions to some extent.
It is also interesting to determine the inuence of
awater molecule on the binding prole, which mediates
an interaction between protein and ligand. For this pur-
the ligand leads to the formation of new direct interactions with the
protein and releases water molecules from the protein pocket. Both are
contributions that favor binding. In part, water molecules can mediate
binding to the protein. The complex formed establishes anew surface
that is re-solvated. H-bonds are shown as blue dashed lines, hydrophobic contacts analogously in yellow. (7 https://sn.pub/iyZd6R)
pose, the complex of darestat4.6 in aldose reductase is
again compared with the analogous complex of sorbinil
4.10. The two inhibitors differ in the attached carboxamide group. In darestat4.6, this group forms adirect
hydrogen bond to the NH function of the amide group
of Leu 300 (. Fig.4.11a). The replacement of leucine
for proline is accompanied by the loss of an H-bond.
This results in achange in Gibbs free energy amounting
to adecrease of 7.8 kJ/mol. Sorbinil 4.10 lacks the carboxamide groups (. Fig.4.11d). Interestingly, the free
enthalpy of binding for the Leu 300 →Pro exchange
now remains almost the same. Since sorbinil lacks the
group to form an H-bond with the NH group of Leu 300,
the removal of the NH function in the proline variant is
hardly noticeable. This explains the almost unchanged
binding free energy. Nevertheless, binding to the wildtype enzyme is enthalpically more favorable, but entropically more “expensive” than for the mutant. The crystal
structure indicates awater molecule that mediates an
H-bond between the ether group of sorbinil and the NH
function of Leu 300 (. Fig.4.11d). This results in an
enthalpy gain of −5.1 kJ/mol. At the same time, however,
entrapping awater molecule is entropically unfavorable.
This contribution of 5.9 kJ/mol just compensates for the
enthalpic gain, leaving virtually no afnity benet in ∆G
ab
cd
ef
Chapter • Protein–Ligand Interactions as the Basis for Drug Action
4
. Fig. 4.10 Experimentally determined water structure in the S1
binding pocket of trypsin before ligand binding(a), after binding of
2-aminopyridine (b), aniline (c), benzylamine (d), N-amidinopiperidine(e), and benzamidine(f). In the uncomplexed structure, ve water
molecules are found surrounding the aspartate residue (Asp 189), of
which W1 and W2 are scattered over two orientations. They indicate
dynamic properties of these two water molecules in the pocket (dynamic water molecules are labeled in blue, compare arrows). The bind-
ing of 2-aminopyridine(b) and aniline (c) displaces two of the ve
water molecules in the S1 pocket. In the 2-aminopyridine complex, W1
remains disordered and W2 adopts an ordered arrangement. In the
aniline complex, W2 remains dynamic and W1 adopts an ordered geometry. In the complex with benzylamine(d), both waters are ordered.
In the case of N-amidinopiperidine (e) and benzamidine (f), W1 is
ordered and W2 is displaced from the complex. The results are based
on structural data determined by combining neutron (density in green)
and X-ray scattering experiments (density in orange)