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Yihong Qiu, Coordinator
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Understanding Physicochemical
Properties for Pharmaceutical
Product Development and
Manufacturing—Dissociation,
Distribution/Partition, and Solubility
Deliang Zhou
Welcome to “Product and Process Design.”
“Product and Process Design” discusses scientific and
technical principles associated with pharmaceutical
product development useful to practitioners in validation and compliance. We intend this column to be a
useful resource for daily work applications. The primary
objective for this feature: Useful information.
Information developed during product and process
design is fundamental to drug and product development and all subsequent activities throughout the
product lifecycle. The quality-by-design (QbD) initiative encourages understanding of product and process
characteristics and process control based on risk analysis.
Process design is the first stage of pharmaceutical process validation as described in the recent US Food and
Drug Administration process validation draft guidance.
This stage comprises activities that support product and
process knowledge leading to a consistent manufacturing process throughout the product lifecycle. Product
development work provides much of the information in
the product and process design stage, including active
pharmaceutical ingredient information, dosage form
characteristics, quality attributes, and so on. Work conducted in development supporting product and process
design stage must be based on scientific and technical
principles. Disciplines supporting this work include
basic chemistry and pharmaceutics; pharamcokinetics
For more Author
information,
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including drug absorption, metabolism, distribution,
and excretion; biopharmaceutics information, and so
on. The principles associated with these areas are
broad and complex. Also, the technical language and
mathematics associated therein may be esoteric and
intimidating. This column addresses topics relevant to
product and process design with these difficulties in
mind. It is our challenge to present these topics clearly
and in a meaningful way so that our readers will be able
to understand and apply the principles in their daily
work applications.
Reader comments, questions, and suggestions are
needed to help us fulfill our objective for this column.
Please send your comments and suggestions to column
coordinator Yihong Qiu at [email protected]
or to journal coordinating editor Susan Haigney at
[email protected].
KEY POINTS
The following are key points addressed in this
discussion:
• Understanding the physicochemical properties
of the active pharmaceutical ingredient (API) is
fundamental to pharmaceutical product development, manufacturing, and stability
• Most pharmaceuticals are organic molecules,
many of which are weak acids or weak bases that
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ABOUT THE AUTHOR
Deliang Zhou, Ph.D., is an associated research investigator in Global Formulation Sciences, Global
Pharmaceutical R&D at Abbott Laboratories. He may be reached at [email protected]. Yihong
Qiu, Ph.D., is the column coordinator of “Product and Process Design.” Dr. Qiu is a research fellow
and associate director in Global Pharmaceutical Regulatory Affairs CMC, Global Pharmaceutical R&D at
Abbott Laboratories. He may be reached at [email protected].
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partially dissociate in aqueous solution
• Dissociation may be characterized by acidity dissociation constant (K a) or basicity dissociation
constant (K b), and their corresponding pK a and
pKb
• Buffer systems such as citrate and phosphate
buffers that resist pH changes upon the addition
of small quantities of acid or base are based on
dissociation behavior or weak acids and weak
bases
• Acid/base dissociation has significant influence
on almost all other physicochemical properties including salt formation, API purification,
formulation, processibility, solubility, stability,
mechanical properties, and biopharmaceutical
performance
• Distribution or partition refers to the relative
solubility of a drug between aqueous and nonaqueous immiscible liquids
• Distribution or partition is characterized by logD
or logP
• The distribution/partition phenomenon is important in manufacturing and laboratory extraction
processes to separate organic and inorganic compounds. Partition is the foundation underlying
various chromatographic techniques
• Distribution/partition is important in determining the biological, pharmacological, pharmaceutical, and toxicological activities of a drug
molecule. It is widely used in designing drug
candidates to possess appropriate absorption,
distribution, metabolism, and excretion (ADME)
properties
• n-Octanol/water is the most commonly used system to characterize drug partitioning
• A solution is a single-phase mixture consisting
of two or more components, usually solute and
solvent
• A solution in which the solute is above the solubility limit (supersaturated) is not thermodynamically stable and will ultimately lead to solute
crystallization
• The United States Pharmacopeia has defined
solubility definitions such as very soluble, freely
soluble, slightly soluble, and other terms
• Molecular interactions important in the solution
process include Van der Waals forces, hydrogen
bonds, and ionic interactions
• Temperature greatly influences the solubility of a
drug molecule and is always noted when solubility values are reported
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• The solubility product of an insoluble compound
is the basis for applications of the common ion
effect
•T
he aqueous solubility of a weak acid or base is
strongly pH-dependent due to relative species
abundances at different pH
• The solubility of polymorphs can be significantly
different because they have different free energies.
The metastable polymorph can convert to the
more stable polymorphs through solution-mediated transformation. The solubility difference
among various solid forms is one of the driving
forces responsible for phase changes during pharmaceutical processing.
• Solubility plays a central role in a number of issues
related to drug dissolution, absorption, formulation and process development and manufacture,
as well as stability
• Solubility is of paramount significance to the
biological performance of a drug molecule. The
physiology of the gastrointestinal (GI) tract plays
an important role in drug solubilization, including the pH environment and solubilization by
bile salts. pH change is particularly relevant for
weakly acidic or basic drugs. Solubility changes
along the GI tract could also cause changes such
as formation of salt, free acid or base, and hydrate,
amorphous, and polymorph conversion.
• Solubility is extensively utilized in crystallization and purification of active pharmaceutical
ingredients and intermediates. Experimental
conditions, such as pH, solvent, antisolvent,
temperature, and other factors may be modified
to optimize processes, control solid forms, and
maximize yields.
• Solubility is one of the underlying reasons during
many process-induced phase transformations
such as hydrate formation during wet granulation
and amorphous transition during drying.
INTRODUCTION
Physicochemical properties are the foundation for
pharmaceutical product development, manufacturing, and stability. They also provide insights to the
biopharmaceutical performance of a drug molecule.
A general appreciation of these subjects is important
for scientists and engineers involved in pharmaceutical product development and manufacturing.
This column discusses basic physical properties
including: Dissociation of weak acids and bases,
distribution and partitioning between aqueous and
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Yihong Qiu, Coordinator.
non-aqueous liquids, and solubility. This column
starts with basic definitions, follows with important
considerations, and completes each discussion with
practical relevance. Particular attention is given to
the contributions of dissociation to distribution and
solubility. The significance of these basic properties
is highlighted in the contexts of biopharmaceutics,
pharmaceutical development, manufacturing, and
stability.
DISSOCIATION AND THE DISSOCIATION
CONSTANT
The vast majority of pharmaceuticals are essentially
organic molecules, a large fraction of which are weak
acids or weak bases. Typical acidic groups include
carboxylic acid (–COOH), sulfonic acid (–SO3H),
sulfonamide (–SO2–NH–), and phenols (C6H5–OH).
Primary, secondary, and tertiary amines, as well as
pyridines and other aromatic amines, are the most
frequently encountered basic groups. These molecules tend to dissociate partially in aqueous solutions, where equilibrium exists between the various
ionized and unionized species.
The Ionization Equilibrium
The dissociation of a monoprotic weak acid, HA, can
be represented as follows:
HA
H+ + A–
Similar to chemical equilibrium, the dissociation
equilibrium is defined as follows:
[H+][A–]
Ka =
[HA]
Where Ka is known as the acidity dissociation constant
or simply as the ionization constant, and the bracket [
] represents concentration of a particular species at
equilibrium. The concentrations, instead of activities,
are used throughout because of their simplicity, and
appropriateness for the purpose of this general discussion. Similarly to the notion of “pH”, pKa is more
commonly used to denote the negative logarithm of
the acidity dissociation constant. The pKa value, often
a small positive number, is more convenient because
most Ka values are so small. It is obvious that the acidity decreases with pKa. A typical organic carboxylic
acid has pKa ~ 5, while phenols and sulfonamides are
weaker, with typical pKa values of 8-10.
Similar treatments can be applied to a weak base:
B + H 2O
BH+ + OH–
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In this case, the proton is transferred from the water
to the base, generating a hydroxyl ion. The basicity
dissociation constant is then as follows:
[BH+][OH–]
Kb =
[B]
By convention, the concentration term of water is
dropped because it is in large excess and its concentration does not change appreciably. It is also well
known that water itself is very weakly dissociable to
produce a proton and a hydroxyl ion as follows:
H 2O
H+ + OH–
The self-dissociation constant of water is often designated as Kw = [H+] [OH–], which equals to 1 × 10 –14
at 25°C. Substituting the Kw in the previous equation,
one obtains the following:
Kw[BH+]
Kw
Kb =
=
[H+][B]
Ka
where Ka is the acidity constant of BH+, the conjugate
acid of the base B, as shown below:
BH+
H+ + B
Ka =
[H+][B]
[BH+]
Therefore, the basicity constant of a weak base can
be represented by the acidity constant of its conjugate
acid as: pKa = pKw – pKb, as is conventionally done.
Basicity increases with increasing pK a, contrary to
the acidity.
Acids and bases can be categorized as monoprotic,
diprotic, or polyprotic based on the number of protons
they can accept or donate. Multiple ion equilibriums exist simultaneously for a polyprotic electrolyte,
and each has its own dissociation equilibrium. For
example, phosphoric acid dissociates in the following three steps:
Ka1
H 3PO4
H+ + H 2 PO4 –
H 2 PO4
HPO42–
Ka2
H+ + HPO42–
Ka3
H+ + PO43–
These stepwise pKas for phosphoric acid are 2.21,
7.21, and 12.7, respectively. The acid becomes progressively weaker.
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Figure 1: Species distribution as a function of
pH for succinic acid.
1.0
H2 Suc
0.8
H Suc
–
( )
[A–]
pH = pKa + log
[HA]
2–
Suc
Fraction
0.6
0.4
0.2
0
0
2
4
6
pH
8
10
12
Species distribution and buffer capacity are direct
results from dissociation and are discussed next.
Species Distribution In Solution
In an aqueous solution of a weak electrolyte various
species, such as the unionized parent, the proton, and
the ion, coexist. The relative fraction of each species
depends on its pKa value and the solution pH, which
can be calculated based on the dissociation equilibrium. For example, consider a monoprotic acid HA:
log
( )
[A–]
[HA]
= pH - pKa
Therefore, the ionized species will dominate at
pH > pKa, while the unionized neutral molecule will
dominate at pH < pKa. To preferably target a certain
species, either ionized or unionized, we often follow
the rule of two pH units from pKa, because the ratio is
either 100 to 1 or 1 to 100. This is particularly useful
in separation science such as high performance liquid
chromatography (HPLC), because mixed mode partition can cause peak broadening or even shouldering.
Similar equations can be worked out for bases. An
example species distribution is shown in Figure 1
for succinic acid. The hydrogen succinate displays a
maximum around pH of 5.
Buffer Capacity Of Weak Electrolyte
Solutions
Weak acids and bases are often used as buffering
agents (e.g., citrate buffer, phosphate buffer). A buffered solution, namely, refers to its ability to resist pH
changes upon the addition of small quantities of acid
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or base. Quantitatively, buffer capacity, β, is the ratio
the added strong base (or acid) to the resulted pH
change. For example, consider the following equation for a weak acid:
= pKa + log
[base form]
[acid form]
This equation is known as the buffer equation or
the Henderson-Hasselbalch equation, which describes
how pH changes with the ratio of its species concentrations. The right most part of the equation is more
general as it also applies to a weak base.
It can be demonstrated that the maximum buffer
capacity occurs at pH = pK a, where concentrations of
the acid and base forms are equal. The exact buffer
capacity can be obtained by differentiating the Henderson-Hasselbalch equation:
β = d[base]/dpH = 2.303 Ctot =
[H+]Ka
([H+]Ka)2
This is the exact reason that most buffers are used
around their pKa values in order to achieve their intended function. Unlike the strong acid and base buffers, a
weak acid and base buffer allows the buffering pH to
be maintained while adjusting the buffer capacity by
changing the total buffer concentration, Ctot.
Significance Of Dissociation
Acid/base dissociation is a simple concept, yet has
significant influences on almost all other physicochemical properties. Salt formation is well built upon
the dissociation concept, which has been extensively
exploited during active pharmaceutical ingredient
(API) manufacture for purpose of purification, processibility, and other aspects of pharmaceutical development such as solubility, stability, mechanical properties, and biopharmaceutical performance. A general
rule of thumb for salt formation is that the difference
in pK a between the drug and the counter ion should
be at least three to increase the chance of success.
Dissociation can also significantly impact stability,
because the ionized and unionized species usually
have different intrinsic reactivity—their respective
rates of degradation may be significantly different.
The more stable form of the API can thus be selected
for product development. One or both of these forms
can be catalyzed by the proton or hydroxyl ions, as
evidence from the pH-stability profiles. Buffering
agents have been frequently used to maintain a foriv thome.com
Yihong Qiu, Coordinator.
mulation at the desired pH for optimal stability, which
is particularly important for parenteral and other
liquid formulations.
Figure 2: Example pH - LogD profile for a weak base with
pK a of 9.
4
DISTRIBUTION AND PARTITION
where C o and C w are the equilibrium concentration of
molecule in the oil and in the water, respectively. The
distribution coefficient can also be approximated as
the solubility ratio. Because distribution coefficients
are usually large for organic molecules, they are often
converted to logarithms, known as “LogD”. Following the convention, the natures of the partitioning
phases are often noted, such as LogD n-octanol / water or
LogD n-octanol / pH 7.4 buffer.
The partition coefficient, LogP, is usually reserved
to the partitioning of the neutral or unionized species,
which can be deemed as the intrinsic distribution
coefficient of a molecule (as to be shown, LogD is
pH-dependent with ionizable molecules). Another
term, called cLogP, refers to calculated LogP value
based on molecular fragmentation algorithms.
LogD For Weak Electrolytes
An ionized species has much less affinity to the oil
phase than the corresponding neutral species. It could
be well assumed that only the unionized species distributes in both phases while the ionized species is
concentrated in the aqueous phase. The pH-distribution relationship can then be derived. For example,
the following equation holds for a monoprotic weak
base:
Ka
LogD = LogP + log
[H+] + Ka
Figure 2 shows such a plot for a weak organic base
with a LogP of 4 and a pK a of 9. LogD decreases
linearly with pH with a slope of negative unity at pH
< pK a. It even becomes negative below pH 5, which
renders itself practically not extractable by oil phase
and can be utilized in separation.
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)
2
LogD
When an immiscible liquid, such as hexane or n-octanol, is added to an aqueous drug solution, the drug
molecules will distribute themselves in each of the
two immiscible phases. This phenomenon is known
as distribution or partition. At equilibrium, the ratio
of the concentrations in each phase is constant, which
is a characteristic of the drug molecule itself and the
nature of the two phases. This ratio is called the distribution coefficient of the drug as follows:
D = CO/CW
0
-2
-4
0
2
4
6
8
10
12
pH
Applications Of Distribution/Partition
The distribution/partition phenomenon is common and has been utilized to a great extent by various industrial and laboratory extraction processes.
Organic compounds can be readily separated from
inorganics by extracting with an immiscible solvent.
Separation of organic mixtures can be possible if their
LogD values are sufficiently different, or by manipulating their LogD values via pH modification.
Partition is the foundation underlying various chromatographic techniques. In HPLC, a solute molecule is
repeatedly partitioned between the stationary (column)
and the mobile (eluent) phases. The number of equivalent partitioning process is so large that minute differences in the solute molecules can be magnified and thus
successfully separated. The pH of the mobile phase is an
important factor to optimize the interactions between
solute and the stationary phase and between the solute
and the mobile phase, because ionized and unionized
species have enormously different LogD values. Both
single mode and mixed mode elution have been utilized
in modern HPLC development.
Distribution/partition also plays an important
role in determining the biological, pharmacological, pharmaceutical, and toxicological activities of a
drug molecule.
The distribution/partition characteristics are directly related to the lipophilicity of a molecule. Lipophicility characterizes the affinity of a molecule to lipid,
fat, or oil. It is well known that the cell membrane
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Product and Process Design.
has the phospholipid bilayer structure, which consists
of various phosphoglycerides. There is a similarity
between water/oil partitioning of a molecule and its
ability to penetrate cell membrane. As a perquisite
for a molecule to function biologically, the molecule
itself has to be able to pass the cell membrane and
penetrate into the cell. Indeed, Hansch et al. (1) established a correlation between n-octanol and water
partition coeffienct and biological activities in the era
of quantitative structure activity relationship (QSAR).
Since then, the n-octanol/water distribution/partition
coefficient has been well established in the pharmaceutical industry and other related fields.
In pharmaceutical sciences, distribution/partition
coefficient is now widely used in designing drug candidates to possess appropriate absorption, distribution, metabolism, and excretion (ADME) properties.
The majority and the most convenient route of oral
administration of a drug requires the drug molecule
to be able to cross the gut wall membrane and get
absorbed into the blood stream, and similar requirements hold true for routes of administration other
than intravenous applications. In the absence of carrier mediated absorption (nutrients or nutrient-like
molecules) or paracelluar diffusion (small hydrophilic
molecules), passive diffusion in which the drug crosses the bilayer of the intestinal epithelium remains
the primary mechanism for most drug absorption.
A molecule needs to be sufficiently lipophilic to be
able to cross the gastrointestinal (GI) membrane.
Therefore, an appropriate LogD value is required for
passive permeability in the GI epithelium. After a
drug gets absorbed into the blood stream, the drug
molecules are then carried to the various tissues and
distributed throughout the body. This process can
be deemed similarly as partition. It is generally true
that a hydrophilic drug is distributed primarily in the
plasma, while a lipophilic molecule is more extensively distributed in peripheral tissues and organs. It
has been well known that sufficient lipophilicity is
necessary for a drug molecule to pass the thick lipid
bilayer called blood-brain-barrier (BBB), a tight layer
of glial cells surrounding the capillary endothelial
cells in the brain and spinal cord. Therefore, the distribution coefficient is of vast importance in designing molecules to have pharmacological activities in
the brain. A recent survey (2) on drug metabolism
has also suggested that lipophilic molecules are usually more extensively metabolized in the body, while
hydrophilic ones are often excreted unchanged in
the urine or bile, relating to its ability to access the
enzymes responsible for metabolism. Therefore, the
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LogD/LogP values can provide tremendous insight on
the drug’s relevant fates in the body.
Distribution/partition has also been useful in
environmental sciences. Hydrophobic chemicals
tend to stay in the environment longer and are more
difficult to clean up. They also tend to preferably
get into various living species. Hence, the partition
coefficient is an important consideration in environmental regulations.
Partitioning Systems
For historical reasons, the n-octanol/water partition
coefficient has prevailed compared to other solvent
systems in the pharmaceutical fields. The octanol/
water system was initially thought to model essential
properties of typical biological membrane. However, a thorough literature review has not resulted
in compelling arguments. Partition between water
and other immiscible solvents, such as hexane, cyclohexane, heptane, isooctane, dodecane, hexadecane,
and chloroform have also appeared in the research
literature, but far less frequently.
From a physicochemical property point of view, the
n-octanol molecule is characterized by a hydrophobic backbone (i.e., the C8 group) and a hydrophilic
hydroxyl head. The hydroxyl group can be both a
hydrogen-bond donor and a hydrogen-bond acceptor.
In addition, water-saturated octanol has ~25 mol %
water, that has led to the suggestion of the 1:4 tetrameric structure. However, it was later determined by
X-ray diffraction analysis that a cluster structure is
more evident where ~16 octanol molecules point their
hydroxyl groups to the water cluster by forming an
extensive hydrogen-bonding network (3). Therefore,
a cluster structure’s interaction with a drug molecule
can be complicated and may not capture all the lipid
permeation characteristics of a molecule.
The proposition of using multiple systems to better characterize biological membrane modeling has
also appeared, such as the “critical quartet” system
(4). However, there has not been much development in this area. The quartet system is composed
of n-octanol (a hydrogen-bond acceptor and donor),
chloroform (a hydrogen-bond donor), alkane (an
inert medium), and propylene glycol dipelargonate
(a hydrogen-bond acceptor).
SOLUBILITY
A solution is a single-phase mixture consisting of two
or more components. A binary solution is composed
of only two components, while a ternary solution is
composed of three components. The liquid compoiv thome.com
Yihong Qiu, Coordinator.
Table: USP designations on solubility.
Part of solvent required
for 1 part of solute
Solubility in mg solute
per mL of solvent
USP terminology
Less than 1 part
>1 g/mL
Very soluble
1-10 parts
100-1000 mg/mL
Freely soluble
10-30 parts
33.3-100 mg/mL
Soluble
30-100 parts
10-33.3 mg/mL
Sparingly soluble
100-1000 parts
1-10 mg/mL
Slightly soluble
1000-10,000 parts
0.1-1 mg/mL
Very slightly soluble
More than 10,000 parts
<0.1 mg/mL
Practically insoluble
nent is usually called the solvent, and the dissolved
component is the solute. But the distinction is not
absolute and becomes vague when two liquid components are concerned.
A drug may not be dissolved in a solvent at any
concentration. Indeed, above the solubility limit, a
solution will not be thermodynamically stable. Still
widely employed today as the most reliable method,
solubility can be determined by equilibrating excess
solids with a solvent. Equilibrium is deemed reached
when the solution concentration reaches an asymptote—the dissolving solvent contains the maximum
amount of solute. At this point, the solution is said
to be saturated with the solute and the solution is
known as a saturated solution. The solution concentration equals the solubility in a saturated solution.
In the language of thermodynamics, the free energy
of the solute in solution equals that of the solid solute. A solution is unsaturated when concentration is
less than the solubility, while a supersaturated solution
refers to one whose concentration is higher than the
solubility. One might be curious as how a solution
could become supersaturated because macroscopic
dissolution ceases once its concentration reaches the
solubility. Well, there are many scenarios where this
could happen. For example, temperature can significantly affect solubility. Therefore, one can purposely
create an unsaturated or supersaturated solution by
merely adjusting the temperature of the solution.
A supersaturated solution is thermodynamically
unstable and will lead to solute crystallization when
adequate time is given. The reason is that the free
energy of solute molecules in a supersaturated solution is always higher than those in the solid phase,
and the crystallization process is spontaneous by
going from a higher free energy state to a lower free
energy state. However, rate of crystallization (kinetics) could vary. Crystallization needs to start with
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nucleation, a process where nuclei are formed, that
serves as the sites for crystal growth. A supersaturated
solution may be kinetically stable for a long duration
without crystallization. Parent crystals, dusts, and
other exogenous materials can in many cases serve as
the nucleation sites and are often utilized in laboratory and industrial crystallization processes.
Solubility is often reported in units such as weight
by volume (w/v, e.g., mg/mL), weight-by-weight (w/w),
molar fraction, molarity, or molarlity. These units
can be inter-converted when adequate information is
given. The most frequently used units for drugs are
mg/mL or µg/mL, because solubility usually falls in
these concentration ranges. It should also be noted
that the United States Pharmacopeia (USP) has a
special solubility designation system, as shown in
the Table.
Solubility And Molecular Interactions
Let’s take a look at schematic solution formation process at the microscopic level. According to Hess’s law,
energy change of a process is path independent, only
the initial and final states being of importance. This
path independence is true for all state functions, such
as enthalpy, free energy, and entropy. Therefore, the
solution formation can be deemed to be composed
of the following three essential steps:
• Step 1. A solute molecule is liberated from its
crystal lattice. Work is needed to overcome the
lattice energy. This process could be approximated as the melting of the solute crystal. A high
melting temperature and melting enthalpy usually
imply high lattice energy.
+
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Product and Process Design.
• Step 2. A solvent cavity is created, large enough
to hold the solute molecule. This process needs to
overcome the cohesive attractions among solvent
molecules.
•S
tep 3. Finally, a solution is formed by placing
the solute molecule into the solvent cavity created
in step 2. Energy is released resulting from solventsolute interaction. Hopefully this energy is large
enough to compensate the work needed in the previous two steps. It should be pointed out, though,
that the energy in the first two steps need not be
entirely compensated, because the entropy gained
after mixing is also favorable, and can negate part
of the energy consumed during steps 1 and 2.
+
Tremendous insight can be gained on solubility by
just considering these processes. The solubility of a
compound can be low due to: high melting point;
high melting enthalpy; lack of solute-solvent interaction; and too strong solvent-solvent cohesion.
For a specific compound and a specific solid form,
the melting temperature and enthalpy is given and can
not be changed. However, the solvent-solute interaction and solvent-solvent cohesion can be modulated
and an appropriate solvent system may be designed
to achieve desired solubility. For this reason, the following molecular interactions are important:
•V
an der Waals forces. Van der Waals forces
include both dipole-dipole, dipole-induced dipole,
as well as induced dipole-induced dipole interactions. A molecule having permanent dipole tends to
interact with another dipole by pointing to opposite
pole and this interaction is known as dipole-dipole
or Keesom forces. Permanent dipoles are also able to
induce an electric dipole in nonpolar but polarizable
neighboring molecules, and this interaction is called
the dipole-induced dipole or Debye force. Finally,
transient polarization exists in nonpolar molecules
because of the electronic cloud movements in the
molecule. This induced dipole-induced dipole interaction is called the dispersion or London force. All
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these van der Waals forces decrease quickly with
distance (1/r6) and are usually in the magnitude
of 1 – 10 kcal/mol. On an individual basis, they
are weak and follow the rank order Keesom force
> Debye force > dispersion force. However, they
cannot be ignored for a collection of molecules.
Even the weakest dispersion force is sufficient to
cause the condensation of nonpolar gas molecules
and crystallization of nonpolar liquids.
•H
ydrogen bonds. Hydrogen bonds refer to
the interaction between a hydrogen atom and a
strongly electronegative atom such as fluorine,
oxygen, and nitrogen. The nature of hydrogenbond is largely electrostatic interaction between
the small size and large electrostatic field of the
hydrogen atom, and the strong electronegativity
of the acceptor atoms. Hydrogen bond requires
certain direction, is a fairly strong intermolecular interaction (2-7 kcal/mol), and is always in
addition to the van der Waals forces. It exists in
water, hydrogen fluoride, and ammonium, and
is responsible for their abnormally high boiling
points and high melting points in their corresponding homologous series.
• I onic interactions. Ionic interaction is the electrostatic interaction (attraction and repulsion)
between ions of the opposite or same charges.
They are the primary interactions responsible for
ionic crystals. Solubility of ionizable molecules
also depends partially on the ion-dipole and ioninduced dipole interactions between the solute
and solvent. A polar solvent has high dielectric
constant (e.g., water) that can reduce the ion-ion
interaction. Therefore, ionizable compounds such
as salts usually are more soluble in polar solvents
than in non-polar solvents.
The empirical rule on solubility states “like dissolves like.” The rule itself is a bit vague; however,
one may make judgment calls by examining the
interactions stated previously. A polar molecule is
likely more soluble in polar solvents. When a drug
molecule is capable of forming hydrogen bonding,
a solvent capable of satisfying such requirements is
likely to provide improved solubility. In the absence
of any polar group, a drug molecule is probably more
soluble in a nonpolar solvent such as alkanes and
benzene. A drug molecule that contains both polar
and nonpolar groups may require a solvent or a mixed
solvent capable of providing simultaneous interactions
to both parts of the molecule.
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Yihong Qiu, Coordinator.
Temperature greatly influences the solubility of a drug
molecule; therefore, the temperature is always noted
when solubility values are reported.
The influence of temperature on solubility depends
on the heat of solution. Solubility decreases with
temperature when the heat of solution is negative
(i.e., when the dissolution is an exothermic process).
Solubility increases with temperature if the solution
process absorbs heat (endothermic). This observation is consistent with the Le Chatelier’s principle on
equilibrium. The magnitude of the heat of solution
determines the steepness of solubility change when
temperature is altered. The quantitative relationship
is well captured by the Van’t Hoff equation.
Some molecules have a more complicated temperature-solubility relationship. For example, phenol shows an upper consolute temperature (UCP) of
66.8°C, above which it can be mixed with water at any
ratio. Triethylamine demonstrates a lower consolute
temperature (LCT) in water of ~18°, below which
the two are miscible in all portions. Nicotine-water
system shows both an UCP and LCT. The existence of
UCP and LCT has something to do with the solute-solvent interactions and their temperature dependences.
For example, the hydrogen-bonding interaction could
be destroyed at higher temperature; therefore, an LCT
could form.
Sometimes solute can form complexes with solvent
called solvates. It is then the solvated form that controls the equilibrium solubility. A solvate is usually
less soluble in the same solvent than the unsolvated
form due to thermodynamic reasons. A solvate may
become unstable and dissociate above a certain temperature. Therefore, the temperature-solubility line
may break its trend. Sodium sulfate is an example
of this type of interaction.
Solubility Product And The Common Ion
Effect
When a salt is dissolved, it dissociates into the respective anion and cation. At equilibrium, the product of
the concentrations of these ions is constant. This is
known as the solubility product, often represented as
K sp. For example, the dissolution of silver chloride
is as follows:
AgCl(s)
Ag+ (aq.) + Cl– (aq.)
where s and aq. represent solids and aqueous,
respectively. Its solubility product is as follows:
Ksp = [Ag+][Cl–]
gxpandjv t.com
Figure 3: A pH-solubility profile for a weak acid with pKa of 4
and intrinsic solubility of 1 µg/mL.
1.E+07
1.E+06
1.E+05
Solubility
Effect Of Temperature
1.E+04
1.E+03
1.E+02
1.E+01
1.E+00
0
2
4
6
8
10
pH
Therefore, the solubility of silver chloride can be
reduced if additional chloride (e.g., NaCl) or silver (e.g.,
silver nitrate) ions are introduced into the solution.
This phenomenon is known as the common ion effect.
pH-Solubility Profile
The aqueous solubility of a weak acid or base is strongly pH-dependent due to relative species abundances
at different pH. The concentration ratio between the
ionized and unionized species is related via the Henderson-Hasselbalch equation, as discussed previously.
The overall solubility can then be derived for a weak
acid or base, noting that the unionized species is in
equilibrium with the excess solid (i.e., at the intrinsic
solubility S0). For example, the total solubility of a
weak acid is as follows:
S = S 0 (1 + 10pH–pKa)
A typical pH-solubility plot is shown in Figure 3 for
a weak acid with pK a of 4 and intrinsic solubility of
10 µg/mL. Below pH 4, solubility is essentially constant and is equal to the intrinsic solubility. However,
above pH 4, the solubility increases exponentially
with pH. The solubility will not, however, increase
monotonically without a limit. Salt formation could
also take place; therefore, the solubility reaches a plateau (shown in Figure 3 as the blue dash line), the
value of which is determined by the solubility product
of the salt. The broken red line is just an extension of
the pH-solubility profile if a salt is not formed.
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Product and Process Design.
Solubility Of Polymorphs And Solvates
Significance Of Solubility
Polymorphism is the existence of multiple crystal
forms by the same chemical entity. They exist as different molecular packings in the crystal lattice. The
reasons leading to polymorphism are complex but
are related to the ability of a molecule to efficiently
pack itself in a unit cell with different conformation,
orientation, or hydrogen-bonding motifs. Crystal
hydrate is the incorporation of water molecules in a
drug crystal unit cell. Similarly, a solvate could be
formed. In addition, amorphous form refers to the
non-crystalline solid when the long-range order in
a crystal is destroyed. All these can be collectively
called the solid forms of a drug molecule.
The solubility of polymorphs can be significantly
different because they have different free energies. For
example, the solubility of ritonavir Form II is about
1/4 of that of Form I at 5°C (5). The most stable form
has the least free energy and lowest solubility. The
free energy difference between two polymorphs is
related to the following solubility ratio:
Solubility plays a central role in a number of issues
related to drug dissolution, absorption, formulation
and process development and manufacture, as well
as stability.
Solubility is of paramount significance to the biological performance of a drug molecule. Drug molecules, when presented as oral solid dosage forms,
need first to dissolve in the GI fluid before absorption can take place. The dissolution rate of a solid is
proportional to its solubility, diffusivity, and surface
area. Therefore, the higher the solubility, the faster
a molecule can get into solution, and the higher its
concentration in the GI fluid. In turn, this leads to
a higher rate of absorption, because the passive diffusion is driven by the concentration gradient across
the GI membrane.
The physiology of the GI tract plays an important
role in drug solubilization, including the pH environment and the solubilization by various bile salts. For
example, the human stomach is characterized by a low
pH of 1.5-2 under fed conditions and pH 2-6 under
fasted conditions. The pH increases going down to
duodenum (pH 4.5-5.5), small intestine (pH 6-6.5),
ileum (pH 7-8), and colon (pH~5.5-7). The small
intestine is also characterized by a high concentration
of bile salts and the large surface area attributing to
its villi and micro-villi structures. The small intestine
serves as the primary site of drug absorption. On
the contrary, the large intestine does not possess the
villi structure. It also lacks the fluid to solubilize a
drug and is challenging for drug absorption, except
for certain highly soluble and highly permeable molecules. The change in pH, bile salt concentration,
fluid contents, motility, as well as the residence time,
all affect the solubility and in vivo dissolution of a
drug molecule, and influences its oral absorption.
The pH change is particularly relevant for weakly
acidic or basic drugs. Solubility changes along the
GI tract could also cause solid form changes, such as
salt formation, parent formation, hydrate formation,
as well as polymorphic conversion, all complicating
the drug absorption process.
Appropriate solubility and solvent system are also
required to develop other formulations such as parenteral,
subcutaneous, transdermal, and aerosol formulations.
Solubility is extensively utilized in crystallization
and purification of active pharmaceutical ingredients
and/or their intermediates. Experimental conditions,
such as pH, solvent, antisolvent, temperature, and
other factors may be modified individually or in com-
ΔG(I – II) = RT ln(SI/SII)
The metastable polymorph can convert to the
more stable polymorphs through solution-mediated
transformation, where the solution concentration is
higher than the solubility of more stable form thus
providing the thermodynamic driving force for conversion. The ritonavir Form I to Form II transition in
the liquid-filled capsules is such an example. Still,
the amorphous form, being the most energetic, has
the highest solubility.
Solvates usually have less solubility in the same solvent than the unsolvated form because equilibrium
dictates so due to the large excess of solvent. Therefore,
a solvated form is often discovered during the crystallization in a particular solvent. There also exists
a critical solvent activity, above which the solvate is
more stable while below which the unsolvated form
is more stable. Specifically for hydrates, the critical
relative humidity (RH) or water activity comes to play a
significant role, because moisture is ubiquitous during
storage and because water is the most frequently used
solvent in wet granulation. It should also be noted
that a solvate has higher solubility in other miscible
solvents, because the dissociated solvent molecules
drag the drug molecules into solution.
The solubility difference among various solid forms
is one of the driving forces responsible for phase
changes during pharmaceutical processing.
22
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iv thome.com
Yihong Qiu, Coordinator.
bination in order to optimize the process. Solubility
plays an important role in controlling the solid forms
as well as the yields.
Amorphous forms have been extensively utilized
to improve drug bioavailability because they have
higher apparent solubility (6).
Solubility is one of the underlying reasons for
many process-induced phase transformations (7).
For example, hydrate formation can occur during wet
granulation. The hydrate may then partially or fully
dehydrate during drying, resulting in formations of
partially amorphous or less ordered molecular form.
A highly water-soluble, low dose drug may completely
dissolve during wet granulation and turn into amorphous upon drying. All these changes can have profound effects on product attributes such as stability,
dissolution, and in vivo performance. These phase
changes can be scale, time, and equipment dependent
due to their kinetic nature, which is worrisome from
the quality control point of view. The solubility and
solubility differences among the various solid forms
are responsible for, and are the keys to, understanding
and resolving these transformations as they provide
the thermodynamic driving forces for change.
SUMMARY
Dissociation, distribution/partition, and solubility
constitute the fundamental physical properties of a
drug molecule. They influence, directly or indirectly,
almost every aspect in drug development: Absorption,
distribution, metabolism and excretion, formulation development, processing development, stability,
product manufacture, and regulatory concerns. Many
issues during pharmaceutical development can be
related to characteristics in these basic properties in
one way or another. Awareness and knowledge of
these subjects is important to understand and solve
various issues during pharmaceutical development,
manufacturing, and stability.
fraction Analysis,” Journal of Pharmaceutical Sciences, 82,
707-712, 1993.
4.Leahy, D. E.; Morris, J. J.; Taylor, P. J. and Wait, A. R.,
“Membranes and Their Models: Towards a Rational
Choice of Partitioning System,” Pharmacochem. Libr.
16, 75-82, 1991.
5.Bauer, J.; Spanton, S.; Henry, R.; Quick, J.; Dziki, W.;
Porter, W. and Morris, J., “Ritonavir: An Extraordinary
Example of Conformational Polymorphism,” Pharmaceutical Research, 18, 859-866, 2001.
6.Law, D.; Schmitt, E. A.; Marsh, K. C.; et al., “RitonavirPEG 8000 Amorphous Solid Dispersions: In vitro and
in vivo Evaluations,” Journal of Pharmaceutical Sciences,
93, 563-570, 2004.
7.Zhang, G. G. Z.; Law, D.; Schmitt, E. A. and Qiu, Y.,
“Phase Transformation Considerations During Process
Development and Manufacture of Solid Oral Dosage
Forms,” Advanced Drug Delivery Reviews, 56, 371-390,
2004. JVT
ARTICLE ACRONYM LISTING
ADME Absorption, Distribution, Metabolism,
Excretion
API
Active Pharmaceutical Ingredient
BBBBlood-Brain-Barrier
FDA
US Food and Drug Administration
GIGastrointestinal
HPLC High Performance Liquid Chromatography
Ka
Acidity Dissociation Constant
Kb
Basicity Dissociation Constant
LCT
Lower Consolute Temperature
QSAR Quantitative Structure Activity Relationship
RH
Relative Humidity
UCP
Upper Consolute Temperature
USP United States Pharmacopeia
REFERENCES
1.Hansch, C. and Leo, A. Substituent Constants for Correlation Analysis in Chemistry and Biology, Wiley-Interscience,
New York, 1979.
2.Custodio, J. M.; Wu, C.Y. and Benet, L. Z., “Predicting
Drug Disposition, Absorption/Elimination/Transporter
Interplay and the Role of Food on Drug Absorption,”
Advanced Drug Delivery Reviews 60, 717-733, 2008.
3.Franks, N. P., Abraham, M. H., and Lieb, W. R., “Molecular Organization of Liquid n-Octanol: An X-ray Difgxpandjv t.com
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