Martin's Physical Pharmacy and Pharmaceutical Sciences, Sixth Edition

       Martin's Physical Pharmacy and Pharmaceutical Sciences, Sixth Edition

Ever since the First Edition of Martin’s Physical Pharmacy was published in 1960, Dr. Alfred Martin’s vision was to provide a text that introduced pharmacy students to the application of physical-chemical principles to the pharmaceutical sciences. This remains a primary objective of the Sixth Edition. Martin’s Physical Pharmacy has been used by generations of pharmacy and pharmaceutical science graduate students for 50 years and, while some topics change from time to time, the basic principles remain constant, and it is my hope that each edition reflects the pharmaceutical sciences at that point in time.

As with prior editions, this edition represents an updating of most chapters, a significant expansion of others, and the addition of new chapters in order to reflect the applications of the physical-chemical principles that are important to the Pharmaceutical Sciences today. As was true when Dr. Martin was at the helm, this edition is a work in progress that reflects the many suggestions made by students and colleagues in academia and industry. There are 23 chapters in the Sixth Edition, as compared with 22 in the Fifth Edition. All chapters have been reformatted and updated in order to make the material more accessible to students. Efforts were made to shorten chapters in order to focus on the most important subjects taught in Pharmacy education today. Care has been taken to present the information in “layers” from the basic to more in-depth discussions of topics. This approach allows the instructor to customize their course needs and focus their course and the student's attention on the appropriate topics and subtopics.

With the publication of the Sixth Edition, a Web-based resource is also available for students and faculty members (see the “Additional Resources” section later in this preface).

Each chapter begins with a listing of Chapter Objectives that introduce information to be learned in the chapter. Key Concept boxes highlight important concepts, and each Chapter Summary reinforces chapter content. In addition, illustrative Examples have been retained, updated, and expanded. Recommended Readings point out instructive additional sources for possible reference. Practice Problems have been

moved to the Web (see the “Additional Resources” section later in this preface).


Important changes include new chapters on Pharmaceutical Biotechnology and Oral Solid Dosage Forms. Three chapters were rewritten de novo on the basis of the valuable feedback received since the publication of the Fifth Edition. These include Chapter 1 (“Introduction”), which is now called Interpretive Tools; Chapter 20 (“Biomaterials”), which is now called Pharmaceutical Polymers; and Chapter 23 (“Drug Delivery Systems”), which is now called Drug Delivery and Targeting.

Martin’s Physical Pharmacy and Pharmaceutical Sciences, Sixth Edition, includes additional resources for both instructors and students that are available on the book’s companion Web site at

Approved adopting instructors will be given access to the following additional resources:


One of the earmarks of evidence-based medicine is that the practitioner should not just accept the conventional wisdom of his/her mentor. Evidence-based medicine uses the science-tific method of using observations and literature searches to form a hypothesis as a basis for appropriate medical therapy. This process necessitates education in basic sciences and an understanding of basic scientific principles.”1,2 Today more than ever before, the pharmacist and the pharmaceutical scientist are called upon to demonstrate a sound knowledge of biopharmaceutics, biochemistry, chemistry, pharmacology, physiology, and toxicology and an intimate understanding of the physical, chemical, and biopharmaceutical properties of medicinal products. Whether engaged in research and development, teaching, manufacturing, the practice of pharmacy, or any of the allied branches of the profession, the pharmacist must recognize the need to rely heavily on the basic sciences. This stems from the fact that pharmacy is an applied science, composed of principles and methods that have been culled from other disciplines. The pharmacist engaged in advanced studies must work at the boundaries between the various sciences and must keep abreast of advances in the physical, chemical, and biological fields in order to understand and contribute to the rapid developments in his or her profession. You are also expected to provide concise and practical interpretations of highly technical drug information to your patients and colleagues. With the abundance of information and misinformation that is freely and publicly available (e.g., on the Internet), having the tools and ability to provide meaningful interpretations of results is critical.

Historically, physical pharmacy has been associated with the area of pharmacy that dealt with the quantitative and theoretical principles of physicochemical science as they applied to the practice of pharmacy. Physical pharmacy attempted to integrate the factual knowledge of pharmacy through the development of broad principles of its own, and it aided the pharmacist and the pharmaceutical scientist in their attempt to predict the solubility, stability, compatibility, and biologic action of drug products. Although this remains true today, the field has become even more highly integrated into the biomedical aspects of the practice of pharmacy. As such, the field is more broadly known today as the pharmaceutical sciences and the chapters that follow reflect the high degree of integration of the biological and physical-chemical aspects of the field.

Developing new drugs and delivery systems and improving upon the various modes of administration are still the primary goals of the pharmaceutical scientist. A practicing pharmacist must also possess a thorough understanding of modern drug delivery systems as he or she advises patients on the best use of prescribed medicines. In the past, drug delivery focused nearly exclusively on pharmaceutical technology (in other words, the manufacture and testing of tablets, capsules, creams, ointments, solutions, etc.). This area of study is still very important today. However, the pharmacist needs to understand how these delivery systems perform in and respond to the normal and pathophysiologic states of the patient. The integration of physical-chemical and biological aspects is relatively new in the pharmaceutical sciences. As the field progresses toward the complete integration of these subdisciplines, the impact of the biopharmaceutical sciences and drug delivery will become enormous. The advent and commercialization of molecular, nanoscale, and microscopic drug delivery technologies is a direct result of the integration of the biological and physical-chemical sciences. In the past, a dosage (or dose) form and a drug delivery system were considered to be one and the same. A dosage form is an entity that is administered to patients so that they receive an effective dose of a drug. The traditional understanding of how an oral dosage form, such as a tablet, works is that a patient takes it by mouth with some fluid, the tablet disintegrates, and the drug dissolves in the stomach and is then absorbed through the intestines into the bloodstream. If the dose is too high, a lower-dose tablet may be prescribed. If a lower-dose tablet is not commercially available, the patient may be instructed to divide the tablet. However, a pharmacist who dispenses a nifedipine (Procardia XL) extended-release tablet or an oxybutynin (Ditropan XL) extended-release tablet to a patient would advise the patient not to bite, chew, or divide the “tablet.” The reason for this is that the tablet dosage.


For molecules to exist as aggregates in gases, liquids, and solids, intermolecular forces must exist. An understanding of intermolecular forces is important in the study of pharmaceutical systems and follows logically from a detailed discussion of intramolecular binding energies. Like intramolecular bonding energies found in covalent bonds, intermolecular bonding is largely governed by electron orbital interactions. The key difference is that covalency is not established in the intermolecular state. Cohesion, or the attraction of like molecules, and adhesion, or the attraction of unlike molecules, are manifestations of intermolecular forces. Repulsion is a reaction between two molecules that forces them apart. For molecules to interact, these forces must be balanced in an energetically favored arrangement. Briefly, the term energetically favored is used to describe the intermolecular distances and intramolecular conformations where the energy of the interaction is maximized on the basis of the balancing of attractive and repulsive forces. At this point, if the molecules are moved slightly in any direction, the stability of the interaction will change by either a decrease in attraction (when moving the molecules away from one another) or an increase in repulsion (when moving the molecules toward one another).

Knowledge of these forces and their balance (equilibrium) is important for understanding not only the properties of gases, liquids, and solids but also interfacial phenomena, flocculation in suspensions, stabilization of emulsions,.


Thermodynamics deals with the quantitative relationships of interconversion of the various forms of energy, including mechanical, chemical, electric, and radiant energy. Although thermodynamics was originally developed by physicists and engineers interested in the efficiencies of steam engines, the concepts formulated from it have proven to be extremely useful in the chemical sciences and related disciplines like a pharmacy. As illustrated later in this chapter, the property called energy is broadly applicable, from determining the fate of simple chemical processes to describing the very complex behavior of biologic cells.

Thermodynamics is based on three “laws” or facts of experience that have never been proven in a direct way, in part due to the ideal conditions for which they were derived. Various conclusions, usually expressed in the form of mathematical equations, however, may be deduced from these three principles, and the results consistently agree with observations. Consequently, the laws of thermodynamics, from which these equations are obtained, are accepted as valid for systems involving large numbers of molecules.

It is useful at this point to distinguish the attributes of the three types of systems that are frequently used to describe thermodynamic properties. Figure 3–1a shows an open system in which energy and matter can be exchanged with the surroundings. In contrast, Figure 3–1b and c are examples of closed systems, in which there is no exchange of matter with the surroundings, that is, the system’s mass is constant. However, energy can be transferred by work (Fig. 3–1b)or heat (Fig. 3–1c) through the closed system’s boundaries. The last example (Fig. 3–1d) is a system in which neither matter nor energy can be exchanged with the surroundings; this is called an isolated system.

For instance, if two immiscible solvents, water and carbon tetrachloride, are confined in a closed container and iodine is distributed between the two phases, each phase is an open system, yet the total system made up of the two phases is closed because it does not exchange matter with its surroundings.


An atom consists of a nucleus, made up of neutrons (neutral in charge) and protons (positively charged), with each particle carrying a weight of approximately 1 g/mole. In addition, electrons (negatively charged) exist in atomic orbits surrounding the nucleus and have a significantly lower weight. Charged atoms arise from an imbalance in the number of electrons and protons and can lead to ionic interactions (discussed in Chapter 2). The atomic mass is derived from counting the number of protons and neutrons in a nucleus. Isotopes may also exist for a given type of atom. For example, carbon has an atomic number of 6, which describes the number of protons, and there are several carbon isotopes with different numbers of neutrons in the nucleus: 11C (with five neutrons), 12C (with six neutrons), 13C (with seven neutrons), 14C (with eight neutrons), and 15C (with nine neutrons).1 Carbon-13, 13C, is a common isotope used in nuclear magnetic resonance (NMR) and kinetic isotope effect studies on rates of reaction, and 14C is radioactive and used as a tracer for studies that require high sensitivity and for carbon dating. Both 11C and 15C are very short-lived, having half-lives of 20.3 min and 2.5 sec, respectively, and are not used in practical applications.

Molecules arise when interatomic bonding occurs. The molecular structure is reflected by the array of atoms within a molecule and is held together by bonding energy, which relies heavily on electron orbital orientation and overlap. This is illustrated in the Atomic Structure and Bonding Key Concept Box. Each bond in a complex molecule has intrinsic energy and will have different properties, such as reactivity. The properties within a molecule depend on intramolecular interactions, and each molecule will possess net energy of bonding that is defined by its unique composition of atoms. It is also important to note that for macromolecules like proteins or synthetic polymers, the presence of neighboring charges or dipole interactions, steric hindrance and repulsion, Debye forces, and so on, within the backbone or functional groups of the molecule can also give rise to distortions in the primary intramolecular binding energies and can even give rise to additional covalent bonding, for example, disulfide bridges in a protein like insulin.

The more bonding that an atom participates in, the lower is the density of the electron cloud around the nucleus and the greater is the shared distribution of atomic electrons in the interatomic bonded space. For example, because of differences in atomic properties such as electronegativity, the nature of each interatomic bond can vary greatly. Here, the “nature” refers to the energy of the bond, the rotation around the bond, the vibration and motion of atoms in the bond, the rotation of a nucleus of an atom in a bond, and so on.

As an example of the nature of a bond, consider the carbonyl–amide bond between two amino acids in a protein(Fig. 4–1), as described by Mathews and van Holde.3 The oxygen in the carbonyl is very electronegative and pulls electrons toward itself. The lone pair of electrons on the nitrogen remains unbound, but the nitrogen has a considerably lower electronegativity than oxygen. Recall that the covalent bond occurs between the carbon (which also has a low electronegativity) in the carbonyl group and the amine of the next amino acid. However, the electrons in the carbon-oxygen double bond in the carbonyl are pulled much closer to the oxygen due to its greater electronegativity, causing a partially negatively charged oxygen. The carbon then carries a partial positive charge and pulls the lone pair of electrons from the nitrogen. This forms a partial double bond, which causes the nitrogen to carry a partial positive charge, yielding a net dipole (see Permanent Dipole Moment of Polar Molecules section). Interestingly, in extended secondary structures, in particular α-helices, the effective dipole of the peptide bond can have a strong stabilizing effect on secondary structure.4 This reduces rotation around the carbon-nitrogen bond, making only cis/trans isomers allowable. Although this is not discussed here, the student should recall discussions about energetically favored bond rotations in organic chemistry. The trans peptide bond orientation is energetically preferred, with the cis conformer appearing in only specialized cases, in particular, around the conformationally constrained pro-line. As discussed in general biochemistry courses, the secondary structure of a peptide is then determined by rotation around the φ and ψ angles only. Finally, amino acids naturally occur in the l-conformer, with the exception of glycine. This chirality is followed in almost all peptides and proteins, resulting in a profound effect on function. An example of chirality with an l- and a d-amino acid is presented in the Key Concept Box Chirality of Amino Acids as a refresher.


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