Contents
MOLECULES AND MEDICINE
THE WILEY BICENTENNIAL-KNOWLEDGE FOR GENERATIONS
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PREFACE
This book is intended for a broad readership, starting with curious and thoughtful college undergraduates and, reaching beyond, to professionals and researchers in the life sciences and medicine. It is hoped that it will also be useful to the educated lay person with an interest in health and medicine.
An effort has been made to integrate chemistry, biology, drug discovery and medicine in a way that is clear and self-explanatory. Heavy use has been made of chemical structures, since they provide a fundamental key to the language of life and the human activities that flow from it. Our age has seen the rapid evolution of molecular medicine as a critical part of the broader fields of health care and the biochemical basis of human disease. The understanding of human illness at the molecular level has brought, and will bring, great benefit to mankind.
There is a price to be paid in any attempt to understand molecular medicine, because that comprehension requires an ability to decipher chemical structures, which many have regarded as too onerous. One purpose of this book is to demonstrate that an adequate understanding of chemical structures is easily within the reach of most educated people, and well worth the effort. Pages 4–31 of this book aim to provide the insights and background required to appreciate the architecture of therapeutic molecules and their target proteins, as they parade though the subsequent pages of this book.
“Molecules and Medicine” delves into the discovery, application and mode of action of well over one hundred of the most significant molecules now in use in modern medicine. It is limited to centamolecules, i.e. molecules with molecular weights in the hundreds (several hundred times more than a hydrogen atom). The important and rapidly developing area of macromolecular therapy, which involves much larger molecules (macro-molecules), such as biologically active proteins and monoclonal antibodies, is a different story, and another book.
We have tried to minimize the amount of prior knowledge required of the reader by providing much background information, both chemical and biomedical. An effort has also been made to be concise as well as clear. A large amount of material has been compressed into a small space for each therapeutic agent. Generally, each medicine is allotted just one page. One advantage of this modular arrangement is that it facilitates the reading of the book in small installments and also its use as a reference work. The therapeutic agents in this book are arranged in sections according to the type of medical condition they treat.
There are also numerous sections in the book that provide biomedical background. For instance, there are two-page to eight-page summaries of topics such as inflammation, metabolic syndrome, immunology, drug resistance, cancer and neurotransmission. These are placed at strategic locations throughout the book.
The structural representations of proteins in this book are in the public domain and may be downloaded from .
In the process of writing this book we have come to appreciate ever more keenly the enormous amount of human talent and effort that has enabled the extraordinary advances in molecular medicine since its advent several decades ago. To the countless chemists, physicians, biologists, educators and other professionals who have participated in this venture we extend our gratitude and thanks. This book is a tribute to them all.
The writing of this work has also been motivated by the realization that the advance of molecular medicine can be even more remarkable during the coming decades of this century, if a steady flow of dedicated and able young people into all the key areas of research in the life sciences can be maintained. If this book plays just a minor part in enhancing progress in molecular medicine and human well-being, our small effort will have been amply rewarded.
A NOTE ON THE USE OF THIS BOOK
Those who are not familiar with the notation used to describe the structures of organic compounds will profit from a close reading of pages 4–22 of this book. Others can read through this part quickly and proceed to the next section (pages 26–31) that reviews the notation used in the book for the structures of proteins.
The coordinates for each of the protein structures shown in the book are in the public domain and can be accessed electronically.
All the crystal structure files can be downloaded from by entering the four-character PDB ID that is indicated in red at the bottom of the page where the protein is displayed or in the reference section. For example, on page 63 in the entry for sitagliptin (Januvia™), the X-ray crystal structure of sitagliptin bound to the target protein DPP-4 is displayed. The access code (PDB ID) for this crystal structure appears below the picture as 1X70 along with the corresponding reference.
The renderings of the X-ray crystal structure data were developed by the authors using the software PyMol v0.99 (DeLanoSci-entific LLC, ). For a description of the Protein Data Bank, see:
H.M. Berman, J. Westbrook, Z. Feng, G. Gilliland, T.N. Bhat, H. Weissig, I.N. Shindyalov, P.E. Bourne: The Protein Data Bank. Nucleic Acids Research, 28, 235–242 (2000).
The discussion of each of the therapeutic agents in the book is limited to one page, and, consequently, much in the way of detail has been omitted. For this reason, a number of up-to-date references on other aspects of each agent are given both at the bottom of the appropriate page and at the end of each section. The page numbers on which detailed references are listed for each therapeutic agent are highlighted in green at the bottom of the pages. Additional information on the pharmacology and properties of each medicinal agent can be found in:
Goodman & Gilman’s The Pharmacological Basis of Therapeutics. Laurence L. Brunton, John S. Lazo and Keith L. Parker (Editors); (McGrawHill, 11th Edition, 2006).
Other recent texts that provide much useful background information are:
The references that appear in this book can serve as a portal to a great deal of information on the chemistry, biology and medicine relevant to the therapeutic agent and disease area. Additional material can be located by appropriate database- and internet searching (e.g., using SciFinder Scholar™ MEDLINE™ or Google™).
A glossary and an index appear at the end of the book.
There is a website for this book maintained by the authors, the address for which is:
Readers can use this website to provide comments or feedback on “Molecules and Medicine”. The website will also contain certain useful updates. The authors invite the views of non-scientists who have read pages 4–31.
INTRODUCTION
We live in a troubled, but wonderful time. It is our good fortune to witness and benefit from scientific advances that would have been literally unimaginable to our grandparents. However, there are dark clouds on the horizon. The rate of growth of scientific knowledge has been so great as to outstrip the ability of our society to assimilate it, the capacity of the educational system to teach it properly, and the wisdom of government adequately to sustain and apply it. There is widespread indifference to science among the young. Even medical science, which touches the lives of us all, is generally left to the practitioners. Whatever the reason for this disparity between the importance of science and the lack of general public understanding, it is important to address it.
In this book we try to take a few steps in this direction. Specifically, the pages to follow tell the tales of many molecules that can qualify as miracles of modernity. These relatively small, highly-structured clusters of atoms, the principal therapeutic agents of modern medicine, can perform in a way that would have been considered miraculous to our ancestors. Such “miracle molecules” can save countless human lives, prolong human life, alleviate pain and suffering, control cells, tissues and organs millions of times their size, and bring enormous material gains through commercial sales of billions of dollars per year. Such molecules also can serve as tools to probe the molecular nature of life processes and disease states and pave the way for the discovery of other effective medicines.
The molecules at the core of this book have been carefully selected from several thousand therapeutic agents that have been used in medicine at one time or another. The development of each of them, arduous and costly though it might have been, represents an enormously valuable investment with very large and ongoing benefits. In the course of discovering all these wonderfully useful molecules, we have learned more about the discovery process itself and have developed an ever expanding set of new discovery tools. The invention of these new platforms for innovation is being powered by dramatic advances in technology, computing and the underlying chemical and biomedical sciences.
The very next section of this book provides a step-by-step introduction to the understanding of the architecture of organic molecules and the general principles that govern structures of molecules. In addition, we explore the fundamental forces that hold molecules together and that allow them to recognize and bind to one another. The affinity of molecules for one another is central to the biological activity of therapeutic agents and to life itself. The section on how to read the chemical diagrams of small molecules is followed by another tutorial on understanding much larger structures, the proteins of life.
It seems quite possible that, in the next century or so, effective treatments for most illnesses will emerge. Disease, premature death, suffering and pain may no longer be a part of the human condition. Humans will as a matter of course live out a full and healthy lifespan, and then depart with grace and dignity. The famous poem of Lady Gio in “The Tale of the Heike” describes the life process and its end in an eloquent and happy way:
Grasses of the plain,
Springing up and withering,
They all fare alike.
Indeed the lot of all things
Is but to wait for autumn.
An impossible dream? Perhaps, but the immense effort required will be well worthwhile, because the gain will be incalculable. The achievements of modern science and technology provide both encouragement and inspiration.
For instance, we are now able to trace our universe back some 14 billion years to an unbelievably hot object, with a temperature of about 1032 Kelvin (10 followed by 32 zeros), and more than a million times smaller than the period at the end of this sentence. From this inferno of exceedingly small and simple objects, the first elements, hydrogen and helium, formed about a million years later, to be followed by all the other objects of the universe — the chemical elements, stars and galaxies, and an unknown collection of other forms of matter and energy, and finally the earth and life upon it. Surely, a time will come when our knowledge of life, intelligence, disease and health will dwarf that of the present.
Hydrogen, the simplest element, has atomic number one because it contains just one proton in the nucleus and one electron that surrounds it. A hydrogen atom is so reactive that it will totally combine with another hydrogen atom to form the simplest molecule, H2. This process, the simplest of all chemical reactions, takes place extremely rapidly and with the release of much energy because the energy content of H2 is much less than two isolated hydrogen atoms. The chemical equation for the reaction is:
Two hydrogen atoms combine to form one hydrogen molecule (H-H) with the evolution of energy.
In , the dots represent electrons. The chemical bond that holds H2 together is designated by a double dot, or simply, by a line between the hydrogens. Because the electrons are shared equally, the molecule is nonpolar. This type of two-electron bond is called a covalent single bond. Although an electron has a mass, it is very small (9×10−28 g, at rest), and behaves like a wave. In addition, because of the Heisenberg uncertainty principle, its position and momentum cannot both be known precisely. Its location is best described in a probabilistic way as a cloud-like representation ().
Approximate representation of the spatial distribution of electrons in an H2 molecule about the nuclei (red spheres). The outer layer of dots encloses about 90% of the total electron cloud.
Helium (He), atomic number two, has two protons in the nucleus and two surrounding electrons. The atom is stable chemically because it neither combines with itself nor reacts with other elements. The reason for this inertness is that two is the maximum number of electrons that can be accommodated in the available orbital. This orbital allows the electrons to distribute themselves like a cloud with spherical symmetry about the nucleus. This orbital is called the 1s orbital. The other orbitals of He are so high in energy that they are not accessible for chemical bonding with any atom or chemical fragment. Helium is the simplest of the inert elements. These elements share the feature of having the full complement of electrons in the available orbitals.
In H2, each H contributes a 1s atomic orbital (AO), leading to the formation of two molecular orbitals (MO’s), one of lower energy than the AO and the other of higher energy. The two electrons of H2 can occupy the lower energy MO, and so H2 is stabilized relative to two H atoms, as the diagram in illustrates.
Energy diagram for the combination of two H atom 1s AO’s to form two MO’s of diatomic H2. The electrons occupy the lower, bonding MO(1) and are shown as blue arrows. MO(2) is an orbital that is not occupied by electrons because of its high energy. MOs can hold only two electrons. These electrons must have opposite spins.
One simple way of thinking about the electron in the H atom is to consider it like a cloud of gas around the nucleus that is kept in place by the electrostatic attraction between the negative electron and the positively charged proton. Once H2 is formed, the two electrons of H2 become delocalized over a greater volume around the two-proton axis of the diatomic molecule, which leads to great stability. For a simple analogy, recall that the rapid expansion of a gas lowers its temperature (i.e., energy content). Conversely, compression of a gas raises its temperature/energy content.
The energy of the H2 molecule is at a minimum when the nuclei are separated by 0.75×10−8 cm or 0.75 Angströms (Å). That bond length is determined by a balance between two factors:
In general, chemical bonds between two atoms have characteristic bond lengths (d0), which result from the balancing of these two effects. A typical graph of the stabilizing energy of a bond as a function of bond length is shown in .
Energy of two hydrogens as a function of internuclear distance.
Figure 4. Three different representations of the hydrogen molecule.
A large amount of energy is required to break the H-H bond of H2 — 104 kcal/mol — (a mol is defined as 6.02×1023 atoms, Avogadro’s number). In the pure state the H2 molecule can be heated to 500 °C without any decomposition.
The bonding diagram for H2 allows us to explain why H2 does not react with another H atom to form H3 as a stable molecule. The bonding orbital MO(1) of H2 is filled since no more than two electrons can occupy a MO. No significant bonding can be attained by combining MO(2) of H2 with the AO of the H atom.
Most of the molecules in this book contain hydrogen. In every case each hydrogen forms one and only one electron-pair bond with its partner leading us to the first rule:
Rule #1: H forms just one electron-pair bond — called a covalent single bond.
This rule is useful in deriving the actual structures of molecules. Thus, the formula of water, H2O, and Rule #1 lead to the structure H-O-H, rather than O-H-H, or something else. Similarly, since methane has the composition CH4, we can draw the structure:
There are a number of alternative ways of representing H2. In addition to a line drawing H-H, we can use a ball-and-stick drawing or a space-filling diagram, as follows:
The space-filling representation contains additional information because it tells us about the shape and size of a molecule. These features are exceedingly important in determining how two molecules can fit together or interact and also how close they can get. With molecules, as with macroscopic objects, two things cannot occupy the same space at the same time.
Carbon, the element with atomic number six has six protons in the nucleus and six electrons surrounding it. Two of these fill the low energy 1s orbital and play no role in chemical bonding. The remaining four electrons can form bonds utilizing a spherical 2s orbital and three dumbbell-shaped, mutually perpendicular 2p orbitals, the shapes of which are illustrated in . The boundary of each dumbbell encloses about 90% of the total electron cloud. The three p orbitals are usually designated as 2px, 2py and 2pz because they can be placed along the axes of a rectilinear coordinate system. The sum of the electron densities for all three 2p orbitals is a spherical cloud of larger diameter than the 2s orbital.
Representation of s and p atomic orbitals. The boundary of each sphere and dumbbell encloses about 90% of the total electron cloud.
The 2s orbital can mix with three 2p orbitals to form hybrid atomic orbitals in three different ways. A sequence for generating the three different 2s/2p hybrids is shown in –. The number of hybrid orbitals always equals the number of AOs from which they derive.
Combination of one 2s and one 2px atomic orbital (AO) to form two sp hybrid orbitals.
Figure 7. Combination of two sp hybrid orbitals and a 2py atomic orbital to form three sp2 hybrid orbitals.
Combination of three sp2 hybrid orbitals and a 2py orbital to form four sp3 hybrid orbitals.
If all three of the 2p orbitals and the 2s orbital are hybridized, a set of 4 equivalent orbitals results, the axes of which are directed at the vertices of a tetrahedron (). This arrangement is used in forming methane and other compounds having four atoms attached to carbon (tetracoordinate carbon). The attachment of four hydrogen atoms (4 H) to a carbon atom results in the formation of four electron-pair bonds and the complete filling of the 4 bonding MOs by eight electrons (4 from C and 4 from H).
Formation of methane from four sp3 hybrid orbitals of carbon and four 1s orbitals of hydrogen.
The angle between any two of the C-H bonds in methane is 109.5°, the angle between the center of a tetrahedron and any two of the vertices (). The internuclear distance of each C-H bond is 1.54 × 10−8 cm or 1.54 angströms (Å). A second useful bonding rule emerges from these facts.
Rule #2: Carbon (C) can bond to a maximum of four atoms (tetracoordination). The preferred angle between any two of the bonds is 109.5°. Tetracoordinate carbon utilizes sp3 hybrid orbitals.
The angle between the center of a tetrahedron and any two of its vertices is 109.5°.
Deviation of the bond angles at an sp3-hybridized carbon atom from the preferred value of 109.5° leads to a higher energy (i.e., less stable) structure. The destabilization increases to a value of about 6 kcal/mol for an angle of 90°. This destabilization, called angle strain, influences the chemistry and properties of a compound.
Carbon forms strong bonds to most atoms, including H, oxygen (O), nitrogen (N), chlorine (Cl), and by no means least, to itself. Thus, methane is just the first in a large family of compounds of carbon and hydrogen (hydrocarbons). That family includes the straight-chain saturated hydrocarbons, the first five members of which are shown in .
The simplest straight-chain hydrocarbons.
Carbon chains can also be branched.
Figure 11a. Two simple branched hydrocarbons, isobutane and trimethylpentane.
There is a simpler notation for depicting the structures of carbon compounds in which the hydrogens are omitted. This shorthand notation leaves it to the reader to add the number of hydrogens corresponding to tetracoor-dination ().
Simplified notation of carbon compounds in which the hydrogens are omitted.
Carbon can form rings by bonding with itself in a cycle as well as chains. The simplest members of the family of cyclic hydrocarbons are shown in .
The first four members of cyclic hydrocarbons: cyclopropane, cyclobutane, cyclopentane and cyclohexane.
Carbon can form compounds in which three atoms are linked to it using hybrid orbitals generated from the combination of the 2s atomic orbital with two of the 2p atomic orbitals. The orbitals of trigonally hybridized carbon are shown in .
Side-on and Top-on views of the orbitals of a trigonally hybridized (sp2-hybridized) carbon atom.
Two of the simplest carbon compounds that involve tricoordinate (trigonally hybridized) carbon atoms are formaldehyde (H2C=O) and ethylene (H2C=CH2). These are planar molecules that contain a double bond to carbon as well as single (electron-pair) bonds to the hydrogens ().
Formaldehyde and ethylene contain trigonally hybridized carbon.
Bonding to tricoordinate carbon utilizes three sp2 hybrid orbitals and the remaining 2p AO. These orbitals allow the derivation of the correct geometry of molecules with tricoordinate carbon. For example, the planar structure of ethylene results because overlap of the p-orbitals is maximum when they are parallel, as shown in .
Formation of ethylene by the combination of two carbon atoms and four hydrogen atoms.
The linkage between the two carbons of ethylene is called a double bond because it involves four electrons. The double bond can be represented by two lines, as in the drawing in , or by σ and π bonds as shown in . The π bond, formed by the side-by-side combination of two parallel p atomic orbitals (shown in blue in ), has two lobes, one above and one below the molecular plane. The σ bond, formed by the combination of two colinear sp2 orbitals, is symmetric about the C-C axis (axial symmetry).
Figure 15b. π-Bond of ethylene.
Rule #3: Tricoordinate carbon is connected to each of the three attached atoms in a planar arrangement. The bonding involves three in-plane hybrid sp2 orbitals and an orthogonal p atomic orbital.
Dicoordinate carbon compounds utilize two sp orbitals formed from the hybridization of the 2s orbital with one 2p orbital, and also the remaining two 2p orbitals. The orbitals for this type of carbon are shown in ; note that the angle between the two sp orbitals is 180°. The py and pz orbitals that are not involved in hybridization remain unchanged.
Two views of the orbitals of a dicoordinate (sp-hybridized) carbon atom.
Two simple examples of dicoordinate carbon compounds are hydrogen cyanide and acetylene, both of which possess triple bonds and linear geometry because of the 180° angle between the sp orbitals of carbon ().
The two simplest dicoordinate carbon compounds, acetylene and hydrogen cyanide.
The triple bond consists of six electrons; two of these are in an axially symmetric MO formed from the combination of two sp AOs. The remaining 4 electrons are in two bonding π-MOS formed from overlap of the four 2p AOs.
The linear structure of acetylene follows from the use of the sp hybrid orbitals and p orbitals of each carbon and two H 1s orbitals to assemble the molecule, as shown in .
Formation of acetylene by the combination of two sp-hybridized carbon atoms and two hydrogen atoms.
An electron cloud representation of the two π bonds of acetylene is shown in . These two π-bonds and the sp-sp σ bond of acetylene hold six electrons and constitute a C-C triple bond.
The two π-bonds of acetylene.
Rule #4: Dicoordinate carbon forms bonds to the two attached atoms in a colinear arrangement. The bonding involves two colinear sp orbitals and two p atomic orbitals at carbon.
Carbon dioxide (CO2) is another linear molecule in which carbon is sp-hybridized ().
Structure and shape of carbon dioxide.
Most of the common elements that make life possible fall within the first three rows of the Periodic Table of Elements. These are shown in along with the corresponding atomic numbers. The atomic number of an atom is identical to the number of protons in the nucleus or the number of orbiting electrons.
A portion of the Periodic Table of Elements.
All the elements shown in combine with hydrogen, with the exception of the inert gases He, Ne and Ar. Some examples of the simplest of these are the following ().
Simple compounds of hydrogen and non-carbon elements.
The dot pairs in the above structures represent electrons in the outer valence shell that are not needed in bonding. The structure of water, for example, involves an sp3 hybridized oxygen atom connected to two hydrogen atoms. The single bonds to the hydrogens use up two out of the six available electrons of an O atom. The remaining four oxygen electrons are located in the remaining (nonbonding) sp3 orbitals. Since the four sp3 orbitals are filled by eight electrons, no further electrons, for instance from a hydrogen atom (H·), can be added. (Reminder: each orbital can hold only two electrons.)
Carbon can also bond to most of the elements, for instance replacing hydrogen in the compounds shown in .
Figure 22. Simple compounds of carbon with non-carbon elements.
Rule #5: Carbon can bond to itself to form either straight or branched chains or rings. Carbon can also bond to many other atoms.
The metallic elements at the left of the Periodic Table lose an electron very readily and tend to form positively-charged ions (cations), rather than covalently bonded compounds. The elements F, Cl, Br and I at the right of the Periodic Table, in contrast, have high electron affinity and readily accept an electron to form negative ions (anions). In the case of fluoride ion (F−) the available orbitals are filled, as with the inert gas neon (Ne) with which it is isoelectronic (i.e., both F− and Ne have a total of 8 outer shell electrons). The bonding between sodium and chlorine, for example, is essentially electrostatic, and the bond is described as ionic in character. It is the extreme of the covalent bond of H2 which involves no charge separation. Sodium chloride (NaCl, salt) in solid form is a crystalline structure containing Na+ and Cl−ions in an indefinitely repeating lattice in which each Na+ is surrounded by 6 Cl−, and vice versa. It is so stable that the melting point of salt is about 800 °C. The energy that holds Na+Cl− in the crystal lattice is 187 kcal/mol, much greater than the H-H covalent bond energy (104 kcal/mol).
Hydrogen chloride (HCl) is a gas at room temperature, in contrast to the ionic solid sodium chloride. The bonding in H-Cl is best described as a covalent bond with appreciable (but far from full) ionic character or charge separation. The electron pair between H and Cl is not shared equally. It is a polarized molecule with more electron density at the Cl end and less at the H end ().
The electron density around HCl. In this computer-generated electron-density map, the blue color represents the lowest electron density whereas the red color represents the highest electron density.
In aqueous solution HCl ionizes to form a hydrated proton and a hydrated chloride ion. Thus, it is a strong acid.
Equation 2. Dissociation of HCl in water.
The polarization of the bond in gaseous hydrogen chloride, often indicated using the notation Hδ+-Clδ-, is a consequence of greater electron affinity of a chlorine atom as compared to a hydrogen atom. Expressed in another way, chlorine is more strongly electron-attracting than hydrogen.
Polarization of covalent bonds is very common. Four examples are shown in .
Electron density maps of four simple compounds. The highest electron density is shown in red whereas the lowest electron density is shown in blue.
Oil and water do not mix because neither can dissolve the other. The former is essentially hydrocarbon-like and nonpolar, whereas water is polarized with the oxygen relatively negative and the two hydrogens positive.
The polarity of the O-H bonds in water causes the boiling point of water (100 °C) to be much higher than that, for instance, of methane (CH4, −161 °C) which has about the same size, or ammonia (NH3, −33 °C). The O-H bond polarity of water causes molecules of H2O to associate with one another, primarily because of electrostatic forces, forming an extended three-dimensional network, a small part of which is shown in .
An extended three-dimensional network of molecules in liquid water involves “hydrogen bonds” (blue dashes) as shown.
The bonds between molecules of H2O in the liquid — called “hydrogen bonds” — are much weaker than the H—H bond (104 kcal/mol), or the C-H bond in methane (105 kcal/mol). The bond dissociation energy of the covalent H-O bond in gaseous H2O is about 117 kcal/mol, whereas the energy of attraction of an H in water with an O atom of a neighboring water molecule is about 6 kcal/mol. Intermolecular hydrogen bonds between water molecules in the liquid are about 90% electrostatic and 10% covalent. The hydrogen bonds in liquid water stabilize it by an energy of cohesion that is responsible for the unusually high heat of vaporization (9.7 kcal/mol, or 538 kcal/liter).
Another unique property of water is the existence of ionized species. In pure, neutral water, the concentration of the hydrated proton H3O+(H2O)n and hydrated hydroxide ion HO−(H2O)n are each ca. 10−7 mol/liter. These species are essentially hydrogen-bonded clusters.
The polarity of water makes it a good solvent for polar ionic molecules because water can form electrostatic or hydrogen bonds to the dissolved species. For example:
Equation 3. Formation of hydrated ions in water.
One of the simplest indications that solutions of NaCl or HCl in water contain ions is their high electrical conductivity. Pure water is only a weak conductor of electricity because the concentrations of the ions H3O+ and HO− are only 10−7 mol/liter. Seawater is a much better conductor because it contains 0.1 mol/liter of Na+ and Cl− ions.
Water is unique as a solvent.
Despite the fact that solid NaCl is bound in the crystal lattice with an energy of 187 kcal/mol, it dissolves in water to give solutions of hydrated Na+ and Cl−. The reason for this is that the energy of solvation of these ions in water (191 kcal/mol) overcomes the high lattice energy by 4 kcal/mol. The sum of energies of solvation by water of a proton (269 kcal/mol) and of a chloride ion (89 kcal/mol) are greater than the dissociation energy of gaseous hydrogen chloride (358 kcal/mol vs 103 kcal/mol), and so it is clear why HCl is both soluble in water and fully ionized. The take-home lesson is that solvation by water strongly stabilizes both positive and negative ions.
The high solvation energy of sodium chloride in water is largely electrostatic. Solvated Na+ is surrounded by a cluster of at least six H2O molecules with oxygen in proximity to Na+. Solvation of Cl− similarly involves a cluster of H2O molecules with hydrogen in proximity to Cl−.
There are also attractive forces that operate between nonpolar molecules such as straight-chain hydrocarbons. These intermolecular attractions, sometimes called van der Waals forces, are very much weaker than covalent or ionic bonds, or even hydrogen bonds. An instructive example is the attraction between two atoms of the inert gas argon (Ar) which has been measured as ca. 0.28 kcal/mol. This attraction arises not from covalent bonding (because the atomic orbitals of Ar are filled), but from fluctuations in electron density around Ar that create transient imbalance, with one side of the atom being more negative than the other. This polarity induces an opposite distribution of charge on a nearby Ar atom, and the result is a transient attraction ().
Electron density fluctuation in Ar(1) induces opposite electron density in Ar(2), leading to net attraction between them.
The attraction between two adjacent non-polar molecules increases in proportion to the area of contact and is usually on the order of ca. 1 kcal per square Å of close contact. One manifestation of van der Waals attraction is the steady increase in boiling point temperatures (in °C) with increasing molecular size for the series of straight-chain hydrocarbons ().
Boiling points (red) of straight-chain hydrocarbons (blue).
Although van der Waals attractions between molecules are weaker than hydrogen bonding or electrostatic interactions, they can become significant when two nonpolar molecular surfaces are complimentary in shape of sizeable area. These forces play a major role in determining three-dimensional protein geometry and specificity of drug action. In addition, this type of interaction is what makes it possible for gecko lizards to walk across smooth ceilings or vertical walls.
Compounds containing a carbon-carbon double bond within the structure show characteristic chemical behavior. For instance, ethylene, 1-pentene and cyclohexene all react with H2 in the presence of finely divided nickel (Ni) as a catalyst to form products in which hydrogen has been added to the sp2 carbons (). Such addition reactions to C=C are so common that these compounds are called unsaturated.
Reaction of compounds containing a C=C double bond with hydrogen gas (H2) in the presence of Ni catalyst.
There are many other characteristic reactions of the C=C subunit, often called an olefinic functional group (or alkene). In addition, there are many other types of functional groups in organic compounds. A tabulation of some of the most common ones are shown in and
A small sample of functional groups in organic compounds.
A small sample of functional groups in organic compounds.
Many of the compounds shown in and are familiar to most people, especially ethanol (alcohol), butanethiol (butylmercaptan, the essence of the odor of skunks), and acetic acid (vinegar). They are also widely useful articles of commerce. For instance ethylene is the building block from which giant molecules of the plastic polyethylene are made.
Often, combinations of directly connected functional groups occur in molecules. Some examples of such combinations are shown in .
Compounds containing a combination of functional groups shown in and .
Functional groups containing sulfur and phosphorus can also exist in states of higher oxygenation since these elements have five 3d-orbitals available for bond formation in addition to the 3s and three 3p orbitals. Some examples of these more highly oxidized functional groups are shown in .
Compounds containing highly oxidized functional groups derived from sulfur and phosphorus.
The classification of structural subunits as functional groups is a very useful technique for organizing the enormous amount of information that is encompassed by organic chemistry, the chemistry of the vast family of carbon compounds (carbogens). In this section the carboxylate functional group is examined in some detail using a few of the smaller molecules of the class (carboxylic acids) as representative.
Formic acid, the simplest carboxylic acid, is a proton donor. In aqueous solution it ionizes partially to hydrated negative formate ion and hydrated H3O+, as indicates.
Dissociation of formic acid in water.
The reaction is rapid (submillisecond time scale) and reversible, i.e., the components are in very fast equilibrium with one another. One reason for the stability of formate ion is that the negative charge is spread equally between the two oxygens. The carbon is sp2 hybridized and there are two delocalized π-bonding MOs as expressed in formulas A, B and C in .
Three different representations of the delocalization of electrons in formate ion. In formula A, atomic orbitals combine to form two 3-center bonding MOs. In formula B, the delocalization over O-C-O is illustrated with dotted lines and the oxygen atoms share the negative charge equally. Formula C is the computer-generated electron density map of formate ion in which red indicates high electron density.
Delocalization of electrons or charge in organic molecules is, in general, strongly stabilizing because it leads to a lower energy structure than the hypothetical electron-localized version(s).
Formic acid is the acidic ingredient that causes the immediate sting in the bite of a bee or hornet. It can be neutralized by a base such as ammonia (NH3) or sodium hydroxide (NaOH) to form formate salts. (See .)
Two simple salts of formic acid, ammonium and sodium formate.
The conversion of a carboxylic acid to a carboxylate salt by treatment with a base is a general property of this functional group class. As might be expected, the salts are generally much more soluble in water than the corresponding carboxylic acids. Solutions of carboxylic acids in water are acidic, just as for formic acid, although the degree of acidity varies from one compound to another. Acetic acid, CH3COOH, is less acidic than formic acid. Trifluoroacetic acid, CF3COOH, is a much stronger acid than either formic or acetic acid. The reason for this is that fluorine is powerfully electron attracting, and considerable negative charge is delocalized to the three fluorines in CF3COO−, conferring extra stabilization. shows the electron density maps for acetate, formate and trifluoroacetate.
Electron density maps for acetate, formate and trifluoroacetate anions, displayed in order of increasing stability.
The subfragment RCO of carboxylic acids is called an acyl group (the R of RCO can be any carbon group, ).
Formation of carboxylic acid derivatives. The acyl group is highlighted with the yellow box.
There are many reactions of carboxylic acids that form compounds of structure RCOX which are called carboxylic acid derivatives. Some examples are shown in .
Examples of simple carboxylic acid derivatives.
These compounds can all be made by standard reactions from carboxylic acids (RCOOH).
The essentially infinite diversity of organic compounds is made possible in part because the various functional groups can occur together in all possible numbers and combinations. Several important specific examples are shown in and .
Polyfunctional carboxylic acids.
Polyfunctional carboxylic acids.
Sulfuric acid is a strong acid that dissolves in water to form the hydrated negative ions (anions) bisulfate and sulfate.
Dissociation of sulfuric acid in water.
Similarly, sulfonic acids, such as methane-sulfonic acid, ionize completely in water and form salts with bases.
Equation 6. Neutralization of methanesulfonic acid with aqueous sodium hydroxide to form sodium methanesulfonate.
Long chain sulfonate salts such as C18H37SO3−Na+ (sodium octadecane sulfonate, see ) are important detergents (e.g., Tide) that interact by van der Waals forces to attract greasy deposits and carry them into a water wash because of the strong aqueous solvation of the sulfonate ion.
Line drawing and space-filling representation of sodium octadecane sulfonate (C18H37SO3Na).
The negative charge in bisulfate, sulfate and sulfonate ions is spread out over oxygens in delocalized and highly stabilized molecular orbitals.
Figure 39. Computer-generated electron density maps of bisulfate, sulfate and methanesulfonate ions.
The behavior of phosphoric acid is similar. Phosphoric acid reacts with sodium hydroxide in water to form soluble mono-, di- or trisodium salts.
Equation 7. Neutralization of phosphoric acid with sodium hydroxide to form phosphate salts.
Phosphate esters can be formed from the attachment of phosphate to a hydroxyl group in an organic molecule (e.g., alcohols), as shown in the examples that follow.
Figure 40. Simple esters of phosphoric acid.
Phosphorylation of hydroxyl groups in proteins is important in living organisms as a key reaction in chemical signaling.
Benzene is a remarkably stable hydrocarbon of formula C6H6. The six carbons are held together in a planar hexagonal arrangement as shown in . That geometry corresponds to sp2 hybridization of the carbons in the ring.
Structure of benzene.
Although analogy with ethylene (see page 8) might suggest structure A for benzene, it has been shown experimentally that the six C-C bonds in benzene are equal in length and that it is a delocalized π-electron structure. Quantum chemical calculations show that the six electrons not used for C-H and C-C σ-bonding lie in three bonding orbitals formed by the overlap of the p AOs on each carbon. A MO energy diagram is shown in .
I: Energy levels of the six π-MOs of benzene that result from the combination of the six 2pz atomic orbitals of the six carbons in the ring. II: An MO picture of the π-electron clouds above and below the benzene ring. III: A picture of the π-electron clouds above and below the benzene ring.
Each carbon contributes a p-AO to form six π-MOs. Three of these will be of lower energy than the original pz atomic orbitals and three will be higher (see diagram I. in ). There are six electrons available for π-bonding. Since these will be accommodated in the lower energy orbitals MO (1) and MO (2) and MO (2’), the resulting π-delocalized structure will be stabilized. In fact, simple calculations show that the delocalized structure B is more stable than the localized three C=C structure A (see ) by about 35 kcal/mol.
Despite the fact that structure A is less realistic than structure B in , it is commonplace to draw the structure of benzene and related compounds with alternating double bonds. The structures of benzenoid compounds in the later sections of this book are drawn with the traditional double/single bond notation.
The delocalization of the six π-electrons of benzene over the whole ring in low-energy orbitals bestows special properties and huge importance to this structural unit. Typically the C=C unit tends to undergo reactions in which groups add to the two carbons of the double bond. However, benzene generally reacts with these same reagents differently. For instance, 1-pentene reacts with chlorine gas (Cl2) to form 1,2-dichloropentane, but benzene reacts to form chlorobenzene and HCl as shown in .
Reaction of 1-pentene and benzene with chlorine gas (Cl2). In the case of benzene one of the hydrogen atoms (H) is replaced with a chlorine (Cl) atom — a process which is known as substitution.
Most reactions of benzene are of the substitution type (replacement of atoms) with retention of the very stable π-system, whereas a double bond in a typical non-benzenoid hydrocarbon undergoes addition reactions.
There are innumerable chemical reactions that replace the hydrogens of benzene by some other groups, and countless thousands of substituted benzene compounds can be made (synthesized). The six carbons of the benzene ring allow different positioning of groups. For instance, there are three distinctly different dichlorobenzenes, as indicated in .
Three distinctly different position isomers of dichlorobenzene.
These compounds, called position isomers, have the same molecular formula (C6H4Cl2) but different physical and chemical properties.
A closed circuit of six π-electrons confers stabilization of the ring systems of many compounds other than benzene. The rings can be a different size or contain non-carbon atoms, especially O, N and S. One or more such rings can be connected to one another. Some examples of important ring systems are shown in and .