CHAPTER OUTLINE
1.2 The Nature of Covalent and Ionic Bonds
Box 1.1—Drawing organic molecules
1.3 Polarization of Covalent Bonds
1.4 Atomic Orbitals and Valence Bond Theory
1.5 Hybridization of Orbitals and Tetrahedral Carbon
1.6 Hybrid Orbitals of Oxygen and Nitrogen and Common Functional Groups
Box 1.2—Functional groups containing Phosphorus or Sulfur
1.8 Heteroaromatic Ring Systems in Drug Structures
1.1 Introduction
In this chapter, we will review fundamental concepts of chemical structure and bonding in the organic molecules that make up drugs and their biological targets. By “organic,” we mean molecules that are constructed primarily from the element carbon (C). Carbon exhibits striking versatility in its ability to form various different bonding arrangements with other carbon atoms as well as with other biologically relevant elements such as nitrogen (N), oxygen (O), sulfur (S), and phosphorus (P). It is this versatility that allowed carbon-based life to emerge on our planet. Thus, to understand the molecules of life—proteins, lipids, nucleic acids, hormones, etc.—and the drugs that interact with them, we must start with a solid understanding of structure and bonding in organic molecules. In this chapter, we will begin by contrasting the nature of ionic and covalent bonding and will describe the polarization of covalent bonds. We will then dive deeper into the nature of the covalent bond, discussing atomic and molecular orbitals, the “hybridization” of orbitals, and aromaticity. Finally, we will review some important functional groups and organic ring systems that figure prominently in the structures of biological molecules and drugs.
In the chapters that follow we will learn more about the intermolecular interactions, mostly non-covalent, that govern the binding of a drug molecule to its intended (and sometimes unintended) biological targets. For now, it is important to recognize that a drug molecule’s particular structure—its shape and the nature and connectivity of its atoms—determines what biological activities it will have. If a molecule’s structure leads to interactions in the body that correct an abnormality, restore normal function of a cell, or kill a pathogenic or cancerous cell, a new medicine is born. The seemingly endless ways in which organic molecules can be assembled has allowed scientists to create our current pharmacopeia and affords confidence that still more new medicines will be developed to address currently unmet medical needs.
1.2 The Nature of Covalent and Ionic Bonds
Atoms are comprised of a nucleus containing positively charged protons and uncharged neutrons surrounded by negatively charged electrons. On account of their very low mass, electrons behave as both particles and waves. The peculiar wave-like nature of the electron is what prevents this negatively charged particle from simply “falling” into the positively charged nucleus, to which it is clearly attracted. Wave-like electrons are spatially confined to specific atomic “orbitals” surrounding the nucleus. While atomic and molecular orbitals (Sections 1.4 and 1.5) underlie our current understanding of chemical bonding, their existence was hinted at much earlier by a certain periodicity in the chemical reactivity of the elements. It was this observation that allowed Mendeleev to construct his periodic table of the elements. A partial periodic table including just the first three “periods” (rows) of elements most relevant to organic chemistry is provided here (Figure 1.1).
Figure 1.1 Periodic table of the first 18 elements (atomic number Z = 1 through 18). Groups (columns) 1–8 represent the “main group” elements and are the elements most relevant to organic chemistry and drug structures. Electronic configurations are provided in condensed format, with configuration of valence electrons shown explicitly and inner sphere electrons indicated by the corresponding noble gas configuration, either [He] or [Ne].
The periodic table arranges the elements in order of increasing number of protons (atomic number, Z) and by “groups” (columns) of elements with similar chemical reactivity. This periodicity led to an understanding of chemical reactivity and bonding as being related to the filling of electron “shells” surrounding the nucleus. To understand why chemical bonds form at all, it is useful to consider those few elements that generally do not form bonds—the noble gases. Found at the far right-hand side of the periodic table, noble gases such as helium (He), neon (Ne), and argon (Ar) are “nobly unreactive” because their outermost electron shell is perfectly filled. If helium requires only two electrons to complete its outermost shell, then neon and argon require an additional 8 and 16 electrons, respectively, to do so. The driving force for chemical bonding can thus be understood as a desire of atoms to achieve perfectly filled electron shells (a noble gas “configuration”) by forming bonds to other atoms. This can be achieved in one of two ways—by the exchange of electrons in an ionic bond or by the sharing of electrons in a covalent bond.
The chemistry of carbon involves covalent bonding and so we will discuss ionic bonding only briefly here. Common table salt (sodium chloride, Na+Cl−) provides the most familiar example of an ionic bond between two atoms. Looking at the periodic table we see that both sodium and chlorine are just one column away (and thus one electron away) from a noble gas configuration. Transfer of an electron from sodium to chlorine produces a sodium cation (Na+) and chloride anion (Cl−), each with the electronic configuration of neon (i.e., a filled outer electron shell). The “bond” in Na+Cl− can be thought of as the electrostatic attraction between the sodium and chloride ions. The benign, unreactive nature of Na+Cl− can be contrasted with elemental sodium metal (Na), which reacts violently with water, and elemental chlorine gas (diatomic Cl2), which was used as a warfare agent in World War I.
Carbon does not form ionic bonds because achieving a noble gas configuration would require that it acquire and stabilize four additional electrons, resulting in a tetra-anion with an overall charge of −4. Small atoms such as C, N, and O are not capable of existing in such highly charged states. Instead, carbon achieves a noble gas configuration by forming four covalent bonds. Each bond comprises two electrons, one provided by the carbon atom and one provided by its bonding partner. With four bonds of two electrons each, a carbon atom has obtained the eight electrons (an octet) required to exactly fill its outermost electron shell. While we commonly shown bonds as simple lines, the chemist Gilbert N. Lewis developed a notation in which a bond is shown as a pair of dots, meant to represent the pair of shared electrons that make up the bond. Lewis structures can be used to show not only single bonds but also double and triple bonds, as illustrated (Figure 1.2). While this notation has clear limitations for drawing larger molecules, we still use Lewis notation to show and keep track of nonbonded lone pair electrons.
Figure 1.2 Ethane, ethylene, and acetylene shown as Lewis drawings and as line drawings.
Since carbon must form four bonds to achieve a noble gas configuration, we say that carbon has a valence of four. By inspecting the periodic table (Figure 1.1), we can furthermore predict that nitrogen should have a valence of three and oxygen a valence of two, since nitrogen and oxygen will require three or two additional shared electrons, respectively, to achieve a noble gas configuration. Hydrogen is only one column removed from helium in the first row of the periodic table and so it has a valence of one. Similarly, the halogens (Cl, Br, I) have a valence of one since this group (column) is immediately adjacent to the noble gasses and thus is just one shared electron away from a filled shell.
Even with this rather crude notion of filling electron “shells,” we can already make sense of a great variety of organic compounds formed from combinations of C, N, O, and H. Some biologically relevant molecules are shown (Figure 1.3) using Lewis structures to illustrate bonding and the filling of electron shells for H (two electrons required) and C, N, and O atoms (eight electrons required). Note that all the bonding and nonbonding electrons associated with a given atom count toward the total shared electron count. Thus the triple bond in hydrogen cyanide (HCN) contributes six shared electrons to both the C and N atoms. These six electrons, when combined with a pair of electrons in the H–C bond and the nonbonded electron pair on the nitrogen atom, produce a total electron count of eight for both C and N (Figure 1.3).
Figure 1.3 Structures of simple organic molecules shown as line drawings and complete Lewis structures.
It’s a good idea to become proficient in drawing Lewis structures as this approach helps us understand the locations of bonded and nonbonded electrons and reinforces the idea that bonds are comprised of pairs of shared electrons. Of course, using Lewis structures for drug-sized molecules is not practical and so chemists have developed short-hand notations for drawing chemical structures. These are reviewed in Box 1.1 and this standard notation will be used throughout most of this text.
Box 1.1 Drawing organic molecules.
Chemists have adopted a drawing convention that avoids the need to explicitly show hydrogen atoms or even write a “C” for each carbon atom. A carbon atom is implicit at each “joint” in a structure, or at an unlabeled terminus. Hydrogen atoms are similarly implicit—each carbon is assumed to contain as many bound hydrogens as necessary to achieve tetravalency. Atoms other than carbon and hydrogen are shown explicitly, as are hydrogen atoms on non-carbon atoms (e.g., the hydroxyl group –OH in 1-butanol). It is also helpful to show hydrogen atoms explicitly on certain functional groups such as aldehydes. Common aromatic rings like phenyl and pyridine are best depicted with alternating double and single bonds.
1.3 Polarization of Covalent Bonds
In our discussion of covalent bonding in the previous section, we described the electrons involved in a covalent bond as being shared between the two atoms involved in the bond. If the bonded atoms are identical then the electrons in that bond will indeed be shared equally. However, when two different atoms form a covalent bond, the electrons in the bond will usually not be shared equally between the bonded atoms and the bond is said to be polarized. Polarization of covalent bonds occurs because certain atoms have more power to pull electrons toward their nucleus than others. Generally, atoms located further to the right in a period (row) of the periodic table exert a stronger pull on electrons and are said to be more electronegative. Fluorine for example is more electronegative than carbon, and oxygen is more electronegative than nitrogen. We can illustrate the polarization of a C–F bond in one of two ways, as shown below. The δ+ nomenclature indicates a partial positive charge and the δ− a region of partial negative charge. This polarization of the C–F bond (with greater electron density on fluorine) can also be illustrated using the special arrow shown below at right. Both of these notations will be used in subsequent sections and chapters of this text.
To a first approximation, we can estimate electronegativity using the concept of effective charge, which is equal to the total positive charge of the nucleus minus the negative charge of the non-valence (“inner shell”) electrons. For example, lithium (Li) has an atomic number of three (Z = 3), and thus three protons in the nucleus and a nuclear charge of +3. Lithium has a single valence electron and two inner shell electrons so the effective charge of lithium is +1 (3 − 2 = 1). Being in the same row of the periodic table as lithium, fluorine also has only two inner shell electrons. With an atomic number 9 however, fluorine has an effective charge of +7 (9 − 2 = 7). Thus, if one negatively charged electron of Li is experiencing the pull of a single positive charge from the nucleus, an electron from F is experiencing a pull that is seven times greater.
Effective charge is useful for estimating relative electronegativity for elements in the same period (row) of the periodic table, but it is less predictive when comparing atoms from different periods and different groups, like sulfur and nitrogen. In these cases, the Pauling electronegativity scale becomes indispensable. Devised by Linus Pauling, the table assigns each atom an electronegativity coefficient, and the covalent bond is always polarized in the direction of an atom with a larger coefficient (Table 1.1). From the Pauling electronegativity scale we see that nitrogen (Pauling coefficient of 3.0) is more electronegative than sulfur (2.5). We will frequently refer to the electronegativity scale in subsequent chapters as this concept is very powerful in helping to understand chemical reactivity and intermolecular interactions of functional groups.
Table 1.1 Pauling Electronegativity Scale for Selected Elements Most Relevant to Organic Chemistry and Drug Action.
1.4 Atomic Orbitals and Valence Bond Theory
The concept of valence and the Lewis view of covalent bonding is useful to help us understand why elements like H, C, N, and O combine in various ways in organic molecules. Unfortunately, this view fails to explain many other important features of organic molecules, such as the three-dimensional arrangement of bonds and the fact that rotation about C–C single bonds is generally facile while rotation about C–C double or triple bonds is not. In this section we will introduce the concept of the atomic orbital as well as valence bond theory, in which covalent bonds are understood as arising from the “overlap” of atomic orbitals to form molecular orbitals. At least notionally, the overlap of atomic orbitals to form bonds can be equated with the sharing of electrons as posited in the Lewis description of the covalent bond.
Quantum mechanics is the field of physics that deals with matter and energy at very small scales, where the dual wave-particle nature of matter becomes important. According to quantum mechanics, electrons do not circle the nucleus in a fixed orbit like a planet around its sun, but rather are “spread out” in three-dimensional space around the nucleus as defined by specific solutions to the Schrödinger equation.
Hψ = Eψ
Each solution to this equation is associated with a particular wave function (ψ), also called an atomic orbital. The easiest way to visualize an atomic orbital is to consider its probability density, the square of the wave function (ψ2), which corresponds to the probability that an electron will be found in a particular region of space surrounding the nucleus. The lowest energy atomic orbital for the hydrogen atom (and by extension all other atoms) is the 1s orbital, which has a spherical probability density (Figure 1.4) and can accommodate at most two electrons, provided they have opposite “spin” as dictated by the Pauli exclusion principle. The filling of a 1s orbital with two electrons is the more accurate picture of what is going on with He and its filled electron “shell.” Next highest in energy is the 2s orbital, which is also spherical but with its electrons, on average, spending more time further from the nucleus. Next higher in energy are three energetically equivalent 2p orbitals, often denoted 2px, 2py, and 2pz. The p orbital has a bilobed or “dumbbell” shaped probability density, with a node of zero probability at the nucleus, where the wave function changes sign. The three p orbitals are oriented along different axes when shown on a typical coordinate system (Figure 1.5). Each 2p orbital can accommodate up to two spin-paired electrons, for a total of six 2p electrons.
Figure 1.4 Boundary surface of 1s and 2s atomic orbitals. The spherical surfaces shown represent the boundary within which the probability of finding an electron is high (>90%). (Reproduced, with permission, from Carey FA, Giuliano RM. Organic Chemistry. 9th ed. New York: McGraw-Hill Education; 2014.)
Figure 1.5 Boundary surface of 2p atomic orbitals. The 2p orbital has a node at the nucleus. Three 2p orbitals are oriented along the three axes of a typical three coordinate system. (Reproduced, with permission, from Carey FA, Giuliano RM. Organic Chemistry. 9th ed. New York: McGraw-Hill Education; 2014.)
At this point, the power of quantum mechanics to describe the physical world of atoms should be apparent. Specific solutions to the Schrödinger equation provide discrete energy states (e.g., 1s, 2s, 2p orbitals) that are consistent with and help explain the particular arrangement of elements in the empirically derived periodic table. As atomic number increases, electrons are added to atomic orbitals in the order 1s, 2s, 2p, 3s, 3p, and so on according to the relative energies of the atomic orbitals, as determined by solutions to the Schrödinger equation. Let us then revisit the electronic configuration of the noble gas neon (Ne, Z = 10), using an atomic orbital diagram, with relative energies of the atomic orbitals displayed on either a vertical or horizontal axis (Figure 1.6). To complete the electronic configuration of Ne, we add 10 electrons sequentially to the 1s, then 2s, and finally the three 2p orbitals, using an up or down arrow to indicate electron spin (and being sure to show paired electrons with opposite spin). As expected for a noble gas, each orbital is filled with exactly two electrons, producing perfectly filled 1s, 2s, and 2p orbitals. We say that the electronic configuration of Ne is 1s2 2s2 2p6. The electronic configuration for all elements in the first three periods of the periodic table is shown in Figure 1.1 and in tabular format for the first 12 elements in Table 1.2.
Figure 1.6 Electronic configuration of the noble gas neon (Ne) shown in two versions of an energy diagram. While the horizontal orientation at right is useful in the way it recalls the periodic table, it incorrectly suggests that the three 2p orbitals are of different energies. In fact the energies of the three 2p orbitals are equal, as is accurately captured in the vertically displayed diagram at left.
Table 1.2 Electron Configuration of the First 12 Elements.
Now consider the ionic bond in Na+F−. We can examine the electronic configurations of Na (1s2 2s2 2p6 3s1) and F (1s2 2s2 2p5) as illustrated in energy diagrams (Figure 1.7). Sodium has a single unpaired electron in a higher energy 3s orbital, meaning that on average this electron spends more time further away from the nucleus as compared to electrons in a 2p orbital. In transferring this 3s electron to a fluorine atom a cationic sodium ion (Na+) and a fluoride anion (F−) are produced, each with a new electronic configuration of 1s2 2s2 2p6—the same electronic configuration as Ne.
Figure 1.7 Electronic configurations of sodium and fluorine in the ground state. Transfer of an electron from sodium to fluorine produces a pair of ions, Na+ and F−, each with the same electron configuration as the noble gas neon (1s2 2s2 2p6).
The valence bond description of the covalent bond involves the mathematical combination of two wave functions (i.e., the “overlap” of atomic orbitals) to produce two new molecular orbitals (Figure 1.8). This is most simply illustrated for the formation of two new molecular orbitals (MOs) by the combination of two hydrogen 1s atomic orbitals (AOs). One of the new molecular orbitals is lower in energy than the 1s atomic orbitals while the other is higher in energy. Since one electron is contributed by each of the two 1s AOs, we will have two electrons in total to occupy the new MOs of the H–H molecule. These electrons will naturally occupy the MO with lower energy, which is called a bonding orbital since its filling represents the formation of a stable bond (being lower in energy than either of the 1s AO). The higher energy MO is called an antibonding orbital since it is higher in energy than the AOs and filling it would not be expected to result in the formation of a stable bond. Note that electrons fill MOs in pairs with opposite spin, since the exclusion principle applies to MOs as well as to AOs. Just as valence bond theory can explain the formation of the molecule H–H, it can also explain why the molecule He–He is not observed. With two electrons contributed from each 1s orbital of He, a total of four electrons would need to be placed into the MOs of He–He. This would involve filling both the bonding and antibonding orbitals and any energetic benefit accomplished by filling the former would be more than offset by filling the latter.
Figure 1.8 The combination of two partially filled hydrogen 1s orbitals leads to two new molecular orbitals, one a bonding MO and the other an antibonding MO. The two electrons fill the bonding orbital, leading to a stable covalent bond in H–H (H2, molecular hydrogen).
A second important tenant of valence bond theory is that stronger bonds are those in which orbital overlap is maximized. Orbital overlap is best visualized using the probability density or boundary surface representations of the AOs involved. From this perspective the combination of two 1s orbitals is equivalent to bringing two spheres together until their surfaces intersect with a circular cross-section. We might expect orbital overlap to be maximal when this circular cross-section is greatest. For the H–H bond in H2, this occurs when the nuclei of the two hydrogen atoms are separated by a distance of about 74 picometers or 0.74 Ångstroms (Å). The two 1s orbitals have been replaced by a bonding MO that is egg-shaped, with a circular cross-section and the highest probability of finding electron density between the two hydrogen nuclei. This type of MO is known as a sigma (σ) orbital and the resulting bond a σ bond. Note that rotation of a σ bond does not change the extent of orbital overlap (the cross-section remains circular), and thus σ bonds generally undergo free rotation.
The formation of a σ orbital as described above results from the “in-phase” combination of two 1s atomic orbitals. The corresponding higher-energy antibonding orbital is called a σ* (“sigma star”) orbital and arises from “out-of-phase” combination of the 1s atomic orbitals. Note that the antibonding σ* orbital has a node of zero electron density between the two H atoms, whereas electron density is maximal at this same location in a bonding σ orbital (Figure 1.9).
Figure 1.9 Graphical illustration of the formation of a bonding σ orbital and an antibonding σ* orbital by the combination of two 1s atomic orbitals of hydrogen. While the dumbbell shape of the σ* orbital resembles a p orbital, these must not be confused. The σ* orbital is a molecular orbital with a node between two different atoms whereas the p orbital represents electron density surrounding a single atom. (Reproduced, with permission, from Carey FA, Giuliano RM. Organic Chemistry. 9th ed. New York: McGraw-Hill Education; 2014.)
1.5 Hybridization of Orbitals and Tetrahedral Carbon
With an understanding of atomic orbitals and valence bond theory we might hope we could explain bonding in simple organic molecules. However, if we examine the electronic configuration of carbon (Figure 1.10) we quickly discover an apparent problem. Carbon has a total of four valence electrons in its 2s and 2p orbitals, but the only unpaired electrons are the two found in 2p orbitals. We can imagine how each 2p orbital might combine with a 1s orbital of an H atom to form two filled MOs (two C–H σ bonds). However, it’s not obvious how the other two electrons in the 2s orbital could be engaged in new bonds since they are already paired in a relatively low energy AO. Moreover, the geometrical arrangement of p orbitals about the carbon nucleus (Figure 1.5) would predict that two C–H σ bonds should be separated by an angle of 90°. However, we know from experimental data that carbon forms four bonds, not just two, and that typical bond angles in carbon-based molecules are ~180°, ~120°, or ~109.5°.
Figure 1.10 Energy diagram showing the relative energies of atomic orbitals for carbon. Valence electrons available for bonding include two electrons in a 2s orbital and two unpaired electrons in 2p orbitals.
Linus Pauling proposed a solution to this problem by suggesting that the valence 2s and 2p orbitals on carbon might “mix” to form new hybrid orbitals with different energies and geometries depending on how the s and p orbitals are combined. While this proposal was made in an effort to rationalize experimental observations, quantum mechanical calculations do in fact support the notion of hybrid atomic orbitals formed by mixing s and p orbitals. Here, we will use energy diagrams and boundary surface illustrations to describe this “hybridization” of carbon. There are three ways in which carbon can be hybridized—by mixing the single 2s orbital with either one, two, or all three of the 2p orbitals (Figure 1.11). The result of mixing one 2s and one 2p orbital is a pair of sp hybrid orbitals, each with equal s and p “character.” This leaves the remaining two p orbitals unchanged (unhybridized) and so we can say that sp hybridized carbon consists of two sp hybrid orbitals and two 2p orbitals. If instead we mix the single 2s orbital with two 2p orbitals, the result is three sp2 hybrid orbitals and a single 2p orbital. Finally, if we mix the 2s orbital with all three 2p orbitals, we obtain four sp3 hybrid orbitals and no unhybridized p orbitals.
Figure 1.11 Energy diagram illustrating three different ways of hybridizing carbon by mixing the valence 2s orbital with either one, two, or all three of the valence 2p orbitals. The resulting forms of hybridized carbon each have four unpaired electrons and are thus capable of forming four bonds with other atoms.
Several points are worth noting at this point. Most importantly, you may have noted in Figure 1.11 that each of the three hybridization schemes results in four orbitals, each with a single unpaired electron. This nicely fits with the known valence of carbon and makes it quite easy to see how these four orbitals might be combined with other atoms to form molecular orbitals (and four bonds). Another important point is that the number of new hybrid orbitals formed in each case exactly matches the number of s and p orbitals used for hybridization. Thus for sp hybridization we combined one s and one p orbital to produce two sp orbitals. Finally, we should note that hybridization occurs because it ultimately leads to molecular orbitals (and bonds) with favorable energies. In other words, hybridization of atomic orbitals is a phenomena of atoms in molecules, where orbital overlap leads to the formation of bonds.
The most useful aspect of hybridization is that it allows us to rationalize the experimentally observed geometries of tetravalent carbon. Thus, sp3-hybridized carbon as in methane (CH4) comprises four sp3 hybrid orbitals, each pointing toward the corners of a tetrahedron (bond angle ~109.5°). Having 25% s character and 75% p character, the sp3 orbital takes on the dumbbell shape of a p orbital, but with one lobe much larger in size than the other (Figure 1.12). Recalling the tenants of valence bond theory, we would say that methane is formed by combining four sp3 orbitals on carbon with the 1s orbitals of four hydrogen atoms. Each hydrogen 1s orbital overlaps with one of the four sp3 orbitals, forming four C–H σ bonds. Overlap occurs on the larger lobe of the sp3 orbital since this maximizes orbital overlap, resulting in a stronger bond. The C–C bonds in related hydrocarbons such as ethane and propane are simply σ bonds formed by overlap of carbon sp3 hybrid orbitals.
Figure 1.12 Bonding in methane (CH4) involves the orbital overlap of half-filled sp3 hybrid orbitals on carbon with half-filled 1s orbitals on hydrogen. (Reproduced, with permission, from Carey FA, Giuliano RM. Organic Chemistry. 9th ed. New York: McGraw-Hill Education; 2014.)
Next, consider sp2-hybridized carbon and the structure of ethylene (C2H4). Each of the two carbon atoms in ethylene comprises three sp2 orbitals lying in the same plane and pointing toward the vertices of an equilateral triangle, a so-called “trigonal-planar” arrangement with bond angles of ~120°. The lone p orbital on each carbon atom is exactly orthogonal to the plane of sp2 hybrid orbitals. The C–H bonds in ethylene are σ bonds formed by end-on overlap of carbon sp2 hybrid orbitals with hydrogen 1s orbitals (Figure 1.13). The double bond in ethylene has two components. The first is a normal σ bond formed by end-on overlap of sp2 orbitals on the two carbon atoms. The second component involves side-on overlap of the unhybridized p orbitals on the two carbon atoms, resulting in what is called a π bond. The π electrons in the π bond of ethylene reside above and below the plane formed by the σ bonds, as illustrated (Figure 1.13). This same plane is a node of the p orbital, where the probability of finding a π electron is zero. Unlike in a σ bond, rotations about the axis of a π bond would result in reduction and ultimately loss of orbital overlap. Thus, the need to maintain side-on overlap of p orbitals in forming a π bond explains why rotation about double bonds is generally forbidden.
Figure 1.13 Bonding in ethylene illustrating the orbital interactions involved. The double bound comprises a σ bond (MO) and a π bond (MO), with two electrons in each molecular orbital. (Reproduced, with permission, from Carey FA, Giuliano RM. Organic Chemistry. 9th ed. New York: McGraw-Hill Education; 2014.)
Finally, we consider sp hybridized carbon in the molecule acetylene (C2H2). In sp hybridized carbon, two sp orbitals project outward in a linear arrangement, 180° opposed from one another. If an sp hybrid orbital were arbitrarily aligned on the x-axis of a three-coordinate system, then the remaining two unhybridized p orbitals would reside, one each, on the y-axis and z-axis. The C–H bonds of acetylene are σ bonds formed from overlap of the carbon sp orbital and 1s hydrogen orbital. The triple bond of acetylene has three components, one σ bond formed by end-on overlap of sp orbitals, and two orthogonal π bonds formed by side-on overlap of the two sets of p orbitals (Figure 1.14).
Figure 1.14 Bonding in acetylene based on sp hybridized carbon. The triple bond comprises one σ bond and two π bonds. (Reproduced, with permission, from Carey FA, Giuliano RM. Organic Chemistry. 9th ed. New York: McGraw-Hill Education; 2014.)
As a final note, remember that the combination (overlap) of two atomic orbitals, whether they are hybrid orbitals or not, must produce exactly two new MOs—a bonding MO (σ or π orbital) as well as an antibonding MO (σ* or π* orbital). In all of the examples provided above, two half-filled AOs were combined, contributing a total of two electrons to the new bond and thus exactly filling a σ or π orbital and forming a stable covalent bond.
1.6 Hybrid Orbitals of Oxygen and Nitrogen and Common Functional Groups
Hybridization of orbitals is important in other main group elements as well, and we will discuss in this section nitrogen and oxygen, which aside from carbon and hydrogen are the most commonly encountered atoms in organic chemistry and drug structure (common functional groups of sulfur and phosphorus are described in Box 1.2). Nitrogen lies next to carbon in the periodic table, with an atomic number of 7 (Z = 7) and an electronic configuration of 1s2 2s2 2p3. We can mix the valence 2s and 2p orbitals of nitrogen just as we did with carbon, resulting in sp, sp2 and sp3 hybridization (Figure 1.15). Note that in each of these arrangements nitrogen has three unpaired electrons while a fourth orbital harbors a pair of electrons, called a lone pair. Thus we predict that nitrogen should form a total of three bonds comprising σ bonds or some combination of σ and π bonds, depending on the form of hybridization. The presence of lone pair electrons on nitrogen atoms has profound implications for the reactivity, intermolecular interactions, and acid-base chemistry of nitrogen containing molecules, as we will see throughout this text.
Figure 1.15 Energy diagram illustrating the mixing of nitrogen valence orbitals into sp, sp2, and sp3 hybridized forms. With one additional valence electron as compared to carbon, hybridized forms of nitrogen all possess one orbital containing a lone pair of electrons, shown in red. Common functional groups containing sp, sp2, and sp3 hybridized nitrogen, respectively, are shown at the bottom of the figure.
The nitrile functional group (CN) is the major example of sp hybridized nitrogen in organic chemistry. By analogy with the linear triple bond in acetylene, the nitrogen atom in a nitrile forms one σ bond and two π bonds with an sp hybridized carbon atom. The nonbonding sp orbital on nitrogen points opposite the C–N bond, and contains a lone pair of electrons (represented using Lewis nomenclature as a pair of dots, Figure 1.15). Like sp2 carbon in ethylene (Figure 1.13), nitrogen sp2 orbitals are disposed in a trigonal-planar arrangement, with an orthogonally located p orbital containing one unpaired electron. Again, the key difference with nitrogen is that one of the sp2 orbitals bears a lone pair of electrons. This form of nitrogen is found in functional groups containing C=N double bonds, such as in imines, oximes, and hydrazones (discussed in Chapter 7) and in a wide variety of aromatic heterocycles (Section 1.8). Finally, sp3 hybridized nitrogen is found in common “saturated” amines, which are commonly encountered in drug structures because they can contribute both to target binding, and in their protonated (charged) form can improve aqueous solubility. Nitrogen with sp3 hybridization has tetrahedral geometry, with the lone pair occupying one of the four corners of a tetrahedron.
