Bond Angles - Localized Chemical Bonding - Introduction - March's Advanced Organic Chemistry: Reactions, Mechanisms, and Structure, 7th Edition (2013)

March's Advanced Organic Chemistry: Reactions, Mechanisms, and Structure, 7th Edition (2013)

Part I. Introduction

Chapter 1. Localized Chemical Bonding

1.K. Bond Angles

The bond angles of sp3 carbon should be the tetrahedral angle 109°28′ when the four atoms or groups are relatively small and identical, as in methane, neopentane, or carbon tetrachloride. As atoms or groups become larger, bond angles are distorted to accommodate the larger size of the attached units. In most cases, the angles deviate a little from the pure tetrahedral value unless two or more units are very large. Molecular models 7–9 illustrate this phenomenon. The H–C–H bond angles in methane (7) are calculated for the model to be 109°47′, whereas the Br–C–H bond angle in 8 is calculated to be 108.08° and the Br–C–Br bond angle in 9 is calculated to be 113.38°. Note that the C–Br bond length is longer than the C–H bond lengths. As the bond angles expand to accommodate the larger atoms, the H–C–H bond angles in 8 and 9 must compress to a smaller angle. In 2-bromopropane, the bromine atom also has a methyl group (compare with bromomethane 8 where Br competes with H) and, the C–C–Br angle in 2-bromopropane is 114.2°.85

Variations are generally found from the ideal values of 120° and 180° for sp2 and sp carbon, respectively. These deviations occur because of slightly different hybridizations; that is, a carbon bonded to four other atoms hybridizes one s and three p orbitals, but the four hybrid orbitals thus formed are generally not exactly equivalent, nor does each contain exactly 25% s and 75% p character. Because the four atoms have (in the most general case) different electronegativities, each makes its own demand for electrons from the carbon atom.86 The carbon atom supplies more p character when it is bonded to more electronegative atoms, so that in chloromethane, for example, the bond to chlorine has somewhat > 75% p character, which of course requires that the other three bonds have somewhat less, since there are only three p orbitals (and one s) to be divided among the four hybrid orbitals.87 Of course, in strained molecules (e.g., 3–6), the bond angles may be greatly distorted from the ideal values (also see Sec. 4.Q).

For molecules that contain oxygen and nitrogen, angles of 90° are predicted from p2 bonding. However, as seen in Section 1.B, the angles of water and ammonia are much larger than this, as are the angles of other organic molecules that contain oxygen and nitrogen (Table 1.6).8892 In fact, they are much closer to the tetrahedral angle of 109°28′ than to 90°. These facts have led to the suggestion that in these compounds oxygen and nitrogen use sp3bonding. Using the hybridization model, these atoms are said to form bonds by the overlap of two (or three) p orbitals with 1s orbitals of the hydrogen atoms, which means that they hybridize their 2s and 2p orbitals to form four sp3orbitals and then use only two (or three) of these for bonding with hydrogen, the others remaining occupied by unshared pairs (also called lone pairs). If this description is valid, and it is generally accepted by most chemists today,93it becomes necessary to explain why the angles of these two compounds are in fact not 109°28′ but a few degrees smaller. One explanation that has been offered is that the unshared electron pair actually has a greater steric requirement (see Sec. 4.Q) than the electrons in a bond, since there is no second nucleus to draw away some of the electron density and the bonds are thus crowded together. However, most evidence is that unshared pairs have smaller steric requirements than bonds94 and the explanation most commonly accepted is that the hybridization is not pure sp3. As seen above, an atom supplies more p character when it is bonded to more electronegative atoms. An unshared pair may be considered to be an “atom” of the lowest possible electronegativity, since there is no attracting power at all. Consequently, the unshared pairs have more s and the bonds more p character than pure sp3 orbitals, making the bonds somewhat more like p2 bonds and reducing the angle. However, these arguments ignore the steric effect of the atoms or groups attached to oxygen or nitrogen. As seen in Table 1.6, oxygen, nitrogen, and sulfur angles generally increase with decreasing electronegativity of the substituents. Note that the explanation given above cannot explain why some of these angles are greater than the tetrahedral angle.

Table 1.6 Oxygen, Sulfur, and Nitrogen Bond Angles in Some Compounds.

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