Summary

Understanding organic chemistry means knowing not just what happens but also why and how it happens at the molecular level. In this chapter, we’ve reviewed some of the ways that chemists describe and account for chemical reactivity, thereby providing a foundation for understanding the specific reactions that will be discussed in subsequent chapters.

Organic molecules often have polar covalent bonds as a result of unsymmetrical electron sharing caused by differences in the electronegativity of atoms. A carbon–oxygen bond is polar, for example, because oxygen attracts the shared electrons more strongly than carbon does. Carbon–hydrogen bonds are relatively nonpolar. Many molecules as a whole are also polar, owing to the presence of individual polar bonds and electron lone pairs. The polarity of a molecule is measured by its dipole moment, μ.

Plus (+) and minus (–) signs are often used to indicate the presence of formal charges on atoms in molecules. Assigning formal charges to specific atoms is a bookkeeping technique that makes it possible to keep track of the valence electrons around an atom and offers some clues about chemical reactivity.

Some substances, such as acetate ion and benzene, can’t be represented by a single line-bond structure and must be considered as a resonance hybrid of two or more structures, none of which would be correct by themselves. The only difference between two resonance forms is in the location of their π and nonbonding electrons. The nuclei remain in the same places in both structures, and the hybridization of the atoms remains the same.

Acidity and basicity are closely related to the ideas of polarity and electronegativity. A Brønsted–Lowry acid is a compound that can donate a proton (hydrogen ion, H⁺), and a Brønsted–Lowry base is a compound that can accept a proton. The strength of a Brønsted–Lowry acid or base is expressed by its acidity constant, K_a, or by the negative logarithm of the acidity constant, pK_a. The larger the pK_a, the weaker the acid. More useful is the Lewis definition of acids and bases. A Lewis acid is a compound that has a low-energy empty orbital that can accept an electron pair; Mg²⁺, BF₃, AlCl₃, and H⁺ are examples. A Lewis base is a compound that can donate an unshared electron pair; NH₃ and H₂O are examples. Most organic molecules that contain oxygen and nitrogen can act as Lewis bases toward sufficiently strong acids.

A variety of noncovalent interactions have a significant effect on the properties of large biomolecules. Hydrogen bonding—the attractive interaction between a positively polarized hydrogen atom bonded to an oxygen or nitrogen atom with an unshared electron pair on another O or N atom, is particularly important in giving proteins and nucleic acids their shapes.