Summary
The purpose of this chapter has been to get you up to speed—to review some ideas about atoms, bonds, and molecular geometry. As we’ve seen, organic chemistry is the study of carbon compounds. Although a division into organic and inorganic chemistry occurred historically, there is no scientific reason for the division.
An atom consists of a positively charged nucleus surrounded by one or more negatively charged electrons. The electronic structure of an atom can be described by a quantum mechanical wave equation, in which electrons are considered to occupy orbitals around the nucleus. Different orbitals have different energy levels and different shapes. For example, s orbitals are spherical and p orbitals are dumbbell-shaped. The ground-state electron configuration of an atom can be found by assigning electrons to the proper orbitals, beginning with the lowest-energy ones.
A covalent bond is formed when an electron pair is shared between atoms. According to valence bond (VB) theory, electron sharing occurs by the overlap of two atomic orbitals. Bonds that have a circular cross-section and are formed by head-on interaction are called sigma (σ) bonds; bonds formed by sideways interaction of p orbitals are called pi (π) bonds.
In the valence bond description, carbon uses hybrid orbitals to form bonds in organic molecules. When forming only single bonds with tetrahedral geometry, carbon uses four equivalent sp3 hybrid orbitals. When forming a double bond with planar geometry, carbon uses three equivalent sp2 hybrid orbitals and one unhybridized p orbital. When forming a triple bond with linear geometry, carbon uses two equivalent sp hybrid orbitals and two unhybridized p orbitals. Other atoms such as nitrogen, phosphorus, oxygen, and sulfur also use hybrid orbitals to form strong, oriented bonds.
Organic molecules are usually drawn using either condensed structures or skeletal structures. In condensed structures, carbon–carbon and carbon–hydrogen bonds aren’t shown. In skeletal structures, only the bonds and not the atoms are shown. A carbon atom is assumed to be at the ends and at the junctions of lines (bonds), and the correct number of hydrogens is supplied mentally.
Why You Should Work Problems
There’s no surer way to learn organic chemistry than by working problems. Although careful reading and rereading of this text are important, reading alone isn’t enough. You must also be able to use the information you’ve read and be able to apply your knowledge in new situations. Working problems gives you practice at doing this.
Each chapter in this book provides many problems of different sorts. The in-chapter problems are placed for immediate reinforcement of ideas just learned, while end-of- chapter problems provide additional practice and come in several forms. They often begin with a short section called “Visualizing Chemistry,” which helps you see the microscopic world of molecules and provides practice for working in three dimensions. After the visualizations are many further problems, which are organized by topic. Early problems are primarily of the drill type, providing an opportunity for you to practice your command of the fundamentals. Later problems tend to be more thought-provoking, and some are real challenges.
As you study organic chemistry, take the time to work the problems. Do the ones you can, and ask for help on the ones you can’t. If you’re stumped by a particular problem, check the accompanying Study Guide and Student Solutions Manual for an explanation that should help clarify the difficulty. Working problems takes effort, but the payoff in knowledge and understanding is immense.
Figure 1.28 The enzyme HMG–CoA reductase, shown here as a so-called ribbon model, catalyzes a crucial step in the body’s synthesis of cholesterol. Understanding how this enzyme functions has led to the development of drugs credited with saving millions of lives. (credit: image from the RCSB PDB (rcsb.org) of PBD ID 1HW9 (E.S. Istvan, J. Deisenhofer) (2001) Structural mechanism for statin inhibition of HMG-CoA reductase Science 292: 1160–1164/RCSB PDB, CC BY 1.0)
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, Ka, or by the negative logarithm of the acidity constant, pKa. The larger the pKa, 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.
