Structure and Bonding in Organic Compounds

Organic chemistry began to emerge as a science about 200 years ago. By the late eighteenth century, substances had been divided into inorganic and organic compounds. In those days, early in the history of organic chemistry, inorganic compounds were isolated from mineral sources, and organic compounds were obtained only from plants or animals. Organic compounds were more difficult to study in the laboratory and decomposed more easily than inorganic compounds. The differences between inorganic and organic compounds were attributed to a “vital force” that was required for the synthesis of organic compounds. It was believed that organic compounds could not be synthesized in the laboratory without the vital force. However, by the middle of the nineteenth century, chemists had learned how to work with organic compounds in the laboratory and how to synthesize them.

The organic compounds we will discuss throughout this text contain carbon and a few other elements, such as hydrogen, oxygen, and nitrogen. We will also examine compounds containing sulfur, phosphorus, and halogens. Many, more exotic, organic compounds are also known, and organic compounds have been made that contain virtually every element in the periodic table.

The molecule shown below is terfenedine, an antihistamine whose formula is C32H41NO2. The structure of terfenedine is an example of the amazing variety of structures of organic compounds. They are everywhere in nature, including interstellar space. No known living organism can exist without organic compounds, and synthetic organic compounds are an integral part of the objects we use every day. Their importance cannot easily be exaggerated.

1.1 Brief Review of Atomic Structure

The physical and chemical properties of a molecule depend on the bonds that hold it together. And these bonds depend on the electron configurations of its atoms. Therefore, we will review some of the electronic features of atoms and the periodic properties of the elements before describing bonding and its relation to structure in organic compounds.

Atomic Structure

Each atom has a central, small, dense nucleus that contains protons and neutrons, which are embedded in a sea of electrons. The atomic number, which equals the number of protons in the nucleus, determines the identity of an atom. Since atoms have an equal number of protons and electrons and are electrically neutral, the atomic number also equals the number of electrons in an atom.

The elements in the periodic table are arranged by atomic number. The elements are arrayed in horizontal rows called periods and vertical columns called groups. In this text, we will emphasize hydrogen in the first period and the elements carbon, nitrogen, and oxygen in the second period. The electronic structure of an atom determines its chemical reactivity.

Atomic Orbitals

The electrons in an atom occupy atomic orbitals, which are designated by the letters s, p, d, and f. Each orbital can contain a maximum of two electrons. An atomic orbital is a mathematical equation that describes the energy of an electron. The square of the equation for the atomic orbital defines the probability of finding an electron within a given region of space.

Orbitals are grouped in shells of increasing energy, designated by the integers 1, 2, 3, 4, …, n. These integers are called principal quantum numbers. Each shell contains a unique number and type of orbitals. The first shell contains a single 1 s orbital. The second shell contains one 2 s orbital and three 2 p orbitals. Each orbital can contain no more than two electrons, and two electrons in any orbital must have opposite spin. We need to consider only the orbitals of the first three shells for the elements commonly found in organic compounds.

All s orbitals are spherically symmetrical (Figure 1.1a). The 2 s orbital is larger than the 1 s orbital. A 2 s orbital is farther from the nucleus, and it has a higher energy than a 1 s orbital. The three p orbitals in a shell are not spherically symmetrical. Electron density in each p orbital is concentrated in two regions or lobes—one on each side of the nucleus. The two lobes together are the orbital. The shapes of the p orbitals are shown in Figure 1.1b. The p orbitals are often designated as px, py, and pz. They are mutually perpendicular to one another, and they are aligned along the x, y, and z axes. Although the orientations of the px, py, and pz orbitals differ, the electrons in each p orbital have equal energies.

Orbitals of the same type within a shell constitute a group called a subshell. For example, an s subshell has one orbital and can contain only two electrons. In contrast, a p subshell, which begins in period two, contains three p orbitals and can contain a total of six electrons.

Electrons are distributed in subshells to give an electron configuration that has the lowest energy. The order of increasing energy of subshells is 1 s < 2 s < 2p < 3 s < 3p for elements of atomic number less than 18. For any subshell, the lowest energy state is the arrangement that maximizes the number of electrons having the same spin. This generalization is Hund’s Rule. This means that electrons first occupy orbitals one at a time within subshells before pairing in a common orbital. Table 1.1 shows the atomic numbers and electron configurations for the first two periods in the periodic table.

Valence Shell Electrons

Electrons in filled, lower energy shells of atoms have no role in determining the structure of molecules, and they do not participate in chemical reactions because they are held too tightly to the nucleus. Only the higher energy electrons, which are located in the outermost shell, called the valence shell, participate in chemical bonding. These are the valence electrons. For example, the single electron of the hydrogen atom is a valence electron. The number of valence electrons for the common atoms contained in organic molecules is given by their group number in the periodic table. Thus, carbon, nitrogen, and oxygen atoms have four, five, and six valence electrons, respectively. With this information we can understand how these elements combine to form organic compounds.

1.2 Atomic Properties

The elements in the periodic table are arranged by atomic number. The elements are arranged in horizontal rows called periods and vertical columns called groups. The physical and chemical properties of an element can be estimated from its position in the periodic table. Two properties that help us explain the properties of organic compounds are the atomic radius and electronegativity.

Atomic Radius

The overall shape of an atom is spherical, and its volume depends both on the number of electrons and on the energies of the orbitals the electrons occupy. The sizes of some atoms, expressed as the atomic radius, in picometers (pm,10− 12 m), are given in Figure 1.2 in a greatly abbreviated periodic table that shows the atoms we will most commonly encounter in our...