CHEMISTRY THE CENTRAL SCIENCE

8 BASIC CONCEPTS OF CHEMICAL BONDING

CHEMICAL BONDING AS ART. The Atomium is a 110-m-high steel sculpture commissioned for the 1958 World's Fair in Brussels. The nine spheres represent atoms, and the connecting rods evoke the chemical bonds holding them together. One sphere sits in the center of a cube formed by the other eight, a common arrangement of the atoms in metallic elements, such as iron.

WHAT'S AHEAD

8.1 LEWIS SYMBOLS AND THE OCTET RULE

We begin with descriptions of the three main types of chemical bonds: ionic, covalent, and metallic. In evaluating bonding, Lewis symbols provide a useful shorthand for keeping track of valence electrons.

8.2 IONIC BONDING

We observe that in ionic substances the atoms are held together by the electrostatic attractions between ions of opposite charge. We discuss the energetics of forming ionic substances and describe the lattice energy of these substances.

8.3 COVALENT BONDING

We examine the bonding in molecular substances in which atoms bond by sharing one or more electron pairs. In general, the electrons are shared in such a way that each atom attains an octet of electrons.

8.4 BOND POLARITY AND ELECTRONEGATIVITY

We define electronegativity as the ability of an atom in a compound to attract electrons to itself. In general, electron pairs are shared unequally between atoms with different electronegativities, leading to polar covalent bonds.

8.5 DRAWING LEWIS STRUCTURES

We see that Lewis structures are a simple yet powerful way of predicting covalent bonding patterns in molecules. In addition to the octet rule, we see that the concept of formal charge can be used to identify the dominant Lewis structure.

8.6 RESONANCE STRUCTURES

We observe that in some cases more than one equivalent Lewis structure can be drawn for a molecule or polyatomic ion. The bonding description in such cases is a blend of two or more resonance structures.

8.7 EXCEPTIONS TO THE OCTET RULE

We recognize that the octet rule is more of a guideline than an absolute rule. Exceptions to the rule include molecules with an odd number of electrons, molecules where large differences in electronegativity prevent an atom from completing its octet, and molecules where an element from period 3 or below in the periodic table attains more than an octet of electrons.

8.8 STRENGTHS OF COVALENT BONDS

We observe that bond strengths vary with the number of shared electron pairs as well as other factors. We use average bond enthalpy values to estimate the enthalpies of reactions in cases where thermodynamic data are unavailable.

WHENEVER TWO ATOMS OR IONS are strongly held together, we say there is a chemical bond between them. There are three general types of chemical bonds: ionic, covalent, and metallic (FIGURE 8.1). We can get a glimpse of these three types of bonds by thinking about the simple act of using a stainless-steel spoon to add table salt to a glass of water. Table salt is sodium chloride, NaCl, which consists of sodium ions, Na+, and chloride ions, Cl-. The structure is held together by ionic bonds, which are due to the attractions between oppositely charged ions. The water consists mainly of H2O molecules. The hydrogen and oxygen atoms are bonded to one another through covalent bonds, in which molecules are formed by the sharing of electrons between atoms. The spoon consists mainly of iron metal, in which Fe atoms are connected to one another viametallic bonds, which are formed by electrons that are relatively free to move through the metal. These different substances—NaCl, H2O, and Fe metal—behave as they do because of the ways in which their constituent atoms are connected to one another.

What determines the type of bonding in any substance? How do the characteristics of these bonds give rise to different physical and chemical properties? The keys to answering the first question are found in the electronic structure of the atoms involved, discussed in Chapters 6 and 7. In this chapter and the next, we examine the relationship between the electronic structure of atoms and the ionic and covalent chemical bonds they form. We will discuss metallic bonding in greater detail in Chapter 12.

8.1 LEWIS SYMBOLS AND THE OCTET RULE

The electrons involved in chemical bonding are the valence electrons, which, for most atoms, are those in the outermost occupied shell.  (Section 6.8) The American chemist G. N. Lewis (1875-1946) suggested a simple way of showing the valence electrons in an atom and tracking them during bond formation, using what are now known as either Lewis electron-dot symbols or simply Lewis symbols.

The Lewis symbol for an element consists of the element's chemical symbol plus a dot for each valence electron. Sulfur, for example, has the electron configuration [Ne]3s2 3p4 and therefore six valence electrons. Its Lewis symbol is

The dots are placed on the four sides of the symbol—top, bottom, left, and right—and each side can accommodate up to two electrons. All four sides are equivalent, which means that the choice of on which sides to place two electrons rather than one electron is arbitrary. In general, we spread out the dots as much as possible. In the Lewis symbol for S, for instance, we prefer the dot arrangement shown rather the arrangement having two electrons on three of the sides and none on the fourth.

The electron configurations and Lewis symbols for the main-group elements of periods 2 and 3 are shown in TABLE 8.1. Notice that the number of valence electrons in any representative element is the same as the element's group number. For example, the Lewis symbols for oxygen and sulfur, members of group 6A, both show six dots.

GO FIGURE

Which of these three bond types do you expect to see in CO2(g)?

FIGURE 8.1 Ionic, covalent, and metallic bonds. Different types of interactions between atoms lead to different types of chemical bonds.

GIVE IT SOME THOUGHT

Are all these Lewis symbols for Cl correct?

The Octet Rule

Atoms often gain, lose, or share electrons to achieve the same number of electrons as the noble gas closest to them in the periodic table. The noble gases have very stable electron arrangements, as evidenced by their high ionization energies, low affinity for additional electrons, and general lack of chemical reactivity. (Section 7.8) Because all the noble gases except He have eight valence electrons, many atoms undergoing reactions end up with eight valence electrons. This observation has led to a guideline known as the octet rule: Atoms tend to gain, lose, or share electrons until they are surrounded by eight valence electrons.

An octet of electrons consists of full s and p subshells in an atom. In a Lewis symbol, an octet is shown as four pairs of valence electrons arranged around the element symbol, as in the Lewis symbols for Ne and Ar in Table 8.1. There are exceptions to the octet rule, but it provides a useful framework for introducing many important concepts of bonding.

TABLE 8.1 • Lewis Symbols

GO FIGURE

Do you expect a similar reaction between potassium metal and elemental bromine?

FIGURE 8.2 Reaction of sodium metal with chlorine gas to form the ionic compound sodium chloride.