8.1 ∣ Lewis Symbols and the Octet Rule

• 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.

• 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 sides for placement of two electrons rather than one electron is arbitrary.

The Octet Rule

• Atoms tend to gain, lose, or share electrons until they are surrounded by eight valence electrons.

• In a Lewis symbol, an octet is shown as four pairs of valence electrons arranged around the element symbol.

8.2 ∣ Ionic Bonding

• Ionic substances generally result from the interaction of metals on the left side of the periodic table with nonmetals on the right side.

• Electron transfer to form oppositely charged ions occurs when one atom readily gives up an electron (low ionization energy) and another atom readily gains an electron (high electron affinity).

• Ionic substances possess several characteristic properties.

Energetics of Ionic Bond Formation

• The formation of sodium chloride from sodium and chlorine is very exothermic.

• When a nonmetal gains an electron, the process is generally exothermic, as seen from the negative electron affinities of the elements.

• The principal reason ionic compounds are stable is the attraction between ions of opposite charge.

• A measure of how much stabilization results from arranging oppositely charged ions in an ionic solid is given by the lattice energy, which is the energy required to completely separate one mole of a solid ionic compound into its gaseous ions.

• The energy released by the attraction between ions of unlike charge more than makes up for the endothermic nature of ionization energies, making the formation of ionic compounds an exothermic process.

• For a given arrangement of ions, the lattice energy increases as the charges on the ions increase and as their radii decrease.

Electron Configurations of Ions of the s- and p-Block Elements

• The energetics of ionic bond formation helps explain why many ions tend to have noble-gas electron configurations.

• The second electron removed would have to come from an inner shell of the sodium atom, and removing electrons from an inner shell requires a very large amount of energy.

• Similarly, adding electrons to nonmetals is either exothermic or only slightly endothermic as long as the electrons are added to the valence shell.

Transition Metal Ions

• Because ionization energies increase rapidly for each successive electron removed, the lattice energies of ionic compounds are generally large enough to compensate for the loss of up to only three electrons from atoms.

• Most transition metals, however, have more than three electrons beyond a noble-gas core.

• Thus, in forming ions, transition metals lose the valence-shell s electrons first, then as many d electrons as required to reach the charge of the ion.

8.3 ∣ Covalent Bonding

• A chemical bond formed by sharing a pair of electrons is a covalent bond.

• By using quantum mechanical methods analogous to those used for atoms, we can calculate the distribution of electron density in molecules.

Lewis Structures

• The formation of covalent bonds can be represented with Lewis symbols.

• In forming the covalent bond, each hydrogen atom acquires a second electron, achieving the stable, two-electron, noble-gas electron configuration of helium.

• While these structures show circles to indicate electron sharing, the more common convention is to show each shared electron pair or bonding pair, as a line and any unshared electron pairs (also called lone pairs or nonbonding pairs) as dots.

• For nonmetals, the number of valence electrons in a neutral atom is the same as the group number.

Multiple Bonds

• A shared electron pair constitutes a single covalent bond, generally referred to simply as a single bond.

• When two electron pairs are shared by two atoms, two lines are drawn in the Lewis structure, representing a double bond.

• A triple bond corresponds to the sharing of three pairs of electrons.

• Because each nitrogen atom has five valence electrons, three electron pairs must be shared to achieve the octet configuration.

• Nitrogen is a diatomic gas with exceptionally low reactivity that results from the very stable nitrogen–nitrogen bond.

• As a general rule, the length of the bond between two atoms decreases as the number of shared electron pairs increases.

8.4 ∣ Bond Polarity and Electronegativity

• Bond polarity is a measure of how equally or unequally the electrons in any covalent bond are shared.

• A nonpolar covalent bond is one in which the electrons are shared equally.

• In a polar covalent bond, one of the atoms exerts a greater attraction for the bonding electrons than the other.

Electronegativity

• Electronegativity is defined as the ability of an atom in a molecule to attract electrons to itself.

• We use a quantity called electronegativity to estimate whether a given bond is nonpolar covalent, polar covalent, or ionic.

• The American chemist Linus Pauling (1901–1994) developed the first and most widely used electronegativity scale, which is based on thermochemical data.

Electronegativity and Bond Polarity

• A nonpolar covalent bond results when the electronegativities of the bonded atoms are equal.

• The electrons are shared unequally—the bond is polar.

• In general, a polar covalent bond results when the atoms differ in electronegativity.

Dipole Moments

• A molecule such as HF, in which the centers of positive and negative charge do not coincide, is a polar molecule.

• Polarity affects several macroscopic properties in the lab and in life. Polar molecules attract each other at their negative and positive ends. Ions attract polar compounds. Polar molecules attract positive and negative ions.

• Whenever two electrical charges of equal magnitude but opposite signs are separated by a distance, a dipole is established.

• The quantitative measure of the magnitude of a dipole is called its dipole moment.

Comparing Ionic and Covalent Bonding

• To understand the interactions responsible for chemical bonding, it is advantageous to treat ionic and covalent bonding separately.

• Covalent bonding causes compounds to behave like molecules, with low melting and boiling temperatures and nonelectrolyte behavior in water. Ionic bonding produces brittle, high-melting solids with extended lattice structures that behave strongly as electrolytes in water.

• The simplest approach is to assume that the interaction between a metal and a nonmetal is ionic and that between two nonmetals is covalent.

8.5 ∣ Drawing Lewis Structures

• Lewis structures can help us understand the bonding in many compounds and are frequently used when discussing the properties of molecules.

How to Draw Lewis Structures

1. Sum the valence electrons from all atoms, taking into account overall charge.

2. Write the symbols for the atoms, show which atoms are attached to which, and connect them with a single bond (a line, representing two electrons).

3. Complete the octets around all the atoms bonded to the central atom.

4. Place any remaining electrons on the central atom.

5. If there are not enough electrons to give the central atom an octet, try multiple bonds.

Formal Charge and alternative Lewis Structures

• The formal charge of any atom in a molecule is the charge the atom would have if each bonding electron pair in the molecule were shared equally between its two atoms.

• If we can draw several Lewis structures for a molecule, the concept of formal charge can help us decide which is the most important, which we shall call the dominant Lewis structure.

How to Identify the Dominant Lewis Structure

1. The dominant Lewis structure is generally the one in which the atoms bear formal charges closest to zero.

2. A Lewis structure in which any negative charges reside on the more electronegative atoms is generally more dominant than one that has negative charges on the less electronegative atoms.

• Although the concept of formal charge helps us to arrange alternative Lewis structures in order of importance, it is important to remember that formal charges do not represent real charges on atoms.

8.6 ∣ Resonance Structures

• Because each oxygen atom contributes 6 valence electrons, the ozone molecule has 18 valence electrons. This means the Lewis structure must have one O=O single bond and one O=O double bond to attain an octet about each atom.

• The placement of the atoms in these two alternative but completely equivalent Lewis structures is the same, but the placement of the electrons is different; we call Lewis structures of this sort resonance structures.

• For some molecules or ions, all possible Lewis structures may not be equivalent; in other words, one or more resonance structures are more dominant than others.

Resonance in Benzene

• Resonance is an important concept in describing the bonding in organic molecules, particularly aromatic organic molecules, a category that includes the hydrocarbon benzene.

• Benzene is commonly represented by omitting the hydrogen atoms and showing only the carbon–carbon framework with the vertices unlabeled.

• The bonding arrangement in benzene confers special stability to the molecule. As a result, millions of organic compounds contain the six-membered ring characteristic of benzene.

8.7 ∣ Exceptions to the Octet Rule

Odd Number of Electrons

• In the vast majority of molecules and polyatomic ions, the total number of valence electrons is even, and complete pairing of electrons occurs.

Less than an Octet of Valence Electrons

• A second type of exception occurs when there are fewer than eight valence electrons around an atom in a molecule or polyatomic ion.

More than an Octet of Valence Electrons

• Molecules and ions with more than an octet of electrons around the central atom are often called hypervalent.

• There are Lewis structures where you might have to choose between satisfying the octet rule and obtaining the most favorable formal charges by using more than an octet of electrons.

• Other researchers claim that the bond lengths in the ion are more consistent with the right structure being dominant. This disagreement is a convenient reminder that, in general, multiple Lewis structures can contribute to the actual electron distribution in an atom or molecule.

8.8 ∣ Strengths and Lengths of Covalent Bonds

• The stability of a molecule is related to the strengths of its covalent bonds.

• As the number of bonds between the carbon atoms increases, the bond length decreases and the bond enthalpy increases. That is, the carbon atoms are held more closely and more tightly together. In general, as the number of bonds between two atoms increases, the bond grows shorter and stronger.