INTRODUCTION AND SECTION 8.1

In this chapter we have focused on the interactions that lead to the formation of chemical bonds. We classify these bonds into three broad groups: ionic bonds, which result from the electrostatic forces that exist between ions of opposite charge; covalent bonds, which result from the sharing of electrons by two atoms; and metallic bonds, which result from a delocalized sharing of electrons in metals. The formation of bonds involves interactions of the outermost electrons of atoms, their valence electrons. The valence electrons of an atom can be represented by electron-dot symbols, called Lewis symbols. The tendencies of atoms to gain, lose, or share their valence electrons often follow the octet rule, which can be viewed as an attempt by atoms to achieve a noble-gas electron configuration.

SECTION 8.2

Ionic bonding results from the transfer of electrons from one atom to another, leading to the formation of a three-dimensional lattice of charged particles. The stabilities of ionic substances result from the strong electrostatic attractions between an ion and the surrounding ions of opposite charges. The magnitude of these interactions is measured by the lattice energy, which is the energy needed to separate an ionic lattice into gaseous ions. Lattice energy increases with increasing charge on the ions and with decreasing distance between the ions. The Born-Haber cycle is a useful thermochemical cycle in which we use Hess's law to calculate the lattice energy as the sum f several steps in the formation of an ionic compound.

SECTION 8.3

A covalent bond results from the sharing of electrons. We can represent the electron distribution in molecules by means of Lewis structures, which indicate how many valence electrons are involved in forming bonds and how many remain as unshared electron pairs. The octet rule helps determine how many bonds will be formed between two atoms. The sharing of one pair of electrons produces a single bond; the sharing of two or three pairs of electrons between two atoms produces double or triple bonds, respectively. Double and triple bonds are examples of multiple bonding between atoms. The bond length decreases as the number between the atoms increases.

SECTION 8.4

In covalent bonds, the electrons may not necessarily be shared equally between two atoms. Bond polarity helps describe unequal sharing of electrons in a bond. In a non polar covalent bond the electrons in the bond are shared equally by the two atoms; in a polar covalent bond on of the atoms exert a greater attraction for the electrons than the other.
Electronegativity is a numerical measure of the ability of an atom to compete with other atoms for the electrons shared between them. Fluorine is the most electronegative element, meaning it has the greatest ability to attract electrons from other atoms. Electronegativity values range from 0.7 for Cs to 4.0 for F. Electronegativity generally increases from left to right in a row of the periodic table and decreases going down a column. The difference in the electronegativity difference, the more polar the bond.
A polar molecule is one whose centers of positive and negative charge do not coincide. Thus, a polar molecule has a positive side and a negative side. This separation of charge produces a dipole, the magnitude of which is given by the dipole moment, which is measured in debase (D). Dipole moments increase with increasing amount of charge separated and increasing distance of separation. Any diatomic molecule X--Y in which X and Y have different electronegativities is a polar molecule.
Most bonding interactions lie between the extremes of covalent and ionic bonding. While it is generally true that the bonding between a metal and a nonmetal is predominantly ionic, exceptions to this guideline are not uncommon when the difference in electronegativity of the atoms is relatively small or when the oxidation state of the metal becomes large.

SECTION 8.5 AND 8.6

If we know which atoms are connected to one another, we can draw Lewis structures for molecules and ions by a simple procedure. Once we do so, we can determine the formal charge of each atom in a Lewis structure, which is the charge that the atom would have if all atoms had the same electronegativity. In general, the dominant Lewis structure will have low formal charges with any negative formal charges residing on more electronegative atoms.
Sometimes a single dominant Lewis structute is inadequate to represent a particular molecule by using two or more resonance structures for the molecule. The molecule is envisioned as a blend of these multiple resonance structures. Resonance structures are important in describing the bonding in molecules such as ozone, O3, and the organic molecule benzene, C6H6.

SECTION 8.7

The octet rule is not obeyed in all cases. Exceptions occur when (a) a molecule has an odd number of electrons, (b) it is not possible to complete an octet around an atom without forcing an unfavorable distribution of electrons, or (c) a large atom is surrounded by a sufficiently large number of small electronegative atoms that it has more than an octet of electrons around it. Lewis structures with more than an octet of electrons are observed for atoms in the third row and beyond in the periodic table.

SECTION 8.8

The strength of a covalent bond is measured by its bond enthalpy, which is the molar enthalpy change upon breaking a particular bond. Average bond enthalpies can be determined for a wide variety of covalent bonds. The strengths of covalent bonds increase with the number of electron pairs shared between two atoms. We can use bond enthalpies to estimate the enthalpy charge during chemical reactions in which bonds are broken and new bonds formed. The average bond length between two atoms decreases a the number of bonds between the atoms decreases as the number of bonds between the atoms increases, consistent with the bond being stronger as the number of bonds increases.

INTRODUCTION AND SECTION 9.1

The 3D shapes and sizes of molecules are determined by their bond angles and bond lengths. Molecules with a central atom A surrounded by 'n' atoms B, denoted AB'n', adopt a number of different geometric shapes, depending on the value of 'n' and on the particular atoms include. In the overwhelming majority of cases, these geometries are related to five basic shapes (linear, trigonal pyramidal, tetrahedral, trigonal bipyramidal, and octahedral).

SECTION 9.2

SECTION 9.3

The dipole moment of a polyatomic molecule depends on the vector sum of the dipole moment associated with the individual bonds, called the bond dipoles. Certain molecular shapes, such as linear AB2 and trigonal planar AB3, assure that the bond spills cancel, producing a non polar molecule, which is one whose dipole moment is 0. In other shapes , such as bent AB2 and trigonal pyramidal AB3, the bond dipoles do not cancel and the molecule will be polar (that is, it will have a nonzero dipole moment).

SECTION 9.4

Valence-bond theory is an extension of Lewis's notion of electron-pair bonds. In valence-bond theory, covalent bonds are formed when atomic orbitals on neighboring atoms overlap one another. The overlap region is one of low energy, or greater stability, for the two electrons because of their simultaneous attraction to two nuclei. The greater the overlap between two orbitals, the stronger will be the bond that is formed.

SECTION 9.5

To extend the ideas of valence-bond theory to polyatomic molecules, we must envision missing s, p, and sometimes d orbitals to form hybrid orbitals. The process of hybridization leads to hybrid atomic orbitals that have a large lobe directed to overlap with orbitals on another atom to make a bond. Hybrid orbitals can also accommodate nonbonding pairs. A particular mode of hybridization can be associated with each of three common electron-roman geometries (linear = sp; trigonal planar = sp2; tetrahedral = sp3).

SECTION 9.6

SECTION 9.7

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Chapter 8 & 9 Overview/Summary