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Basic Concepts of Chemical Bonding

Basic Concepts of Chemical Bonding. 8.1 Chemical Bonds, Lewis Symbols, and the Octet Rule

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Basic Concepts of Chemical Bonding

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  1. Basic Concepts of Chemical Bonding • 8.1 Chemical Bonds, Lewis Symbols, and the Octet Rule • Three different types of chemical bonds are introduced: ionic, covalent, and metallic. Ionic and covalent bonds are the focus of this chapter. Here, you will see that atoms tend to lose, gain, or share electrons to attain a noble-gas electron configuration. This observation is summarized in the octet rule. • 8.2 Ionic Bonding • This section focuses on the bonding due to electrostatic attraction between oppositely charged particles. The principle of electrostatic attraction will be reviewed with an activity in which you can vary the charges on, and distance between, two particles to see what effect each has on the magnitude of the force between them. The term lattice energy is introduced as a quantitative measure of the attractive forces between ions in a solid. • 8.3 Sizes of Ions • Electron configurations and electrostatic forces are used to explain the differences in size between atoms and their corresponding ions. You will watch a movie of the electron-loss and electron-gain processes that details how each affects size. • 8.4 Covalent Bonding • Some atoms share electrons to achieve a noble-gas electron configuration. A chemical bond formed by sharing a pair of electrons is called a covalent bond. Electrostatic forces again are used to explain the length of the H–H bond. • 8.5 Bond Polarity and Electronegativity • This section introduces the concept of electronegativity and illustrates the trends in electronegativity contained in the periodic table. Electronegativity differences give rise to polar bonds, which are bonds in which electron density between atoms is not shared equally. You will learn to calculate a dipole moment—a quantitative measure of the polarity of a bond. An activity will represent bond dipoles as vectors for bonds between atoms that you choose. • 8.6 Drawing Lewis Structures • The rules for representing molecules and polyatomic ions with Lewis structures are introduced here. You will determine which bonds in a molecule are single bonds, and which are double or triple bonds. A movie explains formal charges, a method of keeping track of electrons that helps you determine the best Lewis structures. Nomenclature is revisited briefly from the standpoint of using the covalent nomenclature convention for compounds that contain metals. • 8.7 Resonance Structures • Often, there are two or more equally correct Lewis structures for the same species. Those that differ only by placement of electrons are called resonance structures. None of the resonance structures truly represents the molecule accurately. The bond lengths and strengths are intermediate between the possible resonance structures. • 8.8 Exceptions to the Octet Rule • There are three types of exceptions to the octet rule. All three are discussed in this section and examples are given for each. You will again use formal charges to decide which structures are most feasible. • 8.9 Strengths of Covalent Bonds • Bond enthalpy is defined in this section, and tabulated values of average bond enthalpies are used to calculate enthalpies of reaction. The strength of a bond and its length are shown to have a reciprocal relationship.

  2. Chemical Bonds, Lewis Symbols, and the Octet Rule • An ionic bond results from the powerful electrostatic forces that exist between oppositely charged ions. Ionic substances form readily between elements from the far left of the periodic table (metals) and elements from the far right of the periodic table (nonmetals). Metals tend to lose electrons to form positively charged cations, while nonmetals tend to gain electrons to form negatively charged anions. Ionic substances are solids at room temperature. • When atoms are similar in their tendencies to lose or gain electrons, they share electrons to form a covalent bond. The most familiar examples of covalent bonds are found in the interactions of nonmetals. Substances held together by covalent bonds can be solids, liquids, or gases at room temperature. • Atoms in metallic solids such as copper, iron, and aluminum are held together by metallic bonding. In this type of bonding, valence electrons of the metal atoms are free to move throughout the metal solid. • octet rule. Atoms tend to gain, lose, or share electrons until they are surrounded by eight valence electrons.

  3. Ionic Bonding • A measure of the magnitude of the electrostatic attractive forces in ionic compounds is the lattice energy. The lattice energy is defined as the energy required to convert a mole of an ionic solid into its constituent ions in the gas phase. Although it is not possible to measure the lattice energy directly, it is possible to calculate it with Hess's law, using measurable quantities and the Born-Haber cycle.

  4. Born-Haber cycle

  5. Size of Ions

  6. Lewis Structures A single line denotes the sharing of two electrons in a single bond. Lewis structures can also be used to show multiple bonding in a molecule or polyatomic ion.A double line means that two pairs of electrons are shared in a double bond, and a triple line means that three pairs of electrons are shared in a triple bond.

  7. Bond Polarity and Electronegativity

  8. Drawing Lewis Structures • Sum the valence electrons from all atoms. (Use the periodic table to help you do this.) For an anion, add an electron to the total for each negative charge. For a cation, subtract an electron for each positive charge. Don't worry about keeping track of which electrons come from which atoms. Only the total number is important. • Write the symbols for the atoms to show which atoms are attached to which, and connect them with a single bond (a dash, representing two electrons). Chemical formulas are often written in the order in which the atoms are connected in the molecule or ion. When a central atom has a group of other atoms attached to it, the central atom is usually written first. In other cases you may need more information before you can draw the Lewis structure. • Complete the octets of the atoms bonded to the central atom. (Remember, however, that hydrogen can have only two electrons.) • Place any leftover electrons on the central atom, even if doing so results in more than an octet. • If there are not enough electrons to give the central atom an octet, try multiple bonds. Use one or more of the unshared pairs of electrons in the atoms bonded to the central atom to form double or triple bonds.

  9. Lewis Structures (continued) It is sometimes possible to write more than one correct Lewis structure for a molecule. When two or more correct Lewis structures differ only by placement of electrons, they are called resonance structures.

  10. Exceptions to the Octet Rule • 1. Molecules with an odd number of electrons • 2. Molecules in which an atom has less than an octet • 3. Molecules in which an atom has more than an octet BF3

  11. Strengths of Covalent Bonds

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