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7-2 OBJECTIVES

7-2 OBJECTIVES. Define covalent bond. Give examples of substances that contain covalent bonds. Distinguish between a molecular formula, an empirical formula, and a structural formula. Describe covalent bonds in terms of Lewis dot structures and the octet rule.

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7-2 OBJECTIVES

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  1. 7-2 OBJECTIVES • Define covalent bond. • Give examples of substances that contain covalent bonds. • Distinguish between a molecular formula, an empirical formula, and a structural formula. • Describe covalent bonds in terms of Lewis dot structures and the octet rule. • Explain the difference between single, double, and triple covalent bonds. • Compare and contrast polar and nonpolar covalent bonds. PPowell 05

  2. Covalent Bond • Formed when two atoms share a pair of electrons • How are covalent bonds different from ionic bonds? • Molecule: two or more atoms bonded covalently OBJ: Define covalent bond. OBJ: Give examples of substances that contain covalent bonds.

  3. Molecules and Their Formulas • Molecular formula: tells how many atoms are in one molecule of the compound • molecular formula for table sugar: C12H22O11 • Does NOT show you which atoms are bonded to which • This means you can have the same molecular formula for different compounds OBJ: Distinguish between a molecular formula and a structural formula.

  4. F F Shared electrons Structural Formulas • Show ratio of different kinds of atoms in a compound • Show which atoms in a compound are bonded to each other OBJ: Distinguish between a molecular formula and a structural formula.

  5. Molecular vs. Empirical Formulas • Empirical formulas give you the simplest ratio of the atoms in the compound • Ex: glucose • Mol. formula: C6H12O6 • Emp. formula: CH2O • This means different compounds can share the same empirical formula • Ex: lactose (a type of sugar in milk) • Mol. formula: C3H6O3 • Emp. formula: CH2O OBJ: Distinguish between a molecular formula and an empirical formula.

  6. F F F F Covalent Bond Describing Covalent Bonds • Formation of covalent bonds allows atoms to satisfy the octet rule • Fluorine atoms have 7 valence electrons • Two F atoms each share an e- with the other in a covalent bond • Now each F atom has a complete octet OBJ: Describe covalent bonds in terms of Lewis dot structures and the octet rule.

  7. F F F F Describing Covalent Bonds • When drawing Lewis structures, use a dash to represent a covalent bond: • This is a single bond • But atoms can share more than one pair of electrons… becomes OBJ: Describe covalent bonds in terms of Lewis dot structures and the octet rule.

  8. Double and Triple Bonds • Double covalent bond: made up of two pairs of shared e- (one pair per bond) • Triple covalent bond: made up of three pairs of shared e- (one pair per bond) OBJ: Explain the difference between single, double, and triple covalent bonds.

  9. Drawing Lewis Dot Structures For Molecules • Add up the valence e- from all the atoms • Write the symbols for the atoms to show which atoms are attached to which, and connect them with a single bond • Complete the octets of the atoms bonded to the central atom • Place any leftover electrons on the central atom, even if it results in more than an octet • If there are not enough electrons to give the central atom an octet, try multiple bonds OBJ: Describe covalent bonds in terms of Lewis dot structures and the octet rule.

  10. e- e- e- e- e- e- e- e- e- e- e- e- H C O O H H H H C O H C O C H H Molecular Formula: H2CO 1. 12 valence e- (check the dot diagrams for yourself) 2. Write symbols, connect with single bonds 2a. Notice that C needs 2 e- to make the octet, and O will need 6 e- BUT you only have 6 e- left 3. Use the remaining e- to complete the octets To satisfy the octet rule for both O and C, there must be a double bond. OBJ: Explain the difference between single, double, and triple covalent bonds.

  11. F F F B Exceptions to the Octet Rule 1.Atoms with less than an octet • Example: Boron OBJ: Describe covalent bonds in terms of Lewis dot structures and the octet rule.

  12. S F F F F Exceptions to the Octet Rule • Atoms with more than an octet -Example: Sulfur Extra e- fill the 3d orbitals of atoms that fall into this category OBJ: Describe covalent bonds in terms of Lewis dot structures and the octet rule.

  13. N O N O Exceptions to the Octet Rule 3. Atoms with an odd number of electrons • Example: Nitrogen + OBJ: Describe covalent bonds in terms of Lewis dot structures and the octet rule.

  14. Properties of Covalent Bonds • Polar covalent bond: (aka polar bond) a bond between two atoms with different electronegativities • Uneven sharing causes the more electronegative atom to be slightly negative in charge • The less electronegative atom will be slightly positive in charge OBJ: Compare and contrast polar and nonpolar covalent bonds.

  15. More About Polar Bonds • Figure 7-20, p. 242—draw it! • Water has polar bonds • Slight charges are called partial charges • Partial charges symbolized with Greek letter delta () • Arrows can also be used to show the uneven sharing of electrons—point toward more electronegative atom, crossed at the other end OBJ: Compare and contrast polar and nonpolar covalent bonds.

  16. Properties of Covalent Bonds • Nonpolar covalent bonds: (aka nonpolar bonds) a bond between two atoms of similar or equal electronegativities • Examples: H2, N2, O2 OBJ: Compare and contrast polar and nonpolar covalent bonds.

  17. Can you now meet the 7-2 OBJECTIVES? • Define covalent bond. • Give examples of substances that contain covalent bonds. • Distinguish between a molecular formula, an empirical formula, and a structural formula. • Describe covalent bonds in terms of Lewis dot structures and the octet rule. • Explain the difference between single, double, and triple covalent bonds. • Compare and contrast polar and nonpolar covalent bonds.

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