CHAPTER 6 REVIEW Teacher Notes and Answers 1.In ionic bonding, large numbers of oppositely charged ions join because of mutual electrical attraction. In covalent bonding, atoms join by sharing electron pairs. In metallic bonding, atoms join through an attraction to a sea of valence electrons. 2.The electrons in a covalent bond occupy overlapping orbitals; each electron is free to occupy either of the orbitals, but both are more likely to be in the space between the nuclei of the bonded atoms. 3.A multiple bond is needed when there are not enough valence electrons to complete octets by adding unshared pairs. 4.Ionic compounds have higher melting points and boiling points than molecular compounds do, and they do not vaporize at room temperature. 5.Properties include hardness, brittleness, and electrical conductivity in the molten state. 6.Metals are better heat conductors than ionic or molecular compounds. Metals are more easily deformed and are better electrical conductors in the solid state. Metals are also shiny. 7.According to VSEPR theory, the shapes of molecules are classified based on the number of bonding electron pairs and lone pairs that surround a molecule’s central atom. 8.Intermolecular forces are the forces of attraction between molecules. 9.H and I: 2.4 − 2.0 = 0.4, polar-covalent, I; S and O: 3.5 − 2.5 = 1.0, polar-covalent, O; K and Br: 2.8 − 0.8 = 2.0, ionic, Br; Si and Cl: 3.0 − 1.8 = 1.2, polar-covalent, Cl; K and Cl: 3.0 − 0.8 = 2.2, ionic, Cl; Se and S: 2.5 − 2.4 = 0.1, nonpolar-covalent, S; C and H: 2.4 − 2.0 = 0.4, polar-covalent, C 10.K and Cl, K and Br, Si and Cl, S and O, H and I and C and H, Se and S 11.Bonding is stronger between the ions in sodium chloride because its lattice energy is greater (more negative). Greater lattice energy indicates stronger ionic bonding. 12a. Cl 12b. Ca 12c. C 12d. P F FCF 13a. F 13b. H Se H H H C Cl 13c. H 13d. N N 14a.trigonal-pyramidal 14b.bent or angular 14c.bent or angular 15a.toward F 15b.toward Cl 15c.toward I 16a.nonpolar 16b.polar 16c.polar 16d.nonpolar 16e.polar 16f.polar 17a.polar 17b.nonpolar 17c.nonpolar 17d.polar 17e.nonpolar 18a. Cl S Cl bent or angular I PI 18b. I trigonal-pyramidal 18c. Cl O Cl bent or angular HNH 18d. Cl trigonal-pyramidal Cl Cl Si Cl Br 18e. tetrahedral 18f. O N Cl bent or angular - 18g. ON O O O O S O O 18h. trigonal-planar 2- tetrahedral CHEMICAL BONDING 1 CHAPTER 6 REVIEW 1. Identify and define the three major types of chemical bonding. 2. Describe the general location of the electrons in a covalent bond. 3. In writing Lewis structures, how is the need for multiple bonds generally determined? 4. In general, how do ionic and molecular compounds compare in terms of melting points, boiling points, and ease of vaporization? 5. List three general physical properties of ionic compounds. 6. How do the properties of metals differ from those of both ionic and molecular compounds? 7. How is VSEPR theory used to classify molecules? 8. What are intermolecular forces? 2 CHAPTER 6 9. Complete the table with respect to bonds formed between the pairs of atoms. Bonded atoms Electronegativity difference Bond type More-negative atom H and I S and O K and Br Si and Cl K and Cl Se and S C and H 10. List the bonding pairs from Question 9 in order of increasing covalent character. 11. The lattice energy of sodium chloride, NaCl, is –787.5 kJ/mol. The lattice energy of potassium chloride, KCl, is –715 kJ/mol. In which compound is the bonding between ions stronger? Why? 12. Use electron-dot notation to illustrate the number of valence electrons present in one atom of each of the following elements. a. chlorine, Cl c. carbon, C b. calcium, Ca d. phosphorus, P 13. Draw Lewis structures for each of the following molecules. a. CF4 c. CClH3 b. H2 Se d. N2 CHEMICAL BONDING 3 14 According to the VSEPR theory, what molecular geometries are associated with the following types of molecules? a. AB3 E b. AB2 E2 c. AB2 E 15. For each of the following polar molecules, in which direction is the dipole oriented? (Which end would the dipole arrow point toward?) a. H–F b. H–Cl c. H–I 16. Determine whether each of the following bonds would be polar or nonpolar. a. H–H c. H–F e. H–Cl b. H–O d. Br–Br f. H–N 17. On the basis of individual bond polarity and orientation, determine whether each of the following molecules would be polar or nonpolar. a. H2O c. CF4 b. I2 d. NH3 e. CO2 18. Draw Lewis structures for each of the following molecules or ions, and then use VSEPR theory to determine the geometry of each. 4 a. SCl2 e. SiCl3 Br b. PI3 f. ONCl c. Cl2 O g. NO 3 d. NH2 Cl h. SO24 CHAPTER 6
© Copyright 2024