Name:______Date: ______Period: ______

Chapter 6 Review 1: Chemical Bonding

Short Answer: Answer the following questions in the space provided

  1. Identify the major assumption of the VSEPR theory that is used to predict the shape of atoms.
Electrons in molecules repel each other______

______

  1. In water, two hydrogen atoms are bonded to one oxygen atom. Why isn’t water a linear molecule?

Water has two unshared electron pairs. It is an AB2E2 (bent) molecule; lone pairs occupy space.______

  1. What orbitals combine together to form sp3 hybrid orbitals around a carbon atom?

S, P, P, and P combine to form sp3 (four) making possible the tetrahedral geometry.__

  1. What two factors determine whether or not a molecule is polar?

Polarity and orientation. Bond polarities can be additive causing molecules as a whole to be polar or bond dipoles can cancel one another (molecular polarity=0)

  1. Arrange the following types of attractions in order of increasing strength, with 1 being the weakest and 4 the strongest.

__3__ covalent

__4__ ionic

__2__ dipole-dipole

__1__ London Dispersion

  1. How are dipole-dipole attractions, London dispersion forces, and hydrogen bonding similar?
Dipole-dipole- force of attraction between polar molecules

London dispersion: attraction from creation of instantaneous dipoles.

Hydrogen bonding: H attracted to lone pair of electronegative atoms.

All dipole-dipole attraction between molecules (intermolecular forces)

  1. Complete the following table:

Formula _____ Lewis Structure______Geometry______Polar

H2S

CCl4

BF3

H2O

PCl5

BeF2

SF6

Name: ______Date:______Period ______

Chapter 6 Review: Chemical Bonding

Short Answer: Answer the following questions in the space provided

  1. Name the type of energy that is a measure of strength for each of the following types of bonds:

__lattice energy (Kj/mol)__ a. ionic bond

___bond energy (Kj/mol)__ b. covalent bond

heat of vaporization (Kj/mol)c. metallic bond

  1. Use the electronegativity values shown in Figure 5-20 on page 151 of the text, to determine whether each of the following bonds is nonpolar covalent, polar covalent, or ionic.

____4.0-2.1=1.9 Ionic__ a. H-F ___2.1-2.1=0, N.P.C.___ d. H-H

_____3.0-0.9=2.1 Ionic_ b. Na-Cl __2.5-2.1=0.4 P.C._____ e. H-C

___3.5-2.1=1.4_P.C.__ c. H-O ___3.0-2.1=0.9__P.C.___ f. H-N

  1. How is a hydrogen bond different from an ionic or covalent bond?

Ionic- large number of oppsitely charge ions join because of mutual electrical attraction

Covalent-atoms: join by sharing electron pairs

Hydrogen bond- intermolecular forces, type of dipole-dipole interaction, weaker than forces in ionic or covalent bond. Strength: Ionic > covalent>H-bonds

  1. H2S and H2O have similar structures and their central atoms belong to the same group. Yet H2S is a gas at room temperature and H2O is a liquid. Use bonding principles to explain why this is true.
H-S electronegativity difference: 2.5-2.1=0.4

H-O electronegativity difference 3.5-3.1=1.4

H-S- barely polar v.s. H-O very polar

The stronger the bond the greater the difference in electronegativity

  1. Why is a polar-covalent bond similar to an ionic bond?
Difference in electronegativity
  1. Draw a Lewis structure for each of the following formulas. Determine whether the molecule is polar or non polar.

______POLAR______a. H2S

:S:

H H

______POLAR______b. COCl2

______POLAR______c. PCl3

______POLAR______d. CH2O