AP Chemistry by Satellite Lectureguide
Student Edition
Chemical Bonds
Chapter 8
Objectives
Following your study of this chapter, you should be able to
1a. Complete the following table
b) In general terms summarize the relationship between the number of valence
electrons in a noble gas element and its chemical reactivity.
c) Complete the following table
How many valence electrons are present in an alkali metal? How many electrons must an alkali metal gain or lose to attain a noble gas electron configuration? Which is more likely, gain or loss of electrons?
1d. Complete the following table
How many valence electrons are present in an alkaline earth metal? How many electrons must an alkaline earth metal gain or lose to attain a noble gas electron configuration? Which is more likely, gain or loss of electrons?
e) Complete the following table
How many valence electrons are present in oxygen? How many electrons does oxygen gain or lose to attain a noble gas electron configuration? Which is more likely, gain or loss of electrons?
f) Complete the following table
How many valence electrons are present in a halogen? How many electrons must a halogen gain or lose to attain a noble gas electron configuration? Which is more likely, gain or loss of electrons?
2a. Define the terms cation, anion and ionic bond.
b) Complete the following table.
3. Describe some of the characteristic physical properties of ionic compounds.
4. Draw the Lewis structures for the following atoms or ions.
5. Explain the octet rule.
6a. Given the following list of cations and anions, write all possible chemical
formulas.
6b. Supply the missing information below:
7a. Define the term lattice energy and list factors that effect the magnitude of the lattice energy in a chemical compound.
b) Explain the magnitudes of these lattice energies in terms of ion sizes and ionic
charge.
8. What class(es) of element(s) typically unite to form compounds that contain ionic bonds?
9a. How does loss or gain of electrons affect the radius of an atom? Explain the
magnitude of observed changes in terms of effective nuclear charges on the
valence electrons.
10. Define the term isoelectronic and cite an example of isoelectronic species.
11. If ionic bonds form as a result of electron transfer, how do covalent bonds form?
What class(es) of element(s) typically unite to form compounds that contain
covalent bonds?
12. Describe characteristic physical properties of compounds that contain covalent
bonds.
13. Define the term electronegativity. Use the periodic table to display general trends in electronegativity within periods and groups.
14. Explain the difference between the models for polar bonds and nonpolar bonds. Cite at least one example of each type of bond in your explanation. How do electronegativity differences between bound atoms correlate with bond polarity?
15a. List the steps used to draw the Lewis electron-dot structure for a covalent compound.
16a. Describe the difference among the models for single, double and triple bonds. (Note: Include such properties as bond length and bond energy.)
Problem Set #12
AP Chemistry by Satellite
ALL work must be shown in all problems for full credit.
PS12.1. Write the electron configurations for each of the following atoms or ions.
a) Na
b) Mn2+
c) Br-
d) Kr
e) Pb2+
PS12.2. Predict the formula of the ionic compound formed between the following pairs of elements.
a) Na and Br2
b) Al and O2
c) Ba and S
d) Fe and Cl2
PS12.3. Explain why Mg2+ is smaller than S2-. Explain why Mg is larger than S.
PS12.4. Which of the following species form an isoelectronic group?
N3-, Cl-, Ne, Mg2+, Se2-, H+
PS12.5. Which of the following salts has the largest lattice energy? Explain.
LiF, LiCl, LiBr, LiI
PS12.6. Write the equation which describes the reaction which is associated with the lattice energy of an ionic compound such as MgO.
PS12.7. Predict whether the following compounds are ionic or covalent.
SiCl4, MgBr2, PH3, NH4Cl, HCl, Al2O3
PS12.8. Arrange the following elements from smallest to largest electronegativity.
O, Al, Ga, I, H, Na
PS12.9. Write the Lewis structures for the following ions or molecules.
a) HBr
b) PCl3
c) SO32-
9d. ClO3-
e) C2H4
f) CH2Cl2
g) Cl2CO
h) HCN
17. Cite at least three different examples of compounds that violate the octet rule.
18a. Define the term resonance, and explain when it can be used.
19a. Define the term bond energy (bond dissociation energy.)
Given the information in the table below;
b) Explain the observed relationship between bond length and bond energy in the three examples of carbon-carbon bonds and in the three examples of carbon- oxygen bonds. Which is stronger and why?
19c. Compare the bond strengths in a dihydrogen molecule and a chlorine molecule. Which is stronger and why?
20a. Write the mathematical equation for estimating enthalpies of reaction from bond energies.
21. Define the term oxidation number.
22. List the rules used to assign oxidation numbers for elements in compounds and
illustrate an example for each rule.
23. Briefly state rules for naming binary ionic and binary covalent compounds. Give several examples showing how the rules are applied.
24. List differences between the physical properties of ionic metal oxides and those
of the covalent nonmetal oxides.
25. Use chemical equations to illustrate differences between the reactions with water
and ionic and covalent oxides. Which oxides produce acid solutions and which
produce basic solutions?
Problem Set #13
AP Chemistry by Satellite
ALL work must be shown in all problems for full credit.
PS13.1. Write the Lewis structures for the following ions or molecules. (If a molecule cannot be adequately represented by a single diagram, include all resonance structures.)
a) CO32-
b) NO2-
c) HCO2-
d) N2O
PS13.2. Use bond dissociation energies to estimate the
for HCl(g).
PS13.3. Use bond dissociation energies to estimate the enthalpy of the reaction for
PS13.4. Use bond dissociation energies to estimate the enthalpy of the reaction for
PS13.5. Use bond dissociation energies to estimate the enthalpy of the reaction for
PS13.6. Determine the oxidation state of the boldfaced elements in each of the
following:
a) Ca3P2
b) SO42-
c) K2Cr2O7
d) Na2O2
e) FePO4
PS13.7. Give the name or chemical formula for each of the following substances:
a) lead(II) nitrate
b) dinitrogen pentoxide
c) Na2CO3
d) FeCl3
e) HCl
f) P4O6
g) chromium(III) sulfate
PS13.8. Predict the products of the following reactions.
Microcomputer software
Introduction to General Chemistry by Stan Smith, Ruth Chabay and Elizabeth Kean
Drill-and-practice software
$500 (10-disk set)
Falcon Software
P.O Box 200
Wentworth, NH 03282
1-603-764-5788
Diskette #3 Chemical Formulas and Equations
Diskette #6 Chemaze
Computer Aided Instruction for General Chemistry by William Butler & Raymond Hough
Drill-and-practice software
$40 (4-disk set)
John Wiley & Sons, Inc.
605 3rd Avenue
New York, NY 10158
(this software may not be available)
Diskette #1 Periodic Properties of the Elements
Diskette #2 Nomenclature and Oxidation Numbers