Homework Chapter 6
1. What is the main distinction between ionic and covalent bonding?
2. How is electronegativity used in determining the ionic or covalent character of the bonding between two elements?
3. What type of bonding would be expected between the following atoms?
a) H and F
b) Cu and S
c) I and Br
4. List the three pairs of atoms referred to in the previous question in order of increasing ionic character of the bonding between them.
5. What is a chemical bond?
6. Identify the three major types of chemical bonding.
7. a) What is the meaning of the term polar, as applied to chemical bonding.
b) Distinguish between polar covalent and nonpolar covalent bonds.
8. In general, what determines whether atoms will form chemical bonds?
1. Define molecule.
2. a) What determines bond length?
b) In general, how are bond energies and bond lengths related?
3. Describe the general location of the electrons in a covalent bond.
4. Use electron configuration and orbital notation to illustrate the bonding in each of the following molecules:
a) chlorine, Cl2;
b) oxygen, O2;
c) hydrogen fluoride, HF
1. As applied to covalent bonding, what is meant by an unshared or lone pair of electrons?
2. Describe the octet rule in terms of noble gas configurations and potential energy.
3. Determine the number of valence electrons in an atom of each of the following elements by showing the electron configuration of each element and drawing a circle around the valence electrons.
4. When drawing Lewis structures, which atom is usually the central atom?
5. Use electron dot notation to illustrate the number of valence electrons present in one atom of each of the following elements:
6. Use electron dot structures to demonstrate the formation of ionic compounds involving the following elements:
a) Na and S
b) Ca and O
c) Al and S
7. Draw Lewis structures for each of the following molecules:
a) contains one C and four F atoms
b) contains two H and one Se atom
c) contains one N and three I atoms
d) contains one Si and four Br atoms
e) contains one C, and Cl, and three H atoms
1. Write the structural formula for methanol, CH3OH.
2. Draw the Lewis structure for each of the following:
1. Draw Lewis structures for each of the following molecules:
1. Define the following
a) bond length
b) bond energy
2. State the octet rule.
3. How many pairs of electrons are shared in the following types of covalent bonds?
a) a single bond
b) a double bond
c) a triple bond
4. Draw the Lewis structures for the following molecules:
1. Give two examples of an ionic compound.
2. Use electron dot notation to demonstrate the formation of ionic compounds involving the following:
a) Li and Cl
b) Ca and I
3. Distinguish between ionic and molecular compounds in terms of the basic units that each is composed of.
4. Compound B has lower melting and boiling points than compound A. At the same temperature, compound B vaporizes faster and to a greater extent than compound A. If one of these compounds is ionic and the other is molecular, which would you expect to be molecular? ionic? Explain the reasoning behind your choices.
5. a) Define an ionic compound.
b) In what form do most ionic compounds occur?
6. a) What is a formula unit?
b) What are the components of one formula unit of CaF2?
7. a) What is lattice energy?
b) In general, what is the relationship between lattice energy and the strength of ionic bonding?
8. a) In general, how do ionic and molecular compounds compare in terms of melting points, boiling points, and ease of vaporization?
b) What accounts for the observed differences in the properties of ionic and molecular compounds?
c) Cite three physical properties of ionic compounds.
9. a) What is a polyatomic ion?
b) Give two examples of polyatomic ions.
c) In what form do such ions often occur in nature?
10. 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? Explain why.
1. Use VSEPR theory to predict the molecular geometry of the following molecules:
2. What two theories can be used to predict molecular geometry?
1. Use VSEPR theory to predict the molecular geometries of the molecules whose Lewis structure are given in the practice problem on page 187 of your text. Explain your answer based on VSEPR theory.
2. Draw the Lewis structure, and use the VSEPR theory to predict the molecular geometry of the following molecules:
3. a) How is the VSEPR theory used to classify molecules?
b) What molecular geometry would be expected for F2 and HF?
4. According to the VSEPR theory, what molecular geometry's are associated with the following types of molecules?
1. What are some factors that affect the geometry of a molecule?
2. Explain what is meant by sp3 hybridization.
1. What type of intermolecular force contributes to the high boiling point of water? Explain.
2. Describe the role of each of the following in predicting molecular geometry's:
a) unshared electron pairs
b) double bonds
3. a) What are intermolecular forces?
b) In general, how do these forces compare in strength with those in ionic and metallic bonding?
c) Where are the strongest intermolecular forces found?
4. What is the relationship between electronegativity and the polarity of a chemical bond?
5. a) What are dipole dipole forces?
b) What determines the polarity of a molecule?
6. a) What is meant by an induced dipole?
b) What is the everyday importance of this type of intermolecular force?
7. a) What is hydrogen bonding?
b) What accounts for its extraordinary strength?
8. What are London dispersion forces?
9. For each of the following polar molecules, indicate the direction of the resulting dipole:
10. On the basis of individual bond polarity and orientation, determine whether each of the following molecules would be polar or nonpolar: