CHEMICAL BONDING
TEST REVIEW
1. Understand that when forming chemical bonds atoms are attempting to form more stable
electronic configurations.
2. Understand the concept of electronegativity.
3. Understand that differences in electronegativity in covalent molecules causes dipoles and
some ionic character in covalent compounds
4. Understand that ionic bonding and covalent bonding are at two ends of a sliding scale of
bond type.
5. Understand that polarization caused by small highly charged cations leads to ionic
compounds exhibiting some covalent character.
6. Understand the concept of ionic bonding and the nature of the ionic bond.
7. Understand the concept of covalent bonding and nature of the covalent bond.
8. Be able to draw Lewis structures.
9. Understand the concept of resonance related to Lewis structures.
10. Understand the concept of formal charge related to Lewis structures.
11. Be able to predict the shape of, and bond angles in, simple molecules and ions using VSEPR
theory.
12. Understand the concept of the coordinate covalent bond related to Lewis structures.
13. Be able to predict the shapes of simple molecules and ions using Lewis structures.
14. Understand when molecules exhibit polarity.
15. Describe the formation of sigma and pi bonds and determine sigma and pi bonds in any given
molecule.
16. Describe covalent bond formation in terms of overlap of atomic orbitals.
17. Understand and be able to identify different types of orbital hybridization.
REVIEW QUESTIONS
1.
Explain lattice energy and indicate which substance has the most: RbI, MgO, CaBr2.
2.
Rank the substances in increasing ionic character order: NaI, KF, KCl, KI, NaCl
3.
The C-F bond in CF4 has an electronegativity difference of 1.53 which makes it a polar
covalent bond. How can the CF4 molecule have polar covalent bonds and dipoles, but have a
nonpolar geometry?
4. Draw the resonance structures for ClO3-.
5.
Determine the formal charges on the atoms in OH-1
6. 1989 D
CF4
XeF4
ClF3
(a) Draw a Lewis electron-dot structure for each of the molecules above and identify the shape
of each.
(b) Use the valence shell electron-pair repulsion (VSEPR) model to explain the geometry of each
of these molecules.
7. Draw the Lewis Structure, predict the molecular structures, describe the bond angles, and
describe the bonding in terms of the hybrid orbitals for the central atom for the following:
a. PBr2Cl3
b. PI5
c. H2Se
d. AlCl3
e. H3O+
8. PCl5 has the shape of a trigonal bipyramid whereas IF5 has the shape of a square pyramid.
Account for this difference.
9. Complete the following:
Hybrid Type
a.
Geometry
linear
b. dsp3
c.
Hybrid Type
d.
Geometry
octahedral
e. sp2
tetrahedral
10. Deduce the hybridization of the central atom, the geometry, and the molecular polarity of
each of the following molecules:
a. NH3
b. C2H4
c. ClO3-1
11. Draw the Lewis structure for Benzene, C6H6. Label all pi and sigma bonds. Give numbers
of each. Explain how the pi bonds form.
12. Use simple structure and bonding models to account for each of the following.
a. The bond length between the two carbon atoms is shorter in C2H4 than in C2H6.
b. The H-N-H bond angle is 107.5º, in NH3.
c. The bond lengths in SO3 are all identical and are shorter than a sulfur-oxygen single bond.
d. The I3- ion is linear.
13. 1996 D
Explain each of the following observations in terms of the electronic structure and/or bonding of
the compounds involved.
(b) Molecules of AsF3 are polar, whereas molecules of AsF5 are nonpolar.
(c) The N-O bonds in the NO2- ion are equal in length, whereas they are unequal in HNO2.
(d) For sulfur, the fluorides SF2, SF4, and SF6 are known to exist, whereas for oxygen only OF2 is
known to exist.