Chemical Bonding 10
Chapter Overview
Understanding how atoms bond gives us the power of predicting chemical behavior.
Drawing Lewis structures and predicting the resulting molecular shapes is discussed. The
chemical interactions based on shape, including electronegativity and polarity, are also
explained.
Lecture Outline
10.1 Bonding Models and AIDS Drugs
10.2 Representing Valence Electrons with Dots
Learning Objective: Write Lewis structures for elements.
A. Valence electrons are outermost shell electrons
10.3 Lewis Structures for Ionic Compounds: Electrons Transferred
Learning Objective: Write Lewis structures for ionic compounds.
Learning Objective: Use the Lewis model to predict the chemical formula of an ionic
compound.
10.4 Covalent Lewis Structures: Electrons Shared
Learning Objective: Write Lewis structures for covalent compounds.
A. Two bonded nonmetals share electrons such that both get an octet (or duet)
B. Both species get “credit” for all electrons in bond
C. Two species can share two, four, or six electrons
2. Four electrons shared is a double bond
10.5 Writing Lewis Structures for Covalent Compounds
Learning Objective: Write Lewis structures for covalent compounds.
A. Write the correct skeletal structure for the molecule
B. Calculate the total number of valence electrons
C. Distribute electrons among atoms giving each an octet (or duet)
D. If any atoms then lack an octet, form double or triple bonds as necessary
E. Exceptions to the octet rule
1. Duet rule for hydrogen
10.6 Resonance: Equivalent Lewis Structures for the Same Molecule
Learning Objective: Write resonance structures.
10.7 Predicting the Shapes of Molecules
Learning Objective: Predict the shapes of molecules.
A. VSEPR theory
1. Electron groups repel each other
2. Electron geometry
a. Linear
B. Predicting structures
2. Count the number of electron groups
4. Determine electron geometry and molecular geometry
10.8 Electronegativity and Polarity: Why Oil and Water Don’t Mix
Learning Objective: Determine whether a molecule is polar.
A. Electronegativity
B. Bond dipoles
Chemical Principle Teaching Ideas
Lewis Structure
All covalent bonds involve the sharing of electrons. By understanding how the atoms
bond to each other, the students can begin to understand why species react the way they do.
Molecular Shapes
The most effective way for students to remember the various molecular shapes is to
Electronegativity
Some atoms hold on to electrons tighter than others. In some interactions, the bonding is
therefore uneven. Atoms are involved in a sort of tug-of-war with the electrons. A purely
Skill Builder Solutions
10.2. NaBr is an ionic compound, so Na donates the 1 valence shell electron it has to bromine,
which then has an octet in its valence shell. Sodium has a +1 charge and Br has a –1
10.3. Since Mg has a +2 charge and N has a –3 charge, the molecular formula is Mg3N2. The
10.4. Carbon monoxide has a total of 4 + 6 = 10 valence electrons. The skeletal structure is
C-O, and then we add electrons around the outer atoms, giving them octets. We can start
with :
C-O
:, but carbon does not have an octet, so we must form a triple bond with the
10.5. There are a total of 12 valence electron in this species. Following the symmetry
guidelines, and placing 2 electrons in for each bond, we get
10.6. The species has 7 electrons coming from the Cl and 6 coming from the O atom. This
makes a total of 13, but one more comes from the –1 charge of the ion, for a total of 14.
The two species share one pair of electrons, to give each an octet. The Lewis structure is
10.7. The base Lewis structure is [:
O
–
N
–
O
O
O
O
O
:]–. The nitrogen has only 6 electrons around it,
so it wishes to make a multiple bond with one of the oxygen atoms. It does not matter
10.8. The central nitrogen has three groups of electrons around it, two of which are bonds and
10.9. The central sulfur atom has four groups of electrons around it: one a lone pair and
10.10. a. Because two iodine atoms have the exact same electronegativity, neither is stronger
than the other. Therefore, the bond is pure covalent.
10.11. CH4 has a tetrahedral electron geometry and a tetrahedral molecular geometry. Since all
of the bonds are of the same slight polarity in terms of electronegativity difference and
Suggested Demonstrations
Blow up four equally sized balloons and tie the knots together. The resulting structure is
tetrahedral in geometry, and you can explain how the balloons try to get as far apart as possible.
Then pop one of the balloons to show how three orbitals (balloons) orient themselves. Then pop
another balloon and explain the resulting structure change.
Have a few students (of various sizes) come to the front of the room and have them make various
molecular geometries by holding arms in various orientations. This is an effective method for
showing bond dipoles, dipole moments, and polarity.
Guided Inquiry Ideas
Below are a few example questions that students answer in the guided inquiry activities provided
in the Guided Activity Workbook.
Which atom do you think is central in carbon dioxide? Why?
Carbon dioxide has two groups of electrons surrounding the carbon atom. Why is the OCO bond
angle in CO2 180°?