Chapter 9, Section 10.1 and Bond
Types Lab Test

Electron Dot / Lewis Structures
◦ Be able to draw and analyze them.

Geometries:
◦ Linear, bent, trigonal planar, tetrahedral,
trigonal pyramidal, trigonal bipyramidal,
octahedral
◦ Bond angles for above geometries
◦ Polarity of the molecule
◦ IMF the molecule will have
Chapter 9, Section 10.1 and Bond
Types Lab Test

Hybridization
◦ Hybrid orbitals involving s and p
◦ Sigma and pi bonds
◦ Orbital overlap

Molecular orbital diagrams
◦ Realize that this is another method for
visualizing molecules
◦ Uses orbital notation similar to what is used
for atoms

Would either KBr or C7H8 be soluble in
water? Use intermolecular forces to
explain.
◦ Would either be soluble in hexanes?

How do IMFs relate to the following:
◦
◦
◦
◦
◦
◦
◦
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Viscosity
Vapor pressure
Boiling point
Freezing point
State at room temperature
Solubility of a substance
The bond types lab that we did
Intramolecular bonds
What is wrong with the localized electron
model? What other model improves upon
it?
 What is delocalization? How can we show
this is a molecule like benzene? Or NO3-?


What is PES? What can we learn from PES?
Website with some PES practice:
http://www.chem.arizona.edu/chemt/Flash/ph
otoelectron.html

Use this spectra for the questions in the
next slide (these are the types of questions
on PES you may have to answer on the
test)
a.
b.
c.
d.
e.
f.
Assuming that the PES data shows ALL of the electrons
present in the atom, identify the element.
Still applying the assumption in a., which specific electrons
are associated with the peak at 126 eV?
What does a relatively low value of energy tell you about
the relative position of the electrons within any atom?
Which peak in the spectrum represents electrons that are
closest to the nucleus?
What is the relevance of the relative height of the peaks at
5.31 eV and 9.07 eV?
Why is there such a large difference in energy between
the peaks at 0.74, 5.31 and 9.07 eV, and the peak at 126
eV?
More PES Practice
http://www.adriandingleschemistrypages.c
om/ap/pes-a-first-attempt-at-creating-aquestion/
 http://apchemresources2014.weebly.com/
uploads/9/7/6/4/9764824/photoelectron_s
pectroscopy_info.pdf

Sections 10.8-10.9

Use the worksheet on heating curves and
phase diagrams to help you study. There
will be a minimal amount of information
on the test over this.

Concepts from the Bond Types Lab:
◦ How to determine the type of bond present
in a compound.
◦ Properties that each type of compound has
including solubility, melting point, appearance,
conductivity, etc. What is the chemical reason
behind these properties?
◦ What role to IMFs play in this? Can you
determine the type of IMF based on
properties?
◦ Lab techniques
AP Multiple Choice Review
From chapter 9: All
 From chapter 10: 1-7, 10-13, 16-20

Using principles of chemical bonding and/or
intermolecular forces, explain each of the
following.
(a) Xenon has a higher boiling point than neon
has.
(b) Solid copper is an excellent conductor of
electricity, but solid copper chloride is not.
(c) SiO2 melts at a very high temperature,
while CO2 is a gas at room temperature, even
though Si and C are in the same chemical
family.
(d) Molecules of NF3 are polar, but those of
BF3 are not.
Explain each of the following in terms of
atomic and molecular structures and/or
intermolecular forces.
(a) Solid K conducts an electric current,
whereas solid KNO3 does not.
(b) SbCl3 has measurable dipole moment,
whereas SbCl5 does not.
(c)The normal boiling point of CCl4 is 77ºC,
whereas that of CBr4 is 190ºC.
(d) NaI(s) is very soluble in water, whereas
I2(s) has a solubility of only 0.03 gram per 100
grams of water.
Which of the following substances has the
greatest solubility in C5H12(l) at 1 atm?
(A) SiO2(s)
(B) NaCl(s)
(C) H2O(l)
(D) CCl4(l)
(E) NH3(g)