Name_________________
Unit 8
Atomic Theory and
Periodicity
-1-
Name_________________
Chemistry Unit 8: Atomic Theory and Bonding
Assignment
WB Page Number
Podcast 8.1
Video Clip + Discussion with
your teacher: Mindwalk
Worksheet A
Podcast 8.2
Black Box Lab
Demo: Spectral Tubes +
Discussion with Teacher
Demo: And Then There Was
Light + Discussion with your
teacher
Take Home Lab: Let there be
Light
Lab: Flame Tests
Worksheet B
Podcast 8.3 (CB 15-27)—This is
a long one
Lab-Activity: Periodic Trends
Worksheet C
Worksheet D
Lab: Model Building
Podcast 8.4 (CB 29-35)
Worksheet E
Molecular Modeling Computer
Activity
Worksheet F
Unit 8 Exam
Online + CB Pages
In Class
Can Only
be done in
Class
Online + CB Pages
Teacher Handout
In Class
X
X
In Class
X
Pg 7
Pg 8-9
Pg 16-17
Online + CB Pages
Pg 11-12
Pg 18
Pg 19-20
Teacher Handout
Online and CB Pages
Pg 21
Online: See Teacher
Pg 22
In Class
-2-
X
Done?
Name_________________
Take Home Lab
Parent/Student Experiment
Title: And Then There Was Light....
Subject/Concept: Chemistry - Photon Emission
Purpose:
 The purpose of this activity is to observe the emission of photons in your own home!
Don’t worry, this happens all the time!
Materials:
 several commercial bandage strips (3” x .75” with pull-apart packaging - no strings!)
 CURAD™, KING SOOPERS™, OR SAFEWAY™ brands work well
 regular Wintergreen LifeSavers™ candies or Wintergreen LifeSavers™ Holes
Procedure:
In an absolutely pitch dark room (bathrooms often work), do the following:
 Pull apart the bandage strip packaging with very quick pulls of about a half inch or so.
You should see the emission of a small purple photon!
 While your partner looks on, crush the wintergreen candy between your teeth and your
partner will see the emission of a small photon!
Questions:
1. What is the source of the emitted photon?
2. Make a drawing of the Rutherford-Bohr model of an atom showing the movement
of an electron during the process of light emission.
For Credit:
To receive credit, complete the questions for this lab on a separate piece of paper. Also,
your parent or guardian must write a short note confirming that you performed the experiment
for them and explained the results to their satisfaction using the concept of photon emission and
electron energy levels. Attach your note to the back of this sheet.
-3-
Name_________________
Flame Test Lab
Chemists began studying colored flames in the 18th century and soon used "flame
tests" to distinguish between some elements. Different elements burn with different
colored flames. Although some of the flames you will be seeing will appear similar in
color, their light can be resolved (separated) with a prism into distinctly different
bands of colors on the electromagnetic spectrum (ROYGBIV). These bands of colors
are called atomic line spectra, and they are UNIQUE to each element. Niels Bohr
studied the line spectrum for hydrogen, and wondered what the specific line
spectrum had to do with the structure of the atom. He postulated that an electron
can have only specific energy values in an atom, which are called energy levels. Bohr
believed that the energy levels for electrons were quantized, meaning that only
certain, specific energy levels were possible. How does an electron move between
energy levels? By gaining the right amount of energy, an electron can move, or
undergo a transition, from one energy level to the next. We can explain the emission
of the light by atoms to give the line spectrum like this:
1. An electron in a high energy level (excited state) undergoes a transition to a
low energy level (ground state).
2. In this process, the electron loses energy, which is emitted as a photon (a
particle which behaves like a wave)
3. The energy difference between the high energy level and the low energy level
is related to the frequency (color) of the emitted light.
Pre-lab questions:
1. Bohr's important discovery was that energy levels of electrons are quantized (only
existing in certain, specific levels). In what year was this discovery made?
_____________
2. What happens to an electron when energy is added?
3. What is released when an electron loses energy?
4. What determines the frequency (color) of photons?
5. Why do you think the frequencies (color) for a specific element is always the same?
Procedure: In this lab, you will be observing the colors of the flames for 7 different
elements: lithium, sodium, potassium, calcium, strontium, barium, and copper.
Each element is dissolved in a solution of its chloride salt. There is a different solution
at each lab station. You will go around to all 7, perform the flame test, and make
CAREFUL observations of the colors. You will then be given an unknown solution, for
-4-
Name_________________
which you will have to use your notes below to determine which unknown you were
given.
Post- Lab Questions:
1. If you had 2 colors that seemed identical, how could you tell them apart more
accurately?
2. Albert Einstein determined this equation:
energy (in joules) of a photon is equal to Planck's constant times the frequency of
the light:
E = h •  • Frequency () has units of 1/sec (which is a Hertz, or Hz)
• Planck's constant (h) = 6.63 x 10-34 J·sec
a) If the frequency of a red spectrum line is at 1.60 x 1014 Hz, how much energy does
each photon of this light have?
b) If the frequency of a violet spectrum line is at 2.50 x 1014 Hz, how much energy
does each photon of this light have?
c) On the far ends of the visible spectrum of light, there exists ultraviolet (UV)
radiation and infrared (IR) radiation.
- UV radiation is dangerous. UV radiation is located just past violet on the spectrum.
 IR radiation is harmless. It is located just past red on the spectrum.
 Based on what you calculated in parts a & b, explain -why- UV is more
dangerous than IR:
-5-
Name_________________
Lab: A Black
Box
Overview:
As humans, we rely on our senses to tell us about the world around us. Does that mean
we can only understand those things we can see, feel, hear, taste, or touch? Of course
not. Sometimes we can sense things indirectly. For example, let's say you're alone in
the house, and leave a juicy cheeseburger on the table for a few minutes while you
answer the phone. You return five minutes later and find that the cheeseburger is gone.
You also notice your dog contentedly sitting under the table. Perhaps it looks at you and
burps. It's pretty obvious that the dog must have eaten the cheeseburger since
cheeseburgers don't walk away by themselves. You didn't sense the event directly, but
based on everything you've experienced about the world before allows you to logically
figure out what must have happened. Scientists must also use indirect observations to
answer many questions about things that are impossible or difficult to observe directly.
This is certainly true about our understanding of atoms.
In science, the term "black box" is used to describe something when we understand
how it behaves (and can even make predictions about what it will do in a given
situation), but are not able to see exactly what is going on inside it to make it behave the
way it does. Atoms are a great example of a scientific "black box."
We are doing this lab because, before we begin to study atoms, it's a good idea to
experience what it's like to investigate a "black box." The 4 boxes you'll be using are
literally boxes, but you'll follow the same basic methods that scientists use to study
anything that cannot be directly observed. Here's how you to proceed with this lab:
1. Make and record as many observations about the box as you can without removing
any rods or looking inside. It is helpful to team up with another lab pair and their box so
that you can compare notes. Draw a sketch of what you think the inside of the box is
like based on these observations.
2. Decide which rod you would like to remove, and write down a prediction of what will
happen when you do this. Try it, and record what happens.
3. You may then use another lab pair's box to remove a second rod. Record what
happens. Remember that each time you remove a rod you are possible changing the
contents of the box in a way that cannot be reversed. Take your time, and only remove
one rod at a time.
-6-
Name_________________
4. Draw a final sketch of what you think the inside of the box looked like before you
removed any rods. Continue with the next box until you have developed models for all
four boxes.
Discussion Notes:
a) Discuss how this activity relates to what scientists do in real-life, giving examples.
b) Discuss your confidence in the models of the boxes you have developed. Would you
be willing to publish the results? Do you think it is possible that your model will change if
new information is discovered? How is this similar to the model of the atom that
scientists have developed?
-7-
Name_________________
Model Building
Using Styrofoam balls and toothpicks, build each shaped structure that is taught in the podcasts.
Credit is earned when you show these to your teacher and answer the following questions:
1. What is VSEPR?
2. What is the pattern on angles for those atoms with 4 shells (or clouds) as you go from 4
atoms connected to 2 atoms connected?
3. Why do you think this pattern occurs?
4. How does the shape of the molecule affect the polarity of the molecule?
5. Using your models, explain to your teacher why water is a polar molecule.
6. Using your models, explain why carbon tetrachloride is a non-polar molecule.
-8-
Name_________________
Molecular Shapes
“clouds”
2
Linear Diatomic
Polarity depends upon
electronegativity
difference
Polar if >0.5
Nonpolar if <0.5
Linear Triatomic, Usually nonpolar
CO2, HCN
3
Trigonal Planar: BF3, SO32-, NO3120˚ Usually nonpolar
In molecules where
the outside
molecules are
different, shapes
that tend to be
nonpolar usually
become polar.
Remember to count
the number of
“clouds” of electrons,
not the actual number
of electrons. A
double or triple bond
counts as one
effective pair.
Bent, 12O˚ Usually polar
NO2-
Also: If there ever
is a two molecule
atom (diatomic) that
molecule’s polarity
depends upon the
electronegativity
difference of the
atoms
4
Tetrahedral; 109˚: Usually
nonpolar CH4, CF4
Pyrimidal: 107˚ Usually
polar: NH3, PCl3
Bent: 104.5˚
Usually polar: H2O, OF2
Molecular Shapes
“clouds”
2
Linear Diatomic
Polarity depends upon
electronegativity
difference
Polar if >0.5
Nonpolar if <0.5
Linear Triatomic, Usually nonpolar
CO2, HCN
3
Trigonal Planar: BF3, SO32-, NO3120˚ Usually nonpolar
Remember to count
the number of
“clouds” of electrons,
not the actual number
of electrons. A
double or triple bond
counts as one
effective pair.
Bent, 12O˚ Usually polar
NO2-
4
Tetrahedral; 109˚: Usually
nonpolar CH4, CF4
In molecules where
the outside
molecules are
different, shapes
that tend to be
nonpolar usually
become polar.
Pyrimidal: 107˚ Usually
polar: NH3, PCl3
-9-
Bent: 104.5˚
Usually polar: H2O, OF2
Also: If there ever
is a two molecule
atom (diatomic) that
molecule’s polarity
depends upon the
electronegativity
difference of the
atoms
Name_________________
Periodicity Graphs
Directions: Construct three computerized graphs using the data from the table below.
Graph 1:
electronegativity (on y-axis) versus atomic number (on x-axis)
Graph 2:
first ionization energy (on y-axis) versus atomic number (on x-axis)
Graph 3:
atomic radius (on y-axis) versus atomic number (on x-axis)
On each graph, use a connected line between data points. Also, the title of each graph must also include
your initials. For each graph, include a short explanation which:
(A) defines the property on the y-axis,
(B) discusses the general trend of the property across the rows of the periodic table
(Periodic Trends), and
(C) discusses the general trend of the property down the columns of the periodic table
(Group Trends).
More information on this subject can be found in Chapter 13 of your text book.
Periodic Trends - Element Data Table
Element
Atomic Number Electronegativity
First Ionization Atomic Radius
Symbol (use on x-axis) (no units)
Energy (kJ/mole)
(picometers)
----------------------------------------------------------------------------------------------------------------------------------H
1
2.1
1312
He
2
- -(skip)
2371
Li
3
1.0
520
-----------------------------------------------------------------Be
4
1.5
900
B
5
2.0
800
C
6
2.5
1086
-----------------------------------------------------------------N
7
3.0
1402
O
8
3.5
1314
F
9
4.0
1681
-----------------------------------------------------------------Ne
10
- -(skip)
2080
Na
11
0.9
495.8
Mg
12
1.2
737.6
-----------------------------------------------------------------Al
13
1.5
577.4
Si
14
1.8
786.2
P
15
2.1
1012
-----------------------------------------------------------------S
16
2.5
999.6
Cl
17
3.0
1255
Ar
18
- -(skip)
1520
-----------------------------------------------------------------K
19
0.8
418.8
Ca
20
1.0
589.5
------------------------------------------------------------------
- 10 -
37
50
140
90
80
77
71
66
64
70
157
136
143
118
109
103
91
94
196
174
Name_________________
Periodic Trends Graph
- 11 -
Name_________________
WS A: History of the atom and Electron Configurations
1. Identify the two particles found in the nucleus of an atom.
2. What is an electron?
3. Identify the scientists (Thompson, Bohr, Rutherford, or Newton) who proposed each of the
models illustrated below:
4. What specific evidence (from the gold foil experiment) led Rutherford to come to the
each of the following conclusions?
a. The nucleus occupies very little space in the atom.
b. An atom is made of mostly empty space.
c. The nucleus is positively charged.
- 12 -
Name_________________
5. For each of the elements listed below, complete the following:
a. long hand electron configuration
b. orbital diagram (boxes with arrows)
c. short hand electron configuration (begins with a noble gas)
1. Li
2. Na
3. K
4. B
5. Al
6. Ne
7. Ar
8. Mg
- 13 -
Name_________________
9. P
10. Ni
11. Zn
12. Br
13. H
14. He
15. Ca2+
16. N3-
17.
Na1+
- 14 -
Name_________________
WS B: Light and Light Equations
1. If an electron goes from level 4 to level 2 what happens? Be specific.
2. Fill in the following table
Violet
Energy (Joules)
6.3 x 10-19 J
Blue
Wavelength-λ
(meters)
Green
Yel
Frequency –ν (s-1)
2.4 x 10-7 m
100 s-1
1.5 x 10-14 J
10
Orange
2.2 x 1013 s-1
525 nm
- 15 -
Red
Color of
Light/type of
electromagnetic
radiation
Name_________________
The diagram above represents the spectra of three different elements. Explain the
following:
a. Why are they different
b. Why are they not a continuous series of colors (ROYGBIV)?
c. How are these used in the field of Astronomy?
- 16 -
Name_________________
Part C: Periodicity
1. List the following atoms in order of increasing electronegativity:
a. Cr, Ni, Ga, K
b. P, As, F, Hg, Fr
2. List the following atoms in order of increasing atomic radius:
a. Cr, Ni, Kr, Ga, K
b. P, As, F, Hg, Fr
3. List the following atoms in order of increasing ionization energy:
a. Cr, Ni, Kr, Ga, K
b. P, As, F, Hg, Fr
4. Why are alkali metals stored in kerosene or mineral oil? Why are they not allowed to sit
out in the air?
5. The Mg+2, and the Na+1 ions each have ten electrons surrounding the nucleus. Which ion
would you expect to have the smaller radius?
- 17 -
Name_________________
Worksheet D: Bonding Introduction
1. What type of atoms combine to form a covalent bond?
2. What type of atoms combine to form a ionic bond?
3. What type of atoms combine to form a metallic bond?
4. Give two examples of a covalent compounds?
5. Give two examples of a ionic compounds?
6. Give two examples of a metallic compounds?
7. Describe how a covalent bond forms between two atoms.
8. How does a covalent bond differ from an ionic bond?
9. _________________is defined as the energy required to break the chemical bond
between two atoms and separate them.
10. _________________is the tendency of an atom to attract bonding electrons to itself when
it bonds with another atom.
11. _________________is the attraction between two atoms in which bonding electrons are
shared _________________between two atoms
12. In general, if the difference in electronegativity between two atoms is zero the bond
formed is _________________
13. If the electronegativity difference between two atoms is between 0.5 and 2.1 the bond
formed is _________________
- 18 -
Name_________________
14. If the electronegativity difference between two atoms is greater than 2.1, the bond is
_________________
15. In an ionic bond, the valence electrons are ___________.
16. In a metallic bond the valence electrons form a ______________________________
17. Rank the bonds (ionic, covalent, metallic) in order from strongest to weakest.
18. Classify each of the following compounds as either: Ionic, Covalent, Metallic.
a.
b.
c.
d.
e.
f.
g.
h.
i.
j.
k.
l.
m.
n.
o.
p.
q.
r.
s.
H2O
NaCl
MgSO4
CsCl
Fe
Hg
He
Ca3(PO4)2
NH4Cl
NH3
P2O5
Ag
AgNO3
AgCl
Titanium
Barium Phosphate
Sulfur Dioxide
Bromine
Tungsten V Bromide
______________
______________
______________
______________
______________
______________
______________
______________
______________
______________
______________
______________
______________
______________
______________
______________
______________
______________
______________
- 19 -
Name_________________
WS E: Lewis Structure Worksheet: (do on your own paper: this will take quite a bit of
paper)
 Draw the Lewis Structures
 Determine the shape
 Determine the polarity of the molecule
1. HCl
2. Br2
3. SeBr2
4. CF4
5. PI3
6. O2
7. N2
8. H2
9. OI2
10. CS2
11. SiBr4
12. F2
13. HCN
14. NH4 +
15. NO2 –
16. SO3
17. SO4218. NO3-
19.
20.
21.
22.
23.
24.
25.
26.
27.
28.
PO33CNCO2
CO
I2
CO32SO2
OCNSCNO3
For the following structures: You do
NOT have to determine the shapes
29. H3CCOOH
30.CH3CH2OH
31. H3COCH3
32. H3CCH3
33. H2CCH2
34. HCCH
- 20 -
Name_________________
WS F: Intermolecular Forces
Substance #1
Predominant
Intermolecular Force
Substance #2
(a) HCl(g)
I2
(b) CH3F
CH3OH
(c) H2O
H2S
(d) SiO2
SO2
(e) Fe
Kr
(f) CH3OH
CuO
(g) NH3
CH4
(h) HCl(g)
NaCl
(i) SiC
Predominant
Intermolecular Force
Substance with Higher
Boiling Point
Cu
2. Rank the following substances in order from lowest to highest melting point.
CO2, NaCl, Ag, H2O, He, HBr
3. Rank the following substances in order from lowest to highest freezing point.
H2O, Ca3(PO4)2, Cr, C2H6, OF2
4. Rank the following substances in order from highest to lowest boiling point.
Cl2, Ne, Ca, Cr(OH)3, CH3CH2OH, Diamond
- 21 -
Name_________________
WS A: History of the atom and Electron Configurations
1. Identify the two particles found in the nucleus of an atom.
Proton--Neutron
2. What is an electron?
Negatively Charged Particle in the Nucleus
3. Identify the scientists (Thompson, Bohr, Rutherford, or Newton) who proposed each of the
models illustrated below:
a.
b.
c.
d.
Bohr
Dalton
Rutherford
Thompson
5. What specific evidence (from the gold foil experiment) led Rutherford to come to the
each of the following conclusions?
a. The nucleus occupies very little space in the atom.
Only a few of the alpha particles hit the nucleus
b. An atom is made of mostly empty space.
Only a few of the alpha particles hit the nucleus
c. The nucleus is positively charged.
Alpha Particles are positively charged so they deflected.
- 22 -
Name_________________
5. For each of the elements listed below, complete the following:
d. long hand electron configuration
e. orbital diagram (boxes with arrows)
f. short hand electron configuration (begins with a noble gas)
2
2
6
1s 2s 2p 3s2 3p6 4s1
1. Li
2
1s 2s1
[He] 2s1
1s
2s
2. Na
1s2 2s2 2p6 3s2 3p1
1s
2s
[Ne] 3s1
2p
3. K
1s2 2s2 2p6 3s2 3p64s1
1s
2s
3s
[Ar] 4s1
2p
3s
4s
4. B
1s2 2s2 2p1
1s
[He] 2s2 2p1
2s
2p
- 23 -
3p
Name_________________
5. Al
1s2 2s2 2p6 3s2 3p1
1s
[Ne] 3s2 3p1
2a
p
2s
3s
3p
6. Ne
1s2 2s2 2p6
1s
[He] 2s2 2p6
2s
7. Ar
1s2 2s2 2p6 3s2 3p6
1s
2p
[Ne] 3s2 3p6
2s
2p
3s
2p
3s
3p
8. Mg
1s2 2s2 2p6 3s2
1s
[Ne] 3s2
2s
9. P
1s2 2s2 2p6 3s2 3p3
1s
2s
[Ne] 3s2 3p3
2p
3s
- 24 -
3p
Name_________________
10. Ni
1s2 2s2 2p6 3s2 3p64s13d8
1s
[Ar] 4s13d8
2s
4s
2p
3s
3p
3s
3p
3s
3p
3d
11. Zn
1s2 2s2 2p6 3s2 3p64s13d10
1s
[Ar] 4s13d10
2s
4s
2p
3d
12. Br
1s2 2s2 2p6 3s2 3p64s13d10 4p5
1s
2s
4s
[Ar] 4s13d10 4p5
2p
3d
4p
13. H
1s1
- 25 -
Name_________________
1s
14. He
1s2
1s
15. Ca2+
1s2 2s2 2p6 3s2 3p6
1s
16. N31s2 2s2 2p6
1s
Na1+
1s2 2s2 2p6
[Ne] 3s2 3p6
2s
2p
3s
[He] 2s2 2p6
2s
2p
17.
1s
[He] 2s2 2p6
2s
2p
- 26 -
3p
Name_________________
WS B: Light and Light Equations
Name ____________________
1. If an electron goes from level 4 to level 2 what happens? Be specific.
2. Fill in the following table
Violet
Blue
Green
Yel
Orange
Energy (Joules)
Wavelength-λ
(meters)
Frequency –ν (s-1)
6.3 x 10-19 J
8.28x10-19 J
6.62x10-34 J
1.5 x 10-14 J
1.99 x 10-26 J
1.48 x 10-20 J
3.78 x 10-19 J
3.15x10-7 m
2.4 x 10-7 m
3.00x106 m
1.32x10-11 m
10
1.36x10-5 m
525 nm
9.5 x 1014 s-1
1.2 x 1015 s-1
100 s-1
2.3 x 1019 s-1
3.0 x 107 s-1
2.2 x 1013 s-1
5.7 x 1014 s-1
- 27 -
Red
Color of
Light/type of
electromagnetic
radiation
UV
UV
Long Radio
X-Ray
Radio
IR
Green
Name_________________
The diagram above represents the spectra of three different elements. Explain the
following:
d. Why are they different
Each atom has different orbital levels. They each have the same orbitals (1s, 2s, 2p, etc). But
each has a different value for energy. This then translates into different energy levels. The
lines are made when electrons “fall” from a higher level to a lower level. They are different
because the “height” of the levels is different in each atom
e. Why are they not a continuous series of colors (ROYGBIV)?
They are not continuous because the electrons can only have certain allowable energy levels.
The electrons are said to be quantized.
f. How are these used in the field of Astronomy?
In astronomy these are used to identify the elements in the stars.
- 28 -
Name_________________
Part C: Periodicity
Name ____________________
1. List the following atoms in order of increasing electronegativity:
a. Cr, Ni, Kr, Ga, K
K < Cr < Ni < Ga < Kr
b. P, As, F, Hg, Fr
Fr < Hg < As < P < F
2. List the following atoms in order of increasing atomic radius:
c. Cr, Ni, Kr, Ga, K
Kr < Ga < Ni < Cr < K
d. P, As, F, Hg, Fr
F < P < As < Hg < Fr
3. List the following atoms in order of increasing ionization energy:
e. Cr, Ni, Kr, Ga, K
K < Cr < Ni < Ga < Kr
f. P, As, F, Hg, Fr
Fr < Hg < As < P < F
4. Why are alkali metals stored in kerosene or mineral oil? Why are they not allowed to sit
out in the air?
When they are exposed to air they lose their one valence electron and react.
5. The Mg+2, and the Na+1 ions each have ten electrons surrounding the nucleus. Which ion
would you expect to have the smaller radius?
The Mg+2 ion is smaller because it has one more proton which causes the electron
cloud to be held more tightly, thus making it smaller.
- 29 -
Name_________________
Worksheet D: Bonding Introduction
Name ____________
1. What type of atoms combine to form a covalent bond?
a. Non-Metal to Non-Metal
2. What type of atoms combine to form a ionic bond?
Metal to Nonmetal
3. What type of atoms combine to form a metallic bond?
Metal-Metal
4. Give two examples of a covalent compounds?
CO, H2O, C6H12O6, etc
5. Give two examples of a ionic compounds?
NaCl, Ca3(PO4)2, NH4Cl
6. Give two examples of a metallic compounds?
Iron, Copper, Zinc, Brass
7. Describe how a covalent bond forms between two atoms.
Valence Electrons Share
8. How does a covalent bond differ from an ionic bond?
Ionic Bonds: valence electrons are transfered
9. Bond Energyis defined as the energy required to break the chemical bond between two
atoms and separate them.
10. electronegativity is the tendency of an atom to attract bonding electrons to itself when it
bonds with another atom.
11. Dipole Forcesis the attraction between two atoms in which bonding electrons are shared
unevenly between two atoms
12. In general, if the difference in electronegativity between two atoms is zero the bond
formed is Non-Polar
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Name_________________
13. If the electronegativity difference between two atoms is between 0.5 and 2.1 the bond
formed is Polar
14. If the electronegativity difference between two atoms is greater than 2.1, the bond is
Ionic
15. In an ionic bond, the valence electrons are Transferred
16. In a metallic bond the valence electrons form a Sea of valence electrons
17. Rank the bonds (ionic, covalent, metallic) in order from strongest to weakest.
Covalent > Ionic > Metallic
18. Classify each of the following compounds as either: Ionic, Covalent, Metallic.
a.
b.
c.
d.
e.
f.
g.
h.
i.
j.
k.
l.
m.
n.
o.
p.
q.
r.
s.
H2O
NaCl
MgSO4
CsCl
Fe
Hg
He
Ca3(PO4)2
NH4Cl
NH3
P2O5
Ag
AgNO3
AgCl
Titanium
Barium Phosphate
Sulfur Dioxide
Bromine
Tungsten V Bromide
Covalent
Ionic
Ionic
Ionic
Metallic
Metallic
None
Ionic
Ionic
Covalent
Covalent
Metallic
Ionic
Ionic
Metallic
Ionic
Covalent
Covalent
Ionic
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Name_________________
WS E: Lewis Structure Worksheet:
 Draw the Lewis Structures
 Determine the shape
 Determine the polarity of the molecule
Name _________
1.HCl
Linear Diatomic
Polar
H Cl
1.
2.Br2
Br Br
2.
Linear Diatomic
Non Polar
3.SeBr2
Bent 104.5
Polar
Se
3.
Br
Br
4.CF4
- 32 -
Name_________________
F
4.
Tetrahedral
Non Polar
C
F
F
F
5.PI3
5.
P
I
I
I
Pyrimidal
Polar
6.O2
6.
O
O
Linear Diatomic
Non Polar
7.N2
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Name_________________
7.
N
N
Linear Diatomic
Non Polar
H
Linear Diatomic
Non Polar
8.H2
8.
H
9.OI2
O
9.
I
I
10.
CS2
10.
S
Bent 104
Polar
C S
Linear Triatomic
Non Polar
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Name_________________
11.
SiBr4
Br
Si
11.
Br
12.
12.
14.
Br
F2
F
13.
13.
Br
Tetrahedral
Non Polar
F
Linear Diatomic
Non Polar
HCN
H
C
Linear
Triatomic
Polar
N
NH4 +
- 35 -
Name_________________
H
N
14.
H
H
H
Tetrahedral
Non Polar
15.
NO2 –
- 36 -
Name_________________
15.
O
N
O
Bent 120
Polar
16.
SO3
- 37 -
Name_________________
O
16.
O
S
O
Trigonal Planar
Nonpolar
17.
SO42-
- 38 -
Name_________________
2-
O
17.
O
S
O
O
Tetrahedral
Nonpolar
18.
NO3-
- 39 -
Name_________________
-
O
18.
O
N
O
Trigonal Planar
Nonpolar
19.
PO33-
- 40 -
Name_________________
3P
O
O
Pyrimidal
Polar
20.
CN-
C
N
Linear Diatomic
Polar
21.
CO2
- 41 -
O
Name_________________
21.
O C
O
Linear Triatomic
NonPolar
22.
CO
C
O
Linear Diatomic
Polar
23.
I
I2
I
Linear Diatomic
Non Polar
- 42 -
Name_________________
24.
CO32-
O
24.
O
C
O
Trigonal Planar
Nonpolar
25.
SO2
Bent 120
Polar
- 43 -
Name_________________
26.
26.
OCN-
O C
N
1-
Linear Triatomic
Polar
27.
27.
SCN-
S
C
N
1-
Linear Triatomic
Polar
28.
28.
O3
O O
O
- 44 -
Bent 120
Nonpolar
Name_________________
For the following structures: You do NOT have to determine the shapes
29.
H3CCOOH
H O
H C C
O H
H
30. CH3CH2OH
31.
H3COCH3
- 45 -
Name_________________
H
H
H C O C H
H
H
32.
33.
H3CCH3
H2CCH2
- 46 -
Name_________________
34.
HCCH
- 47 -
Name_________________
WS F: Intermolecular Forces
Substance #1
Name _______________
Predominant
Intermolecular Force
Substance #2
(a) HCl(g)
Dipole
I2
(b) CH3F
Dipole
CH3OH
(c) H2O
H-bonding
(d) SiO2
Predominant
Intermolecular Force
Substance with Higher
Boiling Point
LDF
HCl
H-Bond
CH3OH
H2S
Dipole
H2 O
Covalent Network
SO2
Dipole
SiO2
(e) Fe
Metallic
Kr
LDF
Fe
(f) CH3OH
H-Bond
CuO
Ionic
CuO
(g) NH3
H-Bond
CH4
LDF
NH3
(h) HCl(g)
Dipole
NaCl
Ionic
NaCl
(i) SiC
Covalent Network
Metallic
SiC
Cu
2. Rank the following substances in order from lowest to highest melting point.
CO2, NaCl, Ag, H2O, He, HBr
He < HBr < H2O < Ag < NaCl
3. Rank the following substances in order from lowest to highest freezing point.
H2O, Ca3(PO4)2, Cr, C2H6, OF2
C2H6 < OF2 < H2O < Cr < Ca3(PO4)2
4. Rank the following substances in order from highest to lowest boiling point.
Cl2, Ne, Ca, Cr(OH)3, CH3CH2OH, Diamond
Diamond > Cr(OH)3 > Ca > CH3CH2OH > Cl2 > Ca
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