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Chemistry Unit Five (Chapter 7, 8 and 14)
States of Matter: Bonding
Unit Five Enduring Understanding
1. The behavior of matter is explained by the kinetic-molecular theory.
2. Chemical and physical properties are determined by inter- and intra-molecular forces.
3. Matter is constantly changing according to physical laws.
Chapter 7, 8 and 14 Essential Questions:
1. How can the octet rule be used to explain ionic and covalent bonding?
2. Why is the shape of a molecule important?
Readings for This Unit: Readings in BOLD are REQUIRED for everyone!!!
If you earned under 70% on your last exam, you should do all readings and take notes.
Review These:
(7-1) “Lewis Dot Diagrams” (p231-232) and “The Octet Rule” p 227-229
(7-1) “Ionic Bonding (p225-226 and p231-234)
(7-2) “Covalent Bonding” (p236- 239 and p241-242)
Read These:
(8-1) “The Shape of Small Molecules” (p255-262)
(8-2) “Polarity” (p 266-271)
(14-1) “Condensed States of Matter” (p 457-466)
(14-2) “Properties of Liquids” (p467-470)
(14-3) “The Nature of Solids” (p471-478)
INSIDE FRONT COVER –
INTENTIONALLY LEFT BLANK
Page 2
Unit Five, Chapters 7, 8, 14 Topics – Bonding
Reference Chapter 7, 8and 14 in your textbook
Of Chemical Bondage
Ionic, Covalent, Metallic…it’s all about electrons!
(7-1) Describe properties of ionic compounds.
(7-2) Define and describe properties of molecular compounds.
(7-1) Draw Lewis dot diagrams to show the valence electrons of an atom
(7-2) Explain the difference between single, double, and triple covalent bonds.
(7-2) Distinguish between a molecular formula and a structural formula.
The Shape of Things to Come
Are you square (or trigonal planar?)
(8-1) Define the VSEPR theory and explain its relationship to the shape of molecules.
(8-1) Name and describe the five common shapes of small molecules.
(14-3) Relate structure and bonding to the properties of metallic, molecular, ionic, and
covalent-network solids.
Polarity
It’s not just about cold-lovin’ bears!
(7-2) Compare and contrast polar and non-polar bonds.
(8-2) Determine whether a molecule is polar considering the polarity of its bonds
and the shape of the molecule.
(14-1) Describe the different types of intermolecular forces, and explain how they
influence properties of the different states of matter.
Somebody Bring Me Some Water
Water –it’s wet, it’s wacky, it’s what chemistry is all about!
(14-2) Define and explain the relationship of viscosity and surface tension to
intermolecular forces.
(14-2) Describe some of the unusual properties of water and relate those properties to
hydrogen bonding.
Key areas we’ll cover:
Chapters 7, 8 and 14
 Ionic vs. Covalent Bonding
 Lewis Dot Structures (they’re back!)
 VESPR
 Bond Polarity
What we Already Covered
Chapter 13
 Kinetic Molecular Theory
 Atmospheric Pressure
 The “Classic” Gas Laws
Page 3
Part Two: Chapters 7, 8, 14 States of Matter,
Bonding
Page 4
Video Companion Sheet
Chemistry: Chemical Bonding and Atomic Structure
Metallic Bonds
1. What do we call the powerful electrical forces which hold atoms together?
2. True or False: The physical and chemical properties of elements and compounds depend
on the nature of bonds and weak forces that hold atoms and molecules together?
3. The diameter of the atom is________________ times the size of the nucleus.
4. How are metallic bonds held together?
5. Define:
a) Elements -
b) Compounds
6. Which elements are involved in the bonding of atoms of metals?
7. A metallic bond results from a sea of negatively charged valence electrons flowing among
regularly arranged atoms which have lost their permanent outer electrons. This gives the
atom a positive charge (it becomes an ion), The attraction of the ____________________ ions
and the sea of _____________________________ bonds metallic elements together.
8. What are the characteristics of metallic bonds?
Page 5
Ionic Bonds
9. How are ionic bonds formed?
10. What are the characteristics of ionic bonds?
Covalent Bonds
11. What type of elements combines to form this kind of bond?
12. What is a molecule?
13. How are covalently bonded atoms held together?
14. What are two of the characteristic properties of covalently bonded molecules?
15. What is meant by the “Octet Rule?”
16. What is unique about carbon atoms?
17. What is true about the charge distribution of polar bonds?
18. Covalent bonds between like atoms which are electrostatically balanced are
____________________________________ (polar or non-polar; pick one!)
19. What is “electronegativity?”
20. If the difference in electronegativity between two atoms is great, the bond formed is
__________________________________.
21. If the difference in electronegativity between two atoms is great, the bond formed is
__________________________________.
Page 6
Name: _________________________________________________________________________ Hour ___________
Skill Builder 1: Ionic Compounds: Valence Electrons and Lewis Dot
Structures
1. Write down three characteristics that are unique to ionic compounds.
___________________________________________________________________________________________________
___________________________________________________________________________________________________
_______________________________________________________________________________________
2. What types of elements come together to form an ionic bond (examples: metal + metal,
non-metal + non-metal, etc.)?
_______________________________________________________________________________________________
3. a) What happens to the valence electrons of the elements that form an ionic bond? Why
do these elements form an ionic bond?
___________________________________________________________________________________________________
___________________________________________________________________________________________________
______________________________________________________________________________________
Come up with your own unique analogy of this relationship (example: two toddlers
want one toy (electrons); the bigger and tougher toddler (non-metal) snatches the toy
away from the smaller one (metal) and does not share it).
___________________________________________________________________________________________________
___________________________________________________________________________________________________
_______________________________________________________________________________________
4. What is a cation? ________________________ What is an anion?___________________________
Complete the sentence: Elements with _____ valence electrons or fewer will form
cations/anions (circle one). Elements with _____ or more valence electrons will form
cations/anions (circle one).
5. Fill in the following chart as shown in the example below:
Element
Example:
Mg
Metal or
nonmetal?
Valence
electrons
Metal
2
Lewis dot Charge of
structure
the
resulting
ion
Mg
Na
O
Zn
Page 7
+2
Anion or
cation?
Cation
Element
Metal or
nonmetal?
Valence
electrons
Lewis dot Charge of
structure
the
resulting
ion
Anion or
cation?
F
Tin (IV)
Bi
Ne
Iron (III)
Al
Cs
6. Draw the Lewis dot structures of the following ionic compounds. Then, using a different
colored pen, show how one element “steals” the other’s electrons, resulting in two ions.
(Hint: Some of the compounds may require multiple numbers of one type of element - be
sure to draw in the extra element if needed)
Example: Na
Mg
Li
Cl  Na + Cl
-
O
K
F
Cu (I)
P
O
7) Does a physical “bond” hold the two elements together in an ionic bond? If not, what
force holds these elements together?
Page 8
Skill Builder 2: Lewis Dot Structures
Draw the Lewis structure for all of the following elements:
C
H
F
Cl
O
S
Si
Na
K
Mg
Al
N
Draw the Lewis Structure for each of the following ionic compounds:
KF
AlF3
Draw the Lewis Diagram for the compound that results from the combination of the
following elements:
strontium, sulfur
potassium, oxygen
Page 9
Chem-is-try
Chapter 7: Chemical Formulas
Name ___________________________________
Period _____
Happy Surprise! Naming is BACK! 
Ionic or
Molecular?
Name
Formula
Ions
(if ionic)
ionic
magnesium nitrate
Mg(NO3)2
Mg2+ NO3-
Na2SO4
Ba(ClO3)2
NH4C2H3O2
aluminum nitride
sulfur hexafluoride
nitrogen monoxide
OF2
P2O5
FeSO4
silver nitrate
potassium oxide
dichlorine monoxide
CaCO3
LiOH
Ga2S3
N2O5
Page 10
Skill Builder 3: More Fun With Lewis Structures!
For each of the following compounds or ions, draw the Lewis structures
1)
HF
2)
H2O
3)
CF2S
4)
BH3
5)
CBr4
6)
P2H4
Page 11
13)
7)
NF3
8)
CO2
9)
O2
10)
NH41+
11)
OF2
12)
CH2BrF
CO32-
Page 12
Name: ____________________________________________________________________________ Date: ___________________ Hour ____________
# electron pairs
around central
atom
# bonding pairs
(groups)
# lone
pairs
Sketch it!
Bond
Angle
Molecular
Shape
Skill Builder 4: Valence Shell Electron Pair Repulsion (VSEPR) Theory Worksheet)
Lewis Structure
F2
BeCl3 (B needs 4
total val. e−)
BCl3 (B needs 3
total val. e−)
NO2−
Page 13
Lewis Structure
CH4
NH3
H2O
# electron pairs
around central
atom
# bonding pairs
(groups)
# lone
pairs
Sketch it!
Bond
Angle
Molecular
Shape
The electrons around the atoms in a molecule repel each other. They move to be as far apart as possible while still maintaining the bonding within the molecule.
The procedure for using the model is as follows:
1) Determine the correct Lewis structure for the molecule. If it is a diatomic (has only two atoms) it is linear. If it has 3 or more atoms continue with step 2.
2) Count the number of electron groups around the central atom. A group of electrons is a bond, a nonbonding electron pair, or occasionally an unpaired
nonbonding electron. Each triple or double bond counts as only one group for the purposes of this model.
3) Based on this number of groups around the central atom the molecule falls into one of three basic categories. Within each category there are a number of
different names for the shapes depending upon the number of atoms and nonbonding groups around the central atom.
Page 14
Name _________________________
Date _______________
Hour _______
Skill Builder 5: Molecular Shape and Bond Angles
Directions: For each of the following covalent compounds provide the Lewis Structure, Molecular Shape, and Bond
Angles.
#
1
Formula
NF3
2
CN1-
3
NO21-
4
CO32-
5
OF2
6
BI3
7
O2
Lewis Structure
Molecular Shape
Page 15
Bond Angle
8
SeCl2
9
NH41+
10
CO2
11
PH3
12
CH2Cl2
13
CH2O
14
NH2Cl
Page 16
Name: _______________________________________________________ Date ______________ Hour _______
Skill Builder 6: Lewis Structures and Polar Bonds Worksheet
Directions: For each of the following molecules, construct their Lewis Structures and label the
POLAR bonds present (do not label non-polar bonds). See table of electronegativity for help.
1. H2O
6. OCS
2. BI3
7. PH3
3. NCl3
8. CF4
4. O2
9. SCl2
5. SiO2
10. GeBr4
Page 17
Periodic Table and Table of Electronegativities
Page 18
Name ________________________________________________________ Date ______________ Hour __________
Skill Builder 7: Polarity and Intermolecular Forces Worksheet
Draw the Lewis structure, determine what shape the molecule is, if the molecule is polar, and
what dominant intermolecular forces would be present for each of the following chemicals:
Formula
Lewis
Structure
Molecular
Shape
CH4
PF3
N2
BF3 (exception
to octet rule)
H2O
CO2
CH2Cl2
Page 19
Polar or
Nonpolar
molecule
Dominant
Intermolecular
Force
Name __________________________________________________
Date ________________ Hour ________
Skill Builder 8: Intermolecular Forces
Directions: For each of the following molecular substances, draw lewis structures depicting the intermolecular
forces present. This only needs to be done for the dipole-dipole and hydrogen bonding intermolecular forces. This is
not necessary for the molecules utilizing dispersion intermolecular forces. You should also identify which
intermolecular force(s) are present. Feel free to use the molecular models to help if necessary.
Substance
CO2
Lewis Structures showing the intermolecular force
NCl3
HCl
CH3Cl
HCN
CH4
CH2O
C2H2
NP
H2S
BF2H
Page 20
Intermolecular Force(s)
Name __________________________________________________
Date ________________ Hour ________
Skill Builder 9: Intermolecular Forces (14-1)
1. Describe the characteristics of a solid. (p459 & 471)
2. Describe the characteristics of a liquid. (p 458 & 471)
3. Describe the characteristics of a gas. (p 458)
4. According to the kinetic-molecular theory, what determines the state of matter at room
temperature? (p460)
5. What are the three characteristics of bonding? (p 460)
6. Define intramolecular forces. (p462)
7. Define intermolecular forces. (p462)
8. What is the attractive force of a substance that exhibits a dispersion force? (p464)
9. What is the attractive force of a substance that exhibits a dipole-dipole force? (p465)
10. Explain the attractive force of a hydrogen bond. Be specific. (p465)
Page 21
Name __________________________________________________
Date ________________ Hour ________
Skill Builder 10: Properties of Liquids (14-2)
1. Define “viscosity.” Give an example that will help explain the properties of liquids and
amorphous solids (remember magma?). (p467)
2. List and describe the six special properties of water. (p469)
•
•
•
•
•
•
Page 22
Name __________________________________________________
Date ________________ Hour ________
Skill Builder 11: The Nature of Solids (14-3)
1. List the physical properties of solids. (p473)
2. Define a metallic bond and list the properties of metallically bonded structures. (p474)
3. List the properties of molecular solids. (p476)
4. List the properties of ionic solids. (p476)
5. Describe a covalent-network solid. (p476)
Page 23
Ionic Bonds
Guided Notes: Bond Characteristic Properties
Metallic Bonds
Hydrogen
“Bonds”
Intermolecular Forces
Dipole
Interactions
Van der Waals forces
Dispersion
Forces
Network
Solids
Chapters 7, 8 & 14: Chemical Formulas and Bonding, Molecular Shapes and Liquids and Solids.
Pages 227 and 241-242 in chapter 7 Chemical Formulas and Bonding and pages 460-466 and 473-476 in chapter 14 on Liquids and Solids will be
very useful to you. This worksheet will help you with our lab in this unit.
Properties
State of Matter
(Room Temp.)
Melting Point
Solubility-Polar
Solvent
SolubilityNonpolar
Solvent
Conductivity
(in solid form)
Conductivity
(in solution)
Volatility
Other
Page 24
Part Two Chapters 7, 8 and 14: States of Matter,
Bonding
Page 25
Lab 1: Project: Bonding, Bonding, Bonding
"Dr. Linus Pauling is the man for me / He makes violent changes in my chemistry /
Oh, my, when he rolls his eyes / All my atoms ionize." -Song lyrics from "The Road to Stockholm." 1954
Your task is to create something original that shows me you understand the difference between
ionic bonding, covalent bonding, and metallic bonding.
You will create 3 separate comic strips demonstrating each type of bonding.
Each of the comic strips needs to be 3-5 panels.
The following grade rubric will be used to evaluate your creation.
IONIC BONDING
o You turned in something
o Comic is 3-5 panels
o Electron transfer is clear
o +/- Attraction is clear
COVALENT BONDING
o You turned in something
o Comic is 3-5 panels
o The need for e- is
clearly shown
o The sharing of e- is clear
METALLIC BONDING
o You turned in something
o Comic is 3-5 panels
o Spacing of nuclei is clear
o The mobility of e- is clear
0
0
0
0
1
1
1
1
2
2
0 1
0 1
0 1
2
0 1
2
0
0
0
0
1
1
1
1
CREATIVITY
o Wow!
o Original
o Straight from lecture!
o Straight from a book!
4
3
2
1
ERRORS
o Error free
o 1-2 errors
o 3-4 errors
o 5-6 errors
o More than 6 errors
4
3
2
1
0
WORKSMANSHIP
o Excellent
o Good
o Slapped together
during studyhall
o Slapped together
during passing period
2
2
3
2
1
0
Total ___/29
Page 26
Page 27
CH4
N2
4
5
(nitrogen)
(methane)
PH3
(phosphorous
trihydride)
(water)
H2O
3
2
HBr
1
(hydrogen
bromide)
Molecule
#
Sketch the
Ball-and Stick Model
Structural
Formula
Shape
Bond
Angles
Polar
Bonds?
The purpose of this lab is to practice writing Lewis Structures, and identify molecular shapes, and
determine bond angles and polarity of the bonds and overall molecules..
Lab 2: Molecular Model Kits, Page 1
Polar
Molecul
e?
Name: ____________________________________________________________________ Hour: _________
H2CO
C2H2*
6
7
Page 28
(hydrogen
peroxide)
H2O2
(ydriogen
cyanide)
Sketch the
Ball-and Stick Model
Structural
Formula
Shape
*There are TWO possible shapes for this molecule. Please identify BOTH.shapes and bond angles.
10
HCN
9
(chloromethane)
CH3Cl
8
(ethylene)
(hydroxylmethylene)
Molecule
#
Lab 2: Molecular Model Kits, Page 2
Bond
Angles
Polar
Bonds?
Polar
Molecul
e?
Name: ____________________________________________________________________ Hour: _________
Name _________________________________________
Date _______________
Hour ________
Lab 3: Intramolecular and Intermolecular Forces
“The senses have been conditioned by attraction to the pleasant and aversion to the unpleasant:
a man should not be ruled by them; they are obstacles in his path.” - Bhagavad Gita
Intra- vs InterWrite as many words as you can starting with intra- as a prefix (intramolecular doesn’t count!)
Write as many words as you can starting with inter- as a prefix (intermolecular doesn’t count!)
Based on your answers, what does the prefix intra- mean? _____________________
Based on your answers, what does the prefix inter- mean? _____________________
So, what kind of force is an intramolecular force? ____________________________
So, what kind of force is an intermolecular force? ____________________________
When applying energy (ie. heat), which force do you think is weakened first? Explain why you
think your answer is correct.
Which force do you think is responsible for predicting states of matter (solid, liquid, or gas)?
Page 29
Intermolecular Forces
One type of intermolecular force…
What is a dipole?
Partially positive charge is represented by ________ and partially negative charge is represented
by ___________________________________________.
How do we determine which side of a molecule is partially negative and which is partially
positive?
You and your lab partner should make a molecular model of OCl2. When you have it, you’re your
teacher check. Draw a Lewis structure of your model below, including partial charges:
The other group at the same lab table should have the same molecule as you. These molecules are
going to interact with one another in some way. Using both of your models, orient the molecules
so that they interact (hint: think about partially positive and partially negative charges). When
you have it, have your teacher check. Draw Lewis structures of these two models interacting
below, including partial charges:
Also, draw this sketch on the whiteboard provided so we can compare.
The force occurring between these molecules is an example of an intermolecular force.
These intermolecular forces are called ______________________.
What kind of molecules will use these kinds of intermolecular forces? _______________
Page 30
A second kind of intermolecular force…
You and your lab partner should make a molecular model of H 2O. When you have it, have your
teacher check. Draw a Lewis structure of your model below, including partial charges:
The other group at the same lab table should have the same molecule as you. These molecules are
going to interact with one another in some way. Using both of your models, orient the molecules
so that they interact (hint: think about partially positive and partially negative charges). When
you have it, have our teacher check. Draw Lewis structures of these two models interacting
below, including partial charges:
Also, draw this sketch on the whiteboard provided so we can compare.
These intermolecular forces are very similar to the intermolecular force on the previous page.
What is different about them?
There are three very electronegative atoms commonly found in molecules utilizing this
intermolecular force. They are _____________, ______________, and _____________.
These intermolecular forces are called ___________________. In spite of its misleading name, it is not
a covalent bond…it is an intermolecular force.
The boiling point of water is 100 oC. The boiling point of oxygen dichloride is much less. What
does this tell you about the relative strength of these new intermolecular forces compared to the
dipole-dipole forces?
Page 31
A third kind of intermolecular force…
You and your lab partner should make a molecular model of Cl2. When you have it, have your
teacher check. Draw a Lewis structure of your model below:
You might have noticed that this molecule does not have partial charges. There is an
intermolecular force present, we just have to think about how it happens.
Think about electrons. They are always moving. Is there a way a molecule can become a
temporary dipole? How can a molecule become a temporary dipole?
Draw a Lewis structure of your newly formed temporary dipole (include partial charges):
The other group at the same lab table should have the same molecule as you. These molecules are
going to interact with one another in some way. Your temporary dipole will cause, or induce
another molecule to become another temporary dipole. Using both of your models, orient the
molecules so that they interact. When you have it, have your teacher check. Draw three sets of
lewis structures of these two models interacting below:
Two normal Cl2 molecules
One temporary Cl2 dipole with a normal Cl2 molecule
Temporary Cl2 dipole causing an induced Cl2 dipole
Also, draw this sketch on the whiteboard provided so we can compare.
Page 32
These intermolecular forces are called ____________________________. All molecular substances have
dispersion forces. So the examples above, OCl2 and H2O, also have dispersion intermolecular
forces as well as dipole-dipole and hydrogen-bonding intermolecular forces respectively.
Something interesting should be considered about molecules utilizing only dispersion forces.
F2 and Cl2 are gases, yet Br2 is a liquid, and I2 is a solid at room temperature. They all are nonpolar
molecules, so they must have only dispersion intermolecular forces. Why do they exist in different
states of matter if they all have the same type of intermolecular force?
Let’s draw a flowchart to help us determine which type(s) of intermolecular forces occur
for specific molecules.
Page 33
Now let’s try an example.
What kind of intermolecular force(s) is/are present for NH3?
Draw Lewis structures of two molecules of NH3 including partial charges if necessary.
Explain your answer.
Now let’s try one more example.
What kind of intermolecular force(s) is/are present for CHF3?
Draw Lewis structures of two molecules of CHF3 including partial charges if necessary.
Explain your answer.
Page 34
Name __________________________________________________________________Date _________ Hour _______
Lab 4: Drops on a Penny Lab
“A penny is a lot of money if have not got a penny.” –Yiddish proverb
Target  - Define and explain the relationship of surface tension to intermolecular forces.
Introduction – The forces within water that are responsible for surface tension originate at the
molecular level. As you know, a water molecule is polar. The oxygen atom has a partial negative
charge, and each hydrogen atom has a partial positive charge. As a result, electrical attractions
occur between the oxygen atom of one molecule and the hydrogen atom of another molecule.
These intermolecular attractive forces are called hydrogen bonds.
Procedure:
1. Predict how many drops of water a Heads-Up penny will hold
and record in the table below.
2. Count how many drops it will actually hold.
3. Record your data.
4. Dry the penny off each time & repeat 3 more times.
5. Average the number of drops and record.
6. Draw what the penny looked like with the drops on it.
7. Repeat using the soap solution.
Results:
Trial
Prediction
#1
#2
#3
# of Pennies
of Water
# of Pennies
of Soap
Solution
Penny Drawing with Water Drops
Page 35
#4
Average
Analysis Questions:
1. Describe the shape of the water on the penny & explain why the drops form that shape.
2. What happened when the water finally flowed off the penny? Explain in terms of the forces
involved.
3. Did the penny hold the same number of drops of soap solution & the number of drops of
water? _______________ Why or why not?
Page 36
Lab 5: Solids Lab
“Oh that this too, too solid flesh should melt…” –Hamlet (William Shakespeare)
Purpose: To classify a number of compounds into groups based on their physical properties.
Theory: Compounds can be categorized according to their physical properties. These properties can
be determined by subjecting the compounds to certain physical tests. These tests include electrical
conductivity, solubility in polar and nonpolar solvents, melting point, hardness, flexibility, volatility.
In this lab experience, you will classify eight compounds. It will be your responsibility to determine
what type of substance each material is. Once you have determined the type of substance (Ionic, Polar
Covalent, Non-polar covalent, Covalent Crystal or Metallic), then provide a paragraph explaining
why you classified the substance into that group.
Procedure: Subject each of the compounds to all of the following seven tests:
1. Volatility: Smell (WAFT!) the compound. A strong smell is the mark of a high
volatility.
2. Electrical Conductivity of Solid: Test the electrical conductivity of the solid
with a conductivity tester. If the light goes on or the buzzer buzzes, the compound
conducts electricity.
3. Solubility: Take a couple crystals of the material in a small test tube and see if it
dissolves in water, which is a polar solvent. If the material does dissolve in water,
then test it for solution conductivity (see next step of directions.)
Front Desk Demo: The crystals that did not dissolve in water will be on display at
the teacher’s desk in mineral oil. Observe if the crystals dissolved in mineral oil,
which is non-polar. If the material dissolved in mineral oil, then test the solution for
conductivity.
4. Electrical Conductivity of Solution: If the solid dissolves in a solvent, test the
electrical conductivity of that solution. (If a solid doesn’t dissolve in a solvent, this
test doesn’t apply.) Make sure you test the solvent for electrical conductivity to make
sure that any conductivity seen is due to the compound, not the solvent.)
5. Hardness: Try to grind a few crystals of the solid using a mortar and pestle.
Describe what happens to the compound.
6. Melting Point: Place a small amount of the solid on an evaporating dish. Place the
evaporating dish on a hot plate. If the compound melts quickly, it indicates a low
melting point.
Discussion/Conclusion:
The substances tested represent ionic bonds, metallic bonds or covalent (molecular) bonds. For
covalent compounds, you must identify whether the intermolecular bonding is network solid,
Van der Waals: dispersion forces or dipole interaction. Determine which bond type each of your
tested materials represents. Describe the characteristics you used to determine their bond type.
That is, summarize each substances physical properties. Describe what factors might contribute to
different classification systems or for errors in classifying substances within your system.
Page 37
Page 38
XXXX
XXXX
Water
Mineral Oil
I
H
G
F
E
D
C
B
A
Volatility
Substance
Conductivity
XXXX
XXXX
Solubility
in water
XXXX
XXXX
Solubility in
mineral oil
Solids Lab Data Table
XXXX
XXXX
Conductivity
of solution
XXXX
XXXX
Hardness
XXXX
XXXX
Qualitative
melting point
Lab 6: Striking it Rich!
“Nothing is so hard for those who abound in riches to conceive how others can be in want.” –Jonathan Swift
Seeing is believing – or so it is said. In this lab, the properties of a metal will appear to change.
You will change the appearance of some pennies by heating them with zinc (Zn) metal in a zinc
chloride (ZnCl2) solution.
Data table:
Condition
Appearance
Untreated penny
Penny treated with Zn and ZnCl2
Penny treated with Zn, ZnCl2 and
heated
Procedure:
1.
Wear goggles!
2.
Obtain two old (pre-1982) pennies. Clean using a mixture of salt and vinegar until surface
is shiny. Record appearance in column labeled untreated penny.
3.
If not already present, add 2.0-2.2 grams of zinc to your beaker.
4.
Add approximately 25 ml of 1 M zinc chloride solution to the beaker containing the zinc.
Note: If beaker already contains solution, just add enough to reach 25 ml.
5.
Slide the two pennies into the beaker containing the zinc and zinc chloride solution.
Cover with watch glass and gently heat until boiling.
6.
Continue heating until both pennies have a change in appearance (2-5 minutes).
7.
Fill a small beaker with distilled water.
8.
With forceps or tongs, remove the two pennies from the solution. Remove beaker from
heat. Place both pennies in the beaker of distilled water. Record observations.
9.
Using forceps or tongs, remove the coins from the beaker of water. Rinse under running
water, then dry gently with a paper towel.
10.
Briefly heat one of the coins by placing on the hot surface of the hot plate for 10-20
seconds. Watch for a change in appearance and then remove. Do not overheat.
11.
Immediately immerse the heated coin in the beaker of distilled water. Record
observations.
12.
Remove coin from water and dry.
13.
Turn off hot plate. Do NOT dump the solution.
Questions:
1.
Compare the colors of the coins—untreated, heated in zinc chloride solution only, and
heated on hot plate. How are the coins different than they were originally? How are they
different from one another?
2.
Draw pictures of all three coin conditions at the atomic level which would explain the
difference in their appearance.
Page 39
Part Two Chapters 7, 8, 14: States of Matter,
Bonding
Page 40
Summary Notes on Bonding
Bonding is the foundation of all chemical reactions! How chemists define binds in changing, so some
of this material may go against what you heard before – and it may change again before you hit
college! These notes are designed to supplement your self-inquiry labs and book reading. Remember,
models are simplified versions of reality; just because they can’t explain everything all the times
does not mean they are without value.
Types of Bonds We’ll Cover
Bond
Type
Electronegativity
Difference
Ionic
≥2.0
Polar
Covalent
0.4 – 2.0
Non-Polar
Covalent
0.0 – 0.4
Bond
Description

“Transfer” of electrons from a metal to a non-metal.

Electrons are not truly abandoned by the metal, but
rather they “live” much closer to nucleus of the non-metal.

Unequal sharing of electrons, resulting in a dipole moment

Partial (+) and partial (-) charge imparted on atoms.

Equal sharing of electrons between identical atoms or
atoms with similar electronegativities.
Electronegativity
The relative ability of an atom to attract shared electrons to itself (how strong they “pull”).
o
o
Electronegativity of atoms decreases as you go down a group
Electronegativity of atoms increases as you go across a period.
Bonds: Polar or Non-polar?
A bond’s polarity depends on the difference in electronegativity between the participating
atoms. High difference = polar bond; low/no difference = non-polar bond.
Molecules: Polar or Non-polar?
A good rule of thumb:
o
If the molecule is symmetrical, then it’s non-polar.
o
If the molecule is asymmetrical, it’s most likely polar.
You can also check for things like number of polar vs. non-polar bonds and the absence
or presence of strong functional groups.
Lewis Dot Structures
A model which represents molecules so you can see how the valence electrons are arranged.
There are only three rules for drawing Lewis Dot Structures:
1. Only valence electrons are used, and all appear as shared (covalent) or unshared pairs
2. Duet Rule – Hydrogen, Helium, Lithium and Beryllium are stable in a molecule when
they share or “see” two electrons.
3. Octet Rule – Most other elements need to “see” 8 electrons to be stable in a molecule.
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Lewis Dot Structures
Some molecules (many, actually) have more than one correct Lewis Dot Structure. These are
called “resonance structures.”
To draw Lewis Dot structures with covalent or polar covalent bonds:
1.
Chemistry loves Symmetry! Start by choosing your best “central” atom.
2. Count up every element’s total valence electrons.
3. Draw a basic “single bond” structure, connecting all atoms – make sure you connect as
many atoms as possible to the central atom before connecting them to each other.
4. Do a post-bonding electron inventory, and see how many electrons are still available
(each single bond uses 2 electrons)
5. Distribute all the remaining available electrons around the atoms as unshared pairs.
6. Make sure everyone is happy! (Are the octet or duet rules satisfied?) . If they are,
you’re done! If not, you might have to consider redistributing the electrons in double
or even triple bonds. Some elements will not comply with the octet rule – distribute the
electrons as best you can (Boron, Phosphorus are among the “exception-loving elements!)
7. Check for resonance structures. If more than is correct, draw both!
Molecular Structures
These are three-dimensional arrangements of the atoms in a molecule. The structure requires
the SHAPE and BOND ANGLE.
o
We will study FIVE shapes and SIX different bond angles.
o
VSEPR model tells us structures form to minimize the repulsions between electron pairs,
basically by keeping the pairs as far apart from each other as possible.
To determine molecular structures and bond angles:
1.
Draw the Lewis Dot Structure for the molecule
2. Determine how many electron groups are around the central atom
a.
Unshared pair = 1 group
b. Single bond = 1 group
c. Double bond = 1 group
d. Triple bond = 1 group
3. Determine the number of unshared pairs around the central atom.
4. Use the number of electron groups and the number of unshared pairs to get
the structure and bond angle of your molecule (use the chart below).
# of electron
groups
Unshared
Electron
Pairs
Any structure with 2 elements
2
3
3
4
4
4
0
0
1
0
1
2
Molecular
Geometry
Bond
Angle
Some
Examples
linear
linear
trigonal planar
bent
tetrahedral
trigonal pyramidal
bent
180
180
120
117
109.5
107
105
Cl2. or HBr
CO2
BF3
NO2CCl4
NH3
H2O
Page 42
Let’s See That in 3D!
Molecular formula
Lewis Dot Structure
Structural formula
Molecular shape
CO2
Linear
BCl3
Trigonal Planar
CH4
Tetrahedral
NH3
Trigonal Pyramidal
H2 O
Bent (scenario two)
O3
Bent (scenario one)
Page 43
Need More Help with Lewis Dot Structures?
http://web.chem.ucla.edu/~harding/lewisdots.html
We will use CO2 as our example of a simple method for drawing Lewis dot structures.
While this may not work in all cases, it should be adequate the vast majority of the time.
To see how to do negatively charged and positively charged polyatomic ions, visit the website!
Procedure for Neutral Molecules
1. Determine the total number of valance electrons.
2. Draw a “skeleton” structure of the molecule.
3. Use two valence electrons to form single bonds
between atoms in the skeleton structure.
4. Try to satisfy the octet (or duet) rule for each
atom by distributing the remaining valence
electrons as “non-bonding” electrons.
5.
If any of the octets are incomplete and more electrons remain to be shared, move one
electron per bond per atom to make a double bond.
6. Repeat steps 4 and 5 as needed until all octets are full.
7. Redraw the dots so that electrons on any given atom are in pairs wherever possible.
There are exceptions, of course! Some molecules require triple bonds. Some molecules have Lewis
Structures that refuse to obey the octet rule and have to settle for being “close enough.”
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Notes
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Notes
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Notes
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Notes
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