Monday
March 2nd 3 th
Tuesday
4 th
Wednesday
5 th
Thursday
Notes #1: Bond Types Timed Writing and Properties and
pg 3,
Notes #2: Covalent
Bonding and Lewis
Structures
(pg 5-7)
Hwk: Page 4, 7
Hwk: Finish Pages
4 & 7
Notes #3: Electrical Notes #4: Bond and Molecular
Geometry (pg 8)
Hwk: Page 9
Polarity (pg 10)
Hwk: Page 11
9 th 10 th 11 th 12 th 13 th
16 th
Hwk: Page 13
And REVIEW
Page 14 - 16
17 th
Notes #5: Molecule
Polarity, and
Intermolecular Forces
Covalent Bond and
(pg 1)
Chemthink.com
Molecular Shapes
HWK:REVIEW
Page 14 - 16
18 th
VSPER Lab
REVIEW DUE: Page
14 - 16
19 th 20 th
Friday
6 th
VSPER Lab
HWK:REVIEW
Page 14 - 16

2.
3.
4.
https://www.youtube.com/watch?v=TAVwvOUewLk&list=PLfWuBca_SOr8aVsqv27mpD7UV0aBKEOWU
&index=1
List the properties of Ionic Compounds:
1.
2.
3.
4.
Draw Example:
https://www.youtube.com/watch?v=Oagr9xMAmfY&index=3&list=PLfWuBca_SOr8aVsqv27 mpD7UV0aBKEOWU
List the properties of Ionic Compounds:
The electrons in a metallic compound are able to move throughout the entire compound. The electrons can be described as a "__________________ ____ __________________."
Draw Example:
https://www.youtube.com/watch?v=og8qXtEzrvA&index=4&list=PLfWuBca_SOr8aVsqv27mp
D7UV0aBKEOWU
List the properties of Ionic Compounds:
1.
.
Draw Example:

 Metallic bonds are formed by pooled valence electrons of metallic atoms providing the negative charges to hold positively charged metallic ions together.
This bonding structure provides for relatively high melting points and easy reshaping (bending, flattening). The delocalized electrons provide high electrical conductivity.
 Ionic bonds are formed when metallic atoms donate valence electrons to nonmetallic atoms. The resulting ions have opposite charges and attract each other into rigid lattices. This bonding structure gives high bond strength that provides brittle substances with high melting points and low conductivity. If the lattice is disrupted by being heated or dissolved in water, the ions break apart and find movement easier.
Conductivity of molten or aqueous ions is much higher than that of solids.
 Covalent bonds are formed when two nonmetallic atoms approach and share valence electrons. These are the strongest of all bonds. Covalent networks form when atoms bond each to several others, making an interlocking web of atoms. Covalent networks are very hard to disrupt, giving these substances very high melting points and low conductivity in any state. Molecules form when a few covalent bonds form between a countable number of molecules, as in CO
2
or H
2
O. While the bonds within the molecule are very strong, the molecules are so small that we commonly deal with a very large number of them. One molecule requires little energy to separate from another, so these substances have very low melting points, often below room temperature. Most liquids and gases that we are familiar with are molecular. Because molecules hold their electrons so tightly, molecules also tend to be poor conductors.
Classify the following as characteristics of a compound that contains metallic (M),
ionic (I) or covalent (C) bonding.
_____ 1. Contains a metal and a nonmetal ion
(and possibly polyatomic ions)
_____ 2. Contains 'sea' of delocalized electrons
_____3. The smallest particle is the formula unit
_____ 4. The smallest particle is the molecule
_____ 5. Particles are held together by strong
electrostatic forces
_____ 6. Particles are held together by relatively
weak intermolecular forces
_____ 7. The elements in the compound share
valence electrons
_____ 8. The elements in the compound gain and
lose valence electrons
_____ 9. Are usually soluble in water
_____ 10. Are electrolytes when dissolved in water
or molten (liquid) and conduct electricity
_____11. Have low melting and boiling points
_____ 12. Have high melting and boiling points
(2 answers)
_____ 13. Often exist as gases or vaporize easily at room temperature
_____ 14. Are hard crystalline solids at room
temperature
_____ 15. Malleable and able to be flattened
_____ 16. Are rigid and brittle
_____ 17. An example is MgCl
2
_____ 18. An example is CCl
4
_____ 19. Conduct electricity in solid state
______20. Its structure looks like this:
_____ 21. Its structure looks like this:
_____22. Its structure looks like this:

Ionic Bonding -
Generally occurs between _____________________________.
Can also occur with __________________________.
Involves _________________________ electrons, followed by electrostatic attraction.
Covalent Bonding -
Generally occurs between ________________________________________________
Involves ____________ of electrons, rather than transfer.
Octet Rule - Atoms will acquire, through sharing or transfer, the electron configuration of a
_________________. This happens in order for the atoms to gain stability.
 Most noble gases have ___ valence electrons. o _____ is the exception.
 These elements do not need a full octet: _____________.
Valence Electrons – The electrons in the highest occupied energy level. There are two ways to determine the number of valence electrons.
A.
Using the electron configuration: 1s 2 2s 2 2p 6 3s 2 3p 2 , how many valence electrons are in this element?
B.
Look at the group number to determine valence electrons.
Dot Models - The number of dots is equal to the number of ______________ electrons.
Example: Phosphorus
Examples of Ionic compounds: NaCl MgBr
2
Lewis Structures for Molecules -
Each atom in the molecule is connected by bonds.
Covalent bonds are shared paired of electrons, and they are represented with a
____________.
Pairs that are not shared are called __________________________.
A single bond is created by _______ shared pair of electrons. *Draw it: _____
A double bond is created by _______ shared pairs of electrons. *Draw it: _____
A triple bond is created by _______ shared pairs of electrons. *Draw it: _____

:
1.
Determine the number of electrons pairs (bonds) you must place in the structure
Example: CH
4
 Add up all the valence electrons for all atoms in compound.
 Count up the total number of electrons to make all atoms happy (to fulfill the octet rule…usually 2 or 8).
 Subtract the 2 totals.
 Divide by 2
 This tells you how many bonds (-) connect the central atom.
2.
Draw a skeleton of the molecule, connecting atoms in the molecule with ________________________ (1 shared pair).
 Usually the ____ atom in the molecular formula is
____________.
 __________________ is never central –
______________________________________________
 Nature likes symmetry.
3.
From the total pairs you counted in step 1, _________________ ______________.
This will determine the number of unpaired electrons pairs you have left to distribute.
4.
Distribute remaining pairs as _____________________around the atoms in the molecule.
 Hydrogen always only gets ________ of electrons.
 Each atom should have ________________________________ unless it is an exception to the octet rule.
 If there ___________________________ to fulfill the octet rule, you will need a multiple bond.
Practice:
H
2
O
Bond calculation: Lewis Structure
HCP
Bond calculation: Lewis Structure
CO
2
N
2

Remember: Exceptions to the octet rule:
 Atoms with less than an octet.
Day 2 Homework
Complete the tables below
Compound Bond
Calculation
CCl
4
Lewis Dot
Structure
Drawing
Compound
O
2
Calculation Lewis Dot
Structure
Drawing
SiH
4
H
2
O
NH
3
HCl
BeF
2
HCN
CH
3
Cl
N
2
O

Day #3: Electron and Molecular Geometry
Electron Geometry:
 The geometry based on the number of electron groups on the central atom. o It does not matter if the groups are a single bond, a multiple bond, or lone pairs. o Know the following Geometries:
Electron Geometry: Linear
Number of electron regions on the
Central Atom:
Trigonal Planar Tetrahedral
Picture:
Molecular Shape (Geometry):
• In order to predict the specific molecular shape, we assume the valence electrons
____________ each other. The molecule adopts whichever 3D geometry
________________this repulsion.
• We call this process __________________________________________________.
• Some of the names for the Molecular Geometry are the same as the Electron Geometry
• Know the following Molecular Geometries:
Molecular
Geometry
Number Of Lone
Pairs On The
Central Atom
Linear
Trigonal
Planar
Tetrahedral Pyramidal Bent
Number Of
Atoms Bonded
To The Central
Atom
Bond Angle
Picture

Day 3 Homework: Geometry and VSPER shapes
1.
Explain the VSEPR theory in your own words:
2.
Use the picture to the right to answer the next 3 questions: a.
How many electron groups are in the molecule? b.
What electron geometry will the molecule have? c.
What molecule geometry will the molecule have?
Complete the table below
Formula Total Pairs
Lewis Structure
Sketch
# lone pairs of electrons around central atom
# shared pairs of electrons around central atom
HF
AlBr
3
BF
3
CO
2
Electron
Geometry
C
2
H
4
Molecule
Geometry

In ________________ bonds, shared pairs of electrons are pulled between the nuclei of atoms sharing them. Sometimes electrons are pulled equally and sometimes they are not.
This has to do with electronegativity.
*Recall: electronegativity is the ______________________ electrons.
1.
Nonpolar covalent bonds: When electrons are shared ______________. a.
The molecule will be electrically neutral – little to no difference in electronegativity b.
Ex:
2.
Polar covalent bonds: The electrons are shared ________________due to electronegativity a.
The more electronegative atom will have a _______________ attraction for the bonded electrons and will have a slightly ______________ charge. b.
The less electronegative atom will have a slightly ______________ charge. c.
Ex: HCl – Look up the electronegativities on the chart.
H: ____________, Cl: ___________ o There are two ways to show polarity: HCl has a polar bond or
Ex: Water
Electronegativity
Difference
0.0-0.3
0.4-1.0
1.0-2.0
>2.0
Type of Bond Example
The difference in electronegativities indicates the type bond the atoms will form.
Question: What type of bond will form
between:
 N and H?
 F and F?
 Ca and O?
 Br and Cl?

Day 4: Bond Polarity Homework
1-3 True/False
1) In a polar covalent bond, the more electronegative atom has a slightly positive charge.
T F
2) In general, non-metallic elements have greater electronegativities than metallic elements.
T F
3) If the electronegativity difference between two atoms is greater than 2, they will form an ionic bond.
T F
4) What is the difference between a polar and a non-polar bond? (answer in terms of electrons)
5) If an element with an electronegativity of 0.3 bonds with an element with an electronegativity of
2.9. What type of bond is between these elements?
6) State whether the following contain polar, non-polar or ionic bonds. a) KF e) Na
2
O b) SO
2 f) O
2 c) NO
2 d) Cl
2 g) P
2
O
5
7) Draw the Lewis structure for the following compounds. Label any polar bonds and identify their shape. a) H
2
S d) CHCl
3 b) CS
2 e) OF
2 c) HBr d) CF
4

Molecule Polarity:
If a molecule has all nonpolar bonds then the molecule is __________________.
If a molecule has a polar bond then the whole molecule is usually ______________, but not always.
To determine if a molecule is polar or nonpolar, look at the central atom.
1.
If the central atom has ___ lone pairs and has ________________ of atoms attached to it, then the molecule is ________________.
2.
If the central atom has __ lone pairs but ________ atoms attached to it, the molecule is _____.
3.
If the central atom has lone pairs, the molecule is _______.
*In other words: If the molecular geometry is _____________, the bond polarities cancel, and the molecule is ___________.
Example: carbon dioxide Example: methane Example: water
In a polar molecule, one end has a positive charge and the other has a negative charge. A molecule that has poles is called a ________________________ or a ______________.
Recall: Intermolecular Forces are the ________________________________ neighboring particles.
Dispersion Forces – occurs between
__________molecules,
___________________ charges
Dipole interaction – occurs when ___________
___________ are attracted to one another
Hydrogen bonding – occurs when a __________ atom involved in an
______________________ bond is strongly attracted to an adjacent molecule. This is the
__________ intermolecular force.
Extremely polar bonds:
Hydrogen
Bonding

Day 5 Homework
1) Label the following structures. Also include the bond angles for each structure. a) Which of the above structures would be polar molecule(s) no matter what elements were bonded in it?
2) State whether the following contain polar, non-polar or ionic bonds. a) NaF d) Br
2 b) SeO
2 c) PO
2 e) Li f) N
2
2
O
3) Draw the structural formula for the following compounds. Label any polar bonds and determine whether the molecule itself is polar or non-polar. a) HOOH b) CO
2 c) H
2
0 d) BF
3
4) What are the attractive forces between adjacent water molecules called?
5) In your own words, explain how or why a molecule can have polar bonds, but the molecule is nonpolar.
6) Draw the Lewis structures for CF
4
and CH polar.
2
Cl
2
and then explain which molecule is the more

Study Guide-Covalent Bonding
1.
For each of the following elements, determine the number of valence electrons and draw the Lewis Dot
Diagram (Electron Dot Structure) for that element.
A.
sodium
B.
helium
C.
D.
oxygen neon
2.
Describe each type of bond. For each type, provide at least one example.
A.
nonpolar covalent bond
B.
polar covalent bond C.
ionic bond
3.
Complete the table below.
Type of Bond
Single
Double
Triple
4.
What is electronegativity?
Number of Electrons Number of Electron Pairs

5.
Determine if the following bonds are ionic, polar, or nonpolar using the electronegativity values in the table below.
A.
H – H D.
Na – F
Element Electronegativity
H
Br
C
2.1
2.8
2.5
B.
N – H E.
N - Br N
O
3.0
3.5
C.
C – O F.
C - F
F
Cl
I
4.0
3.0
2.5
6.
For the polar bonds in #5, which bond was the most polar?
Na
P
0.9
2.1
7.
For the POLAR bonds (and only the polar) bonds in #5, assign the positive pole(  +) and negative pole(  +) for the bond.
8.
What shape occurs if there were three atoms and one lone pair of electrons?
9.
What electron geometry does BCl
3
have?
10.
Draw the shape of water. The two sets of lone pairs of electrons cause the atoms to repel and create what shape?
If the lone pairs of electrons were bonded to atoms, what would this shape then be called?
11.
Which two molecular shapes are polar no matter what elements are involved? Explain why.
12.
In the Lewis structure for CCl
4
, how many unpaired electrons would appear on the structure?
13.
According to VSEPR theory, why do molecules adjust their shapes?
14.
Complete the large-ish content frame on the back. 

Compound
A.
H
2
B.
N
2
C.
H
2
S
D.
CO
2
E.
NCl
3
F.
NH
3
G.
PCl
3
H.
CHCl
3
I.
HCN
N-
A-
S-
U-
B-
N-
A-
S-
U-
B-
U-
B-
N-
A-
S-
U-
B-
B-
N-
A-
S-
N-
A-
S-
U-
A-
S-
U-
B-
S-
U-
B-
N-
N-
A-
S-
U-
B-
N-
A-
S-
U-
B-
N-
A-
Covalent Bonding Content Frame
Total Pairs Lewis Structure
(Drawing)
Polar or Non-Polar
Bonds?
(Label each bond with a “P” or “NP”)
(Label each bond with a “P” or “NP”)
(Label each bond with a “P” or “NP”)
(Label each bond with a “P” or “NP”)
(Label each bond with a “P” or “NP”)
(Label each bond with a “P” or “NP”)
(Label each bond with a “P” or “NP”)
(Label each bond with a “P” or “NP”)
(Label each bond with a “P” or “NP”)
Polar or Non-Polar
Molecule?