Unit 8: Bonding
Chemistry—2011
Day
W 3/02
Activity
Nuclear Test
Homework
Th 3/03
Bonding Basics Article (p3-8) + finish
Chernobyl video (shared folder)
Finish Article WS (p7-8)
F 3/04 – M 3/7
Casimir Pulaski Long Weekend
Play outside. PLEASE.
Tu 3/08
Return Nuclear Test. Start Bond Types
Read 12.1 in textbook.
W 3/09
Bond Types- Ionic, Covalent, Metallic Notes Bond Types WS (p17)
(p9-10)
Read section 12.2 in text.
Th 3/10
Electronegativity & Polarity Notes (p11)
Polarity WS (p18)
Bring 2 POST 1982 Pennies
F 3/11
LAB: Golden Penny
Finish Lab. Read 12.3. Play outside!
M 3/14
Lewis Dot Structure Notes (p12)
Lewis Dot WS #1 (p19-20)
Study for Quiz. Read 12.6
Tu 03/15
QUIZ: Bond Types & Polarity
Lewis Dot WS #2 (p21)
W 03/16
VSEPR Model
Molecular Geometry Notes (p13-14)
Lewis/Polar Practice WS (p22-23)
Read 12.9
Th 03/17
LAB: Building Molecules
VSEPR WS #1 (p24-25)
Finish Lab & More VSEPR
VSEPR WS #2 (p26)
Play outside!
M 03/21
Intermolecular Forces Notes (p15-16)
Intermolecular Forces Article (p28-30)
Intermolecular Forces WS (p31)
Study for Quiz
Tu 03/22
QUIZ: Lewis Dot & VSEPR Start IMF Poster
Review Worksheet (p32-34)
W 03/23
Finish IMF Poster. Review
Study for Test
Th 03/24
UNIT TEST
Late Arrival
F 03/18
Progress Reports
F 03/25
ACT Review. Then Spring Break!
Bonding Targets
Textbook
Section
(12.1)
1.
I can define ionic, covalent, and metallic bond types.
(supp.)
2.
I can describe the following properties as they relate to ionic, covalent, and
metallic bonding: melting point, boiling point, solubility, conductivity, and bond
strength.
(12.2.)
3.
I can use the periodic table to predict the bonding type and the polarity of the
bond between two atoms.
(12.3, 12.6)
4.
I can draw the Lewis dot structure of covalently bonded molecules and
polyatomic ions and predict whether the molecule is polar or nonpolar.
(12.9)
5.
I can use VSEPR theory to predict the shape of simple molecules (trigonal
pyramidal, tetrahedral, bent, linear, trigonal planar). Also, I can make Mrs.
Humes proud of me by telling her what VSEPR stands for.
(12.6)
6.
I can predict whether a bond or molecule is polar or nonpolar by comparing
electronegativity values.
(12.6)
7.
I can use the difference in electronegativity to determine whether a bond is ionic,
polar covalent or nonpolar covalent.
(14.3)
8.
I can differentiate between the following intermolecular forces (IMFs):
dipole/dipole, hydrogen bonding and London dispersion forces.
2
Bonding Basics
We just finished the nuclear unit in which the protons and neutrons were the most important subatomic
particles. We now return to chemical reactions where the electron reigns supreme. You must first learn
why atoms bond together. We will use a concept called "Happy Atoms." The idea behind Happy Atoms
is that energy levels (also called atomic shells) want to be filled with electrons. Some atoms have too
many electrons. These atoms tend to be metals and want to give up their electrons. Other atoms are
really close to having a full shell of electrons. These non-metal atoms go around looking for other atoms
that want to give up an electrons.
Let's take a look at some examples of metal atoms. Notice how there are just a few electrons in their
outer shell.
In chemically unreactive atoms, the first shell is filled with 2 electrons, the second is filled with 8
electrons, and the third is filled with 8. You can see that sodium (Na) and magnesium (Mg) have a
couple of extra electrons in their outermost shell. In order to be happy atoms, they have two possibilities.
(1) They can try to get eight electrons to fill up their third shell. Or (2) they can give up a few electrons
and have a filled second shell. For them it's easier to give up a few electrons.
Other atoms are interested in gaining a few extra electrons. Look at these nonmetals. Notice how the
outer shell almost has 8 electrons.
Each of those atoms is looking for 1-2 electrons to make a filled shell. They each have one filled shell
with two electrons but their second shell doesn’t have eight. There are a couple of ways they can get the
electrons. (1) They can share electrons, making a covalent bond. Or (2) they can transfer electrons to
form ions. The oppositely charged ions will attract one another and make an ionic bond.
So if we have a sodium (Na) atom that has an extra electron and a fluorine (F) atom that is looking for
one, an ionic bond can be formed.
Na
F
(transfer e-) Na+
ClThey wind up working together and both wind up happy! Sodium (Na) gives up its extra electron. The
sodium (Na) has a full second shell and the fluorine (F) has a full second shell. Two happy atoms! That's
one way things are able to bond together. The two elements have created an ionic bond.
3
What is Electronegativity?
Definition- Electronegativity is a measure of the tendency of an atom to attract a bonding pair of
electrons to itself.
Fluorine (the most electronegative element) is assigned an electronegativity value of 4.0, and values
range down to cesium (Cs) and francium (Fr) that are the least electronegative at 0.7.
Patterns of Electronegativity in the Periodic Table
The most electronegative element is fluorine. If you remember that fact, everything becomes easy,
because electronegativity must always increase towards fluorine in the Periodic Table.
Note: Ignore the noble gases (group 18). This is because they usually do not form bonds - and if they
don't form bonds, they can't have an electronegativity value.
What Happens if Two Atoms of Equal Electronegativity Bond Together?
Consider a bond between two atoms, A and B with dots representing electrons and the line representing
the region between the 2 nuclei of atoms A and B.
(Note: It's important to realize the electrons are actually moving around all the time within that orbital.)
If the atoms are equally electronegative, both have the same tendency to attract the bonding pair of
electrons, and so it will be found on average half way between the two atoms. To get a bond like this, A
and B would have to be the same atom. For example, you will find this sort of bond in H2 or Cl2
molecules.
This sort of bond could be thought of as being a "pure" covalent bond - where the electrons are shared
evenly between the two atoms. We call this pure covalent a nonpolar covalent bond.
4
What Happens if B is Slightly More Electronegative than A?
B will attract the electron pair rather more than A does:
That means that the B end of the bond has more than its fair share of electrons and becomes slightly
negative. At the same time, the A end (rather short of electrons) becomes slightly positive. In the
diagram, - means "slightly negative" so + means "slightly positive."
Defining Polar Bonds
This is described as a polar bond. A polar bond is a covalent bond in which there is a separation of
charge between one end and the other. In other words, it’s when one end is slightly positive and the other
slightly negative. Examples include most covalent bonds between two different nonmetals. The
hydrogen-chlorine bond in HCl or the hydrogen-oxygen bonds in water are typical polar bonds.
A nonpolar bond is a covalent bond in which there is no separation of charge between one end and the
other. An example is
What Happens if B is a lot More Electronegative than A?
In this case, the electron pair is dragged right over to B's end of the bond. A has lost control of its
electron, and B has complete control over both electrons. Ionic bonds have been formed.
Summary
* No electronegativity difference between two atoms leads to a pure nonpolar covalent bond.
* A small electronegativity difference leads to a polar covalent bond.
* A large electronegativity difference leads to an ionic bond (which is always polar).
5
Polar Bonds: Polar Molecules and Nonpolar Molecules
In any 2-atom molecule like HCl, if the bond is polar, the whole molecule is too. What about more
complicated molecules that contain 3 or more atoms? Chlorine has a greater electronegativity value than
carbon (how did I know that???), so in each of the four carbon to chlorine bonds (C- Cl), the Cl end
carries a partial negative charge, abbreviated -.
Below is a drawing of the molecule carbon tetrachloride, CCl4. Each individual bond is polar.
The molecule as a whole, however, is not polar; it is nonpolar. The entire molecule doesn't have an end
(or a side) that is more negative than any other end. The whole of the outside of the molecule is
somewhat negative, but there is no overall separation of charge from top to bottom, or from left to right.
By contrast, CHCl3 is a polar molecule.
The hydrogen at the top of the molecule is less electronegative than carbon and so is slightly positive.
This means that the molecule now has a slightly positive "top" and a slightly negative "bottom", and so
is overall a polar molecule.
1.
“Bonding Basics,” Andrew Rader Studios 2007. Rader’s Chem4Kids.
http://www.chem4kids.com/files/atom_bonds.html.
2.
Clark, Jim. “chemguide.” 2008. http://www.chemguide.co.uk/index.html#top.
6
BONDING ARTICLE
Name_______________________________
1. How does an atom become a “Happy Atom?”
2. Do atoms in a covalent bond share or transfer electrons from each other?
3. On the periodic table above, draw arrows to show how the electronegativity values increase.
4. What are the electronegativity values for the noble gases?
5. What do the dots represent in this diagram?
6. What happens in a “pure” covalent bond?
7. Give an example of a “pure” covalent bond.
8. What does + represent in this diagram? What does it mean?
9. What does it mean to have a polar bond?
10. What does it mean to have a nonpolar bond?
7
11. Does water have polar or nonpolar bonds?
12. How do ionic bonds form?
13. Determine whether the molecule (or compound) is ionic, polar covalent, or nonpolar covalent.
a.
N—N
b.
H—F
c.
C—O
d.
Na—Cl
e.
Mg—O
f.
Cl—C—Cl
g.
H
|
H—C—H
|
H
h.
F—N—F
|
F
i.
S—F—S
j.
Cl
|
Cl—C—Cl
|
Cl
k.
O—S—O
|
O
l.
K—I
14. From now on we wil NEVER use dashes to represent ionic bonds like we just did in 13e and 13l.
From now on a dash will always indicate 2 shared electrons in a covalent bond.
8
Name__________________________
Bonding Unit
Types of Bonds
Type 1: ___________________________________
- Definition: ________________________________________________________________________
________________________________________________________________________
- Involves _________________________________
- This is how we get molecules!
- _________________ bond- shares ___________________________ of electrons
- _________________ bond- shares ___________________________ of electrons
- _________________ bond- shares ___________________________ of electrons
- Share electrons to be more stable!
- _______________________ Rule- atoms will gain, lose, or share electrons to have
________________________________________________________________________
- same electron arrangement (________________________) as the closest
noble gas (very stable).
EXCEPTIONS to the Octet Rule:
Examples:
Properties:
9
Type 2: ___________________________________
- Definition: ________________________________________________________________________
_______________________________________________________________________
- Involves a ________________________________________________________
- These are called ionic _________________________
- Static electrical attraction creates the bond (+ and – attract)!
- ____________________ will _____________________ their electrons
- They become ______________________
- ____________________ will _____________________ their electrons
- They become ______________________
Examples:
Properties:
Type 3: ___________________________________
- Definition: ________________________________________________________________________
________________________________________________________________________
- Involves _________________________________
- Metal Atoms are more stable when ____________________________________________
________________________________________________________________________
- This is called the “Sea of Electrons” model
- The positive ions remain positive and the loosely-held valence electrons move all around.
Properties:
10
Polarity Notes
Name_________________________
1. Electronegativity (review from periodic trends)
Definition2. Electronegativity Trend on the Periodic Table
3.
Li
F
In the molecule LiF, which element is going to win the tug-of-war for more electrons?
-
4.
Fluorine has a ______________ electronegativity
Lithium has a _______________ electronegativity
Li
5.
F
DipoleNonpolar bondPolar bond-
6.
The _______________ the electronegativity difference the more _______________ the bond.
_______________________: Nonpolar Covalent
_______________________: Polar Covalent
_______________________: Ionic (always _____________)
Is it polar or nonpolar? (Look for a big difference in electronegativity)
11
Cs and Br
As and P
Ca and S
H and O
12
Lewis Dot Notes
Name______________________________
Valence Electrons

The electrons on the _____________________________________________ of an atom.
o The only electrons that ______________________________________________.
o Look on your periodic table and find the pattern!
EX: Lithium Atom
Try some!
C
F-1
S
Drawing Covalent Compounds

Electrons are being SHARED between the atoms.
How to Draw Covalent Bonds
1. Find the number of bonds using the following formula:
b= t-v
2
OR
# of bonds = (w-h)/2
b=
t=
w=
v=
h=
2. Determine the central atom.

An atom wants to be symmetrical

Carbon usually a central atom

Hydrogen can never be a central atom
3. Does the central atom follow the octet rule?
If: YES!
If: NO!
13
- Then
you are all done
VSEPR/Molecular Geometry Notes
- Then you need to draw lone pairs of electrons
Name______________________________
* Valence Shell Electron Pair Repulsion Theory (_______________)
Concept - you can predict the 3D arrangement of the atoms in a molecule based upon
____________________________________________________________________________________
____________________________________________________________________________________
* VSEPR Theory Assumptions
1.
Atoms in a molecule are bound together by ___________________________________.
These are called ___________________________________.
2.
_____________________________________ of bonding pairs of electrons may bind any two
atoms together (double & triple bonding).
3.
Some atoms in a molecule may also possess pairs of electrons _______________________
_______________. These are called ______________________ or
____________________________ pairs.
4.
The bonding pairs and lone pairs around an atom in a molecule are put in positions where their
interactions are _________________________.
5.
Lone pairs occupy _______________________________ than bonding electron pairs.
6.
_____________________ bonds occupy more space than ___________________ bonds.
* VSEPR Theory Predictions
How to predict the shape:
1.
Draw a ______________________________.
2.
Determine the number of valence electrons on _____________________________________.
3.
Determine the number of __________________________ of electrons on the central atom.
4.
Look up the shape on the table (or from memory).
* Predict Polarity by the Shapes:
Is the molecule polar?
If the shape is…
It is always…
If the atoms around the
central atom are
DIFFERENT...
If the atoms around the
central atom are THE
SAME...
14
15
16
Intermolecular Forces
Name___________________
Intermolecular
- _____________ the molecule or between the molecules
1)
2)
3)
(Think international)
Intramolecular
- ______________ the molecule
-
(Think intramural sports)
London Dispersion Forces
-Usually occur between __________________________________________________________
- Temporary dipoles exists when electrons are ________________________________________
- These dipoles induce dipoles on __________________________________________________
Examples:
Look at the boiling points of the noble gases
He: -269 °C
Ne: -246 °C
Ar: -186 °C
Kr: -153 °C
Xe: -107 °C
Rn: -67 °C
- Higher boiling points will have a _______
___________________________________
- Boiling points increase as the number of
___________________________________
Dipole-Dipole Forces
- Occur between _______________________________ that have permanent dipoles.
- Because opposite charges attract, the _____ will attract the ______ of another molecule.
- Allow moleculare substance to exist as solids and liquids.
Examples:
Compare the boiling points of a
polar molecule (SO2) vs. non-polar (CO2).
SO2: -10 °C
CO2: -78 °C
- Higher the boiling point more ___________
_____________________________________
- The more polar the molecule, the ________
____________________________________
17
18
Hydrogen bonding
- Although this is called “bonding” it is an __________________________________
- Relates to molecules that contain _______________________________________
(i.e. molecules with F-H, O-H, N-H)
- Hydrogen bonding is a ___________________________________ but it is special.
- Molecules with these elements _________________________________________ predicted by dipoledipole forces so another force must exist.
Examples
H2O: 100 °C
HF: 30 °C
NH3: -30 °C
Versus
CH4: -190 °C
SiH4: -110 °C
PH3: -90 °C
19
Bonding Types
Name_________________________
Complete the table below regarding Ionic Bonds, Covalent Bonds and Metallic Bonds.
Ionic
Covalent
Metallic
Definition
Composed of
Conducts
Electricity?
Solid Form
Y or N
Solid Form
Y or N
Solid Form
Y or N
Liquid (molten) Form
Y or N
Aqueous Solution
Y or N
Liquid (molten) Form
Y or N
Aqueous Solution
Y or N
Solubility in H2O
Melting Points
Boiling Points
20
Polarity
Name__________________________
Fill in the Blank: Use the words in the word bank to complete each sentence.
Words can be used more than once!
dipole
electronegativity
nonpolar
polar
bond
covalent
ionic
1. A chemical ____________________ represents the force that holds together groups of two or more
atoms and allows them to function as a unit.
2. A(n) ____________________ compound results when a metallic element reacts with a nonmetallic
elements.
3. When electrons in a molecule are shared between atoms, either evenly or unevenly, a(n)
____________________ bond is said to exist.
4. The type of bonding that exists in a Cl2 molecule is ____________________.
5. The type of bonding that exists in an HCl molecule is ____________________.
6. The relative ability of an atom in a molecule to attract electrons to itself is called the atom’s
____________________.
7. When there is no difference in electronegativity between atoms in a molecule, it has a
____________________ bond.
8. When there is a large difference in electronegativity between atoms in a molecule, it has a
____________________ bond.
9. The separation of charge in a chemical bond is called a(n) ____________________.
Answer the following questions.
10. For each of the following sets of elements arrange the elements in order of increasing
electronegativity.
a. Li, F, C
b. I, Cl, F
c. Li, Rb, Cs
d. Rb, Sr, I
e. Ca, Mg, Sr
f. Br, Ca, K
11. Circle which bonds are ionic.
a. K—N
b. Cs—O
c. C—N
d. O—F
12. Indicate whether each of the following bonds would be ionic, nonpolar covalent or polar covalent.
a. K—Cl
b. Br—Cl
c. Cl—Cl
13. Which of the following molecules contain polar covalent bonds?
a. phosphorus, P4
b. oxygen, O2
c. ozone, O3
e. sulfur, S8
f. fluorine, F2
g. iodine monochloride, ICl
d. hydrogen fluoride, HF
h. hydrogen bromide, HBr
14. Circle which is the more polar bond in each of the following pairs.
a. H—F or H—Cl
b. H—Cl or H—I
c. H—Br or H—Cl
15. For each of the following bonds, draw - or + to indicate the direction of the bond dipole.
a. C—F
b. Si—C
c. C—O
d. B—C
21
LEWIS DOT STRUCTURES #1
Name____________________________
1. Draw the Lewis Dot Structure for these single elements.
Be
2. H2S
5. (PCl4)+1
8. BF3
S
As
Ar
3. (ClO4)-1
6. SeO2
9. BeCl2
C
Rb
4. SbH3
7. (NH4)+1
10. HCl
22
11. SO3
14. (IO3)-1
17. (SiO4)-4
12. (PO4)-3
13. C3H8
15. SCl2
16. I2
18. (FCl2)+1
19. HCN
23
Lewis Dot Structures #2
Name__________________________________
Draw the Lewis Dot Structure for the following molecules.
Cl2
CCl4
Br2
H2O
I2
HBr
CH4
H2S
C2H6
PH3
HCl
ClF
CO2
C2H2
H2CO
HCN
C2H4
N2
24
Skill Practice 24
Name: ______________________________
Date: _______________
Hour: _____
1. Draw all of the resonance structures for CO32-.
2. Concerning the structures you drew in question one, what is the bond order for the C—O bonds?
3. Draw the structure for CO2.
4. Comparing the structures you drew for questions 1 and 3, which C—O bonds are the longest:
those in CO32- or those in CO2? Which are the strongest (hardest to break)? Explain your
answers.
5. Fill in the blanks:
A) In general, the stronger the bonds, the ______________ the bonds.
longer or shorter
B) In general, double bonds are _______________ than triple bonds.
longer or shorter
C) Bonds with a low bond energy are _______________ than bonds with high bond energy.
weaker or stronger
25
Skill Practice 26
Name: ______________________________
Date: _______________
Hour: _____
1. What does it mean to say that a bond is polar?
2. Label each of the following bonds as ionic (I), polar covalent (PC) or nonpolar covalent (NC).
_____ Na—Cl
_____ N—O
_____ F—F
_____ S—O
_____ H—C
_____ P—S
_____ Mg—F
_____ P—O
_____ Br—N
3. For each of the sets of bonds, rank them in order from most polar to least polar.
A) F—F, S—O, H—C, P—S
B) H—N, H—O, H—F, H—Cl
C) C—H, C—O, N—O, S—C
D) As—S, P—N, N—N, Cl—C
E) H—F, H—O, Se—Br, Si—Cl
26
VSEPR #1
Name_____________________________________
Complete the chart by drawing the LewisDot Structure and predicting the shape and polarity.
Formula
Lewis Dot
Shape
Polar or
(Name & Drawing)
Nonpolar
H2
N2
H2O
PH3
BH3
NH3
CH4
CO
(NH4)+1
27
28
Formula
Lewis Dot
Shape
(Name & Drawing)
Polar or
Nonpolar
(NO3)-1
(SO4)-2
(PO4)-3
N2O
BF3
H3O+1
COS
HCN
H2O2
29
VSEPR #2
Name________________________________
Draw the Lewis Dot structure, write the name of the shape, and circle whether it’s polar or non-polar.
1. BeO
2. NF3
Shape:
Polar or Non-polar
3. SO3-2
Shape:
Polar or Non-polar
4. HOCl
Shape:
Polar or Non-polar
5. PH3
Shape:
Polar or Non-polar
6. PH4+1
Shape:
Polar or Non-polar
7. CaH2
Shape:
Polar or Non-polar
8. CO3-2
Shape:
Polar or Non-polar
9. NO3-1
Shape:
Polar or Non-polar
10. PO4-3
Shape:
Polar or Non-polar
11. H2SO4
Shape:
Polar or Non-polar
12. AlBr3
Shape:
Polar or Non-polar
Shape:
Polar or Non-polar
Intermolecular Forces Poster
30
Box 1: Title this box “Hydrogen Bromide is a Polar Molecule”. Draw the Lewis Dot Structure for
hydrogen bromide. Label the partial positive, + and partial negative, - ends in your drawing.
Box 2: Title this box “3 Polar Molecules Interact”. Draw 3 hydrogen bromide molecules next to one
another in any logical orientation. Label the partial positive, + and partial negative, - ends in your
drawing. At the bottom of box 2 identify the type of intermolecular attraction between these hydrogen
bromide molecules.
Box 3: Title this box “Water is a Polar Molecule”.Draw the Lewis Dot Structure for water. Label the
partial positive, + and partial negative, - ends in your drawing.
Box 4: Draw 3 water molecules next to one another in any logical orientation. Label the partial positive,
+ and partial negative, - ends in your drawing. At the bottom of box 4 identify the type of
intermolecular attraction between these water molecules.
Box 5: Title this box “Like Dissolves Like.” Draw 3 water molecules and 1 dissolved sodium chloride
compound and one dissolved ammonia, NH3 molecule. Make sure to consider the interactions between
the ions and the partial positive, + and partial negative, - ends of the molecules.
Box 6: Title this box “Intramolecular Forces Are Always Stronger than Intermolecular Forces.” Define
the 3 types of intramolecular forces from Target 1. Identify and define the 4 intermolecular forces from
Target 8. Hint: One of these intermolecular forces is called “ion-dipole.”
31
Chemistry Intermolecular Forces and Water
Attractions Between Polar Molecules:
A molecule is a substance that has several atoms bonded together. Water is a molecule because it has
two hydrogens and one oxygen, and sugar is a molecule because it has lots of carbons, hydrogens and
oxygens. Sometimes, attractions can happen between separate molecules that hold the molecules
together. These bonds between molecules are called van der Waals forces. This attraction is just barely
enough to keep the molecules together.
This is how it works: As you know a particle (atom) has a positive nucleus and many negative pieces
that travel around the nucleus (electrons). Each molecule has a bunch of these particles (atoms) in it. So
when you get a bunch of MOLECULES together, what do you have? Many, many, many particles
(atoms) that have many positive nuclei and negative electrons.
So what does that have to do with attraction?
The many positive portions of one molecule can become attracted to the many negative portions
(electrons) of another molecule. This is just enough to keep the two together. A new compound is NOT
formed, but sometimes these attractions can give some pretty cool results. These are the different kinds
of van der Waals forces:
Hydrogen Bonding: This is the strongest attractive force between separate molecules. This happens
when HYDROGEN is bonded to a small but strong element. Some examples of small but strong
elements are Fluorine, Oxygen, and Nitrogen. This is the force that is so important in water!
So what do hydrogen bonds do anyway? Well, for one thing, they cause the compound ot have unusually
high boiling and freezing points when compared to molecules of similar size that do not have hydrogen
bonding.
32
Dipoles: The dipoles occur when particles with different strengths bond to form a new substance
(molecule) that has some weak particles (atoms) and some strong particles around the center. The new
substance has a positive and negative end (dipoles).
DIPOLE EXAMPLE:
CH3Cl
electronegativities:
C: 2.5
Cl: 3.0
H: 2.1
Arrows indicate direction
for dipole, which make
slightly positive and slightly
negative (+, -) ends in this
example.
As a general rule, substances that have three particles around the center are ALWAYS dipoles, and
substances with two particles can be dipoles if the central atom is from group 15, 16, or 17 on the
Periodic Table.
Attractions Between Nonpolar Molecules:
London dispersion forces are the weakest of all of the attractions. When there is a substance formed
between particles that are equally as strong as each other the substance is called nonpolar. There is NO
positive or negative end to these molecules. So how can they be attracted to each other at all? Well,
every once in a while, for just a second, the negative parts (electrons) of a molecule pile to one side
(after all the electrons do move around). When this happens the negative parts (electrons) from a
neighbor molecule get repelled and move to one side, and so on, and so on... This causes temporary
dipoles, so the molecules are just barely attracted to each other during tiny intervals of time. As the
number of negative particles in the nonpolar substance increase, the attraction between them increases.
As a rule, substance that only have one element are nonpolar (HONClBrIF), substances with two
different particles may be nonpolar if their central particle is from group 2, 13, or 14 on the Periodic
Table.
33
The Wonderful Qualities of Water!!!
Water is a covalent compound with hydrogen bonding! Since oxygen is “ strong” and hydrogen is
“weak” the oxygen pulls the electrons closer to itself. Since the water molecule is shaped the way it is,
the slight charges do not cancel each other out and water is said to be polar (the side with oxygen is
slightly negative and the side with the hydrogens in slightly positive).
What does this mean?
Well, lots of things, but we will just look at a few. For one thing, water has a high heat of vaporization
and boiling point (it has to be really hot for water to boil or evaporate). This is because the negative side
on one “piece” of water is attracted to the positive side of another. This causes all of the water to want to
stay together, so it takes a LOT of energy (heat) to pull them apart. Aren’t we lucky this is true because
otherwise water just might spend more time in the air than on the ground during those HOT summers!
One thing to remember is that when water evaporates the hydrogens and oxygen do NOT break apart.
Remember, one piece of water is a group of one oxygen and two hydrogens. When water evaporates the
GROUPS separate from each other to make lots of separate pieces of water. We know this because if
you put your hand over a pot of boiling water you can feel the steam, and your hand gets damp, so you
still have water but the pieces are just separated. If all of the hydrogens and oxygens split up, we
wouldn’t have water anymore!
Water also has high surface tension. The reason for this is that the top layers of the water pieces are only
bonded to the other pieces on one side, so they are drawn into the body of the water. This makes the
surface of the water quite strong. Another great thing is that ice has a low density (the mass of ice is
small compared to the amount of space it takes up). What happens is that as the temperature goes down
there isn’t enough energy to break up the attractions between the pieces of water, so even though the
liquid is turning into a solid the density does not increase. This is a BIG deal because if we did not have
this, the ice on lakes in the winter would sink to the bottom and all of the fish would die!
34
35
REVIEW WORKSHEET
BONDING UNIT
Name______________________________
1. What type of attraction (force) exists between water molecules?
2. Metals tend to ________________ electrons to form positive ions. Nonmetals tend to _____________
electrons to form negative ions.
3. Covalent compounds are usually made up of 2 ______________________; ionic compounds are
usually made up of a ________________ and a ___________________.
4. Generally speaking, electronegativity ____________________ as you go up a column and
____________________ as you go from left to right.
5. The greater the difference between 2 atoms in a bond, the _____________ the percentage of ionic
character in the bond.
6. A bond that has an electronegativity difference of zero is said to be ________________ and
_______________. (polar or nonpolar; ionic or covalent)
7. A(n) _______________ bond is a bond formed when 2 bonding atoms share electrons.
8. A molecule of hydrogen shares _____ electrons in order to form a single covalent bond.
9. A molecule of oxygen shares 4 electrons forming a ________________ covalent bond.
10. A molecule of nitrogen shares ______ electrons forming a ________________ covalent bond.
11. Ammonia, NH3, has ______ shared pair(s) of electrons and ______ pair(s) of unshared pairs.
12. Water has _____ shared pair(s) of electrons and _____ pair(s) of unshared electrons.
13. _____________ compounds are good conductors of electricity in their aqueous and molten states.
14. A polar bond is a bond made up of 2 atoms that have much different
___________________________ values.
15. Water is a __________ (polar or nonpolar) molecule and has a __________ molecular shape.
16. Oxygen is a _______________(polar or nonpolar) molecule and has a ______________ molecular
shape.
17. Phosphorus trichloride is a ____________ (polar or nonpolar) molecule and has a
__________________ molecular shape.
18. Carbon tetrachloride is a ______________ (polar or nonpolar) molecule and has a
___________________ molecular shape.
19. ____________________ occurs when 2 or more equally valid electron dot structures can be written
for a molecule.
20. The VSEPR theory states that because electron pairs ____________, molecules adjust their shapes
so that the valence electron pairs are as far apart as possible.
21. A molecule that has 2 poles is commonly called a ___________.
22. The attraction between 2 polar molecules is called ______________________.
23. The attraction between 2 nonpolar molecules is called _______________________.
24. __________________ bonding is a particularly strong type of dipole-dipole attraction in which the
molecules involved contain hydrogen and another atom that has a large electronegativity.
25. When 2 dipoles in a molecule cancel each other, a ____________ (polar or nonpolar) bond is
formed.
26. When dealing with dissolving, use the statement “Like dissolves like.” This means polar molecules
dissolve polar molecules and nonpolar molecules dissolve nonpolar molecules. Name 2 substances that
you think would dissolve easily in water.
__________________________
___________________________
27. Which of the following bonds are polar?
C-H
I-I
H-O
S-F
Cl-Cl
28. Which of the following bonds is/are held together by bonds where electrons are being shared?
36
NaCl
KF
CO2
H2O
37
29. Which of the following are nonpolar molecules?
H2O
CH4
NH3
PCl3
N2
30. Which of the following contains 3 bonding pairs of electrons (triple bond)?
H2O
CH4
NH3
PCl3
N2
31. Which of the following contains at least one pair of unshared electrons on the central atom?
H2O
CH4
NH3
PCl3
N2
32. Which of the following has the greatest attraction for electrons when involved in a chemical bond?
Li
Be
B
C
N
O
F
33. Which of the following would you expect to be the most soluble in water? (“like dissolves like”)
carbon dioxide
carbon tetrafluoride
oxygen gas
ammonia (NH3)
34. Which of the following compounds would have the most covalent characteristics?
CaO
NO2
LiBr
BaCl2
35. Which of the following are linear molecules?
CO2
SiCl4
N2
PH3
36. Which of the following would have resonance structures?
NO3-1
CO3-2
O3
SO4-2
CO
37. For the answers you circled in the previous question, draw the resonance structures!
38. What is hydrogen bonding? Name three elements that can combine with hydrogen to result in
hydrogen bonding.
39. Name the type of bond for each of the following:
KCl
Na2O
H2SO4
O2
H2O
Na (many Na bonded)
40. Which would have the highest boiling point?
NaCl
Br2
H2O
41. Which of the following would dissolve in CH4? (“like dissolves like”)
H2O
NH3
H2SO4
38
Draw the Lewis Dot, predict the shape and predict the polarity.
Formula
Lewis Structure
Shape
Polar or Nonpolar
CS2
PO43-
BI3
PCl3
SCl2
HCN
CHBr3
NHCl2
TeH2
CBr4
CO
NCl3
39
Unit 3: Bonding Target Review
Target 1: I can define ionic, covalent, and metallic bond types (12.1)
Target 3: I can use the periodic table to predict the bonding type
between two atoms (12.2).
1. Define the following terms, make sure to include electron behavior.
a. Ionic bond:
b. Covalent bond:
c. Metallic bond:
2. Fill in the blank
a. Covalent compounds are usually made up of 2 ____; ionic compounds are usually made
up of a ___ and ___.
b. Metals tend to __________electrons to form positive ions. Nonmetals tend to___
electrons to form negative ions.
c. A(n) ___ bond is formed when 2 bonding atoms share electrons.
d. A(n) ___ bond is formed when a cation is attracted to an anion.
e. Ionic compounds generally melt at ____ temperatures than covalent compounds.
f. ____ compounds are good conductors of electricity in their aqueous and molten states.
g. In this type of bond, electrons from a sea of electrons ___________________.
3. Identify the bond type
a. KCl
b. Na2O,
c. H2SO4
d. O2
e. H2O
f. Na (many Na atoms bonded together)
4. Which of the following compounds would have the most covalent characteristics?
a. CaO
b. NO2
c. LiBr
d. BaCl2
Target 2: I can describe the following properties as they relate to ionic, covalent, and
metallic bonding: melting point, boiling point, solubility, conductivity, and bond
strength (notes).
5. Fill in the table
Bond type
Ionic
Covalent
Metallic
Melting point
Boiling Point
Solubility in water
Conductivity
Bond Strength
Polarity
40
6. What makes metals such good conductors of electricity?
a. They have delocalized electrons.
b. The metallic bonds are weak enough to be temporarily broken – this allows the atoms of
the metal to move about freely.
c. The ductility of the cations combine with the malleability of the anions in order to form
electrons that move from lattice point to lattice point. The lattice point movement
(including movement of the neutral ions) allow for a current to flow.
d. Metals are generally more dense than nonmetals. This extreme density allows for free
electron movement.
e. All of the above!!!
7. Ionic compounds generally melt at ________________ temperatures than covalent compounds.
8. ______________ compounds are good conductors of electricity in their aqueous and molten
states.
Target 4: I can draw the Lewis dot structure of covalently bonded
molecules and polyatomic ions 12.3 & 12.6).
9. What is the formula for finding the number of bonds within a molecule?
10. Draw a LDS structure of the following atoms
a. C
c. S
b. Na
d. Xe
11. Draw a LDS structure for the following diatomic atoms
a. H2
b. O2
c. F2
d. N2
Target 5: I can use the VSEPR theory to predict the shape of simple
molecules (trigonal pyramidal, tetrahedral, bent, linear, trigonal
planar) (12.9).
12. Which shapes have at least one lone pair of electrons on the central atom?
41
Target 6: I can predict whether a bond or molecule is polar or nonpolar
by comparing the electronegativities (12.6).
Target 7: I can use the difference in electronegativities to determine
whether a bond is ionic, polar covalent, or nonpolar covalent (12.6).
13. Which VSEPR shapes are always polar? Why are they always polar?
14. Water is a ______________ (polar or nonpolar) molecule and has a _____________ molecular
shape.
15. Oxygen is a ______________ (polar or nonpolar) molecule and has a _____________ molecular
shape.
16. Phosphorus trichloride is a ______________ (polar or nonpolar) molecule and has a
_____________ molecular shape.
17. Carbon tetrachloride is a ______________ (polar or nonpolar) molecule and has a
_____________ molecular shape.
18. ________________ occurs when 2 or more equally valid electron dot structures can be written
for a molecule.
19. Generally speaking, electronegativity _________________ from top to bottom within a family
and _________________ from left to right across a period.
20. The greater the difference in electronegativity between 2 atoms in a bond, the _____________
the percentage of ionic character in the bond.
21. A bond that has an electronegativity difference of zero is said to be __________ and
_____________. (polar or nonpolar; ionic or covalent)
22. A polar bond is a bond made up of 2 atoms that have much different ______________ values.
23. Which of the following has the greatest attraction for electrons when involved in a chemical
bond?
a. Li
b. Be
c. B
d. C
e. N
f. O
g. F
24. Which of the following bonds are polar bonds?
a. C---H
b. I---I
c. H---O
d. S---F
e. Cl---Cl
25. When 2 polar bonds in a molecule cancel each other, a _____________ (polar or nonpolar)
molecule is formed.
42
Target 4: I can draw the Lewis dot structure of covalently bonded
molecules and polyatomic ions 12.3 & 12.6).
Target 5: I can use the VSEPR theory to predict the shape of simple
molecules (trigonal pyramidal, tetrahedral, bent, linear, trigonal
planar) (12.9).
Target 6: I can predict whether a bond or molecule is polar or nonpolar
by comparing the electronegativities (12.6).
Target 7: I can use the difference in electronegativities to determine
whether a bond is ionic, polar covalent, or nonpolar covalent (12.6).
26. For each of the following . . .
A- Find the number of Bonds
B- draw the dot structures
C- predict the shape
D- predict whether the molecule will be polar or nonpolar
BI3
NBr3
SiCl4
SiS2
Cl2O
CCl2Br2
NO3-
SO42-
P2
27. Which of the following would show resonance?
a. NO3-1
b. CO3-2
c. O3
d. SO4-2
e. CO
28.
29.
30.
43
31. For the answers that you circled in #1, draw the resonance structures!
Target 8: I can differentiate the following intermolecular forces: dipole,
hydrogen bonding, London dispersion forces (14.3).
32. The attractions between 2 separate polar molecules are called __________________ attractions.
33. The attractions between 2 separate nonpolar molecules are called ___________ ___________
forces.
34. ________________ bonding is a particularly strong type of dipole-dipole attraction in which the
molecules involved contain hydrogen and another atom that has a large electronegativity.
35. What type of IM forces exist between carbon dioxide molecules?
I.
dispersion forces
II.
dipole-dipole interactions
III.
hydrogen bonding
a. I only
b. II only
c. I and II only
d. III only
e. I, II, and III
d. III only
e. I, II, and III
36. What type of IM forces exist between ammonia molecules?
I. dispersion forces
II. dipole-dipole interactions
III. hydrogen bonding
a. I only
b. II only
c. II and III only
37. What does the term “like dissolves like” mean?
38. Name 2 substances that you think would dissolve easily in water.
39. Which of the following would you expect to be the most soluble in water?
a. carbon dioxide
b. carbon tetrafluoride
c. oxygen gas
d. ammonia
44