STUDENT NOTES Pre-AP Chemistry
NAME______________________________________
UNIT 6 NOTES: BONDING
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PERIOD___________
STUDENT OBJECTIVES: Your fascinating teachers would like you amazing learners to be able to…
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List the types of ions involved in ionic bonding.
Describe the basic structure of ionic crystals.
Predict the presence of ionic bonding based up the types of elements involved.
Calculate the likelihood of ionic bonding based on differences in electronegativity.
Determine the number of valence electrons in an atom.
Draw the Lewis dot structure of atoms, ions and ionic compounds.
Predict the oxidation numbers of elements in the “s” and “p” blocks of the periodic table.
Memorize the common oxidation numbers of common transition metals.
Analyze typical physical properties of ionic properties based upon the predicted structure.
Predict the presence of covalent bonding based upon the types of atoms involved in the molecule.
Calculate the bond polarity in a covalent bond using differences in electronegativity.
Memorize the names and formulas of the diatomic elements.
Illustrate the polarity of a bonding using the symbols — and + to show the electron distribution in the bond.
Draw Lewis dot structures for covalent molecules and polyatomic ions that have a central atom.
Draw Lewis dot structures for central atoms that have incomplete and expanded octets.
Explain the concept of resonance in covalent bonding and use Lewis dot structures to support your
explanation.
Use Valence Shell Electron Pair Repulsion (VSEPR) theory to predict the three dimensional shape of
covalent molecules with a central atom.
Predict the molecular polarity of a covalent molecule from its VSEPR structure.
Explain the common physical properties of covalent molecules based upon the predicted structure.
Outline the structure and function in metallic bonding.
Define the type of intermolecular forces involved in ionic and covalent substances.
Predict the type of intermolecular forces involved in ionic and covalent substances.
Compare and contrast the properties of covalent and ionic substances.
STUDENT NOTES Pre-AP Chemistry
I.
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INTRODUCTION – THE THREE TYPES OF BONDS
We have THREE types of bonds that we will discuss that vary, depending on those elusive little
electrons!
II.

Ionic bonding electrons are EXCHANGED: Metals are LOSERS and non-metals are WINNERS in the
electron world!

Covalent bonding electrons are SHARED: sometimes evenly and sometimes one element gets a little
greedy with the electrons!

Metallic bonding electrons live in a little “SEA” where they move from metal ion to metal ion,
swimming between the ions like fish!
IONIC BONDING STRUCTURE
a. General Structural Unit : Ionic structure has alternating positive and negative charges in a rigid
________________. Electrons are ___________________ from one species to another, not shared.
STUDENT NOTES Pre-AP Chemistry
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b. Prediction by Substance Type: If you see the following types of particles present in the formula of
a compound, we usually predict the compound to be ionic.
FYI…
Cation (+ ion, lost
Metal
Metal
Ammonium
Ammonium
e-‘s
) Anion (− ion, gained
Non-metal
Polyatomic
Non-metal
Polyatomic
e-‘s)
Polyatomic Ion: 2 or more
elements grouped together
that have a charge.
ex: Nitrate – NO3-1
Ammonium – NH4+1
c. Prediction by Electronegativity Calculation: Electronegativity (the pull an atom has to gain an
electron) can be used to verify whether a predicted bond is actually ionic or not. If the ΔEN is
_________________________, the bond exhibits 50% or more ionic character and definitely considered to
be ionic.
Example 6-1. Each of the following is predicted to have ionic bonds using their particles present as a
predictor. Calculate the difference in electronegativity to determine if the bonds are actually
>50% ionic.
(a) NaCl
(b) BeI2
(c) AlBr3
(d) CuCl2
NOTE: YOU DO NOT NEED TO PERFORM A BOND CALCULATION UNLESS INSTRUCTED! GO BY YOUR
PREDICTIONS!
LEWIS DOT STRUCTURES OF ELEMENTS AND IONS
d. _____________________ Electrons: Outer _____ and _____ electrons – the electrons most likely to be
involved in _________________.
Valence electrons are most helpful for predicting the __________________________ (often referred to as
________________) for elements in the s and p blocks. The oxidation number is a method used for
tracking electrons, and is equal to the ionic charge for cations (+ ions) and anions (— ions).
Below are the common oxidation numbers for the representative (main group) elements:
Group Number # Valence Electrons
Typical Oxidation Number
1
1
+1 (when ________ one e—)
2
2
+2 (when ________ two e—)
13
3
+3 (when ________ three e—)
14
4
Non-metals: -4 (when ________ four e—)
Metals: +2 when they lose their p2 electrons
+4 when they lose their s2p2 electrons
15
5
Non-metals: -3 when with metals
Metals: +3 when they lose their p3 electrons
+5 when they lose their s2P3 electrons
16
6
-2 (when ________ two e—)
17
7
-1 (when ________ one e—)
As atoms, we
will lose & gain
electrons to
achieve full
sublevels as
ions!
ALSO: The oxidation number for pure elements (elements by themselves shown without any kind of
charge) is always ________________.
Example 6-2. Notice that we have not discussed any common charges for Group 18 (Noble/Inert Gas)
elements. Why do you think this is so?
STUDENT NOTES Pre-AP Chemistry
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e. Transition Metals
The transition elements are notoriously independent! Some follow valence electron predictions
and some don’t. Sometimes, you have to use the other ion’s oxidation number in a compound to
figure out the transition element’s oxidation number… or be told what its oxidation number will be.
There are a few to memorize… I will refer to them as the “staircase”
Glad I know
my oxidation
numbers so
I can make
my
compounds
correctly!
Al
0, +3
Ag
0, +1
Zn
0, +2
Ga
0, +3
Cd
0, +2
In
0, +3
Example 6-3. Label the Periodic Table below with the number of valence electrons and the common oxidation
numbers for the representative elements, as well as for zinc, cadmium, and silver.
f.
LDS of ATOMS: To draw the LDS (Lewis Dot Structure) for atoms, simply imagine a square around
the element symbol and then place the valence electrons around the symbol. Remember Hund’s
Rule: put one in each spot, then double up!
LDS of IONS: (1) Determine how many valence electrons the ion will have! Most metals will lose
all their valence (“s” and “p”) electrons, or you will be told what the charge will be.
Non-Metals will gain enough electrons to fill their “p” sublevel.
(2) Place brackets around your Lewis Dot Structure
(3) Add the charge as a superscript on the upper right side outside of the brackets.
ELEMENT
LEWIS DOT
ION
LEWIS DOT
Tin

 Sn 

Tin(II)
(Sn2+)
[• Sn •]2+
Tin(IV)
(Sn4+)
[ Sn ]4+
NOTICE A PATTERN? Tin has two possible charges/oxidation numbers. How were these indicated in the
name?
STUDENT NOTES Pre-AP Chemistry
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Example 6-4. Draw the Lewis dot structures for the following atoms and ions.
ELEMENT
LEWIS DOT
ION
Lithium
Lithium ion
Boron
Boron ion
Oxygen
Oxide ion
Bismuth
Bismuth (+3) ion
Antimony
Antimony (+5) ion
Phosphorus
Phosphide ion
Iodine
Iodide ion
Xenon
Helium
(THIS ONE IS TRICKY…)
LEWIS DOT
Example 6-5. NOTICE A PATTERN? How did the suffix/ending of a non-metal change when it became an
ion?
STUDENT NOTES Pre-AP Chemistry
III.
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LEWIS DOT STRUCTURES OF IONIC COMPOUNDS
(1) Draw the Lewis dot structure for the positive ion with charge and brackets.
(2) Draw the Lewis dot structure for the negative ion with charge and brackets.
(3) Continue to draw positive and negative ions until the charges cancel themselves out to zero. It is proper notation
to alternate positive and negative ions.
IONIC COMPOUND
LEWS DOT
A GLIMSE OF THE FUTURE:
WRITE THE FORMULA UNIT!
magnesium chloride
MgCl2
Example 6-6. Draw the Lewis Dot Structure for the following ionic compounds:
IONIC COMPOUND
LEWS DOT
A GLIMSE OF THE
FUTURE: WRITE THE
FORMULA UNIT!
sodium chloride
calcium fluoride
aluminum oxide
LEGGETT PreAP Bonding 6-4 Ionic Properties (4:05)
IV.
https://vimeo.com/53288361
http://youtu.be/bGjlZeem9Lk
IONIC COMPOUND PROPERTIES
Ionic bonds have the following characteristics:
PROPERTY
High melting (Tm) and boiling (Tb) points – most are __________________ at
room temperature
Brittle and Cleave when struck
Many (but not all) are _______________in water
Non-conductive as __________________
Conductive in _____________and _____________ (dissolved in water) states
EXPLANATION
In order for a phase change to occur, strong
ion-ion attractions that are present must be
broken. This takes a significant amount energy!
Caused by ion-ion repulsion
When soluble, ions are attracted to the partial
charges on water
Movement of electrons require mobile ions ionic solids are rigid
Ions become mobile and allow electron
movement
STUDENT NOTES Pre-AP Chemistry
LEGGETT PreAP Bonding 6-5 BOND POLARITY (9:34)
V.
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https://vimeo.com/53288362
http://youtu.be/PUNoV6zFosY
COVALENT BONDING STRUCTURE
a. General Structural Unit: Sometimes the more electronegative element is not ____________enough to
actually take another atom’s electron away. The electrons of both atoms are placed between the
two and are _____________, which forms a covalent bond. Covalent bonds are found in molecular
compounds and within (not between) polyatomic ions.
The sharing that takes place can happen __________________ (Non-Polar Covalent) or _____________________
(Polar Covalent).
b. Prediction by Substance Type: Covalent bonds are always between two _______________________
elements!
c. Prediction by Electronegativity Calculation:
1. Non-polar covalent: electrons are shared approximately _________________and ΔEN < 0.3
Now is a good time to memorize the _________________. These are elements found in nature with 2 atoms bonded
together. They are the most common examples of non-polar covalent molecules.
Hydrogen (H2), Nitrogen (N2), Oxygen (O2), Fluorine (F2), Chlorine (Cl2), Bromine (Br2), Iodine (I2)
A trick for memorizing them is the phrase “I Bring Clay For Our New House.”
2. Polar Covalent: electrons are __________________ shared and 0.3 ≤ ΔEN < 1.7
When you determine that a bond is “polar”, you indicate its direction of pull (also known as
electron distribution) by drawing an arrow OVER the bond pointing to the atom which is the
most electronegative. This shows that the electron is being shared MUCH closer to the more
electronegative element, causing that side of the molecule to become “partially” negative.
Because there is a shift in charge within the molecule it causes the molecule to become
“partially” positive on the other side. Since these are not true charges, we cannot use a + or –
sign. Instead, we use the lower-case Greek letter delta + to indicate “partially positive” or —
to indicate “partially negative”.
STEPS FOR INDICATING DIRECTION OF PULL (POLARITY) OF A BOND…
(1) Label each element with its electronegativity (EN) value.
(2) Place an arrow above the bond pointing to the more EN element.
(3) Place a — next to the more EN element, and a + next to the less
EN element. This shows the electron distribution in the bond!!!
H ─ O
Example 6-7. Show the direction of pull (electron distribution) for the bonds within F2 and HCl.
Example 6-8. How many polar bonds and how many non-polar bonds are there in C2Cl6?
STUDENT NOTES Pre-AP Chemistry
LEGGETT PreAP Bonding 6-6 Covalent Lewis Dot (10:34)
VI.
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https://vimeo.com/53288363
http://youtu.be/kY_xQwAVwe0
LEWIS DOT STRUCTURES OF COVALENT MOLECULES
We aren’t very interested in which element brought which
electron to the table. Think of the process like a potluck: each
element brings its valence electrons.
Guidelines:
(1) Add up the total number of valence electrons in the particle. THIS IS
THE MAXIMUM NUMBER OF DOTS YOU ARE ALLOWED TO DRAW!!!
(2) Decide which atom is the central atom—this is usually the atom present in fewest number, or, if there is the same
number of all atoms, it is the _______ _____ electronegative. However, hydrogen and fluorine can never be central
atoms! Write down the central atom, and place all of the other elements around it.
(3) Form covalent bonds between the central atom and the peripheral atom using __________ electrons per bond.
Sometimes a covalent bond is designated with “ ─ “ (dash) which indicates 2 electrons being shared.
(4) Start filling up atoms to satisfy the ______________ Rule, which states that most elements need _______ electrons to be
stable!
There are a few exceptions to the Octet Rule, WHICH YOU MUST MEMORIZE:

H only want 2 electrons (one bond)

Be only wants 4 electrons (two bonds)

B is satisfied with 6 electrons (three bonds)
Is the MAGIC NUMBER!
(Well, except for H, Be, B, and
 Sometimes S and P have more than 8 electrons (called an “Expanded Octet”)
sometimes S & P…..)
(5) Place any left-over electrons on the central atom in pairs. If the central atom has extra pairs of electrons beyond the
typical “4”, draw a line from the extra pair to the central atom… this is called an “Expanded Octet”.
While S and P are the most common two elements to have expanded octets, other elements can occasionally be forced
to have them as well. This includes Group 18 elements!
(6) After placing all of the electrons, if one of the atoms does not have a complete octet,
atoms will share a pair of electrons to make a double bond.
RANDOM THOUGHT: We are going to have you write a general formula for each of these. Trust us…you will
understand the purpose in just a little while! Use the letter “A” to represent the central atom, the letter “B” to
represent peripheral atoms, and the letter “X” to represent non-bonded electron pairs on the central atom. This is
the “ABX” Formula.
Example 6-9. Draw the LDS for the following compounds containing single bonds.
MOLECULE
H2O
HI
LEWIS DOT STRUCTURE
GENERAL
“ABX”
FORMULA
STUDENT NOTES Pre-AP Chemistry
MOLECULE
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LEWIS DOT STRUCTURE
GENERAL
“ABX”
FORMULA
BeCl2
BI3
NCl3
CF2H2
LEGGETT PreAP Bonding 6-7 Covalent Lewis Dot (11:26)
http://vimeo.com/53528962
http://youtu.be/4Zqh9olD2Mw
Example 6-10. Draw the LDS for the following compounds containing double or triple bonds.
MOLECULE OR
ION
LEWIS DOT STRUCTURE
GENERAL
“ABX”
FORMULA
N2
Don’t forget to
address the
charge shown!
CO32−
CO2
When you have
only carbons and
hydrogens, the
carbons will align
themselves in a
central chain and
the hydrogens
will go around
that chain!
C2H4
xxxxx
STUDENT NOTES Pre-AP Chemistry
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Example 6-11. Draw the LDS for structures that violate the Octet Rule (Expanded Octets). Put any extra
electrons on the central atom.
MOLECULE OR
ION
LEWIS DOT STRUCTURE
GENERAL
“ABX”
FORMULA
BrF5
ClF3
SF4
LEGGETT PreAP Bonding 6-8 Resonance (6:17)
http://vimeo.com/53528961
http://youtu.be/XA6idB8LRB8
a. LDS for ______________________ Structures: Having resonance structures means that there is a double bond
that can be in more than one possible location… but in actuality, the extra bond is shared equally among the
multiple locations!
For CO3-2, there are actually three equivalent ways to draw the structure. The final structure is actually in
between.
When drawing, you would show this….
But in actuality, the real structure is like this…
STUDENT NOTES Pre-AP Chemistry
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Questions to ask yourself when determining if a structure shows resonance:
(1) Do I have any double bonds?
(2) Can the double bond be moved to other locations?
Example 6-12. Draw Lewis structures for NO3─ ion. Don’t forget brackets and charges, since it is an ion! Give
the general “ABX” structure.
Example 6-13. Which of these examples has resonance structures?
b. LDS for _________________________________: You might have noticed that polyatomic ions tend to have covalent
bonds within them. However, these polyatomic ions can form ionic bond between the polyatomic anion
and other cations (either a metal ion or ammonium). Therefore, the resulting structure will have the LDS
shown as covalent within the polyatomic ion, and the LDS shown as ionic overall.
Example 6-14.
Draw the LDS for the following ionic compounds containing polyatomic ions.
Na2SO4
LEGGETT PreAP Bonding 6-10 VSEPR (6:13)
VII.
NH4NO3
http://vimeo.com/53528959
http://youtu.be/BXF7SqPni9A
VALENCE SHELL ELECTRON PAIR REPULSION (VSEPR) THEORY – FOR
COVALENTS
When we draw a Lewis Dot Structure, we are representing the molecule in a 2-D way… but in actuality, the molecule is 3-D!
The VSEPR Theory refers to the actual 3-D shape that a covalent molecule has. The VSEPR Theory is based on the premise that
the valence electrons of each peripheral atom will repel each other strongly, and therefore cause the peripheral atoms to
move as far from each other as possible. Placing the molecules as far away from each other as possible minimizes the
electrostatic repulsion between them.
The VSEPR Theory assigns a shape “name” to a covalent molecule based on how many bonded atoms and how many nonbonded electron pairs are present around the central atom of the molecule. You must memorize these shape names, and know
how to assign them! We will be using the “ABX” Structure to help us do this!
STUDENT NOTES Pre-AP Chemistry
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Example 6-15. Fill in the following chard about VSEPR shapes. Wait on the last column… we will finish it soon!
Example
BeCl2
BF3
SiH4
PH3
H2O
AsBr5
SeF6
Lewis Dot Structure
# bonded
atoms on
CA
# nonbonded pairs
of e- on CA
General
ABX
structure
VSEPR
Shape
Name
General
Polarity
of Shape
STUDENT NOTES Pre-AP Chemistry
http://vimeo.com/54341122
LEGGETT PreAP 6-11 Molecular Polarity (9:02)
VIII.
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http://youtu.be/WRTxUY4L0aI
COVALENT MOLECULE PROPERTIES & POLARITY OF COVALENT MOLECULES
In order to discuss covalent compound properties, we need to discuss the two types of covalent molecules –
Non-Polar Covalent Molecules and Polar Covalent Molecules. MOLECULE POLARITY IS DIFFERENT THAN
BOND POLARITY!!! You can have polar bonds in a molecule, but the overall molecule can still be
non-polar!
a. ________________________ Molecules: Central atom has no non-bonded pairs of electrons and is
surrounded by the same atoms. KEY: there is not preference for electron density on one part of
the molecule compared to any other. We say the molecule is symmetrical around the central atom.
Having this structure means that there is no “pull” holding one molecule to another molecule,
meaning that the molecules are held together very weakly. We will discuss more about these
“intermolecular forces” later.
This gives Non-Polar Covalent Molecules the following properties:
 Low melting (Tm) and boiling points (Tb)
 Most are not solids, but rather liquids or gases at room temperature
 Does not conduct electricity in ANY state (solid, liquid, or gas) as there is no pathway for mobile
electrons
 Insoluble in water and other polar solvents, but soluble in other non-polar solvents
 Exist as a MOLECULE, not as ions
b. __________________ Molecules: Central atom has non-bonded pair(s) of electrons OR is surrounded by
different atoms. We say the molecule is unsymmetrical around the central atom.
Being polar means that there is some pull between the molecules, which holds the molecules
together (intermolecular forces).
This gives Polar Covalent Molecules the following properties:
 Higher Tm & Tb than non-polar molecules, but much lower than ionic compounds (we’ll discuss
more later!)
 Most are liquids at room temperature (a few with large mass are solids)
 Does not conduct electricity in ANY state (solid, liquid, or gas) as there is no pathway for mobile
electrons
 Typically soluble in water and other polar solvents, but not soluble in non-polar solvents
 Exist as a MOLECULE, not as ions
The presence of polar bonds in a molecule MAY OR MAY NOT cause the entire molecule to be polar. The
basic question to ask yourself when determining polarity is: DOES THE MOLECULE LOOK THE
SAME ALL THE WAY AROUND THE CENTRAL ATOM? If “yes”, the molecule is
___________________. If “no”, the molecule is ________________. (Note: We only look at the sides that have
something there… either a bonded atom or a non-bonded pair of electrons.)
(1) All four sides have the same
element? NON-POLAR MOLECULE.
(2) Three sides have one element, and
the fourth side has a different
element? POLAR MOLECULE.
(3) Some sides have an element, and the
other side(s) has a non-bonded
electron pair(s)? POLAR MOLECULE.
3.0
2.1
2.5
2.5
3.0
3.0
3.0
3.0
3.0
STUDENT NOTES Pre-AP Chemistry
U N I T 6 | Page 14
ALSO: The reason why you saw the trends of solubility (as listed above) is that LIKE DISSOLVES LIKE…
polar substances dissolve in polar solvents (like water). Non-polar substances dissolve in non-polar
solvents (like oil).
Example 6-16. How many polar bonds are in CF4? Is the molecule itself
polar? Would it more likely be soluble in water or oil?
Example 6-17. How many polar bonds are there in CO2? Is the molecule
itself polar? Would it more likely be soluble in water or oil?
Example 6-18. How many polar bonds are there in ammonia (NH3)? Is the molecule itself polar? Would it
more likely be soluble in water or oil?
LEGGETT PreAP 6-12 IMF (10:12)
IX.
http://vimeo.com/54341121
http://youtu.be/1bQ2LbRQ1WU
METALLIC BONDS
Structure of Metallic Bonds: These bonds result when metal atoms donate their valence electrons to an
______________________________ that binds the atoms together. This actually forms a very strong bond!
Having this “sea of electrons” where electrons can move around gives
metallic bonds some very specific properties.
Properties of Metallic Bonds
(1) Have extremely high Tm & Tb due to the strong bonds present
(2) Conduct electricity in their solid and liquid (molten) forms, as electrons
are free to move
(3) Can conduct heat very well, again because of all of the movement from
the sea of electrons
(4) Have __________________ (shine) due to light reflecting off the sea of
electrons
(5) Are _______________________ (can be hammered into a sheet)
(6) Are ____________________ (can be pulled into a wire)
While metallic bonds are present in pure-element metals, they can also be present when you have a homogeneous
mixture (solution) of different metal elements.
A mixture of metal elements is referred to as an ____________________.
Interesting Tidbit: Some common examples of alloys include…
(1) Bronze: Cu and Sn
(2) Brass: Cu and Zn
(3) Sterling Silver: Ag and Cu
(4) Steel: Fe, C, and Cr
STUDENT NOTES Pre-AP Chemistry
X.
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INTERMOLECULAR FORCES (IMF)
__________________________ bonding is what occurs between two atoms. When we have called a compound ionic or covalent, we
have been referring to its intramolecular bonding.
However, for compounds to remain in a solid or liquid state, there must be some sort of pull between the molecules. This pull
between the molecules is what we call ______________________________ forces.
In these notes, you’ve seen a lot about the different melting and boiling points of different types of bonds. These temperatures
are greatly influenced by intermolecular forces. The stronger the type of intermolecular force, the higher the melting or
boiling point will be.
The following are the types of intermolecular interactions in order of decreasing strength:
A. Ion-Ion: These are the dominant forces felt within _____________ substances. Ion-ion interactions are very strong,
which is why ionic compounds have high melting & boiling points.
B. Hydrogen Bonding: These are the dominant forces felt in __________________________ molecules in which a H-F, H-O,
or H-N bond is present somewhere in the molecule. This is a special type of dipole-dipole interaction.
C.
Dipole-Dipole: These are the dominant forces felt in _______________________
_________________ molecules that don’t fit into the Hydrogen Bonding category. The
partial negative charge on one molecule can interact with the partial positive charge on
another molecule, causing a pull between the molecules.
D. London Dispersion: These are the dominant forces felt in
______________________________________ molecules. However, ALL
molecules are capable of London Dispersion forces, but other types
of IMF’s may dominate. The electron cloud on non-polar molecules can be temporarily
distorted to cause an instantaneous dipole moment. These instantaneous dipole moments can
interact, but the temporary interactions are very weak. Since it is easier to shift electrons on
bigger molecules, the strength of the London Dispersion interactions increases with size
(molar mass). This means that for all substances – if 2 compounds have the same
dominant type of IMF, the substance with the higher molar mass will have the stronger
IMFs.
However, it is very important to remember that INTRAMOLECULAR bonding (ionic & covalent bonds) is WAY
STRONGER than INTERMOLECULAR forces.
LEGGETT PreAP 6-13 Summary (10:02)
http://vimeo.com/54341123
http://youtu.be/NiNq9eNMTx4
Example 6-19. For each of the following: determine the dominant type of IMF between the molecules or
compounds in the pure state. Then, put them in order of increasing melting/boiling points.
SUBSTANCE
H2S
CH4
LiBr
NH3
CBr4
DOMINANT IMF
RANK of MP/BP
(5=LOWEST)
STUDENT NOTES Pre-AP Chemistry
U N I T 6 | Page 16
BOND COMPARISONS: Let’s Review!
Ionic Bond
General Formula
& Structure
Difference in
Electronegativity
Usually called a…
How can elements
reach 8 electrons to
be stable? (OCTET
RULE)
How do we show their
Lewis Dot Structure?
Basic Properties…
Polar Covalent Bond
Non-Polar Covalent
Bond