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Chemistry 20 –
Unit A – Diversity of Matter and Chemical Bonding
Section 3.1













structural formula
valence electron
quantum mechanics
orbital
valence orbital
bonding electron
lone pair
octet rule
Lewis symbol
electronegativity
covalent bond
ionic bond
metallic bond
Section 3.2









bonding capacity
coordinate covalent bond
empirical formula
molecular formula
Lewis formula
structural formula
stereochemical formula
polar covalent bond
bond dipole
Section 3.3






stereochemistry
VSEPR theory
polar molecule
non-polar molecule
polar covalent bond
bond dipole
Section 3.4








intermolecular bond
intramolecular bond
van der Waals force
dipole-dipole force
London force
momentary dipole
isoelectronic molecule
hydrogen bond
Section 3.5


crystal lattice
covalent network
1
Chapter 3–Understanding Chemical Compounds- Chemical Bonding: page 76
Bonding Theory: A Little History
 Bonding theory begins prior to Mendeleyev’s periodic table

Edward Frankland (1852) -stated each atom had a certain bond capacity

Friedrich Kekule (1858) – first to use a dash between atom to illustrate a bond (structural
formulas)

Jacobus van’t Hoff and Joseph Le Bel (1874) – extended structural formulas to 3 dimensions.

Richard Abegg (1904) – theorized that bonding capacity was related to the atoms electronic
structure. Also suggested that bonding occurred such that an atom could obtain a “noble gas
like” structure. He was the first to bring up the idea of an electron transfer due a chemical
reaction to achieve this stability. Atoms would then be held together by having opposite
charges. (Ionic bond) Gilbert Lewis (1916)- proposed that atoms could achieve a “noble gas
like” stable structure by sharing electrons as well as by transferring them. The electron
involved were called valence electrons. (Covalent bond)

Linus Pauling (1939)explained why certain electron arrangement were more stable covering
both ionic and covalent bonding.
Section 3.1 Bonding Theory and Lewis Formulas:
energy level: a specific energy an electron can have in an atom or ion
orbital: region of space around an atom’s nucleus where an electron may exist. An orbital is not
small particles in orbit but may be thought as a 3-D space where the electron might be.
valence orbital: volume of space that can only be occupied by electrons in an atom’s highest energy
level
valence electron: electrons that occupies the outermost energy level of an atom or ion. The number
of valence electrons can be determined by referring to the group number (Roman Numeral ‘A’) on
the periodic table.
1. Determine the number of valence electrons for the following atoms or ions:
a. chlorine
f. helium
b.
strontium
g.
carbon
c.
aluminum
h.
sulfur
d.
lithium
i.
krypton
e.
nitrogen
j.
chloride
2
Lewis Symbols:
The valence electrons are shown as dots around the central symbol, with the entire diagram showing
(for an atom) a net charge of zero.
Drawing Lewis symbols for main group atoms:

Write the element symbol to represent the nucleus and filled energy levels of the atom

Add a dot to represent each valence electron

Start by placing valence electrons singly into each of four valence orbitals (represented by
the four sides of the element symbol). Element can only have 4 orbitals. The only exception
is H and He which have only one orbital.

Each orbital must be occupied before electrons can be paired up.

(Each orbital can only hold a max of 2 e-. More than 2 result in too high of a repulsion force

If additional locations are required for electrons once each orbital is half-filled; start filling
each of the four orbitals with a second electron until up to eight valence electrons have been
represented by dots. (known as the octet rule)
Period 2 elements of the periodic table:
Group
Number
Atom
Number
of
valence
electrons
Li
Be
B
C
N
O
F
Ne
Lewis
Symbol
An atom with a valence orbital occupied by a single electron can theoretically share that electron
with another atom. Such an electron is, therefore, called a bonding electron. (
)
A full valence orbital, occupied by two electrons, has a repelling effect on electrons in any
nearby orbitals. Two electrons occupying the same orbital are called a lone pair. (
)
Number
of
bonding
electrons
Number
of lone
pairs
3
2. Draw a Lewis Symbol for each of the following:
a.
iodine (I)
f.
argon (Ar)
b.
barium (Ba)
g.
tin (Sn)
c.
boron (B)
h.
selenium (Se)
d.
bismuth (Bi)
i.
chloride (Cl1-)
e.
helium (He)
j.
sulfide (S2-)
3. Complete the table for the following atoms:
Atom
Valence
electrons
Lewis Symbol
Bonding
electrons
Lone Pairs
K
Ca
B
Si
N
O
F
Ar
Rb
As
P
C
4
4. Draw the Lewis Symbol of the following atoms and indicate how many bonding electrons
are present:
a
bromine
e
silicon
b
lithium
f
oxygen
g
polonium
h
germanium
c magnesium
d
bismuth
Electronegativity: page 81
 atoms exhibit (have) different abilities to attract negative valence electrons
 the farther away from the nucleus that negative electrons are, the weaker their
attraction to the nucleus (which is positive due to the presence of protons)
 inner electrons (those closer to the nucleus) shield the valence electrons from the
attraction of the positive nucleus
 the greater the number of positive protons in the nucleus, the greater the attraction for
negative electrons
 combining these three points, it was possible to assign a value to any atom
 the term electronegativity is used to describe the relative ability of an atom to attract a
pair of bonding electrons in its valence level (energy level)
 electronegativity is usually assigned on a developed scale – according to this scale,
fluorine has been given the highest assigned electronegativity of 4.0 – cesium has the
lowest assigned electronegativity of 0.8
 metals tend to have low electronegativities
 non-metals tend to have high electronegativities
5. Using your periodic table, determine the electronegativity of the element a-h of
table for question 4.
a. ____
b. ____
c. ____
d. ____
e. ____
f. ____
g. ____
h. ____
5
Bonding: page 82
A chemical bond is:
Covalent Bonding: page 82
According to Science 10: bond formed when non-metallic atoms share electrons; atoms in a
molecule are bound together by covalent bonds
A covalent bond is:
The key ideas of the Lewis theory of bonding are:

Atoms and ions are stable if they have a noble gas-like electron structure; i.e. a stable octet of
electrons.

Electrons are most stable when they are paired.

Atoms form chemical bonds to achieve a stable octet of electrons.

A stable octet may be achieved by an exchange of electrons between metallic and
nonmetallic atoms.

A stable octet of electrons may be achieved by the sharing of electrons between nonmetallic
atoms.

The sharing of single valence electrons results in a covalent bond.
6
Ionic Bonds: page 83
Table salt is made up of two elements that can be very dangerous, namely, sodium and chlorine.
Combined they form a safe compound.
Sodium and chlorine combine during a reaction that will transfer
one valence electron from sodium to chlorine.
After the transfer sodium and chlorine are no longer neutral.
The sodium and chlorine have becomes ions (atoms with
unequal number of protons and electrons). The chlorine now has
a net charge of -1 and the sodium ion has a charge of +1. Since
objects of opposite charge attract each other, the two ions will be held
tightly together. This type of attraction is called an ionic bond.
In a crystal of sodium chloride, each ion will have six
neighboring ions of opposite charge. This makes ionic bonding a very
strong type of interaction between atoms. These strong bonded ionic
compounds have a high melting point and boiling point.
Ionic bonding
type of bond formed when electrons transfer between metals
and nonmetals
electrostatic attraction between oppositely charged ions
(positive and negative ions) in the crystal lattice of a salt
metals ---------- lose electrons ---------- change into a positive ion
nonmetals ---------- gain electrons ---------- change into a negative ion
due to unlike charges being present : positive ion ------ is attracted by the ----- negative ion
Bonding of NaCl: Lattice Structure and Empirical Formulas
7
Metallic Bonding: page 83
Summary: Bonding Theory (page 84)
 the formation of a chemical bond involves competition for bonding electrons
occupying valence orbitals.
 if the competing atoms have equal electronegativities, the electrons are shared
equally.
 electrons sharing between atoms with high electronegativities results in covalent
bonding.
 electrons sharing between atoms with low electronegativites, often results in metallic
bonding.
 if the competing atoms have unequal electronegativities, the electrons are unequal
covalent bonding; if unequal enough that electrons transfer, the result will be ionic
bonding.
 bonding theory was created by chemists to describe, explain, and predict natural
events and observed properties.
8
Section 3.2 Explaining Molecular Formulas: page 85
Molecular Elements: page 85
hydrogen
H2
bromine
Br2
nitrogen
N2
iodine
I2
oxygen
O2
astatine
At2
fluorine
F2
phosphorus
P4
chlorine
Cl2
sulfur
S8
Molecular Compounds: page 86
Table from page 87: Bonding Capacities of Some Common Atoms
Atom
Number of
valence
electrons
Lewis
Symbol
Number of
bonding
electrons
Bonding
capacity
Electronegativity
carbon
nitrogen
oxygen
fluorine
hydrogen
9
Covalent bonds occur when there is a sharing of single valence (bonding) electrons between
two nonmetallic atoms in order to form an octet (8) of valence (bonding) electrons
(exception to this rule is hydrogen, which will have two valence (bonding) electrons).
 single covalent bond – is formed when 1 pair of single valence (bonding) electrons is
shared
 double covalent bond - is formed when 2 pairs of single valence (bonding) electrons
are shared
 triple covalent bond - is formed when 3 pairs of single valence (bonding) electrons
are shared
 coordinate covalent bond – a covalent bond in which one of the atoms donates
both electrons
.
hydrogen gas, H2(g)
oxygen gas, O2(g)
nitrogen gas, N2(g)
Types of Formulas: page 87
Ionic Compound Formulas:
Lewis formula for ionic compounds:
sodium chloride
potassium fluoride
calcium chloride
barium oxide
10
Molecular Compound Formulas:
Table 2
Names of Types of Formulas for Molecular Compounds
page 88
Example: acetic acid (CH3COOH)
Empirical
Formula
An empirical formula shows the
simplest whole-number ratio of atoms
in the compound. Empirical formulas
are rarely useful for molecular
compounds.
Molecular
Formula
A molecular formula shows the actual
number of atoms that are covalently
bonded to make up each molecule.
A molecular formula often has the atom
symbols written in a sequence that
helps you determine which atoms are
bonded to which.
Lewis Formula * A Lewis formula is also commonly
called a Lewis diagram, or an electron
dot diagram. It uses Lewis symbols to
show electron sharing in covalent
bonds, electron transfer in ionic bonds,
and the formation of stable valence
octets of electrons in molecules.
Structural
Formula
* A structural formula is also
commonly called a structural diagram.
As well as showing which atoms are
bonded, the type of covalent bond is
represented by the number of lines
drawn between atomic symbols.
Stereochemical A stereochemical formula is a structural
formula drawn to try to represent the
Formula
three-dimensional molecular shape.
For much larger molecules, this style of
representation often becomes too
complex to be practical.
* It is strongly recommended that you memorize this terminology information.
11
Determining Lewis Formulas: page 89
Lewis formula predictions are limited to entities with only one central atom. The central
atom is the atom to which all the other atoms – peripheral atoms – are bonded to.
Simple Lewis formulas can be predicted with a series of five steps.
Determine the Lewis Formula and the structural formula for _____________________,
________________________ and _________________________.
1. Count the total valence electrons in the entity by adding the valence electrons of each atom. If
the entity is a polyatomic ion, add (usually) or subtract valence electrons to account for the net
charge, one for each unit of charge.
2. Arrange the peripheral atom symbols around the central atom symbol, and place one pair of
valence electrons ( ) between each peripheral atom and the central atom (bond pairs).
3. Place more pairs of valence electrons (lone pairs) on all the peripheral atoms, to complete their
octets. Recall that a hydrogen atom’s energy level is completed with only two valence electrons.
4. Place any remaining valence electrons on the central atom as lone pairs.
5. If the central atom’s octet is not complete, move a lone pair from a peripheral atom to a new
position between that peripheral atom and the central atom. Repeat until the central atom has a
complete octet. For a molecule, this completes the Lewis formula. If the entity is a polyatomic
ion, place square brackets around the entire Lewis formula, and then write the net charge outside
the bracket on the upper right.
To show the structural formula, omit all lone pairs and replace every bond pair with a line.
Example 1
Example 2
Example 3
E.g. Draw the Lewis diagram for the carbonate ion
12
Draw Lewis diagram for CO321. total e- = C( 4 ) + 3O ( 3 x 6 ) + 2- ( 2 ) = 24 e-
O
2.
C
O
3.
O
4.
C
O
O
O
O
6.
O
C
O
O
C
O
5.
O
2-
C
O
O
O
Resonance structure :
When 2 or more Lewis diagram can describe a formula, these diagrams are called resonance
structures of that formula. The three resonance structure of CO32- is given below
O
O
2-
O
C
C
O
2-
O
O
2-
C
O
O
O
Draw Lewis Formulas and Structural Formulas for the following
13
fluorine gas, F2(g)
hydrogen fluoride
oxygen difluoride
water
ammonia
methane
sulfur hexafluoride
phosphine, PH3(g)
NH3BH3
5. Draw a Lewis formula of the following molecules:
Lewis Symbol
Lewis Symbol
a
HBr(g)
b
carbon
disulfide
c
HOCl
d
sulfur
dichloride
14
e
H2Se(g)
f
NH2Br
g
CH3F
h
carbon
dioxide
i
Silicon
tetrafluoride
j
HI(g)
k
PBrIF
l
diiodine
sulfide
m
CH3SH
o
nitrogen
tribromide
15
Draw a Lewis formula to represent the following ions: Write the formula of the ion before
drawing the Lewis formula.
a.
hydrogen ion
b.
nitrate ion
c.
hydroxide ion
d.
carbonate ion
e.
hydrogen sulfide ion
f.
hydrogen carbonate ion*
* Hint: A hydrogen ion (proton) bonds to any lone pairs from answer (d).
16
Rules for Drawing Lewis Symbols For Molecules That Have More Than One
Central Atom.
Step 1
Determine the atom with the largest bonding capacity. Use The group number
“Roman Numeral” to determine bonding capacity. This is the central atom.
Step 2
Draw the central atom. Draw the valence electrons around the central atom.
Unpaired (bonding) electrons are available for bonding
Step 3
All other atoms are drawn around the central atom pairing up bonding electrons
such that each atom is surrounded by an octet of electrons (except hydrogen)
Step 4
Work though the chemical formula from left to right ensuring that all atom
groups remain intact in the Lewis model.
Step 5
If any electrons are left unpaired after all atoms have been drawn, determine if a
double bond (sharing of two pairs of electron) or a triple bond (three pairs of
electrons) is required to complete the diagram
Example
a. CH3Cl
b. N2Cl2Br2
c. C2H5OH
d. CH3OCH3
e. CH2O
CH3COCH2OH
Draw the Lewis Formula and Structural Formula of the following:
a.
Lewis Symbol / Structural Formula
hydrogen peroxide
b.
Lewis Symbol / Structural Formula
O2F2
17
c.
ethane
d.
C2H4
e.
ethanol
f.
methanethiol
CH3SH(g)
g.
methanol
h.
C2H2
i.
N2H3OF
j.
HCHO
k.
CH3NHCH3
l.
CH3CHOHCH3
n.
N2H4
m. CH3OCH3
18
o.
CH2FCH2I
p.
propane
q
arsenic trihydride
r
PHO
Structural Formulas:
Color
Atom
represented
hydrogen
oxygen
nitrogen
carbon
halogens
sulfur
white ball
red ball
orange ball
black ball
green ball
yellow ball
Bonding
capacity
forms 1 bond
forms 2 bonds
forms 3 bonds
forms 4 bonds
forms 1 bond
forms 2 bonds
Using a molecular model kit – construct models of each of the following molecules and
draw their structural formula.
O2
Cl2
HCl
H2O
NH3
CH4
H2O2
N2H4
19
C2H6
CH3OH
C2H5OH
CH3OCH3
C2H5OCH3
C2H4
CO2
H2CO
Bonding
Capacity
1
Example
Chlorine
2
Oxygen
3
Nitrogen
4
Carbon
Possible combinations
Draw the structural formula of the following without the aid of the molecular model
kit.
1
CH3NHCH3
2
CH3CH2OCOCHO
20
3
C2H5CCNH2
4
C2H2
5
CH3OH
6
CH3CHOHCONH2
7
CH3COOC2H5
8
CH3CCCH3
9
HOF
10
C2H5COOH
11 C2H4
12
C4H10
13 HCHO
14
CS2
15 C2H5CONH2
16
C2H5COCH3
21
17 HNO
18
CH3CHOHCONH2
19 C3H4
20` PCl3
21 CH3CH2COONH2
22
BrCP
23 CHONHOH
24
CO2
25 C3H7OCH3
26
HCC2CH
27
CO(NH2)2
28
C2H2Cl2
29 CH2CHCH2OH
30
CH2COHOCH3
22
Predicting molecular formulas of nonmetallic compounds
 Write a partial chemical equation show only the reactants.(ie. reactant + reactant )
 Predict the formulas of the following products using the fewest number of atoms and the
least number of double or triple bonds.
 Show the Lewis formula of each type of atom involved and then build Lewis formula of
the simplest molecule.
 From the Lewis formula draw the structural formula.
e.g. #1. oxygen gas + fluorine gas 
e.g. #2. phosphorus + hydrogen gas 
1. oxygen gas + hydrogen gas 
2. iodine + bromine 
3. phosphorus + chlorine gas 
4. oxygen gas + chlorine gas 
5. carbon + sulfur 
6. sulfur + oxygen gas 
23
Section 3.3
Molecular Shapes and Dipoles
page 91
Valence shell electron pair repulsion theory ( VSEPR theory )

The shapes of molecules are very important because many of their physical and chemical
properties depend upon the three dimensional arrangements of their atoms.

e.g , The functioning of enzymes, which are substances that control how fast biochemical
reactions occur, requires that there be a very precise fit between one molecule and another.
Enzymes are large molecule of protein, because of their shape will react only with a specific
molecule much like a key will fit into a lock. Even slight alterations in molecular geometry
can destroy this fit and deactivate the enzyme, which in turn prevents the biochemical reaction
involved from occurring.

Very difficult to predict shapes of large molecules

1957 a simpler theory was developed to help explain and predict the stereochemistry of
certain chemical elements and compounds. Called VSEPR theory.
Valence shell electron pair repulsion theory ( VSEPR ) is the best theoretical explanations of
molecular shapes. The general theory is based on the idea that valence shell electron pairs (bonding
electrons and nonbonding lone pairs), being negatively charged, stay as far apart from each other
as possible so that the repulsion between them are at a minimum.
Stereochemistry:
is the study of the 3-D spatial configuration of molecules and how this affects
reactions.
According to VSEPR theory:

only the valence electrons of the central atom(s) are important for molecular shape

valence electrons are paired in a molecule or polyatomic ion

bonded pairs of electrons and lone pairs of electrons are treated approximately equally

Valence electron pairs repel each other electrostatically

the molecular shape is determined by the positions of the electron pairs when they are a
maximum distance apart
24
Using VSEPR to predict the shapes of of the hydrogen compounds of period 2.
BeH2(s), BH3(g), CH4(g), NH3(g), H2O(l), HF(g)
Table 1: Geometry of beryllium dihydride (pages 92 – 93)
Lewis formula
Bond
pairs
Lone
pairs
Total
pairs
General
formula
Explanation:
Consider BeH2 molecule. The Lewis structure is
H Be
Electron pair
arrangement
Stereochemical
formula
H
How are these atoms arranged ? Are they in a straight line or at some angle ?
The answer to this question depends on the arrangement of electron pairs around the central atom.
The central atom is Berylium and there are 2 shared pairs of electrons around Beryllium. The minimum
repulsion will occur, when the electron pairs are on opposite sides of the nucleus. We can represent this
as:
H– Be – H
the molecule should be linear
Be
Table 2: Geometry of boron trihydride
Lewis formula
Bond
pairs
Lone
pairs
Total
pairs
General
formula
Electron pair
arrangement
Stereochemical
formula
Explanation:
H
••
The molecule BH3 . The Lewis structure is
H
B H
The central atom B has 3 bonding pairs, the arrangement which will lead to minimum repulsion is a
triangle with boron in the center.
B
The shape of the molecule is trigonal planar
25
Table 3: Geometry of methane
Lewis formula
Bond
pairs
Lone
pairs
Total
pairs
General
formula
Electron pair
arrangement
Stereochemical
formula
Total
pairs
General
formula
Electron pair
arrangement
Stereochemical
formula
Total
pairs
General
formula
Electron pair
arrangement
Stereochemical
formula
Explanation:
Table 4: Geometry of ammonia
Lewis formula
Bond
pairs
Lone
pairs
Explanation:
Table 5: Geometry of water (page 94)
Lewis formula
Bond
pairs
Lone
pairs
Explanation:
26
Table 6: Geometry of hydrogen fluoride
Lewis formula
Bond
pairs
Lone
pairs
Total
pairs
General
formula
Electron pair
arrangement
Stereochemical
formula
Explanation:
Shapes of Molecules
VSEPR theory describes, explains, and predicts the geometry of molecules by counting
pairs of electrons that repel each other to minimize repulsion. The process for predicting
the shape of a molecule is summarized below:
Step 1: Draw the Lewis formula for the molecule, including the electron pairs around the
central atom.
Step 2: Count the total number of bonding pairs (bonded atoms) and lone pairs of
electrons around the central atom.
Step 3: Double bonds and triple bonds can be treated as though they are single bonds between
neighbouring atoms.
Step 4: When applying VSEPR, we general consider the shape around the central atom.
If there are more than one central atom, there are more than one shape.
Step 5: Refer to Table 7, and use the number of pairs of electrons to predict the shape of the
molecule.
Eg.
Construct the Lewis structure and predict the shape of each of the following.
a. SiF4
b. SO42c. PH3
d. ClO4-
27
Table 7: Using VSEPR Theory to Predict Molecular Shape (page 95)
General
formula*
AX2
Bond
pairs
2
Lone
pairs
0
Total
pairs
2
Geometry**
Stereochemical
formula
Examples
linear
(linear)
AX3
3
0
3
trigonal planar
(trigonal planar)
AX4
4
0
4
tetrahedral
(tetrahedral)
AX3E
3
1
4
trigonal
pyramidal
(tetrahedral)
AX2E2
2
2
4
angular
(tetrahedral)
AXE3
1
3
4
linear
(tetrahedral)
* A is the central atom; X is another atom; and E is a lone pair of electrons
** The electron pair arrangement is in parenthesis
The Multiple Bond in VSEPR Models: page 96
28
Draw structural formulas for the following and determined the shape around their central
atom(s)
E.g.
a. C2H4 (ethane)
b. NO2- (Nitrite ion)
c. Cl2CO
d. C2H2
e. CH3OH
f. CH3COCH3
Summary
Number of
electron pairs
2 pairs
Molecular shape
BeCl2
X
3 pairs
Example
Cl – Be – Cl
linear
planar triangular
(trigonal planar)
120
BF3
F
o
B
F
F
4 pairs
tetrahedral
CH4 , NH4+
H
109.5o
C
H
H
H
29
Bonding
pair
4
Lone
pair
0
Name
& example
Tetrahedral
CH4
Shape
H
109.5o
C
H
H
H
3
1
Pyramidal
NH3
N
H
H
H
2
2
107.3o
Bent
H2O
O
H
104.5o
Molecular Polarity: Dipole Theory
H
page 98
Electronegativity and Bond Polarity
Recall that electronegativity of elements
 increases from left to right within a period

and bottom to top within a group or family on the periodic table

Linus Pauling explained that the difference in electronegativity between atoms was the
reason for the polarity of a covalent bond.

Covalent bonds in which the bonding electrons are unequally shared and unsymmetrical
distributed between the nuclei of two bonded atoms are called covalent bonds.

Such bonds occur between atoms of different electronegativity.

The greater the electronegativity difference, the more polar the bond will be.
30

Pauling reasoned that when two atoms of different electronegativity share electrons, the
attraction for these electron are unequal, therefore the electrons will spend more of their time
closer to the nucleus of higher electronegativity.

In this situation, the bond formed is called a polar covalent bond.
The presence or absence of polar bonds within molecule plays a very important part in determining
the chemical and physical properties of molecules.
Some of these properties are:
 melting points,
 boiling points,
 viscosity and
 solubilities in solvents.
eg.
Cl2
and
HCl
∙∙
∙∙
‫׃‬Cl ‫ ׃‬Cl‫׃‬
¨ ¨
H ‫׃‬Cl‫׃‬
3.2
2.2 3.2
∙∙
¨
3.2
in the middle
nonpolar covalent bond
closer to the Cl
polar covalent bond
Both molecules are formed by sharing one pair of electrons. In the chlorine molecule, the electrons
are shared and distributed equally by both chlorine atoms since both atoms have the same attraction
for the shared pair of electrons(ie. same electronegativity). However in the hydrogen chloride, the
shared electrons are not equally shared by both atoms. The chlorine with a stronger electronegativity
attracts the shared pair of electrons more strongly than does the hydrogen. Since the shared
electrons are displaced towards the chlorine and away from the hydrogen, the chlorine becomes
partially negative(δ-) and the hydrogen partially positive (δ+).
In this situation, a bond dipole occurs between the hydrogen and chlorine and the bond is said to be
polar. Bonds or molecules that have opposite poles are said to be polar. Polar covalent bonds can
be said to have a charge separation or bond dipole. A bond dipole can be represented by an arrow
pointing to the negative side of the bond.
eg.
δ+ H—Cl δ-
δ – Greek lower case delta represents partial charge
2.1 3.0
31
For each of the following:
1. Draw the structural formulas
2. Draw in the bond dipoles
eg H2O
a) HF
b) NH3
c) CH4
d) CO2
e) CH3OH
f) OF2
g) C2H6
h) CH3Cl
O
H
H
Polar Molecule
A molecule is polar if the dipoles do not cancel each other out. If bond dipoles do not cancel each
other out, the molecule is said to be polar.
Using this information, determine if the above molecules are polar or nonpolar.
Predicting Bond Types:
Bond type can be described as belonging to one of three classes:
1. nonpolar covalent:
no difference in electronegativity
2. polar covalent:
a difference in electronegativity exists
3. ionic:
attraction between a metal and a nonmetal
32
Summary of Bonding:
Intramolecular force
ionic bond
Bonding model

involves an electron transfer between a metallic atom and a
nonmetallic atom, resulting in the formation of cations and anions
polar covalent bond

cations and anions attract each other

involves unequal sharing of pairs of bonding electrons by atoms
of two different nonmetallic elements

bonds can involve 1, 2, or 3 pairs of bonding electrons, i.e., single
(weakest), double or triple (strongest) bonds
non-polar covalent

involves equal sharing of electrons
bond

bonds can involve 1, 2, or 3 pairs of electrons, i.e., single
(weakest), double or triple (strongest) bonds
Eletronegativites Revisted
Note the electronegativities of elements from left to right across periods and down groups.
The trends observed for each of these is?

Across a period: the electronegativities generally increase from left to right across a period
with the Group VII element having the highest value for the period.
Why?

Down a group: the electronegativities generally decrease from top to bottom down a group.
Francium is the element with the lowest electronegativity.
Why?
Electronegativity is a numerical value describing the relative tendency of an atom to attract
negative electrons to itself when bonded to another atom.
As we move along a period of the periodic table the nuclear charge increases and so electrons are
attracted more strongly. However, when we start a new period, the electrons in the filled inner shell
are able to shield the outer electrons from the nuclear charge. Consequently, electronegativity
decreases as we move down a group.
The noble gases are an exception to this; they are not given a value. This is because their
exceptional stability makes it difficult to attract electrons (i.e. they form NO bonds).
33
Label the electronegativity, and label the charges (if any) on the ends of the bond. Classify the
following bonds as either ionic, polar covalent, or nonpolar covalent.
a
Al - Si
h
B - Na
b
Ba - O
i
Ca - Cl
c
C-H
j
F-S
d
Li - S
k
Br – Rb
e
Ca - P
l
F-F
f
H – Cl
m
N–O
g
I – Br
n
Mg – S
Polarity
Indicate the polarity (polar / nonpolar) of the following bonds:
a
P – Cl
f
O–H
b
C=O
g
S–C
c
H–F
h
I–I
d
P–H
i
C–C
e
N – Br
j
N=N
1. Molecular Polarity
Draw the structural formulas of the following and using bond dipoles, determine if it is polar
or nonpolar.
1
H2O
2
CO2
3
HI
4
CH3OH
34
5
F2
6
CH3I
7
NI3
8
C3H8
9
C2H6
10
CS2
11
CCl4
12
H2Se
Guidelines for Predicting Polar and Non-polar Molecules
Type
Polar
AB

Description
diatomic compounds
CO(g)
HBx

any molecule with a single H
HCl(g)
AxOH

any molecule with an OH at one end
C2H5OH(l)
OxAy

any molecule with an O at one end
H2O(l), OCl2(g)
NxAy

any molecule with an N at one end
NH3(g), NF3(g)

all elements
Cl2(g), N2(g)

most carbon compounds (including
CO2(g), CH4(g)
Nonpolar Ax
CxAy
Examples
organic solvents, fats, and oils)
35
Polar Substances: page 103
What empirical test can we do to demonstrate the polarity of molecules? (page 98)
36
Section 3.4
Intermolecular Forces
page 105
Types of forces that exist:
Intramolecular Forces
Intermolecular Forces
forces that hold atoms together in a molecule
forces of attraction and repulsion that exist
between molecules
forces of attraction within a molecule
between two atoms
forces of attraction between molecules
ionic
covalent
Van der Waals Forces
transfer or
exchange of
valence electrons
sharing of single
valence electron
single
double
triple
polar

London Dispersion Forces

Dipole-Dipole Forces
Hydrogen Bonding
polar / non-polar
Intermolecular Forces
Intermolecular forces are the forces of attractions that exist between molecules in a compound.
These cause the compound to exist in a certain state of matter: solid, liquid, or gas; and affect the
melting and boiling points of compounds as well as their solubilities of one substance in another.
Solid:
A state of matter in which the matter is not compressible nor does it flow.
Liquid:
A state of matter in which the matter is not compressible but can flow.
Gas:
A state of matter in which the matter is compressible and can flow.
The melting point of a compound is the temperature at which a compound turns from
a solid to a liquid or a liquid to a solid.
The boiling point of a compound is the temperature at which a compound turns from a liquid to a gas
or a gas to a liquid. This temperature is a true measure of the forces of attractions between molecules
as molecules separate from one another when they turn from a liquid to a gas.
The stronger the attractions between particles (molecules or ions), the more difficult it will be to
separate the particles. When substances melt, the particles are still close to one another but the forces
of attraction that held the particles rigidly together in the solid state have been sufficiently overcome
37
to allow the particles to move. When substances boil, the particles are completely separated from
one another and the attractions between molecules are completely overcome.
The energy required to cause substances to melt and to boil, and thus disrupt the forces of attraction,
comes from the environment surrounding the material. If you place a piece of ice in your hand, the
ice will melt more quickly than if it is placed on a cold counter top. The energy required to melt the
ice comes from your hand, your hand gets colder and the ice gets warmer.
Intermolecular forces are generally much weaker than covalent bonds

Only 41 kJ/mol of energy is required to overcome the intermolecular attraction between H2O
molecules in the liquid state (i.e. the energy required to vaporize the sample) Intermolecular forces

However, 286 kJ/mol of energy is required to break the covalent bond between the H and O
atoms in the H2Ol molecule. Intramolecular Forces (ie covalent bonds)
Thus, when a molecular substance changes states, the atoms within the molecule are unchanged
The temperature at which a liquid boils reflects the kinetic energy needed to overcome the attractive
intermolecular forces (likewise, for the temperature at which a solid melts).
Types of attractive forces between neutral molecules
 Dipole-dipole forces

London dispersion forces

Hydrogen bonding forces
Typically, dipole-dipole and dispersion forces are grouped together and termed van der Waals
forces
Dipole-Dipole Forces
A dipole-dipole force exists between polar molecules

Polar molecules attract one another when the partial positive charge on one molecule is near
the partial negative charge on the other molecule

The polar molecules must be in close proximity for the dipole-dipole forces to be significant

In a condensed state such as a liquid, where many molecules are in close proximity, the
dipoles find the best compromise between attraction and repulsion. This means the
molecules orient themselves to maximize the + &
_
interactions and to minimize the + & +
and - & - interactions as shown below.
38
London Dispersion Forces
Nonpolar molecules would not seem to have any basis for attractive interactions.
 However, gases of nonpolar molecules can be liquefied indicating that if the kinetic energy is
reduced, some type of attractive force can predominate.

Fritz London (1930) suggested that the motion of electrons within an atom or non-polar
molecule can result in a transient dipole moment
A Model To Explain London Dispersion Forces:
Helium atoms (2 electrons)
 Consider the particle nature of electrons

The average distribution of electrons around each nucleus is spherically symmetrical

The atoms are non-polar and posses no dipole moment

The distribution of electrons around an individual atom, at a given instant in time, may not
be perfectly symmetrical
o
Both electrons may be on one side of the nucleus
o
The atom would have an apparent dipole moment at that instant in time (i.e. a
momentary or instantaneous dipole)
o
A close neighboring atom would be influenced by this apparent dipole - the electrons
of the neighboring atom would move away from the negative region of the dipole
Due to electron repulsion, a temporary dipole on one atom can induce a similar dipole on a
neighboring atom
 This will cause the neighboring atoms to be attracted to one another

This is called the London dispersion force (or just London force)

It is significant only when the atoms are close together
39
The ease with which an external electric field can induce a dipole (alter the electron distribution)
with a molecule is referred to as the “polarizability” of that molecule

The greater the polarizability of a molecule the easier it is to induce a momentary dipole and
the stronger the dispersion forces

Larger molecules tend to have greater polarizability
o
Their electrons are further away from the nucleus (any asymmetric distribution
produces a larger dipole due to larger charge separation)
o
The number of electrons is greater (higher probability of asymmetric distribution)
In general, we may conclude that the more atoms in a molecule(greater number of electrons),
the greater the opportunity to form an instantaneous dipole. The result is to increase the
attractive forces and raise the boiling point

Dispersion forces are also present between polar/non-polar and polar/polar molecules (i.e.
between all molecules)
Hydrogen Bonding
A hydrogen atom in a polar bond (e.g. H-F, H-O or H-N) can experience an attractive force with a
neighboring electronegative molecule or ion which has an unshared pair of electrons (usually an F, O
or N atom on another molecule)
Hydrogen bonds are considered to be dipole-dipole type interactions

A bond between hydrogen and an electronegative atom such as F, O or N is quite polar:

The hydrogen atom has no inner core of electrons, so the side of the atom facing away
from the bond represents a virtually naked nucleus

This positive charge is attracted to the negative charge (lone pairs)of an electronegative
atom in a nearby molecule
40

Because the hydrogen atom in a polar bond is electron-deficient on one side (i.e. the side
opposite from the covalent polar bond) this side of the hydrogen atom can get quite close to a
neighboring electronegative atom (with a partial negative charge) and interact strongly with it
(remember, the closer it can get, the stronger the electrostatic attraction)
o
Hydrogen bonds vary from about 4 kJ/mol to 25 kJ/mol (so they are still weaker than
typical covalent bonds.
o
But they are stronger than dipole-dipole and or dispersion
forces.
o
They are very important in the organization of biological
molecules, especially in influencing the structure of proteins
Water is unusual in its ability to form an extensive hydrogen bonding network
 As a liquid the kinetic energy of the molecules prevents an extensive
ordered network of hydrogen bonds

When cooled to a solid the water molecules organize into an
arrangement which maximizes the attractive interactions of the hydrogen bonds
o
This arrangement of molecules has greater volume (is less dense) than liquid water,
thus water expands when frozen
o
The arrangement has a hexagonal geometry (involving six molecules in a ring
structure) which is the structural basis of the six-sidedness seen in snow flakes
o
Each water molecule can participate in four hydrogen bonds , one with each nonbonding pair of electrons and one with each H atom
Diagram
41
Higher IMFs result in a strong attractive force between molecules that requires alot of energy to break
resulting in a high boiling point.
-A molecule like water possesses all 3 types of IMF has a a high boiling point of 100°C.
-A molecule like F2 has only LDF has a much lower boiling point of -188°C.
Eg.#1 The boiling point of argon is –189.4ºC.
a. Why is it so low?
b. How does this boiling point help prove that London dispersion forces exist?
c. The boiling point of xenon is –119.9ºC. Why is it higher than argon?
Eg#2 Put the following in order from the lowest to the highest boiling point? Explain your answer.
F2, C2H6, NH3, CH3Cl
How does structure and polarity affect boiling points?
The variation of boiling point is decided by two factors:
1. The molecular size. The larger molecules have a higher boiling point.
2. The intermolecular forces acting upon the molecules. The greater the forces the higher the
boiling point. The more attracted the molecules are to one another the less likely the molecules
will be able to escape from one another and become vapor molecules.
With regard to strength:
Hydrogen Bonding > Dipole-dipole Forces > London Dispersion Forces
42
Relation of Boiling Point to the Number of Electrons
and to the Type of Intermolecular Forces
Complete the following page. The first is given as an example. Note that there are six series or
groups of molecules.
Molecular
Substances
with state
Polarity of
molecule
Number of
electrons
Boiling
point
(oC)
London
Dispersion
Forces
(polar / non-polar)
eg
F2(g)
non-polar
18
-188

1
Cl2(g)
-35

2
Br2(l)
59

3
I2(s)
184

4
ClF(g)
-101
5
BrF(g)
-20
6
BrCI(g)
5
7
ICl(s)
97
8
IBr(s)
116
9
CH4(g)
-162
10
C2H6(g)
-87
11
C3H8(g)
-45
12
C4H10(g)
-0.50
13
C5H12(l)
36
14
CF4(g)
-129
15
CCl4(l)
77
16
CBr4(s)
189
17
CH3F(g)
-78
18
CH3Cl(g)
-24
19
CH3Br(g)
3.6
20
CH3I(l)
43
21
CH3OH(l)
65
DipoleDipole
Forces
Hydrogen
Bonding
43
22
C2H5F(g)
-38
23
C2H5Cl(g)
13
24
C2H5Br(l)
38
25
C2H5I(l)
72
26
C2H5OH(l)
78
1. Which of the following would you expect to have the highest boiling point BrF(g) or C3H8(g).
Account (why) for the difference in boiling points. Check your answer with the table above.
2. The different series of substances give in the pervious table, in general, have increasing boiling
points with increasing number of electrons. Explain this trend of number of electrons and
strength of intermolecular forces.
3. Methanol, CH3OH(l), and ethanol, C2H5OH(l), each have the least number of electrons but the
highest boiling point of their respective series. Account for this.
4. Explain the difference in boiling point between C2H6 and CH3F.
44
PbF2 melts at 855 oC and boils at 1290 oC. PbI2 melts at 402 oC and boils at 954 oC. Why?
Which would you expect to have the higher melting and boiling point: Cl2 or O2? ________ Why?
Which would have the higher boiling point, HF or HBr? ________ Why?
Bonds holding ionic substances (ionic bonds) together are (stronger or weaker) _______________
than bonds holding molecular substances (covalent bonds) together.
Physical Properties of Liquids
Many physical properties of liquids can be explained by intermolecular forces. This includes surface
tension, shape of the meniscus, volatility and capillary action.
Surface Tension
By definition, the molecules of a liquid exhibit intermolecular attraction for one another.
What happens to molecules at the surface in comparison to those in the interior of a liquid?

Molecules in the interior experience an attractive force from neighboring molecules which
surround on all sides
Molecules on the surface have neighboring molecules only on one side (the side facing the interior)
and thus experience an attractive force which tends to pull them into
the interior
45
The overall result of this asymmetric force on surface molecules is
that:
 The surface of the liquid will rearrange until the least
number of molecules are present on the surface

The surface molecules will pack somewhat closer together
than the rest of the molecules in the liquid
o
The surface molecules will be somewhat more
ordered and resistant to molecular disruptions
o
Thus, the surface will seem to have a “skin”
The “inward” molecular attraction forces, which must be overcome to increase the surface area, are
termed the “surface tension”

Liquids that have strong IMF also have high surface tensions. Thus, because water has
hydrogen bonding, it has considerable more surface tension than other liquids

Surface tension determines whether or not droplets of a liquid will bead up, as on a freshly
waxed car, or spread out when placed on a flat surface.
Cohesive forces bind molecules of the same type together(binding of like molecules)
Adhesive forces bind a substance to a surface (binding of unlike molecules)
For example, attractive forces (hydrogen bonding) exists between glass materials (silicon dioxide)
and water.

This is the basis of “capillary” action, where water can move up a thin capillary, against the
force of gravity. Surface tension “pulls” neighboring water molecules along.

The liquid climbs until the adhesive and cohesive forces are balance by the force of gravity
A. When adhesion is greater than
cohesion. The liquid (for example water)
rises in the capillary tube.
B. When cohesion is greater than
adhesion, as it is for mercury, a
depression of the liquid in the capillary
tube results. Note that the meniscus in
the tube of water is concave or rounded
downward, whereas that in the tube of
mercury is convex, or rounded upward.
46
Eg.
Which would have a higher surface tension H2O or C6H14? Why? Would the shape of the
H2O meniscus in a glass tube, be the same or different than C6H14?
Volatility
47
3.5 Structures and Physical Properties of Solids
Solids
 definite shape and volume
 vary in properties such as hardness, melting point, conductivity, characteristics
 fracture – the way solids break. May be curved (conchoidal) as in glass of flint or perfectly
flat.
 Are classified according to their chemical properties as ionic, metallic, molecular or covalent
network solids.
Ionic Solids

attraction of a cation to an anion to form the strongest attractive force known in chemistry

form very rigid lattices with large lattice energies

high melting and boiling points

regular crystal of alternating positive and negative charges.

Extremely common since they are very stable and unreactive. Most metal ions are found in
this form in nature.
Formation of Ionic Compounds
Lewis symbol equation
.....O .
Ca
2+
..O
..
2-
..
..
.Ca .
Balanced chemical equation
2 Ca(s) + O2(g)  2 CaO(s)
A Model for Ionic Compounds
The structure of most binary compounds such
as NaCl, KCl, and KI can be explained by the
close packing of spheres. Generally the larger
ion (usually the anion) is in one of the close
packing arrangements with the smaller cations
fitting within the holes between the closest
packed anions. The packing arrangement is
done in such a way as to minimize anionanion and cation-cation repulsions.
48
The strong attractions of the positive and negative charges hold the crystal rigidly together. Hitting a
metallic crystal with a hammer would give a completely different result compared to hitting an ionic
crystal with the same force. In a metal the atoms would shift their positions with no disruption of the
metallic bond. However in an ionic crystal, movement of the atoms by a little as one ionic diameter
will cause cations previously aligned with anions to be now aligned with other cations and anions to
be aligned with other anions. This would cause large repulsion forces, which would shatter the
crystal as shown below.
Illustration of why
crystals shatter.
Light circles are
cations and dark
circles are anions.
Metallic Crystals
Metals have the physical characteristics of malleability, ductility, conductors of heat and electricity,
durability, and high melting points. The simplest model that explains these observations is the
electron sea model.
For metals the atoms have low electronegativities; therefore the electrons are free to move over
(delocalized) over all the atoms. We can think of the structure of a metal as an arrangement of
positive atom cores in a sea of electrons
To maximize the bonding in a metal it makes sense to pack
as many atoms around each other as possible, maximize the
number of nearest neighbors The electrons are not bound to
any particular atom and are free to move about in the solid.
The mobile electrons can conduct heat and electricity and
the metal ions can easily be hammered into a sheet or pulled
into a wire.
49
Molecular Solids
In network solids, the solid can be considered as one giant molecule. In a molecular crystal, the
attractive forces between them are van der Waals forces and/or hydrogen bonding. Eg. Sulfur
dioxide (SO2), in which the predominant attractive force is a dipole-dipole interaction.
Intermolecular hydrogen bonding is mainly responsible for maintaining the three dimensional lattice
of ice. Some substances that hold molecular crystals with LDF are Ne, Xe, S8, CH4 etc
In general except ice, molecules in a molecular crystal are packed together as closely as their size
and shape will allow. Since the IMF holding these crystals together are very weak, relative to
covalent and ionic bonds, molecular crystals are more easily broken apart than ionic or covalent
crystals. (most melt below 100ºC)
Covalent Network Crystals

Eg. Diamond and quartz

Among the hardest materials on earth

Properties include hard, brittle, high melting points (higher than even ionic crystals).

Other examples include SiC(s) (silicon carbide) used for sandpaper and WC(s) (tungsten
carbide used for tips of saw blades and drill bit

Diamond, Analogy to Methane
In methane the bonding network of the carbon atoms is terminated by the H atoms. Because the H
atoms can only form one covalent bond, which is confined within the molecule, there is no
possibility for extended covalent bonding between different molecules. The inter-molecular bonds
must therefore rely on weaker dispersion forces for their cohesive forces.
In diamond, which can be constructed by replacing the H atoms in methane by other C atoms, the
bonding is not terminated and the C atoms are extended throughout the solid in a tetrahedral
geometry. These are strong bonds, as a result diamond has a very high melting (>3550°C) and
boiling (4827°C) point and is the hardest natural substance known...
50
Methane
Crystal Structure of Diamond
Silicon dioxide forms a covalent crystal, with each silicon atom forming bonds to four oxygen atoms
and each oxygen bonding to two silicon atoms with a tetrahedral geometry. Silicon carbide is
another network crystal similar to diamond with alternating tetrahedral silicon and carbon atoms. It
is very hard and is sometimes used as an industrial substitute for diamonds.
Benzene and Graphite.
The similarity between the bonding in methane and diamond allowed us to interpret the insulating
properties of diamond; in a similar way it is possible to interpret the electronic properties of
graphite. The properties of diamond and graphite are very different, diamond is hard, colorless, and
an insulator; graphite is a well known lubricant, black, and a conductor. These differences arise
from the different bonding and structure. The structure of graphite is based on layers of connected
six-member rings of carbon atoms - the structure of the layers is very similar to that of benzene, see
below. While the bonding in the layers is very strong, only weak van der Waals forces bond the
layers together (thus the effectiveness of graphite as a lubricant).
51
Benzene
Layered structure of Graphite
The slipperiness or lubricating qualities of graphite can be explained by noting that graphite has
very strong bonding within the layers but very little bonding between the layers giving a structure of
flat sheets. This arrangement allows the layer to slide past one another.
52