Chapter 1: Chemical Bonding
1.1 Forming and Representing Compounds
A. The Basics
scientists have studied the way elements and
compounds appear in nature in order to
categorize chemical bonding
 most metals are combined with non-metals in
nature (called ores ) and are solids
 a few metals are found in their pure form
…precious metals
 metals (except Hg) in pure
form are solids
 non-metals combine
with one another to form
solids, liquids or
gases
 the only elements that are never found in
combined form in nature are the noble gases
 atoms gain, lose or share electrons in such a way
that they create a full outer energy level…
called the octet rule
the first energy level can only hold two e
and therefore satisfies the octet rule when it has
two e in it
octet rule is a guideline…not all elements follow it
at all times
 valence e are the electrons in the outermost
energy level of an atom
they are the only electrons involved in
chemical bonding
for representative elements ( groups 1,2 and 13-18)
group number (ignore the “1” in front of groups above
10) tells you the number of valence electrons
period number tells you the number of
energy levels occupied by electrons
for many transition metals , the number of valence
electrons is not as predictable… it depends on the
environment around the ion
eg) iron can be Fe3+ or Fe2+
the ion charge can be used to determine number
of valence e
eg) Fe3+ had 3 valence electrons
B. Electron Dot Diagrams
you can’t see atoms and electrons, therefore it is
convenient to draw models
to show the
structure and formation of chemical bonds
an electron dot diagram is one such model
consists of the symbol for the element with
dots representing the valence e
when drawing the diagrams, look up the number
of valence e , then place dots around the
symbol clockwise for a maximum of four dots
if you have more electrons to place, go back to the
top of the symbol and start pairing up the e

Na


P


Ca





O 

Al


 Cl 

 Si 




Ar


a full orbital is called a lone pair and is
not involved in bonding (at this level)
a half full orbital contains a bonding electron
bonding e


O 

lone pairs
the bonding capacity of an atom is the
maximum number of single covalent bonds
that it can form (equals the number of bonding e )
Try These
Draw the Lewis diagram (electron dot diagram) for each
of the following:
H
C
Mg
He
F
K
Be
S
Br
P
C. Ionic Bonding
an ionic bond is the electrostatic attraction
between oppositely charged ions
most
metals have three or fewer valence e
they tend to lose these electrons and become
positive ions (cations)
+
Na
Na
most non-metals have more than four
valence e
they tend to gain electrons and become
negative ions (anions)
2-
O
O
 after ions form, the attraction between the
positive charge and negative charge draws the
ions together, forming an ionic bond
when drawing the electron dot diagrams for ionic
compounds:
 the number of electrons lost by the metal
must equal the number of electrons gained
by the non-metal
 the net charge on the compound must be
zero
 you may have to have more than one of
the metal and/or non-metal to balance
out the charges
Examples
Na
Cl
NaCl
Mg
O
MgO
Examples
F
Ca
F
CaF2
K
S
K
K2S
Fe
O
O
Fe
O
Mg
Fe2O3
N
Mg
Mg
N
Mg3N2
notice the following about the diagrams:
 the metal has no valence electrons
(since they lose them)
 the non-metal has the valence level filled
 both ions have square brackets and the
charge
 positive charges = negative charges
D. Covalent Bonding
a covalent bond is formed when two nonmetals share a pair of electrons
compounds containing covalent bonds
called molecular compounds
 ions are not formed!!!
electron dot diagrams used to
show molecular compounds
are called Lewis structures
are
instead of transferring electrons, valence electrons are
now shared to satisfy the octet rule
the electrons that are shared are called a bonding
pair
sharing two or three pairs of electrons between two
atoms results in a double or triple bond,
respectively
to draw the structures:
 place the atom with the most bonding
electrons in the centre
 arrange all other atoms around it as
symmetrically as possible
 share electrons to make sure that all atoms
have the octet rule satisfied (remember that
hydrogen only needs two electrons to be
satisfied)
eg) PH3
H
H
H
P
H
H P
H
Try These
Draw the Lewis diagram (electron dot diagram) for
each of the following:
1. HCl
4. NBr3
2. CH4
5. C2H4
3. F2
6. N2
Br
H Cl
H
H C H
H
F
F
H
C
H
N
N
Br
C
Br
H
H
N
E. Structural Formulas
a structural formula (diagram) is another way
of drawing molecules
to draw them, figure out the Lewis structure then
replace all shared pairs of e with a line and
leave off the lone pairs
eg) PH3
Lewis Diagram
H P H
H
Structural Diagram
H P
H
H
Try These
Draw the structural formula for each of the following:
1. HCl
4. NBr3
2. CH4
5. C2H4
3. F2
6. N2
H
Cl
Br
N
Br
H
H
H
C
H
F
Br
H
H
F
H
C
C
H
N
N
F. Metallic Bonding
most metals are solids at
room temperature which
means that there must be
strong attractive
forces holding the atoms
of a pure metal together
metals DO NOT form covalent or ionic bonds
with other metal atoms
in metallic bonding all the atoms share all
the valence e
the valence electrons are delocalized , which means
they are free to move from one atom to another
metallic bonds are made up of a network
of positive metal ions in a “sea” of electrons
a metallic bond is the electrostatic force of
attraction between the positive metal ions
and the negative sea of electrons
this theory helps explain the properties of
metals
eg) good conductors of electricity and heat,
ductility, malleability
Metallic Bond Model
metal
cations
“sea” of
delocalized
electrons
1.2 The Nature of Chemical Bonds
A. Electronegativity
the electronegativity of an element is the relative
measure of the ability of an atom to attract
electrons in a chemical bond
there is an attraction between the nucleus (protons)
of an atom and the valence e in an adjacent
atom
nucleus
electrons
each element is designated a number to represent
how strong it’s nucleus is at attracting
another atom’s valence e
 higher electronegativity means greater attraction
(affinity)
trend on periodic table – electronegativity
decreases down group and increases across
period
since noble gases do not readily react with other
substances, electronegativities have not been
assigned to them
understanding electronegativity has contributed to
the knowledge of bonding in ionic and molecular
compounds
Electronegativity and the Periodic Table
decreases
increases
B. Size & Electronegativity
as you move from left to right across a period, both the
electronegativity, and atomic number increase
however size of the atom decreases
here is why size decreases across a period:
 the size of an atom depends on the radius of
the energy level containing the valence e
 in any given period, the valence e of each atom
occupy the same energy level
 as you move across the period, the atomic
number increases and thus the number of
protons in the nucleus increases
 there is a greater amount of attraction
between the nucleus and e when there
are more protons , therefore the atom is
smaller
Period 2 Elements
3 p+
1 valence e
Li
6 p+
9 p+
4 valence e
7 valence e
C
F
so, the next question is “why does electronegativity
increase when atomic size across a period
decreases?”
the strength of the attraction (and therefore
electronegativity) between oppositely charged
particles depends on two factors:
 the distance between the charges – the
attractive force between opposite charges
decreases with the square of the distance
between them
 the magnitude of the charges – the attractive
force is directly proportional to the
amount of charge
this means that an atom that is small and has
lots of protons (like fluorine) will have a
very large amount of electrostatic attraction
(electronegativity) for the e of another atom
big atoms have lots of protons but they are
shielded by the inner levels of e therefore
have a small amount of attraction
(electronegativity) for the e of another atom
Cs

nucleus of
cesium
valence
electrons
of silicon
F

Si
distance
between
nucleus of
cesium and
valence
electrons of
silicon
nucleus of
fluorine
valence
electrons of
silicon
Si
distance
between
nucleus of
fluorine
and valence
electrons of
silicon
C. Bond Type & Electronegativity
electronegativities can be related to bond types:
ionic bonds occur between metals and non-metals
 metals have low electronegativities and will
lose e
while non-metals have high
electronegativities and will gain e
 the two ions that are formed will attract each
other and form a chemical bond
covalent bonds occurs between non-metallic atoms
 if you look at two atoms that
have the same
electronegativity, like in H2(g),
the two nuclei of the atoms will
attract the electrons
with exactly the same strength
 the electrons are shared
equally between the two
atoms
 when two non-metals that have
different electronegativities
share electrons, the sharing is
no longer equal
 the element with the higher
electronegativity pulls the
e closer to itself
 this results in one end of the bond having a
slightly negative charge () and the other
end of the bond having a slightly positive charge
(+)
+

 bonds that have unequal sharing of electrons
are called polar covalent bonds
 also called bond dipoles since the bonds have
oppositely charged ends
Bond Dipole Arrows
“+” at the end that is +
+
H–F
“arrow” points
towards element with
higher
electronegativity (-)
-
Try These:
Draw the bond dipole arrow, label the + and  ends,
and state the bond type (polar, nonpolar, ionic)
0.4
1.
H – H nonpolar
polar
0.8
2. - N – H +
3. +B – F -
7.
polar
8. +Si – Cl - polar
0.8
4.+ S – O P–H
Cl – Cl
polar
2.0
5.
6. - C – H +
1.3
nonpolar
1.2
polar
nonpolar
9. - O – H + polar
10.
Na – Cl
ionic
you can use the difference in electronegativity
between two atoms to determine bond character
Difference in Electronegativity
3.3
1.7
mostly ionic
0.5 0
polar
covalent
slightly
polar
covalent
non-polar
covalent
as you can see, bond classification is not
simple
bonding is considered a continuum and there
is no clear distinction between ionic and
covalent bonding
Chapter 2: Chemical Bonding
2.1 Three Dimensional Structures
A. Ionic Crystals
ionic compounds have crystal structure
they form so that oppositely charged ions are
as close together as possible
this 3-D array of alternating positive and
negative ions is called a crystal lattice
since all the attractive forces in the lattice are
the same, you cannot call them molecules…
each positive ion is attracted to all of the negative
ions around it (and vice versa)
the chemical formula is the lowest whole
number ratio for that type of crystal
eg) NaCl has a 1:1 ratio of
Na ions to Cl ions
sodium chloride
there are many different crystal shapes and
they all depend on the way the ions pack together
shape also depends on the relative size of the
ions and the charges on the ions
B. Structure of Molecules
in molecular compounds, covalent bonds exist
between specific pairs of atoms
these compounds exist as molecules that have a
given number of atoms and therefore are not
necessarily written with the lowest whole
number ratio
C. VSEPR & Structure of Molecular Compounds
the valence shell electron pair repulsion
(VSEPR) theory states that molecules adjust
their shapes so that valence e- are as far away
from each other as possible
electron pair repulsion is not always equal…
it is greatest between two lone pairs (LP),
less between a LP and a bonding pair (BP),
and lowest between two BP’s
shape is determined around the central atom
shapes can be classified into five categories:
1. linear – central atom is bonded to two other
atoms and has zero lone pairs, or there is
only two atoms in the molecule
eg) CO2(g), HCN(g), HCl(g)



O


C
O 
H
H

Cl



C
N


2. trigonal planar – central atom is bonded to
three other atoms and has zero lone pairs
eg) CH2O(l)
H
H
C

O



3. tetrahedral – central atom is bonded to four
other atoms and has zero lone pairs
H
eg) CH (g)
4
C
H
H
H
4. pyramidal – central atom is bonded to three
other atoms and has one lone pair
eg) NH3(g)
 
N
H
H
H
5. bent – central atom is bonded to two other
atoms and has either one or two lone
pairs
eg) H2O(l), HNO(g)
H
H

O
 
H
O
N

we can use a code to determine the shape of a
molecule around the central atom
the code has two numbers:
1. the number of atoms attached to the central atom
2. the number of lone pairs on the central atom
CH4
eg) NH3(g)
 
H
N
H
H
H
3-1
pyramidal
C
H
H
H
4-0
tetrahedral
Code
Shape
Example
4–0
tetrahedral
CH4
3 – 0 trigonal planar
CH2O
3–1
pyramidal
NH3
2–1
bent
HNO
2–2
bent
H2O
***all other codes are linear
D. Polar Bonds & Polar Molecules
a molecule that contains polar covalent bonds
can be overall nonpolar
the individual bond dipoles are vectors that
can be added to each other
if the bond dipoles are equal in strength and
opposite in direction , they cancel each other
out resulting in a nonpolar molecule
this canceling happens in symmetrical molecules
if the bond dipoles do not cancel , the entire
molecule will have a slightly positive and
slightly negative end… called dipoles
general rules:
 tetrahedral: nonpolar if all atoms attached
have the same pull (in or out), polar if
different atoms attached
 trigonal planar: nonpolar if all atoms
attached have the same pull (in or out), polar
if different atoms attached
 pyramidal: polar as long as there is a
difference in electronegativity between the
atoms
 bent: polar
 linear: polar or nonpolar …look at
electronegativity difference
Examples
1. H2O
2. HCl
O
H
H
H
polar
polar
3. C2H2
4. C2HI
np
np
H
C
Cl
C
nonpolar
H
H
C
C
polar
I
Journal #6

In what ways is a covalent bond like
running in a relay race?
Try These
1. HF
H
2. CH4
F
C
polar
3. N2
H
H
H
H
nonpolar
np
N
N
4. PI3
nonpolar
H
5. C2H6
H
C
H
np
C
H
H
nonpolar
H
P
I
I
polar
I
2.2 Intermolecular Forces
A. Types of Forces
 intramolecular forces are the forces of attraction
within molecules
(eg. ionic or covalent bonding)
 intermolecular forces are the forces of attraction
between molecules
they are the weakest of all forces
 responsible for state, melting point, boiling
point etc.
there are three types of intermolecular forces that
we will look at:
1. Dipole-Dipole Forces
 electrostatic force of attraction between the
dipoles of polar molecules
 slightly negative “poles” attract the slightly
positive “poles” in other molecules and vice versa
-
+
+
+
-
+
+
2. Hydrogen Bonding
this is a special type of dipole-dipole interaction that
is very strong
hydrogen bonding is the attraction between a
hydrogen on one molecule which is bonded
to O, F or N, to the O, F or N of an adjacent
molecule
when hydrogen is bonded to a highly
electronegative element such as O, F or N, the
electrons are pulled far away from it
since hydrogen doesn’t have any other electrons,
its proton is basically exposed


this proton is then able to be attracted not only to the
 pole but also to the lone pairs
H

O

H
3. London (Dispersion) Forces
the attractive force that occurs between all
molecules is called London Dispersion force
the result of the electrostatic attraction of
induced dipoles
electrons in atoms and molecules are always in
constant rapid motion
for brief instances, the distribution of electrons
becomes distorted which produces very weak
dipoles
this temporary dipole induces a dipole in the
adjacent molecule when like charges repel
each other
this process “disperses” throughout the
substance, causing flickering dipoles that
attract each other
even though this force lasts only a moment and
is very weak , the overall effect in a substance is
significant
LD forces are affected by two factors:
1. size of the atoms –
more e means higher probability of creating
temporary dipoles
2. shape of the molecule –
the more contact between molecules, the
higher the force of attraction
Journal #7
If you could ask a question about a “lone
pair”, what would the question be?
Scale of Forces
very
high
very
low
LD
DD
HB
covalent
Intermolecular Forces
(between)
London Dispersion
Dipole – Dipole
Hydrogen Bonding
ionic
network
covalent
Intramolecular Forces
(within)
metallic ** wide range
ionic
covalent
network covalent
eg) diamond, SiC, SiO2
2.3 Relating Structures and Properties
A. States of Matter
the state of a substance depends on the strength
of the attractive forces between its particles
solids have the greatest forces of attraction, liquids
have the next greatest and gases have very few
if any intermolecular attractions between the
particles
B. Melting & Boiling Points
both melting and boiling points are indicators as to
the strength of attractions within or between
molecules
when metals melt or boil, metallic bonds must
be broken
when ionic compounds melt or boil, ionic
bonds
must be broken
when molecular compounds melt or boil,
only intermolecular forces must be broken
(covalent bonds DO NOT break)
these forces are much weaker therefore it takes a
lot less energy to separate the molecules
as you increase the number of intermolecular
forces, the melting and boiling points increase
Order of bp’s
Using the scale of forces you can order
compounds based on their relative bp’s
 Ex. From Highest to Lowest






Network covalent compound (ex. SiO2)
Ionic compound
Molecular compound with HB, DD, LD
Molecular compound with DD, LD
Molecular compound with LD (if 2 molecular
compounds have LD only then bigger molecule or
molecule with more electrons has higher bp)
Strongest to weakest, highest
b.p. to lowest b.p.






Cats
In
My
House
Don’t
Live
= covalent network
= ionic
= metallic
= hydrogen bonds
= dipole-dipole
= london dispersion
Determining the heat of bonding



The amount of energy that is stored in a liquid state is
called the enthalpy of fusion
The amount of energy that is stored in the gaseous
state is called the enthalpy of vaporization.
To calculate enthalpy we use






DH = mHx or DH = nHx
DH = energy (J)
m = mass (g)
n = moles (mol)
Hx = Hf = enthalpy of fusion (J/mol)
= Hvap = enthalpy of vaporization (J/mol)
Example
How much heat is required to melt 30 g of
water, if the enthalpy of fusion is 333 J/g?
DH = mHf
= (30 g )(333J/g)
= 9990 J
= 10 kJ
Specific Heat Capacity







The amount of energy stored in intermolecular bonds
can cause substances to change temperature slowly or
quickly.
The amount of energy required to change the
temperature of 1 gram of a substance 1 degree Celsius
is called the specific heat capacity.
To calculate the energy required to change the
temperature of a substance use:
H = mcDt
H = energy (J)
m = mass (g)
c = specific heat capacity (J/goC)
Dt = change in temperature (o C)
Example
What would be the final temperature of 300
mL of water, at room temperature, that has
4.00 kJ applied to it?
H = mcDt
4000 J = (300 g)(4.19 J/goC)Dt
3.18 = tf – 20oC
Dt = 3.18oC
tf = 23oC
C. Mechanical Properties of Solids
the mechanical properties of solids are determined by
the types of bonds in the substance
delocalized e cause bonds to be non-directional,
which allows a solid to be malleable and ductile
eg) metals
solids that do not have delocalized e have
directional bonds, which causes them to be brittle
and hard
eg) ionic compounds
D. Conductivity
electric current is the directional flow of
electrons or ions
metals are good conductors of electricity because
the delocalized valence e are free to move
solid ionic compounds have valence electrons
that are held solidly in place therefore they
cannot conduct electricity
when ionic compounds melt or dissolve in water,
the ions are able to move past one another
which allows them to carry an electric current
in most molecular compounds, valence electrons
are not free to move through the molecule
therefore they are not able to conduct electricity
when molecular compounds melt or dissolve in
water, they do not form ions and therefore they
do not carry an electric current
How does a gecko climb?