Chapter 8
Chemical Bonding
Bonds
•
•
Forces that hold groups of atoms together and
make them function as a unit.
We will consider three major categories of
chemical bonds: Ionic Bonds, Covalent Bonds,
and Metallic Bonds.
Examples of Bonding Types
Lewis Structures
• Since bonds hold atoms close together, then the valence
electrons are responsible for bonding since they are on
the outside of an atom.
• It has been recognized for a long time that the noble
gases have great chemical stability. With few exceptions
they are unreactive or inert.
• The noble gases have 8 valence electrons with the
exception of He which has 2.
Lewis Structures
He
Ne
Ar
Kr
Xe
1s2
1s22s22p6
1s22s22p63s23p6
1s22s22p63s23p64s23d104p6
1s22s22p63s23p64s23d104p65s24d105p6
Lewis Structures
The electronic configuration of the noble gases is described as
being energetically stable.
We can draw a Lewis diagram to illustrate the number of
valence electrons an atom has.
In a Lewis diagram valence electrons are represented by dots
placed above, below and to the left and right of the atoms
symbol.
e.g. element with 4 valence electrons
E
Lewis Structures
There are two simple rules to keep in mind when drawing
Lewis diagrams:
• Place one dot in each of the four locations before
doubling up.
• There can be only a maximum of 2 dots in any one
location.
E 
E 
E 
E 
Lewis Structures
What is the Lewis diagram for H?
1. First write the electron configuration:
1s1
2. Identify the number of valence electrons.
1 valence electron.
H
For a representative element it is easy to identify the number of
valence electrons as this is equal to the group number.
Lewis Structures
What is the Lewis diagram for S?
1. First write the electron configuration:
[Ne]3s23p4
2. Identify the number of valence electrons.
6 valence electrons
S
Alternatively you can recognize that S is in
group VIA so has six valence electrons
Lewis Structures
What is the Lewis diagram for S?
1. First write the electron configuration:
[Ne]3s23p4
2. Identify the number of valence electrons.
6 valence electrons
S
Alternatively you can recognize that S is in
group VIA so has six valence electrons
Lewis Structures
What is the Lewis diagram for S?
1. First write the electron configuration:
[Ne]3s23p4
2. Identify the number of valence electrons.
6 valence electrons
S
Alternatively you can recognize that S is in
group VIA so has six valence electrons
Lewis Structures
What is the Lewis diagram for S?
1. First write the electron configuration:
[Ne]3s23p4
2. Identify the number of valence electrons.
6 valence electrons
S
Alternatively you can recognize that S is in
group VIA so has six valence electrons
Lewis Structures
What is the Lewis diagram for S?
1. First write the electron configuration:
[Ne]3s23p4
2. Identify the number of valence electrons.
6 valence electrons
S
Alternatively you can recognize that S is in
group VIA so has six valence electrons
Lewis Structures
What is the Lewis diagram for S?
1. First write the electron configuration:
[Ne]3s23p4
2. Identify the number of valence electrons.
6 valence electrons
S
Alternatively you can recognize that S is in
group VIA so has six valence electrons
LEWIS STRUCTURES OF THE ELEMENTS
IA
IIA
SKIP
B’S
IIIA
IVA
VA
VIA
VIIA
VIIA
He
H
Be
B
C
N
O
F
Ne
Na Mg
Al
Si
P
S
Cl
Ar
Li
LEWIS STRUCTURES OF IONS
(AFTER REMOVAL OR ADDITION OF ELECTRONS)
IA
1+
IIA
H
2+
Li
Be
Na Mg
SKIP
B’S
IIIA
B
3+
Al
IVA
VA
VIA
VIIA
VIIA
4-
3-
2-
1-
He
C
N
O
F
Ne
Si
P
S
Cl
Ar
Lewis Structures
The octet rule states that:
“Atoms interact in order to obtain a stable octet of eight valence
electrons”
The octet rule works extremely well at describing the interactions
of the representative elements.
Lewis Structures
One way in which atoms can interact to satisfy the octet rule is by
transferring electrons between each other.
Transferring of electrons results in the atoms acquiring net
positive and negative charges.
When an atom loses or gains electrons a simple ion is formed.
Cations have more protons than electrons and are positive.
Anions have more electrons than protons and are negative.
Ionic Bonds
•
Formed from electrostatic attractions of
closely packed, oppositely charged ions.
•
Formed when an atom that easily loses
electrons reacts with one that has a high
electron affinity.
Ionization A Review
Consider a Na atom what happens if it loses one electron?
I.E.
Na+
Na
[Ne]3s1
1e-
+
[Ne]
11 P and 10 e-
11 P and 11 e-
Consider a Cl atom would you expect it to lose or gain
electrons?
Cl
[Ne]3s23p5
17 P and 17 e-
+
1e-
E.A.
Cl[Ne]3s23p6
17 P and 18 e-
Metals tend to lose electrons forming positively charged ions
called cations.
• A representative metal will lose its group number of
electrons to obtain a stable octet.
Na
→
Na+
Mg
→
Mg2+ +
+
1e- ( Isoelectronic with Ne)
2e-
(isoelectronic with Ne)
What would the charge be of the ion formed by a Li atom?
And which Noble gas is it isoelectronic with?
+1 The ion formed would be Li+
Isoelectronic with He
Noble Stability
Non-metals tend to gain electrons forming negatively charged
ions called anions.
• A representative non-metal will gain (8 - group number)
electrons to obtain a stable octet.
O +
2e-
→
O2- (isoelectronic with Ne)
S +
2e-
→
S2- (isoelectronic with Ar)
What would the charge be of the ion formed by a I atom?
Which Noble gas is it isoelectronic with?
-1 The ion formed would be IIsoelectronic with Xe
Lewis Structure of NaCl
Na+Cl-Na+Cl-Na+ClCl-Na+Cl-Na+Cl-Na+
Forces between oppositely charged ions are called
Ionic bonds. Each ion is surrounded by an octet of
Electrons, thus making the ions stable.
Crystal Lattice of NaCl
Ionic compounds do not exist as discrete molecules. Instead they
exist as crystals where ions of opposite charges occupy
positions known as lattice sites.
Ions combine in the ratio
that results in zero
charge to form ionic
compounds.
Which ions are the
smaller ones?
Crystal Lattice of NaCl
Crystal Lattice of NaCl
Ionic compounds do not exist as discrete molecules. Instead they
exist as crystals where ions of opposite charges occupy
positions known as lattice sites.
Ions combine in the ratio
that results in zero
charge to form ionic
compounds.
Which ions are the
smaller ones? Sodium
Crystal Lattice of NaCl
Sodium Chloride Lattice
Molecular Compounds
In our early lectures we defined a molecule as “as a compound
Made of nonmetals.”
Molecules exist as particles containing the number of atoms
specified by their formula.
e.g. a water molecule is a particle containing 2 hydrogen atoms
and one oxygen atom and has the formula H2O.
Molecular Compounds
Non-metals may also complete their octets by sharing
electrons.
This may occur between non-metal atoms of the same type:
e.g. H2, O2, N2, Cl2, F2, I2, etc
Or between different types of non-metal atoms:
e.g. CO2, H2O, CH4, etc
Covalent Bond Formation
-
-
+
+
Consider two hydrogen atoms separated by a large distance.
Each has 1 electron in a 1s atomic orbital.
Why does the electron stay around the nucleus?
Now lets bring the two atoms together so there orbitals overlap.
-
+
+
-
The atomic orbitals overlap to form a new molecular orbital.
This is a stable configuration as each H atom can have a full 1s
susbshell (like He) where the electrons spend most of their time
shared between the atoms. In this arrangement each nucleus
feels an inwards attraction to the two electrons. This is called
covalent bonding.
-
+
+
-
This new arrangement of protons and electrons is more stable
than separate hydrogen atoms since the attraction of a proton to
two electrons is a stronger attraction compared to one proton to
one electron of a hydrogen atom.
Lewis Structures
• A single bond results when two atoms share
one pair of electrons.
• A lone pair, or unshared pair, of electrons is a
pair of electrons that is not shared.
• A bonding pair of electrons is a pair of
electrons shared between two atoms.
Multiple Bonds
• A double bond results when two atoms share
two pairs of electrons.
• A triple bond results when two atoms share
three pairs of electrons
• Bond length is the distance between the
nuclear centers of the two atoms jointed
together in a bond.
Covalent Bond Formation
We can draw Lewis diagrams showing the
arrangement of valence electrons in covalent
compounds. In these diagrams we represent each pair
of electrons between atoms as a line.
So for the H2 molecule discussed previously the Lewis
diagram would be:
H–H
All other electrons are represented by dots as described
previously.
Molecular Compounds
Draw Lewis Structures of the following molecular
compounds
H
a. H2O
H
H
H O
O
Nonbonding
electons
Note each element has a
Noble gas structure by
electron sharing
b. NH3
H
N
H
HN H
H
H
Covalent bonding e’s
Simplified Lewis Structures
Straight lines are used to indicate a shared
pair, or a covalent bond.
H O
H
Nonbonding electrons
Lewis Structure Construction
Step 1
Step 2
Step 3
Step 4
Step 5
Connect each element with a single line
Use the “P” formula to determine the number of extra bonds
Insert the extra bonds, to make double or triple bonds.
Give each atom an octet of electrons, except hydrogen
Determine the formal charge of each element
P = 8(n-q) +2q - 2(n-1) - v
N = number of atoms in molecule
Q = number of hydrogen atoms
V = total number of valence electrons
Examples: Give Lewis Structures for the following
CO2
H2CO3
SO3
NO2+
Lewis Structure of Carbon Dioxide
First, connect atoms with lines
O
C
O
Second, use “p” formula to determine the number
of extra bonds.
P = 8(n-q) + 2q – 2(n-1) - v
P = 8(3-0) + 2(0) – 2(3-1) - 16
P = 24 + 0 – 4 - 16
P=4
4 extra bonding electrons
2 extra bonds
2 extra lines
Lewis Structure of Carbon Dioxide
Third, add extra lines
O
C
O
O
C
O
O
C
Fourth, give each atom an octet of electrons
O
Lewis Structure of Carbon Dioxide
Third, add extra lines
O
C
O
O
C
O
O
C
O
Fourth, give each atom an octet of electrons
O
C O
O C O
O C O
Lewis Structure of Carbon Dioxide
Third, add extra lines
O
C
O
O
C
O
O
O
C
Fourth, give each atom an octet of electrons
O
C O
O C O
O C O
Fifth, give each an atom a formal charge
If the element owns less than its valence, then it is positive
If the element has more than its valence, then it is negative
O
C
O
O C O
O
C
O
Lewis Structure of Carbon Dioxide
Third, add extra lines
O
C
O
O
C
O
O
O
C
Fourth, give each atom an octet of electrons
O
C O
O C O
O C O
Fifth, give each an atom a formal charge
If the element owns less than its valence, then it is positive
If the element has more than its valence, then it is negative
-
O
C
O
O C O
O
C
O
Lewis Structure of Carbon Dioxide
Third, add extra lines
O
C
O
O
C
O
O
O
C
Fourth, give each atom an octet of electrons
O
C O
O C O
O C O
Fifth, give each an atom a formal charge
-
If the element owns less than its valence, then it is positive
If the element has more than its valence, then it is negative
+
O
C
O
O C O
O
C
O
Lewis Structure of Carbon Dioxide
Third, add extra lines
O
C
O
O
C
O
O
O
C
Fourth, give each atom an octet of electrons
O
C O
O C O
O C O
Fifth, give each an atom a formal charge
-
If the element owns less than its valence, then it is positive
If the element has more than its valence, then it is negative
+
-
O
C
O
O C O
O
C
O
Lewis Structure of Carbon Dioxide
Third, add extra lines
O
C
O
O
C
O
O
O
C
Fourth, give each atom an octet of electrons
O
C O
O C O
O C O
Fifth, give each an atom a formal charge
-
If the element owns less than its valence, then it is positive
If the element has more than its valence, then it is negative
+
+
-
O
C
O
O C O
O
C
O
Practice
Draw the most stable Lewis structure for CO2.
Practice
Draw the most stable Lewis structure for CO2.
O C O
Sulfur Trioxide Lewis Structure
There are actually three possible Lewis structures for SO3.
2+
O
S
O
-
-
O
-
2+
O
O
S
O
-
-
O
2+
-
S
O
O
Each of these three structures is equivalent. We say they are in
“resonance” or that they are “resonance structures”.
Resonance Form Rules
Resonance Form Rules
Resonance Form Rules
Resonance Structures
Practice
Determine the most stable structure for the
phosphite ion by calculating formal charge for each
ion.
33O
O
O P O
O P O
-
Practice Problem
Give the most stable resonance form of H2SO4
Practice Problem
Give the most stable resonance form of H2SO4
Step #1 Connect atoms with single bonds in the most
symmetrical arrangement possible
O
H
O
S
O
O
H
Practice Problem
Give the most stable resonance form of H2SO4
Step #2 Find the number of extra bonding electrons with the
“P” formula P=8(7-2)+2(2)-2(7-1)-32=0 NO extra lines
O
H
O
S
O
O
H
Practice Problem
Give the most stable resonance form of H2SO4
Step #3 Give each atom an octet of electrons
P=8(7-2)+2(2)-2(7-1)-32=0 NO extra lines
O
H
O
S
O
O
H
Practice Problem
Give the most stable resonance form of H2SO4
Step #4 Give each atom a formal charge
P=8(7-2)+2(2)-2(7-1)-32=0 NO extra lines
-
O
H
O
+
S
+
O
O
-
H
Practice Problem
Give the most stable resonance form of H2SO4
Step #4 Give each atom a formal charge
P=8(7-2)+2(2)-2(7-1)-32=0 NO extra lines
-
O
H
O
+
S
+
O
O
H
According to resonance rules this should not be very stable
Practice Problem
Give the most stable resonance form of H2SO4
Step #4 Give each atom a formal charge
P=8(7-2)+2(2)-2(7-1)-32=0 NO extra lines
-
O
H
O
+
S
+
O
O
H
-
Moving electrons around might improve stability
Practice Problem
Give the most stable resonance form of H2SO4
Step #4 Give each atom a formal charge
P=8(7-2)+2(2)-2(7-1)-32=0 NO extra lines
-
O
H
O
+
S
+
O
O
H
-
Moving electrons around might improve stability
Practice Problem
Give the most stable resonance form of H2SO4
Step #4 Give each atom a formal charge
P=8(7-2)+2(2)-2(7-1)-32=0 NO extra lines
-
O
H
O
S
+
O
H
O
This arrangement is better, but still not the best.
Practice Problem
Give the most stable resonance form of H2SO4
Step #4 Give each atom a formal charge
P=8(7-2)+2(2)-2(7-1)-32=0 NO extra lines
-
O
H
O
S
+
O
H
O
This arrangement is better, but still not the best.
Practice Problem
Give the most stable resonance form of H2SO4
Step #4 Give each atom a formal charge
P=8(7-2)+2(2)-2(7-1)-32=0 NO extra lines
O
H
O
S
O
H
O
Since sulfur is a period three element, then it can accommodate
more than eight electrons. This is the most stable arrangement.
Exceptions to the Octet Rule
Some molecules have less than eight electrons in a
Lewis structure.
BF3 is an example
These are known as electron-deficient compounds
Other molecules have more than an octet.
SF6 is an example
Some molecules have an odd number of electrons
when summing up the total number of electrons. These
don’t obey the octet rule.
NO is an example
Free radicals are molecules having an odd number of
valence electrons.
Example
What is the Lewis dot structure of SO32- ion? Does
this ion have resonance structures?
Metallic Bonds
A metallic bond consists of the nuclei of metal
atoms surrounded by a “sea” of evenly spaced
shared electrons. The nuclei is then attracted
to each electron by the same amount, but in
different directions, thus making the nuclei stay
in a fixed position.
Electron Sea Model
Electronegativity
• Electronegativity is a measure of an element’s ability
to attract bonding electrons. It is used for predicting
the degree to which bonding pairs of electrons are
shared unequally.
Electronegativity Trends
Atomic Size Relationship
• Changes in electronegativity are related to
increasing atomic size.
• The size of the valence orbitals increases as
the value of the principle quantum number
(n) increases.
• Therefore, the atomic size increases and
electronegativity decreases as you go down
a group.
Bond Types
Bond Polarity
Bond polarity is a measure of the extent to which
bonding electrons are shared between two atoms in
a covalent bond.
Polar Covalent Bonds
• Compounds that contain two or more different
elements may contain polar covalent bonds.
:C O:
• A polar molecule contains bonds that have an uneven
distribution of charge because electrons in the bonds
are not shared equally by the two atoms.
Ionization vs. Electronegativities
Practice
Which of the following bonds in each pair are
more polar?
C-S or C-O
Cl-Cl or O=O
N-H or C-H
Different Elemental Forms
• Oxygen is found in two forms: oxygen, O2,
and ozone, O3.
• Carbon is found in many different forms:
soot, diamond, graphite, nanotubes, etc.
• These different forms of an element are
known as allotropes.
Evidence for Resonance Structures
• Ozone bond lengths were both found to be
128pm
• A single bond would have a longer bond
length than a double bond, but in ozone we
find the bond lengths are equal meaning they
are a mixture of single and double bonds.
Bond Lengths
Bond length depends on the identity of the atoms as
well as on the number of bonds between them.
The bond order is the number of bonds between two
atoms.
1 for a single bond
2 for a double bond
3 for a triple bond
Covalent Bond Lengths
Examples of Bond Length
Bond Strengths
Bond Energies
• The energy needed to break 1 mole of
covalent bonds in the gas phase is the bond
energy of that bond.
• Breaking bonds consumes energy whereas
forming bonds releases energy.
• Bond energies can be used to estimate Hrxn.
Hrxn = Hbond breaking + Hbond forming
Energy of Reaction Calculation
Calculation ΔH for the following reaction.
CH4 + O2
CO2 + HOH
Energy of Reaction Calculation
Calculation ΔH for the following reaction.
CH4 + 2 O2
CO2 + 2 HOH
First energy must be added to break the reactant
bonds. Bonds to break are 4 C-H bonds and 2 O=O
double bond. Next energy is released when 2 C=O
bonds are formed and 4 H-O bonds are formed.
Energy of Reaction Calculation
Calculation ΔH for the following reaction.
CH4 + 2 O2
CO2 + 2 HOH
First energy must be added to break the reactant
bonds. Bonds to break are 4 C-H bonds and 2 O=O
double bond. Next energy is released when 2 C=O
bonds are formed and 4 H-O bonds are formed.
Bonds broken (endothermic)
4 C-H bonds 4(414) = 1656 kj
Bonds formed (exothermic)
2 C=O 2(799) = 1598 kj
2 O=O bonds 2(498) = 996 kj
2652 kj
4 H-O 4(464) = 1856 kj
3452 kj
ΔH = 2652 -3452 = 800 kj
Example
The End
Review Questions
ChemTour: Bonding
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This ChemTour shows how ionic and covalent bonds form.
ChemTour: Lewis Dot Structures
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Students learn to draw and use Lewis dot structures to
visualize and represent molecular structures and the
locations of valence electrons. Includes Practice Exercises.
ChemTour: The Periodic Table
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This ChemTour offers a guided tour of the trends
summarized by the periodic table (metallic properties,
subshells, electronegativity, and atomic radius), and
explains how to use this tool to predict an element’s
characteristics, including its bonding capacity.
ChemTour: Resonance
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Molecules that cannot be represented by a single Lewis dot
structure are said to exhibit resonance. This ChemTour
describes resonance and explains how to represent it
visually.
ChemTour: Expanded Valence
Shells
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In this ChemTour, students explore the exceptions to the
octet rule, and learn to identify the conditions under which
an element will expand its outer electron shell to hold more
than 8 electrons. Includes Practice Exercises.
ChemTour: Estimating Enthalpy
Changes
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In this ChemTour, students learn how to use average bond
energies to estimate the energy released during a
combustion reaction. Includes Practice Exercises.
The plot to the left shows the
molecular potential energy curve for
N2. Which of the plots below shows
the correct molecular potential energy
curve for O2 (blue solid line)
compared to N2 (red dashed line)?
A)
B)
Molecular Potential Energy of N and O
C)
Consider the following arguments for each answer
and vote again:
A. The O2 double bond is less stable than the N2 triple
bond, so the dissociation energy of O2 should be less
than that of N2.
B. An oxygen atom is smaller than a nitrogen atom, so
O2 should have a smaller bond length than N2.
C. The fact that O2 has two electrons in antibonding Π*
orbitals is reflected in the potential energy curve
always being above zero.
Molecular Potential Energy of N and O
The oxygen molecule, O2, has a bond order of 2 in its
ground state.
How many electrons can be removed from O2 without
altering the bond order? Assume the electrons are always
removed from the highest-occupied molecular orbital.
A) 2
Bond Order of Oxygen
B) 4
C) 6
Consider the following arguments for each answer
and vote again:
A. Removing 2 electrons from antibonding orbitals will
not affect the bond order.
B. To retain a bond order of 2, a total of 2 electrons must
be removed from a bonding orbital and 2 must be
removed from an antibonding orbital.
C. Removing 6 electrons will leave 4 electrons in 2
bonding orbitals.
Bond Order of Oxygen
Which of the following three molecules would be made
less stable by removing a single electron from the
highest-occupied molecular orbital?
A) He2
Stability of He , CN, and NO
B) NO
C) CN
Consider the following arguments for each answer
and vote again:
A. The 4 electrons in He2 are incapable of holding 2
helium atoms together, so removing 1 electron will
only make matters worse.
B. Removing an electron from NO would decrease its
bond order from 2.5 to 2, thus decreasing its
stability.
C. The highest-occupied molecular orbital for CN is a
bonding orbital, so removing an electron will
decrease CN's stability.
Stability of He , CN, and NO
Lewis electron dot structures depict the arrangement of
valence electrons around the atoms of a molecule, as
shown below for water, H2O.
Which of the following is the correct Lewis electron dot
structure for formaldehyde, CH2O?
A)
B)
Lewis Electron Dot Structure of Formaldehyde
C)
Please consider the following arguments for each
answer and vote again:
A. The octet rule is satisfied, and the formal charges of
all of the atoms are zero.
B. The carbon-oxygen triple bond lends extra stability to
this structure.
C. The octet rule is satisfied without the need to form
double bonds.
Lewis Electron Dot Structure of Formaldehyde
How many bonds would it take to
complete the structure of the nitrite
ion, pictured to the left?
A) 0
Bonding Structure of NO -
B) 1
C) 2
Please consider the following arguments for each
answer and vote again:
A. All of the atoms are attached, so more bonds are not
needed.
B. The correct overall charge and the octet rule are
satisfied by 1 single bond and 1 double bond.
C. Both oxygens are double bonded to the nitrogen in
order for their formal charges to be zero.
Bonding Structure of NO -
Which of the following depicts a
molecule that is different from the
one shown to the left?
A)
B)
Molecular Structures of Pentane
C)
Please consider the following arguments for each
answer and vote again:
A. This molecule, known as isopentane, has a unique
T-shaped arrangement.
B. In this molecule, only two carbon atoms are bonded
to the second carbon.
C. This molecule is unbranched, whereas the one
pictured in the question is not.
Molecular Structures of Pentane