We will now relate the

concepts learned in previous

chapters to chemical bonding.
A strong attractive force that exists between
atoms in a molecule.
Three Types of Bonds
 We will focus mainly on ionic and
covalent bonds in this chapter. But
first, let’s discuss Lewis Symbols.

A diagram in which the elemental symbol represents

the nucleus along with all inner electrons, and the
dots represent outer shell electrons.
Lewis symbols are very useful because they portray the
heart of a chemical bond; valence electrons!

Let’s take a look at a few examples:

Bohr Diagram Lewis Diagram

Dot’s represent valence electrons:

**Notice the negative sign on the chloride ion.**

Don’t forget to include the charge when writing
the Lewis symbols for ions.
What would the Lewis symbols be for
elemental sodium and the sodium ion?
 Atoms gain, share, or lose electrons with the goal of
achieving the electron configuration of their neighboring
noble gas. They do so because as we know, the noble gases
are stable. All elements want to be like the NOBLE gases.

How many valence electrons do the

noble gases have?
*So, atoms gain, lose, or share electrons until*
they are surrounded by 8 valence electrons.

--Aren’t there always exceptions!--

These nasty little buggers will show
themselves later! For now, stick to the rule!
 It’s finally time to start talking about
chemical bonding. We will start with:

Let’s start by viewing a movie of a chemical

reaction that is the result of ionic bonding!
Click the icon below.
Sodium chloride is composed of a crystal lattice of
sodium cations (Na+) and chloride anions (Cl-).
Take a closer look at the 3 dimensional array of NaCl.

Display Model in Class

 Can you relate this reaction to the
periodic trends that you learned?
Sodium metal is on the far left of the periodic
table; therefore it has a low ionization energy.
Remember, this is the energy required to
remove an electron. Since the required energy
is so low, sodium is willing to give up an
electron; and thus, become positively charged.

Chlorine, on the other hand, is on the far right

of the periodic table, and therefore has a high
affinity for electrons. It is willing to take the
electron given up by the sodium atom, which
causes it to become negatively charged.
 Let’s look at this from a Lewis symbol perspective.


 
 Cl   Na      
Na       Cl 
    

Notice how the electron is transferred from

the sodium atom to the chlorine atom
resulting in oppositely charged particles.
 The oppositely charged particles that result
are strongly attracted to one another just like
the opposite ends of a magnet.
What is the chemical formula for:
Magnesium Fluoride?

MgF2
 MgF2
Magnesium has two electrons to give, whereas the
fluorines have only one “vacancy” each.
: :
: F . . Mg.

: :
. F:

- 2+ -
: :

: :
[: F: ] Mg [: F: ]
• Consequently, magnesium can accommodate
two fluoride ions.
The electrostatic forces that exist between ions of
opposite charge as a result of the transfer of one
or more electrons from one atom to another.
 Do you think energy would be required to break the
ionic bonds that form the sodium chloride crystal
lattice, or do you think energy would be released if the
bonds of sodium chloride were broken?
 Energy required to completely separate a mole of
a solid ionic compound into its gaseous ions.

NaCl (s)  Na+ (g) + Cl- (g) ΔHlattice = +788 KJ/mol

Energy Involved in Ionic Bonding
• The transfer of an electron from a sodium atom to a
chlorine atom is not in itself energetically favorable;
it requires 147 kJ/mol of energy.
• 496 KJ/mol needed to remove an electron from Na.
• 349 KJ/mol released when Cl gains the electron.
• 496-349 = 147 KJ/mol for the transfer of the
electron.
However, 493 kJ of energy is released when
these oppositely charged ions come together.
An additional 293 kJ of energy is released when
the ion pairs solidify.
The magnitude of the lattice
energy of a solid depends on:
• Charges of ions
• Size of ions
• Type of arrangement
 Q1Q2
E k
d
We won’t be doing any calculations with this
equation, but it does allow us to compare lattice
energies of different ionic compounds.

Q1 and Q2 = charges of particles

d = distance between centers
k = 8.99 x 109 J-m/C2
 Q1Q2
E k
d
This equation tells us that as the charges increase,
and the radii decreases, the lattice energy increases.

Compare the following in terms of

increasing lattice energy:
LiCl, LiF, MgCl2
 LiCl and LiF both have 1+ and 1- charges. Therefore, we
must consider the sizes of the ions.

F- is smaller than Cl-

So the distance between Li+ and F- in the crystal structure will
be smaller than the distance between the Li+ and Cl- ions.

Therefore more energy is needed to separate LiF

LiCl < LiF

The fact that magnesium has a 2+ charge indicates that MgCl2
will have the largest lattice of all three.

LiCl < LiF < MgCl2

 LiCl < LiF < MgCl2

Actual Lattice Energies

MgCl2 = 2326 KJ/mol
LiF = 1030 KJ/mol
LiCl = 834 KJ/mol
 Review: Sizes of Ions
Cations
smaller than parent atom because the loss of electron(s)
increases the effective nuclear charge and decreases
electron-electron repulsion.
Anions
larger than parent atom because the gain of electron(s)
increases electron-electron repulsion causing electrons
to spread out.
 Review: Isoelectronic Series
A group of ions that have the same number of electrons.

Example
Increasing nuclear charge

O2- F- Na+ Mg2+ Al3+

Decreasing ionic radius

Electrons remain constant, while protons increase
At room temperature, are
these two compounds Carbon monoxide
solids, liquids, or gases? CO

What can we conclude

about the amount of heat
needed to boil CO and
CO2? Carbon dioxide
CO2
 (Molecular Compounds)

*A bond formed between two or more atoms

as a result of sharing of electrons.

*Consists entirely of nonmetals.

*Low Melting Points because the bonds

between molecules are not as strong as the
bonds between ions in crystal lattices.

*Can be polar or nonpolar, which we will

discuss shortly.
Covalent bonding explains why
diatomic molecules exist.
H2, N2, O2, F2, Cl2, I2, Br2
 3 Types of Covalent Bonds
Single covalent bond – one pair of electrons are shared

Double covalent bond – two pairs of electrons are shared

Triple covalent bond – three pairs of electrons are shared

Again, we start with Lewis Symbols!
 Orbital Diagram

H 1S _____
= H2
H 1S _____
Which rule that you previously learned allows 1
electron from each atom to share the same orbital?
The electrons that are shared spin in opposite directions as
explained in the Pauli Exclusion Principle.
Both hydrogen atoms now have a filled outer shell.

Both chlorine atoms now have a filled outer shell.

Notice that no ions are formed!

 Structural Formula
Chemical formulas that show the
arrangement of atoms in molecules.

Each dash represents a shared pair of electrons.

What is the structural formula for hydrogen?

A single dash shows that it is a single

covalent bond.
 O + O O O

What is the structural formula for Oxygen?

O=O
 +

Each needs 3 electrons

to fill outer shell.
What is the structural formula for Nitrogen?

:N N:
What do you think the structural formula of methane
would be?

(CH4)
Hint: Start by writing the Lewis symbol for carbon.

H
HC H
H
 H H
HC H H C H
H H
 Before we discuss these Lewis
STRUCTURES in more detail, we
need to discuss bond polarity and it’s
dependence upon electronegativity.
Nonpolar Covalent Bond
electrons are shared equally
 Polar Covalent Bond
one atom in the molecule has a stronger
attraction for the bonding electrons.
the ability of an atom in a molecule to compete for electrons.
Electronegativities according to Pauling Scale
(Relative Scale from 0.7 - 4.0)
 It is the difference in electronegativity
between two atoms that allows us to determine
the polarity and ionic character of a bond.
 We don’t consider any numerical fine line in classifying the
different bonds. Instead, we consider whether a particular bond
has ionic character or covalent character.
Most compounds with a difference > 1.5 have
more ionic character than covalent character.
Use the table below to determine which
of the following bonds are more polar.
H—O or C — Cl
 H—O or C — Cl
3.5 – 2.1 = 1.4 3.0 – 2.5 = .5
 H—O or C — Cl
3.5 – 2.1 = 1.4 3.0 – 2.5 = .5
The symbols used to indicate polarity
are delta plus and delta minus.

H—O
Polarity is also indicated using the following symbol:

Indicates electron density being greater near Cl.

Electron density of 3 types of bonds

 Dipole
A term used to describe the occurrence of a + charge
separated from a - charge
So, it is basically another term for “Polar molecule”.

Dipole Moment
μ
Quantitative measure of a dipole. (reported in debyes: C-m)

μ = Qr
*Gets larger as size of charge and distance between charge
increases.
Writing Lewis Structures
Writing Lewis Dot Structures

The Lewis electron-dot structure of a

covalent compound is a simple two-
dimensional representation of the positions of
electrons in a molecule.
Bonding electron pairs are indicated by either
two dots or a dash.
In addition, these formulas show the positions
of lone pairs of electrons.
Writing Lewis Dot Structures

The following rules allow you to write

electron-dot structures even when the central
atom does not follow the octet rule.

To illustrate, we will draw the structure

PCl 3

of PCl3, phosphorus trichloride.

Writing Lewis Dot Structures

Step 1: Total all valence electrons in the molecular

formula.
PCl3 26 e- total

5 e- (7 e-) x 3

For a polyatomic anion, add the number of

negative charges to this total.
For a polyatomic cation, subtract the number of
positive charges from this total.
Writing Lewis Dot Structures
Step 2: Arrange the atoms radially, with the
least electronegative atom in the center. Draw
a short line between the central atom and each
peripheral atom. (A line equals one pair of
electrons).
Cl Cl
P

Cl
Writing Lewis Dot Structures
Step 3: Distribute the remaining electrons to
the peripheral atoms to satisfy the octet rule.
:

:
:Cl: :Cl :
P
:Cl :
:
Writing Lewis Dot Structures

Step 4: Distribute any remaining electrons to

the central atom. If there are fewer than eight
electrons on the central atom, a multiple bond
may be necessary.
:

:
:Cl: : :Cl :
P
:Cl :
:
Writing Lewis Dot Structures

Try drawing Lewis dot structure for the

following covalent compound.
SCl2 20 e- total
16 e- left
4 e- left
: :

: :

: Cl S : :
Cl :
Writing Lewis Dot Structures

Try drawing Lewis dot formulas for the

following covalent compound.
COCl2 24 e- total
18 e- left
:
:O : 0 e- left

C
: :

: :
:Cl Cl:
Writing Lewis Dot Structures

Note that the carbon has only 6 electrons.

– One of the oxygens must share a lone pair.

COCl2 24 e- total
18 e- left
:
:O : 0 e- left

C
: :

: :
:Cl Cl:
Writing Lewis Dot Structures

Note that the carbon has only 6 electrons.

– One of the oxygens must share a lone pair.

COCl2 24 e- total
18 e- left
:O : 0 e- left

C Note that the

octet rule is
: :

: :
:Cl Cl: now obeyed.
 Delocalized Bonding:
Resonance
The structure of ozone, O3, can be
represented by two different Lewis
electron-dot formulas.

:
:

O or O

: :
: :
: :

: :

O O: :O O

Experiments show, however, that both bonds

are identical.
 Delocalized Bonding:
Resonance
According to theory, one pair of bonding
electrons is spread over the region of all
three atoms.
O
O O
This is called delocalized bonding, in which a
bonding pair of electrons is spread over a
number of atoms.
 Delocalized Bonding:
Resonance
According to the resonance description, you
describe the electron structure of molecules with
delocalized bonding by drawing all of the
possible electron-dot formulas.

:
:

O and O

: :
: :
: :

: :

O O: :O O
These are called the resonance formulas of
the molecule.
 Double headed arrows are used to
indicate resonance structures.

:
:
O O

: :
: :

: :
: :
O O: :O O

If the resonance structure is an ion, brackets are

placed around both structures, and the charge is
included outside of the brackets.
Draw the Lewis Structure for the nitrate ion.

I will draw it on the board when

you are finished.
Exceptions to the Octet Rule
Although many molecules obey the octet
rule, there are exceptions where the
central atom has more than eight
electrons.
Generally, if a nonmetal is in the third period
or greater it can accommodate as many as
twelve electrons, if it is the central atom.
These elements have unfilled “d” subshells that
can be used for bonding.
Nonmetals with atomic # 15 or higher may
consist of structures with exceptions.
Exceptions to the Octet Rule

For example, the bonding in phosphorus

pentafluoride, PF5, shows ten electrons
surrounding the phosphorus.

:
:F:

:
F:
:

: :
:F P
F:
:

:
:F:
:
Exceptions to the Octet Rule

In xenon tetrafluoride, XeF4, the xenon

atom must accommodate two extra lone
pairs. :

:
:F : F:
: :

: :
Xe
:F F:
:
:

:
 Draw the Lewis Structure of the following:

SF4

F
F S F

F
34 e-
 Exceptions to the Octet Rule
There are two other elements where an
exception may occur.

B and Be may consist of less than an octet.

Cl Cl
B
BCl3
24 e- Cl
 Draw the Lewis Structure of the following:

BeF2

F Be F

16 e-
 Formal Charge and Lewis
Structures
In certain instances, more than one
feasible Lewis structure can be illustrated
for a molecule. For example,

H C N: or H N C:
The concept of “formal charge” helps us determine the
structure that is most stable and most likely to form.
The formal charge of an atom is determined
by the following formula:

number of valence electrons

minus

½ number of bonding electrons

minus

number of unshared valence electrons

Formal Charge = VE - 1/2B - U
Formal Charge and Lewis
Structures
The most likely structure is the one with the least
number of atoms carrying formal charge. If they have
the same number of atoms carrying formal charge,
choose the structure with the negative formal charge
on the more electronegative atom. formal
H C N: or H N C: charge

0 0 0 0 +1 -1
In this case, the structure on the left is most
likely correct.
 The following are possible Lewis Structures for carbon dioxide.
Based on formal charge, which structure is most likely the correct structure?

0 0 0

O C O

O C O
-1 0 +1
Bond Length and Bond Order

Bond length (or bond distance) is the

distance between the nuclei in a bond.
•Knowing the bond length in a molecule can
sometimes give clues as to the strength of the
bond in a molecule.
• Covalent radii are values assigned to atoms
such that the sum of the radii of atoms “A” and
“B” approximate the A-B bond length.
Bond Length and Bond Order
Table 8.5 on page 305 lists bond lengths
for some covalent bonds.

We learned about atomic radii in chapter 7.

To predict the bond length of C-Cl, you
simply add the radii of the two atoms.

C Cl
Bond
length
Bond Length and Bond Order
The bond order, determined by the Lewis
structure, is the number of pairs of
electrons in a bond.
Bond length depends on bond order.
As the bond order increases, the bond gets
shorter and stronger.
Bond length Bond energy
C C 154 pm 346 kJ/mol
C C 134 pm 602 kJ/mol
C C 120 pm 835 kJ/mol
 Bond length Bond energy
C C 154 pm 346 kJ/mol
C C 134 pm 602 kJ/mol
C C 120 pm 835 kJ/mol

Since this bond is stronger, it requires

more energy to break this bond.

Bond Enthalpy
Energy needed to break a bond in 1
mole of a gaseous substance.
The strength of a bond is also related to it’s stability.
Consider N2 versus Cl2.

N N Cl Cl
Bond Length Bond Length
1.10 A 1.96 A

Bond Enthalpy Bond Enthalpy

941 KJ / mol 242 KJ / mol

More Stable Less Stable