I. Types of Bonds

1. Ionic bond
Transfer of e- from a
metal to a nonmetal
and the resulting
electrostatic force that
holds them together forms
an ionic compound.
EX: Na+ + Cl-  NaCl
(neutral)
Ionic Bonds: One Big Greedy Thief Dog!
Types of Bonds

2. Covalent bond
Formed from the sharing of
e- pairs between
two or more nonmetals
resulting in a molecule.
EX: H2 + O  H2O

Here are
some
common
examples of
covalent
bonds.
Types of Bonds

3. Metallic bond
Metals bonding with
other metals
do not gain or lose eor share e- unequally.
These bonds are
created from the
delocalized e- that hold
metallic atoms
together.
II. Electronegativity


Remember that electronegativity is the
tendency of an atom to attract an ein a chemical compound.
Electronegativities of an element are
influenced by atomic radii and number of
valence e-
II. Electronegativity
The greater # of valence e-, the
greater the electronegativity (Why?
More attraction to fill octet rule)
 The smaller the atomic radius, the
greater the electronegativity (Why?
Smaller atoms will hold e- closer)

Electronegativity
Flourine (F)
is the most
electronegative
element
Notice that the noble gases are NOT electronegative… Why?
Because they don’t react to make chemical compounds!
Bond Character


Electrons of atoms are exchanged when
the difference between electronegativity
between the atoms is high.
A difference more than 1.7 will USUALLY
produce ionic bonds.
EX: Magnesium + Oxygen
|1.31 – 3.44| = 2.13 difference
What type of bond? IONIC BOND
Bond Character

Electrons are shared when the
electronegativity difference between
atoms is low, 1.69 and below.
EX: Hydrogen and Oxygen
|2.20 – 3.44| = 1.24 difference
What type of Bond?
COVALENT BOND
Bond Character
Calculate e-neg differences and label as ionic
or covalent.
COVALENT
B – P |2.04 – 2.19| = 0.15 (below 1.69)
Mg – N |1.31 – 3.04| = 1.73 IONIC
(above 1.70)
C – Na |2.55 –.93| = 1.62
COVALENT
(below 1.69)
QOD

Name the three
types of bonds
1. Ionic
2. Covalent
3. Metallic

Covalent bonds
_______ electrons.
Share
III. Valence Electrons

The electrons that are involved in bonding are the
outer most electrons.




They are always “s” or “s and p” electrons of the
outermost principal energy level.
These electrons are called valence electrons.
According to the Octet Rule, eight valence
electrons assures stability.
Atoms exchange or share electrons so that they
can have eight valence electrons.
Valence Electrons


Because e- config. demonstrates a
periodic trend, most valence e- can be
determined by placement on periodic
table.
Remember oxidation states (+ or –
charges of atoms), valence e- are the
cause of these states.
Periodic Table - Valence Electrons

Group 1 has 1 valence, Group 2 has 2,
Group 13 has 3, and so on…
8
1
2
3
4 5
6 7
IV. Lewis Dot Structures
To visualize valence e-, we will use Lewis
Dot Structures.
 Step 1: The element symbol represents
the nucleus and all e- except valence.
 Step 2: Write the e- config. From
config., select e- in the outer level.
These e- are the ones with the largest
principal quantum numbers.
Lewis Dot Structures

Step 3: Each “side” of symbol
represents an orbital. Draw dots on
sides as you would fill orbitals. One at a
time, then paired. Start with the bottom
orbital and work clockwise when filling.
Lewis Dot Structures

EX: carbon
step 1: C
step 2: [He]2s22p2
step 3:
C
Lewis Dot Structures

EX: bromine
step 1: Br
step 2: [Ar]4s23d104p5
step 3:
Br
QOD


Which element
demonstrates the
weakest attraction
for electrons, or
lowest
electronegativity?
Na, P, Si, S
Na

When two or more
atoms bond
covalently to form a
neutral particle, the
particle is called a
Molecule
V. IONIC BONDS
Drawing Ionic Bonds

Draw Lewis dots for each element.
Na + Cl


Draw an arrow to show the e- exchange
between atoms.
Resulting in:

Don’t have to draw this step
Na
+
Cl
-
Drawing Simple Ionic Bonds

Now draw them together & show the
ion charge.
Na


+
Cl-
Metals are always first in the formula.
Ionic bonds form because of the
attraction of the opposite charges.
Ionic Bonds –
Exchange of Electrons
Drawing Simple Ionic Bonds


Lewis Dot for Ionic bonds must
continue until the metals have lost all
their electrons and the nonmetals have
eight total electrons.
To accomplish this, you may need to
bring in more metal or non-metal ions.
QOD
Draw the Lewis Dot
Structure for the
following:
Sr + P
Sr (2ve) + P (5ve)

When the
electronegativity
difference between
bonded atoms is
greater than 1.67,
the substance is
classified as
IONIC COMPOUND
VI. COVALENT BONDS
REVIEW

Let’s first review ionic bonding!
In an IONIC BOND,
electrons are lost or gained,
resulting in the formation of IONS
in ionic compounds.
K
F
K
F
K
F
K
F
K
F
K
F
K
F
K
+
_
F
+
K F
_
K
+
_
F
The compound potassium fluoride
consists of potassium (K+) ions
and fluoride (F-) ions
K
+
_
F
The IONIC BOND is the attraction
between the positive K+ ion
and the negative F- ion
COVALENT BONDS



In covalent bonding, atoms still want to
achieve a noble gas configuration (the
octet rule).
But, rather than losing or gaining
electrons, atoms now share an electron
pair
The shared electron pair is called a
bonding pair
Chlorine
forms
a
covalent
bond
with
itself
Cl2
Cl
Cl
How
will
two
chlorine
atoms
react?
Cl
Cl
Each chlorine atom wants to
gain one electron to achieve an octet
Cl
Cl
Neither atom will give up an electron –
chlorine is highly electronegative.
What’s the solution – what can they
do to achieve an octet?
Cl
Cl
Cl Cl
Cl Cl
Cl Cl
Cl Cl
octet
Cl Cl
octet
circle the electrons for
each atom that completes
their octets
Cl Cl
The octet is achieved by
each atom sharing the
electron pair in the middle
circle the electrons for
each atom that completes
their octets
Cl Cl
The octet is achieved by
each atom sharing the
electron pair in the middle
circle the electrons for
each atom that completes
their octets
Cl Cl
This is the bonding pair
circle the electrons for
each atom that completes
their octets
Cl Cl
It is a single bonding pair
circle the electrons for
each atom that completes
their octets
Cl Cl
It is called a SINGLE BOND
circle the electrons for
each atom that completes
their octets
Cl Cl
Single bonds are abbreviated
with a dash
circle the electrons for
each atom that completes
their octets
Cl Cl
This is the chlorine molecule,
Cl2
Draw this molecule in your
notes
circle
the
electrons for
each atom that completes
their octets
O2
Oxygen is also one of the diatomic molecules
O
O
How will two oxygen atoms bond?
O
O
Each atom has two unpaired electrons
O
O
O
O
O
O
O
O
O
O
O
O
O
O
Oxygen atoms are highly electronegative.
So both atoms want to gain two electrons.
O
O
Oxygen atoms are highly electronegative.
So both atoms want to gain two electrons.
O
O
O
O
O O
O O
O O
Both electron pairs are shared.
O O
6 valence electrons
plus 2 shared electrons
= full octet
O O
6 valence electrons
plus 2 shared electrons
= full octet
O O
two bonding pairs,
making a double bond
O O O =O
For convenience, the double bond
can be shown as two dashes.
O =O
This is the oxygen molecule,
O2
Draw this molecule in your notes
this
is so
cool!
!
Drawing Structural Formulas
So far we have looked at F2 and O2. These are
simple structural formulas showing covalent
bonds.
What about those more complex structural
formulas that use covalent bond (AsI3)?
There must be some tricks to drawing these
structural formulas using covalent bonds.
Yes there is… Here are the steps…
Lewis Dot Diagrams of Molecules
1. Count the total number of valence e-.
2. Determine the central atom. The following are
guides:

Often the unique atom (only one of it) is the
central atom.

Or put the least electronegative element in the
middle.
3. Arrange the other atoms around the central
atom creating a skeleton.
Lewis Dot Diagrams of Molecules
4. Connect all bonded atoms in the
skeleton with one bond.
5. Subtract the number of electrons already
used for the single bonds; two for each
bond.
6. Distribute the remaining electrons in pairs
around the atoms, trying to satisfy the
octet rule. Assign them to the most
electronegative atom first.
Lewis Dot Diagrams of Molecules
7. If you run out of electrons before all atoms
have an octet of electrons, you need to form
double or triple bonds.
8. If you have extra electrons and all of the atoms
have an octet, put the extra electrons on the
central atom in pairs.
9. Math column must reach zero electrons
10. Check to make sure all atoms have octet of
electrons. (if not make multiple bonds)
Lewis Dot Diagrams of Molecules
Exceptions to octet:
*Hydrogen can only have two electrons which is
one shared bond. Hydrogen can never have
unshared electrons or multiple shared bonds.
*Aluminum is satisfied with 6 electrons, which is
three shared bonds.
*If the central atom has an atomic number
greater than fifteen, you are allowed to have
more than eight electrons around it.
Lewis Dot Diagrams of AsI3

EX: AsI3
1. Count valence e- [5 + (3 x 7)]= 26
2. Place As (only one As atom and lowest
electronegativity) in center.
3. Place the three Iodines around As.
4. Draw lines (bonds connecting them)
I –As –I
I
Lewis Dot Diagrams of AsI3
5. Subtract # used for bonds (26 – 6) =
20 e-.
6. Place e- around the three I atoms
first because they are the most
electronegative.
I –As –I
I
7. We did not run out of e- so we do
not have any double or triple bonds.
Lewis Dot Diagrams of AsI3
8. We have 2 extra e- (20 starting –
(3x6) = 2. Place them around the
central atom.
I –As –I
I
9. Check each atom to make sure each
has a full octet of electrons.
Lewis Dot Diagram of CH2O
1. Count total valence e(C=4, H=1x2, O=6)= 12 valence e2. Place C in the center.
3. Place the 2 hydrogens and one oxygen
around C.
4. Draw lines connecting H and O to C.
5. Subtract number of bonded e- from
total. (12-6) = 6
Lewis Dot Diagram of CH2O
6. Place e- around the oxygen atom first
because it is the most electronegative.
H
C
H
O
Lewis Dot Diagram of CH2O
7. We ran out of e- before carbon
satisfied its octet. This means that we
will have a double bond between the
carbon and oxygen. (Hydrogen cannot
form double bonds).
8. There will be no unshared e-.
Lewis Dot Diagram of CH2O
H
C
H
O
QOD
Draw the Lewis Dot
Structure for the
following:
1. Na + S
2. H2O
Special Types of Covalent Bonds
1. Diatomic Molecules
 2. Polyatomic Ions


All bonds in diatomic molecules
and polyatomic ions are
covalent bonds
Diatomic Molecules


1. Molecule – two or more atoms
covalently bound together
2. Diatomic Molecule – two of the
same atom bound together.


Br, I, N, Cl, H, O, F
They are called the BIG SEVEN



Color these on your periodic table
These atoms do not exist alone because
they are extremely reactive.
They always come in pairs.
Diatomic Molecules

For example:







Br
I
N
Cl
H
O
F
_____
I2
_____
Cl2
H2
_____
F2
Lewis Dot Diagram of
Polyatomic Ions


Polyatomic ions are groups of covalently bonded
atoms that carry an overall charge.
The atoms in a polyatomic ion are very tightly bound
and are treated as a single entity. Because they do
not come apart in ionic bonding, we often put them
in parenthesis.


Example: (CN)
–
Constructing Dot Diagrams for the polyatomic ions is
the same, except the difference in charge (+ or -)
must be accounted for.
Lewis Dot Diagram of ClO311. Count valence e-, including charge
7 + (3x6) + 1 = 26 total
2. Place Cl in the center.
3. Arrange the three O around it.
4. Draw bonds from O to Cl.
O
Cl
O
O
Lewis Dot Diagram of ClO315. Subtract e- used in bonds (26 – 6) = 20 e-.
6. Place remaining e- around the three oxygens
to satisfy octet.
7. There are two e- left over so we will not have
multiple bonds.
8. Place the last two e- on the central atom.
*For polyatomic ions: place structure in
brackets with the charge indicated on outside
as demonstrated
Lewis Dot Diagram of ClO31-1
O
Cl
O
O
*don’t forget the brackets
QOD
Draw the Lewis Dot Structure for the
following:
ClO4-
Electron Pair Repulsion


Each bond and each unshared pair of eform a charge cloud that repels other
charge clouds.
Electron pairs spread as far apart as
possible to minimize repulsive forces.
VSEPR Theory


Valence shell, electron-pair repulsion
(VSEPR) Theory states that repulsion
between the sets of valence-level
electrons surrounding an atom causes
these sets to be oriented as far apart as
possible.
Using VSEPR, we can predict the
geometries (shapes) of molecules.
VSEPR Theory to Predict
Molecular Geometries





Step 1: Write the Lewis Dot Structure for
the molecule.
Step 2: Represent the central atom in
molecule by the letter A.
Step 3:Represent any atoms bonded to it
by the letter X.
Lone pairs are represented by letter E
Step 4: Now refer to Table 6-5 VSEPR and
Molecular Geometries
Valence Shell Electron Pair Repulsion
Theory
VSEPR theory lets us predict the shape of a molecule based on the
electron configurations of the constituent atoms. It is based on
maximizing the distance between points on a spherical surface.
Number of
Objects
2
3
4
5
6
Geometry
linear
trigonal
planar
tetrahedral
trigonal
bipyramidal*
Octahedral
Lewis Dot Diagram of SO3
EX: SO3
1. Lewis Dot: S has 6 valence and
O has 6 x 3 for a total of 24 e-.
2. Place S in center.
3. Place O around S.
4. Draw bonds connecting S and O.
5. Subtract bonds from total e(24 – 6) = 18 eO

O
S
O
Lewis Dot Diagram of SO3
6. Now put electrons (in pairs)
around the most
electronegative atoms (O)
first, try to satisfy octet rule.
7. There are no remaining e-.
8. Check to see if every atom
has an octet. Sulfur does not
so make one double bond.
O
O
S
O
Predicting Molecular Geometry of SO3
Refer to structure.
2. S will be represented by letter A.
3. O will be represented by letter X.
4. We have AX3. Refer to Table 6-5 and
look for AX3.
5. According to the Table, the molecular
geometry for SO3 (AX3) is Trigonal-planar.
Trigonal-Planar Geometry
O
O
S
O
Molecular
Geometry
of SO3
Molecular Geometry of PCl3
1. Lewis Dot Diagram:
Cl
Cl
P
Cl
2. P is represented by A. The unshared pair is
represented by E.
3. Cl is represented by X.
4. We have AX3E. Refer to Table 6-5.
Trigonal-Pyramidal Geometry
Molecular
Geometry of
PCl3
Molecular Geometry of H2O

Remember, H2O has a Lewis Structure of
H—O—H


The O is represented by A and the two H
are represented by X. There are two
unshared pairs represented by E2.
We have AX2E2. Refer to Table 6-5.
Bent

AX2E2 is bent.
Molecular
Geometry
of H2O
Covalent Bonds:
Polar or Nonpolar



Covalent bonds are created from the
sharing of e-.
Since some atoms are more
electronegative than others, this sharing is
often unequal.
This results in polarity for a molecule
where there are moments when atoms
are partially negative and the others are
partially positive.
Polarity of Water


Take for example H2O. When we draw the
molecular geometry it looks like:
O
H
H
The electronegativities of H and O are
2.20 and 3.50. O is the more
electronegative and will pull the shared emore of the time. We illustrate this by this
symbol
Polarity of Water


The arrow is pointing
from the least
electronegative atoms to
the most electronegative
atoms.
The arrows show
movement of electrons.
Draw the direction of the
arrow in your notes
Polarity of Water


When electrons are being
pulled so that the forces
of the pulls are not even,
the molecules have
dipole moments.
These dipoles give a
partial positive and
negative charge to atoms
in a polar molecule.
O
H
H
Polarity of Water
O
EX: H2O
The oxygen has a partial
negative charge and the
hydrogens have a partial
positive charge.
So……H2O is a polar
molecule.
H
H
Nonpolar


EX: CH4
Carbon is more
electronegative than
Hydrogen, but this
molecule is nonpolar
because the charges are
pulled in opposite
directions equally and that
causes the forces of
electron pull to cancel.
H
H
C
H
H
Determining Polarity
1. General guidelines to determine polarity:
2. If the central atom has 1 or more lone
pair of electrons, the molecule is polar.
3. If there are different types of elements
bonded to a central atom, the molecule is
polar. EX: CH3Cl.
4. If two different atoms are bonded
together, the molecule is polar. EX: HF
In Summary


The tendency of a bonded atom to
attract shared electrons to itself when
bonded to another atom is called
electronegativity (e-neg)
Large differences in e-neg lead to the
formation of ions, atoms that have
gained or lost e-, and then ionic
compounds, compounds bonded ionically
(attraction of opposite charges).
In Summary




Little or no difference in e-neg leads to
covalent bonds (sharing of e-)
Molecules are held together by covalent
bonds.
The structure of a molecule or a polyatomic
ion can be represented by a Lewis Dot
Diagram, which shows the pattern of shared
and unshared pairs of e-.
Two atoms sometimes share more than one
pair of e-, forming double or triple bonds.

THE END!