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
+1
F
-1
The compound potassium fluoride
consists of potassium (K+) ions
and fluoride (F-) ions
+1
K
-1
F
potassium fluoride
The ionic bond is the attraction
between the positive K+ ion
and the negative F ion.
an electrostatic attraction between a pair of
electrons and positively charged nuclei
 atoms are sharing valence electrons

 this is still in order to achieve an noble gas electron
configuration (stable and less energy)
 exists where groups of atoms (or molecules) share
one or more pair/s of electrons

normally happens between nonmetals
Each hydrogen now has the electron configuration
of the nearest noble gas- helium

some atoms form stable molecules without a
stable octet
 boron
 beryllium
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 (3.2)
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
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 (3.4)
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
this
is so
cool!
!
CCl4 - Covalent
Cl
HCl - Covalent
H Cl
MgF2 - Ionic
[Mg]2+[
F
]2–
Cl C Cl
Cl
NH3 - Covalent
H N H
H2O - Covalent
H O H
NaCl - Ionic
[Na] + [ Cl ]–
H
H2 - Covalent
OH– - Covalent
H H
O H
HCl - Covalent CO2 - Covalent Na2O - Ionic
O
O
C
Cl
Cl
H
H
[Na]2+ [ O ]2–
O C O
NH3 - Covalent
O2 - Covalent I2 - Covalent
I I
O O
H N H H N H
O O
H
H
I
O
Covalent
3
H
Al2O3 - Ionic
O O O H C H H
3+
2–
[Al]2 [ O ]3
O O O
H
I
H
C H
H

occurs when one atom donates both of the
electrons that are shared between two atoms
› should know examples: CO, NH4+ and H3O+ and…
Arrow represents the dative bond.
Pointing towards the atom that did NOT contribute
to the bond

more bonds means stronger
› single ⇒

double ⇒ triple
more bonds means shorter length
› triple ⇒ d o u b l e ⇒ s i n g l e
Know difference in strength and length between
single, double, and triple bonds between two
carbon atoms.
Know the bond length difference between the C and
O in the carboxyl group

Draw plausible Lewis dot structures of the
following covalent compounds. Note if any
bonds are dative bonds.
› start with the “main element” and then attach
the others
O2
NF3
H3O+1
N2
H2Se
HCN
NH3
CO2
C2H4
NH4 +1
CO
more than one Lewis structure can be
drawn
 often view certain molecules as if they were
able to resonate (go back and forth)
between two or more different structures

› example: NO31- behaves as if it were a blend of
the three resonance structures


VSEPR Theory YouTube (4:52)
the structure of many molecules is determined
mostly by minimizing electron pair repulsions
› electrons don’t “want” to be near each other

electron domains are places you would find
electrons around a central atom
 lone pairs (unshared)
 pairs of electrons around a central atom
NOT in a bond with another atom
 have more repulsive force than those
found in a bond
 bonding pairs (shared)
 pairs of electrons being SHARED found in
the space between the atoms (can be
single, double, or triple bonds)
the repulsion between electron domains
causes molecular shapes to adjust
 electron pairs arrange themselves
around the central atom so that they are
as far apart from each other as possible

› explains the three dimensional shape of a
molecule
SL
level
HL
level
Copyright © Cengage Learning. All rights reserved
Copyright © Cengage Learning. All rights reserved
Copyright © Cengage Learning. All rights reserved
Shapes for species with 2, 3, and 4
electron domains on the central atom
2 bonding pairs
 0 lone (unshared) pairs
 bond angle is 180º

shape is linear
3 bonding pairs
 0 lone pairs
 bond angle 120º

2 bonding pairs
 1 lone pair
 bond angle <120º
trigonal planar

bent


4 bonding pairs
0 lone pairs
› bond angle 109.5º


3 bonding pairs
1 lone pair
› pushes 3 bonding pairs
tetrahedral
closer together
› bond angle < 109.5º (107º)


2 bonding pairs
2 lone pairs
trigonal pyramid
› these push 2 bonding pairs
even closer together
› bond angle < < 109.5º (105º)
bent
Drawing 3-D molecules

shared bonding electrons pairs are
sometimes pulled (as in a “tug-of-war”)
between atoms
› equal sharing is non-polar
› unequal sharing is polar
 atoms
in the bond pull the shared pair of
electrons equally-- so no polarity
 always the case in diatomic molecules
› HOFBrINCl meaning…
 H2 O2 F2 Br2 I2 N2 Cl2
memorize these
atoms in the bond pull the shared pair of
electrons unequally since they have
different electronegativities
 results in a dipole because it has two poles

 use the symbol + or - for areas that are
slightly positive or negatively charged

or use a vector to point to the negative
part of the molecule
 has
BF3
polar
bonds yet NOT
a polar
molecule
 there is no
“negative end”
or “positive
end” because
of its shape
So why are some molecules polar?
•
more electronegative atoms have a
greater attraction for electrons
• a number is assigned to each element to
quantify its attraction to a pair of electrons
in a shared in bond (example- F is 4.0)
•
•
atoms with the higher electronegativity
give that “side” of the molecule a
slightly negative charge (δ -)
atoms on the “other side” with a lower
electronegativity therefore have a
slightly positive charge (δ +)
covalent, non-polar
covalent, polar
ionic
Ionic and covalent are not separate
“things” but differences in degree
electronegativty
difference
0.0 – 0.3
0.4 – 1.0
1.1 – 1.8
> 1.8
probable type of bond
covalent, nonpolar
covalent, slightly polar
covalent, very polar
ionic
Arrange the following bonds from most to
least polar:
a)
N–F
O–F
C–F
a) C–F, N–F, O–F
b) C–F
N–O
Si–F
b) Si–F, C–F, N–O
c) Cl–Cl,
B–Cl,
S–Cl
c) B–Cl, S–Cl, Cl–Cl
 H2O
 CO2
 CH3Cl
 CCl4
non-polar
polar
Which of the following molecules have a dipole moment
(another way to say polar?
H2O, CO2, SO2, and CH4
S
O
dipole moment
polar molecule
dipole moment
polar molecule
H
O
C
O
H
C
H
H
no dipole moment
nonpolar molecule
no dipole moment
nonpolar molecule