When Atoms Meet
1
Bonds
 Forces that hold groups of atoms
together and make them function
as a unit.
Bonding Forces
 Electron – electron
repulsive forces
 Nucleus – nucleus
repulsive forces
 Electron – necleus
attractive forces
2
Metals and Nonmetals
3
Types of Chemical Bonding
1. Metal with nonmetal:
electron transfer and ionic bonding
4
Three models of chemical bonding
Ionic
Electron transfer
5
Types of Chemical Bonding
1. Metal with nonmetal:
electron transfer and ionic bonding
2. Nonmetal with nonmetal:
electron sharing and covalent bonding
6
Three models of chemical bonding
Ionic
Covalent
Electron transfer
Electron sharing
7
Types of Chemical Bonding
1. Metal with nonmetal:
electron transfer and ionic bonding
2. Nonmetal with nonmetal:
electron sharing and covalent bonding
3. Metal with metal:
electron pooling and metallic bonding
8
Three models of chemical bonding
Ionic
Covalent
Electron transfer
Electron sharing
Metallic
9
Electron pooling
Valence Electrons
• The outer shell electrons of an atom
• Participate in chemical bonding
Group
e- configuration
# of valence e-
1A
ns1
1
2A
ns2
2
3A
ns2np1
3
4A
ns2np2
4
5A
ns2np3
5
6A
ns2np4
6
7A
ns2np5
7
10
Lewis Structures
Developed the idea in
1902.
G. N. Lewis
11
Lewis Dot Symbols
Place one dot per valence electron on each of the four
sides of the element symbol.
Pair the dots (electrons) until all of the valence electrons are
used.
:
Nitrogen, N, is in Group 5A and therefore has 5 valence
electrons.
.
.
.
.
.
. N. .
. N.
: N
N:
.
.
:
12
Lewis Dot Symbols
13
The Octet Rule
Chemical compounds tend to form so that each
atom, by gaining, losing, or sharing electrons, has
eight electrons in its highest occupied energy
level.
The same number of electrons as in the nearest
noble gas
The first exception to this is hydrogen, which
follows the duet rule.
The second exception is helium which does not
form bonds because it is already “full” with its
two electrons
14
Ionic Bond
Li+ F -
Li + F
1 22s22p5
1s22s1s
[He]
1s
1s2[2Ne]
2s22p6
Li+
Li
1s
2s
1s
2p
+
+ F
1s
2s
2p
2s
2p
F1s
2s
2p
15
Electrostatic (Lattice) Energy
Lattice energy (E) is the energy required to completely separate
one mole of a solid ionic compound into gaseous ions.
Q+Q E=k
r
Q+ is the charge on the cation
Q- is the charge on the anion
r is the distance between the ions
Lattice energy (E) increases
as Q increases and/or
as r decreases.
cmpd
MgF2
MgO
LiF
LiCl
lattice energy
2957 Q= +2,-1
3938 Q= +2,-2
1036
853
r F < r Cl
16
Covalent Bond
A chemical bond in which two or more electrons are shared by
two atoms.
How should two atoms share electrons?
F
+
7e-
F
F F
7e-
8e- 8e-
Lewis structure of F2
single covalent bond
lone pairs
F
F
lone pairs
single covalent bond
lone pairs
F F
lone pairs
17
Distribution of electron density of H2
H H
18
Lewis structure of water
H
+
O +
H
single covalent bonds
H O H
or
H
O
H
2e-8e-2eDouble bond – two atoms share two pairs of electrons
O C O
or
O
O
C
double bonds
- 8e8e- 8ebonds
double
Triple bond – two atoms share three pairs of electrons
N N
triple
bond
8e-8e
or
N
N
triple bond
19
Polar Covalent Bond
A covalent bond with greater electron density around
one of the two atoms
electron poor
region
H
electron rich
region
F
e- poor
H
d+
e- rich
F
d-
20
Electron density distributions in
H2, F2, and HF.
21
Electronegativities (EN)
The ability of an atom in a molecule to attract shared electrons to itself
Linus Pauling
1901 - 1994
22
Classification of Bonds
Difference in EN
Bond Type
0
Covalent
2
0 < and <2
Ionic
Polar Covalent
Increasing difference in electronegativity
Covalent
Polar Covalent
share e-
partial transfer of e-
Ionic
transfer e-
23
Classification of Bonds
Classify the following bonds as ionic, polar covalent,
or covalent: The bond in CsCl; the bond in H2S; and
the NN bond in H2NNH2.
Cs – 0.7
Cl – 3.0
3.0 – 0.7 = 2.3
Ionic
H – 2.1
S – 2.5
2.5 – 2.1 = 0.4
Polar Covalent
N – 3.0
N – 3.0
3.0 – 3.0 = 0
Covalent
24
Rules for Writing Lewis Structures
1. Draw skeletal structure of compound showing what
atoms are bonded to each other. Put least
electronegative element in the center.
2. Count total number of valence e-. Add 1 for each
negative charge. Subtract 1 for each positive
charge.
3. Use one pair of electrons to form a bond (a single
line) between each pair of atoms.
4. Arrange the remaining electrons to satisfy an octet
for all atoms (duet for H), starting from outer atoms.
5. If a central atom does not have an octet, move in
lone pairs to form double or triple bonds on the
central atom as needed.
25
Write the Lewis structure of nitrogen trifluoride (NF3).
Step 1 – N is less electronegative than F, put N in center
Step 2 – Count valence electrons N - 5 (2s22p3) and F - 7 (2s22p5)
5 + (3 x 7) = 26 valence electrons
Step 3 – Draw single bonds between N and F atoms.
Step 4 – Arrange remaining 20 electrons to complete octets
F
N
F
F
26
Write the Lewis structure of the carbonate ion (CO32-).
Step 1 – C is less electronegative than O, put C in center
Step 2 – Count valence electrons C - 4 (2s22p2) and O - 6 (2s22p4)
-2 charge – 2e4 + (3 x 6) + 2 = 24 valence electrons
Step 3 – Draw single bonds between C and O atoms
Step 4 - Arrange remaining 18 electrons to complete octets
Step 5 – The central C has only 6 electrons. Form a double bond.
2-
O
C
O
O
27
Resonance
More than one valid Lewis structures can be written for a
particular molecule
The actual structure of the carbonate ion is an average of the
three resonance structures
-
2-
O
C
O
O
-
O
C
O
O
-
-
2-
-
2-
O
C
O
O
28
-
Exceptions to the Octet Rule
The Incomplete Octet
BeH2
BF3
B – 3e3F – 3x7e24e-
Be – 2e2H – 2x1e4e-
F
B
H
F
Be
H
3 single bonds (3x2) = 6
9 lone pairs (9x2) = 18
Total = 24
F
29
Exceptions to the Octet Rule
Odd-Electron Molecules
NO
N – 5eO – 6e11e-
N
O
The Expanded Octet (central atom with principal quantum number n > 2)
SF6
S – 6e6F – 42e48e-
F
F
F
S
F
F
F
6 single bonds (6x2) = 12
18 lone pairs (18x2) = 36
Total = 48
30
Covalent Bond Lengths
Bond
Type
Bond
Length
(pm)
C-C
154
CC
133
CC
120
C-N
143
CN
138
CN
116
Bond Lengths
Triple bond < Double Bond < Single Bond
31
Covalent Bond Energy
The energy required to break a particular bond in one mole of
gaseous molecules is the bond energy.
Bond Energy
H2 (g)
H (g) + H (g)
436.4 kJ
Cl2 (g)
Cl (g) + Cl (g)
242.7 kJ
HCl (g)
H (g) + Cl (g)
431.9 kJ
O2 (g)
O (g) + O (g)
498.7 kJ
O
O
N2 (g)
N (g) + N (g)
941.4 kJ
N
N
Bond Energies
Single bond < Double bond < Triple bond
32
Light-Matter Interactions
3 x 1020
3 x 1016
3 x 1012
3 x 108
3 x 104
Frequency in Hz
Dissociation
Ionization
Vibration
Rotation
33
Vibrational Modes of Water
Infrared light
34
Infrared Spectrum of Water
Absorbance
Liquid
3000
2000
1000
Gas
2000
1500
1000
Wavenumber (cm-1)
Reveal the interactions between molecules and their environments
35
Absorbance
Infrared Spectrum of Caffeine
3000
2000
1000
Wavenumber (cm-1)
Identification of compounds
36
Lab 1
37
Acknowledgment
Some images, animation, and material have been taken from the following sources:
Chemistry, Zumdahl, Steven S.; Zumdahl, Susan A.; Houghton Mifflin Co., 6th Ed., 2003;
supplements for the instructor
General Chemistry: The Essential Concepts, Chang, Raymon; McGraw-Hill Co. Inc., 4th
Ed., 2005; supplements for the instructor
Principles of General Chemistry, Silberberg, Martin; McGraw-Hill Co. Inc., 1st Ed., 2006;
supplements for the instructor
NIST WebBook: http://webbook.nist.gov/
http://www.lsbu.ac.uk/water/vibrat.html
http://en.wikipedia.org/wiki/Caffeine
http://www.wilsonhs.com/SCIENCE/CHEMISTRY/MRWILSON/Unit%204%20Chemical%2
0Bonding%20Powerpoint1.ppt
38