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Bonding: Part Two
Three types of bonds:
•Ionic Bond
transfer valence e-
(NaCl)
•Metallic bond
mobile valence e-
(Fe)
•Covalent bond
shared valence e-
(H2O)
1
Single Covalent Bond
H + H
H-atoms
H H
H2 molecule
Electrons are shared
by the two H atoms
2
Single Covalent Bond
“Structural” Formula
H—H
Single covalent bond (2 shared e-)
3
Ionic vs. Covalent Compounds
Covalent: discrete molecules
H2O, CH4, CO2
Ionic: cations and anions
(occupy crystal lattice points)
”formula unit” NaCl, MgI2
4
Covalent Bond
Covalent bonds usually form
between nonmetal atoms
(Groups 14, 15, 16, and 17).
Each atom tries to attain the econfiguration of a noble gas by
sharing electrons. (“octet rule”)
5
Back to H2
H + H
H-atoms
H H
Each H has econfig. of He
6
Fluorine: F2 or F—F
F + F
F F
Each F has econfig of [Ne]
7
Lone Pairs
Paired valence e- not in the bond
F F
Each F has
an octet
“lone pairs” or
“nonbonding pairs” of e8
Water (H2O)
O H
O H
H
H
Each atom has Noble Gas
configuration
9
You Try It!
Draw dot structures for:
1.Ammonia (NH3)
2.Chlorine gas (Cl2)
3.Methane (CH4)
10
Multiple Bonds
Sometimes more than one pair
of bonding electrons are needed
in a bond to attain a noble gas
configuration.
Double bond: two pairs of eTriple bond: three pairs of e
11
Double Bond: O2
O + O
O O
Still no octet, so
form a double bond
O O
or
O O
12
Triple Bond
Try nitrogen (N2)
N
N
N
N
13
Coordinate Covalent Bond
e
Sometimes both
in the
bond come from just one of
the atoms.
14
Coordinate Covalent Bond
e.g. CO
no octet
O + C
e- pair came
from oxygen
O C
O C
15
Coordinate Covalent Bond
The ammonium ion, NH4+
H
+
H
no
e
+
NH
H
H
+
H NH
H
16
Dot Structure Rules
e.g. NF3
1.Arrange the atoms with least
electronegative element in center.
The central atom is never hydrogen.
F N F
F
17
Dot Structure Rules
2.Count total valence electrons.
Account for charges in
polyatomic ions.
F N F
F
5 + 3(7) = 26
18
Dot Structure Rules
3.Connect atoms with single
covalent bonds. Then complete
octets (H has only 2 not 8).
F N F
F
19
Dot Structure Rules
4.If octet rule is not satisfied for the
central atom, try double or triple
bonds, using lone pairs from
surrounding atoms.
F N F
F
OK
20
Dot Structure HNO3
Step 1. Skeletal structure
O N O H
O
21
HNO3
Step 2. Number of valence e-
O N O H
O
N, O, H  5 + 3(6) + 1 = 24
22
HNO3
Step 3. Add bonds and
complete octets
O N O H
O
Out of
e,
but no octet for N
23
HNO3
Step 4. Add multiple bonds
O N O H
O
24
You Try It !!!
Draw the e- dot structures for:
1.Hydrogen chloride HCl
2.Hydrogen peroxide H2O2
+
3.Hydronium ion H3O
4.Ozone O3
25
Lewis Structures
26
Resonance: Ozone O3
O O O
O O O
Each is called a
“resonance structure”.
The bonds are equal
(~1.5 bond)
27
Resonance: Try It !!!
Draw resonance
structures for SO2
28
Exception to the Octet
Rule: Odd # e
Try to write the dot structure of
nitrogen monoxide.
With an odd number of valence
electrons (11), it is impossible to
have octets around both atoms.
NO is “paramagnetic.”
29
Predicting Molecular Shapes
linear
triatomic
trigonal
pyramid
trigonal planar
bent
triatomic
tetrahedral
others
30
VSEPR Theory
“Valence-shell
e
pair repulsion”
•All valence electron pairs
(bonding & nonbonding pairs)
repel each other.
• Predicts geometry.
31
VSEPR Theory
Methane (CH4) is drawn as:
H
H—C—H
H
or H
H
Actually CH4 is 3-D
CH
H
32
VSEPR Theory
Maximum
repulsion of
e- pairs
109.5o
C
 tetrahedron
33
VSEPR Theory
Ammonia (NH3)
1 nonbonding e- pair
H—N—H
H
3 bonding e- pairs
34
VSEPR Theory
Ammonia
tetrahedral
electrons
trigonal pyramidal
atoms
35
VSEPR Theory
Water (H2O)
2 nonbonding e- pairs
H—O—H
2 bonding e- pairs
36
VSEPR Theory
Water
e-: tetrahedral
atoms: bent linear
37
VSEPR: Rules
1.Draw the Lewis dot structure.
2.Move e pairs (bonding and lone
pairs) as far apart as possible.
3.Treat double and triple bonds as if
they were single bonds.
4.Distinguish between shape of e
pairs and molecular shape
38
Possible Molecular Shapes
linear
bent
tetrahedron trigonal pyramid trigonal planar
See problem set.
39
Predict the Molecular Shape
Hydrogen sulfide: H2S
Carbon tetrachloride: CCl4
Sulfur dioxide: SO2
Sulfur trioxide: SO3
Nitrogen tribromide: NBr3
40
Is Breaking a Bond
Endo- or Exothermic?
atoms
Endothermic! It takes
energy to break a bond.
Energy is given off when
bonds form (exothermic).
molecule
41
Bond Energy
The energy needed to break a
bond is called “bond dissociation
energy” or “bond energy”.
H—H  H + H
DH = +435 kJ (per mol)
42
Bond Dissociation Energy
Bond
H—H
C—H
C—C
C C
C C
Bond Energy Bond Length
(kJ)
(pm)
435
74
393
109
347
154
657
133
908
121
43
Bond Strength vs. Length
Multiple bonds are stronger
than single bonds.
Multiple bonds have shorter
bond lengths! Why?
C—C
weakest,
longest
C C
C C
strongest,
shortest
44
Bond Polarity
Covalent bonds involve sharing e-,
however the two bonded atoms
don’t always share equally.
45
Bond Polarity
How do you know which atom
wins the “tug-of war” for the
bonding electrons?
The more electronegative one.
Regents Table ‘S’
46
Bond Polarity
In some cases neither atom wins.
Both atoms have same electronegativity.
Hydrogen
Nitrogen
Oxygen
Chlorine
“Nonpolar”
47
Polar Molecules
EN = 2.2
H F
d+
EN = 4.0
d-
Bonding electrons shift toward F.
Thus HF is polar.
48
EN = 3.4
O
H
H
EN = 2.2
Water
Each bond in water is
polar, and the overall
molecule is polar
because of its shape
(nonsymmetrical).
Distinguish between bond polarity
and molecular polarity.
49
Nonpolar Molecules
O C O
CO2
Where is the average center of
positive charge, and where is the
center of negative charge?
Even though CO2 has polar bonds,
it is a nonpolar molecule.
50
Summary: Polar Molecules
If only 2 atoms in the molecule:
the molecule is polar if the atoms
have different electronegativities.
H-Cl
polar
Br-Br
nonpolar
51
Summary: Polar Molecules
If more than 2 atoms in molecule:
Draw dot structure
If the central atom has
• a lone pair or
• if the outside atoms are
different, the molecule is polar
(not symmetrical).
52
Summary: Polar Molecules
CH4 is nonpolar (symmetrical)
CH2F2 is polar (nonsymmetrical)
NH3 is polar (nonsymmetrical)
CH4
CH2F2
NH3
53
Polar Molecule: Yes or No?
Symmetrical
 nonpolar
Nonsymmetrical  polar
CHCl3
HI
NI3
Br2
SO3
CI4
54
Bond Polarity & Bond Type
If the difference in electronegativity
is greater than ~2.0, one atom takes
all the bonding e- and the compound
is ionic.
What types of elements would
have a big difference in EN?
55
Difference in Electronegativity
0
1.0
2.0
covalent ionic
3.0
more ionic
more covalent
H2
HCl
nonpolar
NaCl
CsF
polar
56
Identify Bond Types
nonpolar covalent
polar covalent
ionic
for these pairs of atoms:
Cs & F P & O H & Br



http://www.wimp.com/chemistrydogs/
57
Ionic vs. Covalent Compounds
Covalent
molecule
e sharing
nonmetals
D EN
Ionic
formula unit
e transfer
metal +
nonmetal
> 2.0
State @STP
M. P.
solid
high
S, L, G
low
Unit
Bond
Elements
< 2.0
58
59
Warm-up
Write the e- dot symbol
of calcium chloride.
Define the metallic bond.
How does this bonding explain why
metals conduct electricity?
Draw dot structure for methane, CH4.
60
Warm-up
What is the Lewis dot
structure of PH3 ?
61
Warm-up:
Write dot structure:
• ammonium ion
• nitrate ion
62
Warm-up
e
Draw the dot structure for NO2
and draw its resonance structures.
63
Warm-up
Predict the shape of:
SeO3 CS2 PCl3
NO2
64
Warm-up
What is the molecular shape of:
NBr3
SCl2
CO3
-2
65
Warm-up
Write the dot formula for the
bromine atom and the bromide ion.
What is the shape of the nitrate ion?
66
Warm-up
Draw the correct dot structure
and state whether the molecule is
polar or nonpolar?
• SO2
•
PBr3
•
CH2Br2
67
Warm-up
Draw a diagram showing at least
12 atoms of metallic potassium
in the solid state.
Write the dot formula
for sodium oxide.
68