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Title: Chemical Bonding I: Basic Concepts


1
Chemical Bonding IBasic Concepts
  • Chapter 9

2
Intermolecular Forces
Intermolecular forces are attractive forces
between molecules.
Intramolecular forces hold atoms together in a
molecule.
  • Intermolecular vs Intramolecular
  • 41 kJ to vaporize 1 mole of water (inter)
  • 930 kJ to break all O-H bonds in 1 mole of water
    (intra)

Measure of intermolecular force boiling
point melting point DHvap DHfus DHsub
11.2
3
Valence electrons are the outer shell electrons
of an atom. The valence electrons are the
electrons that particpate in chemical bonding.
9.1
4
The Ionic Bond
He
Ne
1s22s1
1s22s22p5
1s2
1s22s22p6
9.2
5
9.1
6
A covalent bond is a chemical bond in which two
or more electrons are shared by two atoms.
Lewis structure of F2
9.4
7
Lewis structure of water


Double bond two atoms share two pairs of
electrons
or
Triple bond two atoms share three pairs of
electrons
or
9.4
8
Lengths of Covalent Bonds
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 lt Double Bond lt Single
Bond
9.4
9
9.4
10
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. Complete an octet for all atoms except hydrogen
  4. If structure contains too many electrons, form
    double and triple bonds on central atom as needed.

9.6
11
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
and complete octets on N and F
atoms.
Step 4 - Check, are of e- in structure equal
to number of valence e- ?
3 single bonds (3x2) 10 lone pairs (10x2) 26
valence electrons
9.6
12
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 2e-
4 (3 x 6) 2 24 valence electrons
Step 3 Draw single bonds between C and O atoms
and complete octet on C and O
atoms.
Step 4 - Check, are of e- in structure equal
to number of valence e- ?
3 single bonds (3x2) 10 lone pairs (10x2) 26
valence electrons
Step 5 - Too many electrons, form double bond
and re-check of e-
9.6
13
Two possible skeletal structures of formaldehyde
(CH2O)
An atoms formal charge is the difference between
the number of valence electrons in an isolated
atom and the number of electrons assigned to that
atom in a Lewis structure.
The sum of the formal charges of the atoms in a
molecule or ion must equal the charge on the
molecule or ion.
9.7
14
-1
1
formal charge on C
4 -
2 -
½ x 6 -1
formal charge on O
6 -
2 -
½ x 6 1
9.7
15
0
0
formal charge on C
4 -
0 -
½ x 8 0
formal charge on O
6 -
4 -
½ x 4 0
9.7
16
Formal Charge and Lewis Structures
  1. For neutral molecules, a Lewis structure in which
    there are no formal charges is preferable to one
    in which formal charges are present.
  2. Lewis structures with large formal charges are
    less plausible than those with small formal
    charges.
  3. Among Lewis structures having similar
    distributions of formal charges, the most
    plausible structure is the one in which negative
    formal charges are placed on the more
    electronegative atoms.

9.7
17
Predicting Molecular Geometry
  1. Draw Lewis structure for molecule.
  2. Count number of lone pairs on the central atom
    and number of atoms bonded to the central atom.
  3. Use VSEPR to predict the geometry of the molecule.

AB4E
AB2E
distorted tetrahedron
bent
10.1
18
Valence shell electron pair repulsion (VSEPR)
model
Predict the geometry of the molecule from the
electrostatic repulsions between the electron
(bonding and nonbonding) pairs.
AB2
2
0
10.1
19
10.1
20
VSEPR
AB2
2
0
linear
linear
AB3
3
0
10.1
21
10.1
22
VSEPR
AB2
2
0
linear
linear
AB4
4
0
10.1
23
10.1
24
VSEPR
AB2
2
0
linear
linear
AB4
4
0
tetrahedral
tetrahedral
AB5
5
0
10.1
25
10.1
26
VSEPR
AB2
2
0
linear
linear
AB4
4
0
tetrahedral
tetrahedral
AB6
6
0
10.1
27
10.1
28
10.1
29
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30
VSEPR
trigonal planar
trigonal planar
AB3
3
0
AB2E
2
1
10.1
31
VSEPR
AB4
4
0
tetrahedral
tetrahedral
AB3E
3
1
10.1
32
VSEPR
AB4
4
0
tetrahedral
tetrahedral
AB2E2
2
2
10.1
33
VSEPR
trigonal bipyramidal
trigonal bipyramidal
AB5
5
0
AB4E
4
1
10.1
34
VSEPR
trigonal bipyramidal
trigonal bipyramidal
AB5
5
0
AB3E2
3
2
10.1
35
VSEPR
trigonal bipyramidal
trigonal bipyramidal
AB5
5
0
AB2E3
2
3
10.1
36
VSEPR
AB5E
5
1
10.1
37
VSEPR
AB4E2
4
2
10.1
38
10.1
39
A resonance structure is one of two or more Lewis
structures for a single molecule that cannot be
represented accurately by only one Lewis
structure.
9.8
40
Exceptions to the Octet Rule
The Incomplete Octet
BeH2
BF3
9.9
41
Exceptions to the Octet Rule
Odd-Electron Molecules
NO
The Expanded Octet (central atom with principal
quantum number n gt 2)
SF6
9.9
42
Classification of bonds by difference in
electronegativity
Difference
Bond Type
0
Covalent
? 2
Ionic
0 lt and lt2
Polar Covalent
9.5
43
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
9.5
44
Polar covalent bond or polar bond is a covalent
bond with greater electron density around one of
the two atoms
electron rich region
electron poor region
e- rich
e- poor
d
d-
9.5
45
Electronegativity is the ability of an atom to
attract toward itself the electrons in a chemical
bond.
Electron Affinity - measurable, Cl is highest
Electronegativity - relative, F is highest
9.5
46
9.5
47
Dipole Moments and Polar Molecules
electron rich region
electron poor region
m Q x r
Q is the charge
r is the distance between charges
1 D 3.36 x 10-30 C m
10.2
48
10.2
49
10.2
50
dipole moment polar molecule
dipole moment polar molecule
no dipole moment nonpolar molecule
no dipole moment nonpolar molecule
10.2
51
10.2
52
Intermolecular Forces
Dipole-Dipole Forces
Attractive forces between polar molecules
11.2
53
Intermolecular Forces
Ion-Dipole Forces
Attractive forces between an ion and a polar
molecule
11.2
54
11.2
55
Intermolecular Forces
Dispersion Forces
Attractive forces that arise as a result of
temporary dipoles induced in atoms or molecules
ion-induced dipole interaction
dipole-induced dipole interaction
11.2
56
Intermolecular Forces
Dispersion Forces Continued
Polarizability is the ease with which the
electron distribution in the atom or molecule can
be distorted.
  • Polarizability increases with
  • greater number of electrons
  • more diffuse electron cloud

11.2
57
What type(s) of intermolecular forces exist
between each of the following molecules?
HBr
HBr is a polar molecule dipole-dipole forces.
There are also dispersion forces between HBr
molecules.
CH4
CH4 is nonpolar dispersion forces.
SO2
SO2 is a polar molecule dipole-dipole forces.
There are also dispersion forces between SO2
molecules.
11.2
58
Intermolecular Forces
Hydrogen Bond
The hydrogen bond is a special dipole-dipole
interaction between they hydrogen atom in a polar
N-H, O-H, or F-H bond and an electronegative O,
N, or F atom.
A B are N, O, or F
11.2
59
Hydrogen Bond
11.2
60
Why is the hydrogen bond considered a special
dipole-dipole interaction?
11.2
61
Types of Crystals
  • Molecular Crystals
  • Lattice points occupied by molecules
  • Held together by intermolecular forces
  • Soft, low melting point
  • Poor conductor of heat and electricity

11.6
62
Types of Crystals
  • Metallic Crystals
  • Lattice points occupied by metal atoms
  • Held together by metallic bonds
  • Soft to hard, low to high melting point
  • Good conductors of heat and electricity

Cross Section of a Metallic Crystal
11.6
63
Types of Crystals
  • Ionic Crystals
  • Lattice points occupied by cations and anions
  • Held together by electrostatic attraction
  • Hard, brittle, high melting point
  • Poor conductor of heat and electricity

CsCl
ZnS
CaF2
11.6
64
Electrostatic (Lattice) Energy
Lattice energy (E) is the energy required to
completely separate one mole of a solid ionic
compound into gaseous ions.
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.
r F lt r Cl
9.3
65
Types of Crystals
  • Covalent Crystals
  • Lattice points occupied by atoms
  • Held together by covalent bonds
  • Hard, high melting point
  • Poor conductor of heat and electricity

diamond
graphite
11.6
66
Types of Crystals
11.6
67
Hybridization mixing of two or more atomic
orbitals to form a new set of hybrid orbitals.
  • Mix at least 2 nonequivalent atomic orbitals
    (e.g. s and p). Hybrid orbitals have very
    different shape from original atomic orbitals.
  • Number of hybrid orbitals is equal to number of
    pure atomic orbitals used in the hybridization
    process.
  • Covalent bonds are formed by
  • Overlap of hybrid orbitals with atomic orbitals
  • Overlap of hybrid orbitals with other hybrid
    orbitals

10.4
68
10.4
69
10.4
70
10.4
71
Formation of sp Hybrid Orbitals
10.4
72
Formation of sp2 Hybrid Orbitals
10.4
73
Count the number of lone pairs AND the number of
atoms bonded to the central atom
of Lone Pairs of Bonded Atoms
Hybridization
Examples
2
sp
BeCl2
3
sp2
BF3
4
sp3
CH4, NH3, H2O
5
sp3d
PCl5
6
sp3d2
SF6
10.4
74
Valence Bond Theory and NH3
N 1s22s22p3
3 H 1s1
If use the 3 2p orbitals predict 900
Actual H-N-H bond angle is 107.30
10.4
75
10.5
76
10.5
77
Sigma (s) and Pi Bonds (p)
1 sigma bond
Single bond
1 sigma bond and 1 pi bond
Double bond
Triple bond
1 sigma bond and 2 pi bonds
s bonds 6
1 7
p bonds 1
10.5
78
The enthalpy change required to break a
particular bond in one mole of gaseous molecules
is the bond energy.
Bond Energy
9.10
79
Average bond energy in polyatomic molecules
9.10
80
Bond Energies (BE) and Enthalpy changes in
reactions
Imagine reaction proceeding by breaking all bonds
in the reactants and then using the gaseous atoms
to form all the bonds in the products.
DH0 total energy input total energy released
SBE(reactants) SBE(products)
9.10
81
9.10
82
DH0 SBE(reactants) SBE(products)
DH0 436.4 156.9 2 x 568.2 -543.1 kJ
9.10
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