Bonding and Molecular Structure:

1
Chapter 9

• Bonding and Molecular Structure
• Fundamental Concepts

2
Chemical Bonding

• Problems and questions
• How is a molecule or polyatomic ion held
  together?
• Why are atoms distributed at strange angles?
• Why are molecules not flat?
• Can we predict the structure?
• How is structure related to chemical and physical
  properties?

3
Structure Bonding
NN triple bond. Molecule is unreactive
White phosphorus is a tetrahedron of P atoms.
Very reactive!
Red phosphorus, a polymer. Used in matches.
Less reactive!
4
Forms of Chemical Bonds

• There are 2 extreme forms of connecting or
  bonding atoms

Ioniccomplete transfer of 1 or more electrons
from one atom to another
Covalentsome valence electrons shared between
atoms
Most bonds are somewhere in between.
5
Ionic Compounds
Metal of low IE
Nonmetal of high EA
2 Na(s) Cl2(g) ---gt 2 Na 2 Cl-
6
Covalent Bonding

• The bond arises from the mutual attraction of 2
  nuclei for the same electrons. Electron sharing
  results.

Bond is a balance of attractive and repulsive
forces.
7
Bond Formation

• A bond can result from a head-to-head overlap
  of atomic orbitals on neighboring atoms.

Note that each atom has a single, unpaired
electron.
Overlap of H (1s) and Cl (3p)
8
Chemical Bonding Objectives

• Objectives are to understand
• 1. valence e- distribution in molecules and ions.
• 2. molecular structures
• 3. bond properties and their effect on
  molecular properties.

9
Electron Distribution in Molecules

• Electron distribution is depicted with
  Lewis electron dot structures
• Valence electrons are distributed as shared or
  BOND PAIRS and unshared or
  LONE PAIRS.

10
Bond and Lone Pairs

• Valence electrons are distributed as shared or
  BOND PAIRS and unshared or LONE PAIRS.

lone pairs
This is called a LEWIS ELECTRON DOT structure.
11
Valence Electrons

• Electrons are divided between core and valence
  electrons
• B 1s2 2s2 2p1
• Core He , valence 2s2 2p1

Br Ar 3d10 4s2 4p5 Core Ar 3d10 ,
valence 4s2 4p5
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Rules of the Game

• No. of valence electrons of a main group atom
  Group number

For Groups 1A-4A, no. of bond pairs group
number.
For Groups 5A -7A, BPs 8 - Grp. No.
Group 5A
Group 3A
13
Rules of the Game

• No. of valence electrons of an atom Group
  number
• For Groups 1A-4A, no. of bond pairs group
  number
• For Groups 5A -7A, BPs 8 - Grp. No.

Except for H (and sometimes atoms of 3rd and
higher periods), BPs LPs 4
This observation is called the OCTET RULE
14
Hydrophobic vs. Hydrophilic

• Hydrophobic - translation of Greek water fear
• Hydrophilic translation of Greek water
  friendship

15
Building a Dot Structure

• Ammonia, NH3
• 1. Decide on the central atom never H.
• Central atom is atom of lowest affinity for
  electrons.
• Therefore, N is central
• 2. Count valence electrons
• H 1 and N 5
• Total (3 x 1) 5
• 8 electrons / 4 pairs

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Building a Dot Structure

• 3. Form a single bond between the central atom
  and each surrounding atom

4. Remaining electrons form LONE PAIRS to
complete octet as needed.
3 BOND PAIRS and 1 LONE PAIR.
Note that N has a share in 4 pairs (8 electrons),
while H shares 1 pair.
17
Lewis Structures

• CH4 methane

Step 1. Central atom C
Step 2. Count valence electrons
C 4
4 x H 4 x 1 4
TOTAL 8 e- or 4 pairs
Step 3. Form bonds
18
Lewis Structures

• C2H6 ethane

19
Multiple Covalent Bonds

• double bond gt 2 pairs shared
• triple bond gt 3 pairs shared
• normally occurs between
• C atoms N atoms O atoms
• a C atom and a N, O or S atom
• a N atom and a O or S atom
• a S atom and an O atom

20
Double and even triple bonds are commonly
observed for C, N, P, O, and S
H2CO
SO3
C2F4
21
Carbon Dioxide, CO2
C

• 1. Central atom _______

4
2
2. Valence electrons __ or __ pairs
3. Form bonds.
This leaves 6 pairs.
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Carbon Dioxide, CO2

• 4. Place lone pairs on outer atoms.

5. So that C has an octet, we shall form DOUBLE
BONDS between C and O.
The second bonding pair forms a pi (p) bond.
23
Steps to form Lewis Electron Dot Structure

• 1. Central atom _______
• 2. Valence electrons __ or __ pairs
• Form bonds.
• Place lone pairs on outer atoms.
• Form multiple bonds as necessary to obey Lewis
  Octet Rule.
• Remember that there are MANY compounds that do
  not obey the Octet Rule.

24
Lewis Structures

• CO carbon monoxide

25
Exceptions to Octet Rule

• NO nitric oxide

26
Exceptions to Octet Rule

• NO2 nitrogen dioxide

resonance
27
Exceptions to Octet Rule

• PF5
• expanded octet

28
Exceptions to Octet Rule

• expanded octet
• SF4

29
Exceptions to Octet Rule

• SF6 expanded octet

30
Formal Atom Charges

• Atoms in molecules often bear a charge ( or
  -).
• The predominant resonance structure of a molecule
  is the one with charges as close to 0 as
  possible.
• Formal charge Group number 1/2 (no. of
  bonding electrons) - (no. of LP electrons)

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Carbon Dioxide, CO2
32
Thiocyanate Ion, SCN-
6 - (1/2)(2) - 6 -1
5 - (1/2)(6) - 2 0
4 - (1/2)(8) - 0 0
33
Thiocyanate Ion, SCN-
Which is the most important resonance form?
34
MOLECULAR GEOMETRY
35
MOLECULAR GEOMETRY

• VSEPR
• Valence Shell Electron Pair Repulsion theory.
• Most important factor in determining geometry is
  relative repulsion between electron pairs.

Molecule adopts the shape that minimizes the
electron pair repulsions.
36
Electron Pair Geometries
37
Sulfur Dioxide, SO2

• 1. Central atom S
• 2. Valence electrons 18 or 9 pairs

3. Form double bond so that S has an octet
but note that there are two ways of doing this.
38
Sulfur Dioxide, SO2

• This leads to the following structures.

These equivalent structures are called RESONANCE
STRUCTURES. The true electronic structure is a
HYBRID of the two.
39
Geometries for Four Electron Pairs
40
Structure Determination by VSEPR

• Ammonia, NH3
• 1. Draw electron dot structure
• 2. Count BPs and LPs 4

3. The 4 electron pairs are at the corners of a
tetrahedron.
41
Structure Determination by VSEPR

• Ammonia, NH3
• The electron pair geometry is tetrahedral.

The MOLECULAR GEOMETRY, the positions of the
atoms, is PYRAMIDAL.
42
Structure Determination by VSEPR

• Water, H2O
• 1. Draw electron dot structure

2. Count BPs and LPs 4
3. The 4 electron pairs are at the corners of a
tetrahedron.
The electron pair geometry is TETRAHEDRAL.
The molecular geometry is BENT.
43
Consequences of H2O Polarity
44
Structures with Central Atoms with More Than or
less Than 4 Electron Pairs
Often occurs with Group 3A elements and with
those of 3rd period and higher.
45
Molecular Geometries for Five Electron Pairs
46
Molecular Geometries for Six Electron Pairs
47
Bond Properties

• What is the effect of bonding and structure on
  molecular properties?

Free rotation around CC single bond
No rotation around CC double bond
48
Bond Order of bonds between a pair of atoms
Acrylonitrile
49
Bond Order

• Fractional bond orders occur in molecules with
  resonance structures.
• Consider NO2-

1
1
The NO bond order 1.5
50
Bond Order

• Bond order is proportional to two important bond
  properties
• (a) bond strength
• (b) bond length

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(No Transcript)
52

• Compare O-O and OO. Is OO expected to be
• stronger, weaker, or the same strength?

53

• Is OO expected to be
• longer, shorter, or the same length?

54
Using Bond Energies

• Estimate the energy of the reaction
• HH ClCl ----gt 2 HCl

HH 436 kJ/mol ClCl 242 kJ/mol HCl
432 kJ/mol
Net ?H S bondsbroken S bondsformed
Sum of H-H Cl-Cl bond energies 436 kJ 242
kJ 678 kJ
2 mol H-Cl bond energies 864 kJ
Net ?H 678 kJ - 864 kJ -186 kJ
?Hfo (HCl(g)) -92.31 kJ/mol or -184 kJ
55
Molecular Polarity
WaterBoiling point 100 C
MethaneBoiling point -161 C
Why do water and methane differ so much in their
boiling points?

• Why do ionic compounds dissolve in water?

56
Bond Polarity

• HCl is POLAR because it has a positive end and a
  negative end.

Cl has a greater share in bonding electrons than
does H.
Cl has slight negative charge (-d) and H has
slight positive charge ( d)
57
Bond Polarity

• Three molecules with polar, covalent bonds.
• Each bond has one atom with a slight negative
  charge (-d) and another with a slight positive
  charge ( d)

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Linus Pauling, 1901-1994

• The only person to receive two unshared Nobel
  prizes (for Peace and Chemistry).
• Chemistry areas bonding, electronegativity,
  protein structure

59
Electronegativity, ?

• ? is a measure of the ability of an atom in a
  molecule to attract electrons to itself.

60
Bond Polarity

• Due to the bond polarity, the HCl bond energy is
  GREATER than expected for a pure covalent bond.

BOND ENERGY pure bond 339 kJ/mol calcd real
bond 432 kJ/mol measured
Difference 92 kJ. This difference is
proportional to the difference in
ELECTRONEGATIVITY, ?.
61
Electronegativity

• Pauling Scale
• relative attraction of an atom for electrons, its
  own and those of other atoms
• same trends as ionization energy, increases from
  lower left corner to the upper right corner
• fluorine E.N. 4.0

62
Electronegativities of the Elements
63
Bond Polarity

• Which bond is more polar (or DIPOLAR)?
• OH OF
• ? 3.5 - 2.1 3.5 - 4.0
• ? 1.4 0.5
• OH is more polar than OF

and polarity is reversed.
64
Molecular Polarity

• Molecules will be polar if
• a) bonds are polar
• AND
• b) the molecule is NOT symmetric

All above are NOT polar
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Polar or Nonpolar?

• Compare CO2 and H2O. Which one is polar?

66
Covalent Bond Properties

• electronegativity
• nonpolar bonds gt diff. EN 0
• polar bonds gt diff. EN gt 0
• ionic bonds gt diff. EN gt 1.5

67
CH4 CCl4 Polar or Not?

• Only CH4 and CCl4 are NOT polar. These are the
  only two molecules that are symmetrical.