Chapter 6 Chemical Bonding and Molecular Structure

1
Chapter 6 Chemical Bonding and Molecular
Structure
2
Bonds

• Forces that hold groups of atoms
• together and make them function
• as a unit.

• Ionic bonds transfer of electrons
• Covalent bonds sharing of electrons

3
Electronegativity

• The ability of an atom in a molecule to attract
  shared electrons to itself.

Linus Pauling 1901 - 1994
4
Table of Electronegativities
5
Covalent Bonds
Polar-Covalent bonds

• Electrons are unequally shared
• Electronegativity difference between .3 and 1.7

Nonpolar-Covalent bonds

• Electrons are equally shared
• Electronegativity difference of 0 to 0.3

6
Polarity

• A molecule, such as HF, that has a center of
  positive charge and a center of negative charge
  is said to be polar, or to have a dipole moment.

7
Bonding Forces

• Electron electron
• repulsive forces

• Proton proton
• repulsive forces

• Electron proton
• attractive forces

8
Bond Length Diagram
9
Bond Energy

• It is the energy required to break a bond.
• It gives us information about the strength of a
  bonding interaction.

10
Electron Dot Notation
11
The Octet Rule
Chemical compounds tend to form so that each
atom, by gaining, losing, or sharing electrons,
has an octet of electrons in its highest occupied
energy level.
Diatomic Fluorine
12
Hydrogen Chloride by the Octet Rule
13
Formation of Water by the Octet Rule
14
Comments About the Octet Rule

• 2nd row elements C, N, O, F observe the octet
  rule.
• 2nd row elements B and Be often have fewer than 8
  electrons around themselves - they are very
  reactive.
• 3rd row and heavier elements CAN exceed the octet
  rule using empty valence d orbitals.
• When writing Lewis structures, satisfy octets
  first, then place electrons around elements
  having available d orbitals.

15
Lewis Structures

• Shows how valence electrons are arranged among
  atoms in a molecule.
• Reflects central idea that stability of a
  compound relates to noble gas electron
  configuration.

16
Completing a Lewis Structure -CH3Cl

• Make carbon the central atom

• Add up available valence electrons

• C 4, H (3)(1), Cl 7 Total 14

• Join peripheral atoms
• to the central atom
• with electron pairs.

H
..
..
..
H
..
C

• Complete octets on
• atoms other than
• hydrogen with remaining
• electrons

..
Cl
..
..
H
17
Multiple Covalent BondsDouble bonds
Two pairs of shared electrons
18
Multiple Covalent BondsTriple bonds
Three pairs of shared electrons
19
Resonance

• Occurs when more than one valid Lewis structure
  can be written for a particular molecule.

• These are resonance structures.
• The actual structure is an average of
• the resonance structures.

20
Resonance in Ozone
Neither structure is correct.
21
Models

• Models are attempts to explain how nature
  operates on the microscopic level based on
  experiences in the macroscopic world.

Models can be physical as with this DNA model
Models can be mathematical
Models can be theoretical or philosophical
22
Fundamental Properties of Models

• A model does not equal reality.
• Models are oversimplifications, and are therefore
  often wrong.
• Models become more complicated as they age.
• We must understand the underlying assumptions in
  a model so that we dont misuse it.

23
Ionic Bonds

• Electrons are transferred

• Electronegativity differences are
• generally greater than 1.7
• The formation of ionic bonds is
• always exothermic!

24
Sodium Chloride Crystal Lattice
Ionic compounds form solids at ordinary
temperatures.
Ionic compounds organize in a characteristic
crystal lattice of alternating positive and
negative ions.
25
Table of ionic radii
26
VSEPR Model
(Valence Shell Electron Pair Repulsion)

• The structure around a given atom is determined
  principally by minimizing electron pair
  repulsions.

27
Predicting a VSEPR Structure

• Draw Lewis structure.
• Put pairs as far apart as possible.
• Determine positions of atoms from the way
  electron pairs are shared.
• Determine the name of molecular structure from
  positions of the atoms.

28
Table VSEPR Structures
29
VSEPR and the water molecule
30
VSEPR and the ammonia molecule
31
VSEPR and a molecule of I3
Which structure is the correct one?
32
VSEPR and Xenon tetrafluoride
Which one will it be???
33
VSEPR and Phosphorus hexachloride
34
Table of dipole moments
35
Hybridization

• The Blending of Orbitals

36
Lets look at amolecule of methane, CH4.
We have studied electron configuration notation
and the sharing of electrons in the formation of
covalent bonds.
Methane is a simple natural gas. Its molecule has
a carbon atom at the center with four hydrogen
atoms covalently bonded around it.
37
Carbon ground state configuration
What is the expected orbital notation of
carbon in its ground state?
Can you see a problem with this?
(Hint How many unpaired electrons does this
carbon atom have available for bonding?)
38
Carbons Bonding Problem
You should conclude that carbon only has TWO
electrons available for bonding. That is not not
enough!
How does carbon overcome this problem so that it
may form four bonds?
39
Carbons Empty Orbital
The first thought that chemists had was that
carbon promotes one of its 2s electrons
to the empty 2p orbital.
40
A Problem Arises
However, they quickly recognized a problem with
such an arrangement
Three of the carbon-hydrogen bonds would
involve an electron pair in which the carbon
electron was a 2p, matched with the lone 1s
electron from a hydrogen atom.
41
Unequal bond energy
This would mean that three of the bonds in a
methane molecule would be identical, because they
would involve electron pairs of equal energy.
But what about the fourth bond?
42
Unequal bond energy 2
The fourth bond is between a 2s electron from
the carbon and the lone 1s hydrogen electron.
Such a bond would have slightly less energy than
the other bonds in a methane molecule.
43
Unequal bond energy 3
This bond would be slightly different in
character than the other three bonds in methane.
This difference would be measurable to a
chemist by determining the bond length and bond
energy.
But is this what they observe?
44
Enter Hybridization
The simple answer is, No.
Measurements show that all four bonds in methane
are equal. Thus, we need a new explanation for
the bonding in methane.
Chemists have proposed an explanation they call
it Hybridization.
Hybridization is the combining of two or more
orbitals of nearly equal energy within the same
atom into orbitals of equal energy.
45
sp3 Hybrid Orbitals
In the case of methane, they call the
hybridization sp3, meaning that an s orbital is
combined with three p orbitals to create four
equal hybrid orbitals.
These new orbitals have slightly MORE energy
than the 2s orbital
and slightly LESS energy than the 2p orbitals.
46
sp3 Hybrid Orbitals
Here is another way to look at the sp3
hybridization and energy profile
47
sp Hybrid Orbitals
While sp3 is the hybridization observed in
methane, there are other types of hybridization
that atoms undergo.
These include sp hybridization, in which one s
orbital combines with a single p orbital.
Notice that this produces two hybrid orbitals,
while leaving two normal p orbitals
48
sp2 Hybrid Orbitals
Another hybrid is the sp2, which combines two
orbitals from a p sublevel with one orbital from
an s sublevel.
Notice that one p orbital remains unchanged.
49
Relative magnitudes of forces
The types of bonding forces vary in their
strength as measured by average bond energy.
Strongest Weakest
Covalent bonds (400 kcal)
Hydrogen bonding (12-16 kcal )
Dipole-dipole interactions (2-0.5 kcal)
London forces (less than 1 kcal)
50
Hydrogen Bonding
Bonding between hydrogen and more
electronegative neighboring atoms such as oxygen
and nitrogen
Hydrogen bonding in Kevlar, a strong polymer used
in bullet-proof vests.
51
Hydrogen Bonding in Water
52
Hydrogen Bonding between Ammonia and Water
53
Dipole-Dipole Attractions
Attraction between oppositely charged regions of
neighboring molecules.
54
The water dipole
55
The ammonia dipole
56
London Dispersion Forces
The temporary separations of charge that lead to
the London force attractions are what attract one
nonpolar molecule to its neighbors.
London forces increase with the size of the
molecules.
Fritz London 1900-1954
57
London Forces in Hydrocarbons