Transcript JVILLAMIL - Orgsites

Chapter 6
Chemical
Bonding
and
Molecular
Structure
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
Electronegativity
The ability of an
atom in a molecule
to attract shared
electrons to itself.
Linus Pauling
1901 - 1994
Table of Electronegativities
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
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.
H
F
+

Bonding Forces
 Electron – electron
repulsive forces
 Proton – proton
repulsive forces
 Electron – proton
attractive forces
Bond Length Diagram
Bond Energy
It is the energy required to break a
bond.
It gives us information about the
strength of a bonding interaction.
Electron Dot
Notation
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
Hydrogen Chloride by the Octet Rule
Formation of Water by the Octet Rule
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.
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.
Completing a Lewis Structure -CH3Cl
Make carbon the central atom
Add up available valence electrons:
Join peripheral atoms
to the central atom
with electron pairs.
H
..
..
Complete octets on
H
atoms other than
hydrogen with remaining
electrons
C
..
H
Total = 14
..
Cl
..
..
C = 4, H = (3)(1), Cl = 7
Multiple Covalent Bonds:
Double bonds
Two pairs of shared electrons
Multiple Covalent Bonds:
Triple bonds
Three pairs of shared electrons
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.
Resonance in Ozone
Neither structure is correct.
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
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 don’t
misuse it.
Ionic Bonds
 Electrons are transferred
 Electronegativity differences are
generally greater than 1.7
 The formation of ionic bonds is
always exothermic!
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.
Table of ionic radii
VSEPR Model
(Valence Shell Electron Pair Repulsion)
The structure around a given atom
is determined principally by
minimizing electron pair repulsions.
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.
Table – VSEPR Structures
VSEPR and the water molecule
VSEPR and the ammonia molecule
VSEPR and a molecule of I3
Which structure is the correct one?
VSEPR and Xenon tetrafluoride
Which one will it be???
VSEPR and Phosphorus hexachloride
Table of dipole moments
Hybridization
The Blending of Orbitals
We have studied electron configuration notation and
the sharing of electrons in the formation of covalent
bonds.
Lets look at a
molecule of methane,
CH4.
Methane is a simple natural gas. Its molecule has a
carbon atom at the center with four hydrogen
atoms covalently bonded around it.
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?)
Carbon’s 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?
Carbon’s Empty Orbital
The first thought that
chemists had was that
carbon promotes one of
its 2s electrons…
…to the empty 2p orbital.
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.
A Problem Arises
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…?
Unequal bond energy
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.
Unequal bond energy #2
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?
Unequal bond energy #3
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.
Enter Hybridization
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.
sp3 Hybrid Orbitals
sp3 Hybrid Orbitals
Here is another way to
look at the sp3
hybridization
and energy profile…
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
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.
Relative magnitudes of forces
The types of bonding forces vary in their
strength as measured by average bond
energy.
Strongest
Covalent bonds (400 kcal)
Hydrogen bonding (12-16 kcal )
Dipole-dipole interactions (2-0.5 kcal)
Weakest
London forces (less than 1 kcal)
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.
Hydrogen
Bonding
in Water
Hydrogen Bonding between
Ammonia and Water
Dipole-Dipole
Attractions
Attraction between
oppositely charged
regions of neighboring
molecules.
The water dipole
The ammonia dipole
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
London Forces in Hydrocarbons