Chemical Bonding

Report
Chemical Bonding
Chemical Bonding




is all about the electrons.
in most of our discussions we will
concentrate on the outer ‘s’ and ‘p’
electrons. These are called the valence
electrons.
since there are 4 of these orbitals in any
quantum it requires 8 electrons (an octet) to
fill them.
the tendency of atoms to try to fill out the
outer ‘s’ and ‘p’ shells is the octet rule.
Chemical Bonding

Elements will combine in order to fill their
valence shell with electrons, like the noble
gases. This can happen in one of two ways:

To share electrons 

The outer orbitals of 2 atoms overlap so that each
atom is in the vicinity of a full set of valence electrons.
This type of bonding is called covalent bonding.
Chemical bonding




Atoms gain or lose electrons to arrive at a full set
of valence electrons.
When atoms gain or lose electrons they become
ions.
Ions are attracted to ions of opposite charge and
repelled by ions of the same charge.
This type of bonding is called ionic bonding
How Many Bonds ???
• the number of bonds made by atoms in
the ‘s’ and ‘p’ blocks of the Periodic Table
is determined by how many electrons they
are away from an octet:
How Many Bonds ???
Group
1
2
Valence
Covalent
Electrons
Bonds
1
1
2
21
13
14
15
3
4
5
31
41
3 (52)
3 (3+)
3 (3-)
16
17
18
6
7
8
2 (62)
1 (72)
0 (82)
2 (2-)
1 (1-)
0
1electron
promotion
2valence
Ionic
Bonds
1 (1+)
2 (2+)
level expansion
• most of our discussion will be centred on
covalent bonding.
Valence level expansion

Some compounds occur which cannot be
easily explained:
PF5, SF6, ClF7, ArF8


in each case the number of chemical bonds
is equal to the number of valence electrons.
this can only happen if electrons are
promoted to a higher energy level. In this
case it is the adjacent ‘d’ orbital.
PF5

the normal orbital diagram looks like
this:

with valence level expansion it looks
like this:
Valence level expansion


includes elements of groups 15 to 18,
from period 3 down; periods 1 and 2
do not have a ‘d’ orbital to promote to.
Please note that valence level
expansion is the exception, not the
rule.
Polar-Covalent Bonding


Covalent bonding implies equal
sharing of electrons.
If sharing is not equal, the electrons in
a bond will spend more time with one
atom than the other.
Polar-Covalent Bonding


The atom where the electrons spend
more time will have a net negative
charge, while the atom at the other
end of the bond will be positive.
This type of bond is polar-covalent.
Electronegativity
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

is a measure of how strongly an atom is
holding on to its valence electrons.
If an atom loses an electron fairly easily it
has a low electronegativity (and tends to be
a cation).
If an atom tends not to lose electrons, but
tends to steal them from other atoms (and
become an anion) it has a high
electronegativity.

To determine what type of bond exists
between two atoms you subtract their
respective electronegativities:


if the electronegativity difference is 0.2
or less, the bond is covalent
if the electronegativity difference is 1.7
or greater the bond is ionic.
Electronegativity


If the electronegativity difference between two
atoms is between 0.3 and 1.6 the bond is polarcovalent.
The greater the electronegativity difference the
greater the ionic character of the bond:
Assignment

Determine the electronegativity difference for each
chemical bond. If the bond is polar covalent draw an
arrow in the direction of the dipole, from positive to
negative:
C-H
END = | 2.5 - 2.1 | = 0.4
polar covalent bond
N-H
P-H
Br - Cl
B-F
Si - Cl
O-H
S-O
Cu - Br
C - Cl
N-I
C-O
N-H
• END = |3.0 – 2.1| = 0.9
• polar covalent bond:
N-H
B-F
• END = |2.0 – 4.0| = 2.0
• ionic bond:
[B]3+[F]1S–O
• END = |2.5 – 3.5| = 1.0
• polar covalent bond:
S-O
P-H
• END = |2.1 – 2.1| = 0.0
• covalent bond:
P–H
Si - Cl
• END = |1.8 – 3.0| = 1.2
• polar covalent bond:
Si - Cl
Cu - Br
• END = |1.9 – 2.8| = 0.9
• polar covalent bond:
Cu - Br
N–I
• END = |3.0 – 2.5| = 0.5
• polar covalent bond:
N-I
Br – Cl
• END = |2.8 – 3.0| = 0.2
• covalent bond:
Br – Cl
O–H
• END = |3.5 – 2.1| = 1.4
• polar covalent bond:
O–H
C – Cl
• END = |2.5 – 3.0| = 0.5
• polar covalent bond:
C – Cl
C–O
• END = |2.5 – 3.5| = 1.0
• polar covalent bond:
C–O
Covalently-Bonded Structures



we now have to consider molecules
made of several atoms.
most of the following discussion will
concern itself with molecules made
with covalent or polar-covalent bonds.
ionic bonds (and others) will return
later in the unit.
Lewis Structures


Lewis structures allow us to predict
how atoms will come together to make
molecules.
Lewis structures are representations of
molecules showing all electrons,
bonding and nonbonding.
Writing Lewis Structures
1.
PCl3
Find the sum of
valence electrons of all
atoms in the
polyatomic ion or
molecule.

5 + 3(7) = 26

If it is an anion, add one
electron for each
negative charge.
If it is a cation, subtract
one electron for each
positive charge.
Writing Lewis Structures
2.
Keep track of the electrons:
26  6 = 20
The central atom is
the least
electronegative
element that isn’t
hydrogen. Connect
the outer atoms to it
by single bonds.
Writing Lewis Structures
3.
Keep track of the electrons:
26  6 = 20  18 = 2
Fill the octets of the
outer atoms.
Writing Lewis Structures
4.
Keep track of the electrons:
26  6 = 20  18 = 2  2 = 0
Fill the octet of the
central atom.
Writing Lewis Structures
5.
If you run out of
electrons before the
central atom has an
octet…
…form multiple bonds
until it does.
Exceptions to Octet Rule

Electron Promotion



group 2 central atom will have 4 electrons
group 13 central atom will have 6 electrons
Valence Level Expansion




group 15 central atom will have 10 electrons
group 16 central atom will have 12 electrons
group 17 central atom will have 14 electrons
group 18 central atom will have 16 electrons
Writing Lewis Structures

Write Lewis Structures for the following
molecules:
F2
PCl3
SeH2
MgH2
SiF4
C 3H 8
C 3H 6
C 3H 4
Coordinate Covalent bonds

is defined as a covalent bond where one
atom provides both of the electrons:
Polyatomic Ions

we can use coordinate covalent bond
theory to explain most ions:
Resonance
This is the Lewis
structure we
would draw for
ozone, O3.
+
-
Resonance

But this is at odds
with the true,
observed structure
of ozone, in
which…


…both O—O bonds
are the same
length.
…both outer
oxygens have a
charge of 1/2.
Resonance


One Lewis structure
cannot accurately
depict a molecule
such as ozone.
We use multiple
structures, resonance
structures, to
describe the
molecule.
Resonance
Just as green is a synthesis
of blue and yellow…
…ozone is a synthesis of
these two resonance
structures.
Resonance


In truth, the electrons that form the second C—O
bond in the double bonds below do not always sit
between that C and that O, but rather can move
among the two oxygens and the carbon.
They are not localized, but rather are delocalized.
Resonance


The organic compound
benzene, C6H6, has
two resonance
structures.
It is commonly
depicted as a hexagon
with a circle inside to
signify the delocalized
electrons in the ring.
Molecular Shapes


The shape of a
molecule plays an
important role in its
reactivity.
By noting the number
of bonding and
nonbonding electron
pairs we can easily
predict the shape of
the molecule.
What Determines the Shape of
a Molecule?


Simply put, electron pairs, whether they be
bonding or nonbonding, repel each other.
By assuming the electron pairs are placed as far
as possible from each other, we can predict the
shape of the molecule.
Valence Shell Electron Pair
Repulsion Theory (VSEPR)
“The best
arrangement of a
given number of
electron domains is
the one that
minimizes the
repulsions among
them.”
Electron-Domain Geometries


All one must do is count the number of
electron domains in the Lewis structure.
The geometry will be that which corresponds
to that number of electron domains.
Molecular Geometries


The electron-domain geometry is often not
the shape of the molecule, however.
The molecular geometry is that defined by
the positions of only the atoms in the
molecules, not the nonbonding pairs.
Molecular Geometries
Within each
electron domain,
then, there might be
more than one
molecular
geometry.
Group 2 Geometries
Central
Atom
Magnesium


Bonding Lone Pair Bond Type Shape Example
Electrons Electrons
2
0
all single
linear
MgI2
one double
linear
MgO
In this domain, there is only one molecular
geometry: linear.
NOTE: If there are only two atoms in the
molecule, the molecule will be linear no
matter what the electron domain is.
Group 13 Geometries
Central Bonding Lone Pair
Atom Electrons Electrons
Boron

3
0
Bond
Type
Shape
Example
all single
trigonal
planar
BI3
one
double
linear
BIO
one triple
linear
BN
because there are no lone pair electrons the
molecular is planar (flat).
Group 14 Geometries
Central Bonding Lone Pair
Atom Electrons Electrons
Carbon
4
0
Bond
Type
Shape
Example
all single
tetrahedral CH4
one
double
trigonal
planar
one triple, linear
or
two
double
COH2
HCN
CO2
Summary for Groups 2, 13 & 14

because there are no lone pair
electrons:




central atom bonded to 1 atom: linear
central atom bonded to 2 atoms: linear
central atom bonded to 3 atoms:
trigonal planar
central atom bonded to 4 atoms:
tetrahedral
Group 15 Geometries
Central
Atom
Nitrogen
Bonding Lone Pair
Electrons Electrons
3
1
Bond
Type
Shape
all single
trigonal
pyramidal
NH3
one
double
angular
NOH
one triple linear
5
0
Example
all single
N2
trigonal
NCl5
bipyramidal
Group 16 Geometries
Central
Atom
Oxygen
Bonding Lone Pair
Electrons Electrons
2
6
2
0
Bond
Type
Shape
Example
all single
angular
H2O
one
double
linear
O2
all single
octahedral OF6
Group 17 & 18 Geometries

group 17



normally makes 1 chemical bond (linear)
with valence level expansion can make 7
bonds (ClF7).
group 18


normally makes no bonds
with valence level expansion can make 8
bonds (ArF8).
Larger Molecules


In larger molecules, it makes more sense to talk
about the geometry about a particular atom rather
than the geometry of the molecule as a whole.
This molecule is tetrahedral about the first carbon,
planar trigonal about the second and angular about
the single-bonded oxygen.
Larger Molecules
This approach
makes sense,
especially because
larger molecules
tend to react at a
particular site in the
molecule.
Using VSEPR Theory

Determine the shape of each
molecule, or atom within the molecule:
F2
PCl3
SeH2
MgH2
SiF4
C 3H 8
C 3H 6
C 3H 4
Polarity


previously we
discussed bond
dipoles.
But just because a
molecule possesses
polar bonds does not
mean the molecule as
a whole will be polar.
Polarity


By adding the
individual bond
dipoles, one can
determine the overall
dipole moment for
the molecule.
In other words we can
see if the dipoles
reinforce each other,
or cancel out.
Polarity
Assignment

Perform the following activities for the
molecules at the bottom of the slide:
1.
2.
3.
4.


Draw a lewis structure
Draw a structural diagram and indicate shape.
Determine the electronegativity difference for
each bond type and determine whether the
bond is covalent, polar covalent or ionic
Determine if they will be polar, or non-polar.
F2 PCl3 SeH2 MgH2
SiF4 C3H8 C3H6 C3H4

Complete “Fun With Balls & Sticks”
Properties of Molecules
Intermolecular Forces
Properties of Molecules
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
are determined by how the molecules
interact with each other.
How they interact is determined by the
forces of attraction between molecules.
These are called intermolecular forces;
the forces which act between molecules, to
draw them together, forming the various
phases of matter.
How do we measure force ?




the best way is by temperature.
temperature is a measure of kinetic energy;
the lowest temperature represents zero
kinetic energy.
for a substance to melt or boil the kinetic
energy must overcome the intermolecular
force.
the higher the melting temperature or boiling
temperature, the greater the intermolecular
force.
Van der Waals Forces


occur between covalently bonded
molecules.
three kinds:



London disersion
dipole-dipole attraction
hydrogen bonding
London Dispersion Forces
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

are the dominant forces between
covalently bonded, non-polar
molecules
based on the formation of
instantaneous dipoles.
the more electrons in a molecule, the
stronger the force.
Relationship of Boiling Point to Number of
Electrons and Molar Mass
Melting Point Boiling Point
# e°C
°C
CH4
C2H6
C3H8
C4H10
C8H18
10
18
26
34
66
- 182
- 183
- 190
- 138
- 57
-161
- 88
- 44
- 0.5
+125
London dispersion forces

are influenced by
shape:

Normal Pentane
(C5H12)



m.p. -130C,
b.p. 36°C
Neopentane (C5H12)


m.p. -20°C,
b.p. 9°C
Dipole-Dipole attraction



is a force which acts between polar
molecules (ex. H2S).
results from the attraction of the opposite
poles of the permanent molecular dipoles.
These substances generally have higher
melting and boiling points than non-polar
molecules with similar molecular weights (or
numbers of electrons).
Ion-Dipole Interactions


A fourth type of force, ion-dipole interactions
are an important force in solutions of ions.
The strength of these forces are what make
it possible for ionic substances to dissolve in
polar solvents.
Dipole-Dipole Interactions
The more polar the molecule, the higher
is its boiling point.
How Do We Explain This?


The nonpolar series
(SnH4 to CH4)
follow the expected
trend.
The polar series
follows the trend
from H2Te through
H2S, but water is
quite an anomaly.
Hydrogen Bonding





a specialized form of dipole-dipole attraction.
It occurs as when O, N, and F are bonded to H, owing to the
large electronegativity difference:
 O - H
3.5 - 2.1 = 1.4
 N - H
3.1 - 2.1 = 1.0
 F - H
4.1 - 2.1 = 2.0
This is a stronger force than standard dipole-dipole
attraction.
Molecules with hydrogen bonding will have boiling points
and melting points quite a bit higher than molecules that
have only dipole-dipole or London dispersion forces.
Hydrogen bonding is responsible for many of the unusual
properties of water.
Hydrogen bonding


is responsible for folding, final
structure and function of proteins
holds the DNA strands together, but
allows them to be unzipped for copying
and gene expression.
Relationship Between Polarity, Force
and Boiling Point
# e- B.P. (°C) Polarity
C2H6 18
- 161
H2S 18
- 60
H2O 10
+ 100
Force
Non-polar London
Dispersion
Polar
Dipole
- Dipole
Very polar Hydrogen Bonds
Ionic Bonding








generally occurs in compounds of metals and non-metals
(salts).
It is the result of the attraction of oppositely charged ions.
The structures formed are very orderly and are given the
name crystal lattice.
Ionic solids are called crystals.
No sharing of electrons occurs between the ions in the
crystal lattice.
As a result, ionic solids are brittle.
Ionic solids conduct electricity only in the molten state, and
not very well.
Ionic solids are characterized by very high melting and
boiling points.
Metallic Bonding






is the bonding which occurs between metals in the Periodic
Table.
It is characterized by close packing of the atoms, with the
electrons delocalized; that is, they are free to jump from
atom to atom, filling unoccupied orbitals.
This free sharing of electrons allows metals to conduct
electricity freely (copper conducts electricity 100 000 times
better than molten NaCl).
The free electrons also act as a lubricant, allowing metal
atoms to slide over one another without affecting the
integrity of the material. Thus metals are malleable and
ductile.
This bond is strong, giving most metals high melting and
boiling points.
This bond is also variable.
Network Covalent Bonding






This is the traditional covalent bond, expanded to 2 or 3
dimensions in a network which is theoretically infinite (much
like an ionic crystal lattice).
Network solids include diamonds, graphite, quartz, and
most rocks.
Because the covalent bond is stronger than any other bond,
the network solid is very hard (diamonds are the hardest
substance known).
Because electrons are held tightly in their bonds, network
solids are brittle, and they do not conduct electricity.
Because the orientation of the atoms is very specific to the
bonding orientation (tetrahedral, planar trigonal), network
solids form distinct crystals.
Because of the strength of the bonds, network solids have
very high melting and boiling points
Summarizing Intermolecular Forces
Intermolecular Forces Affect
Many Physical Properties
The strength of the
attractions between
particles can greatly
affect the properties
of a substance or
solution.
Viscosity
• Resistance of a liquid
to flow is called
viscosity.
• It is related to the ease
with which molecules
can move past each
other.
• Viscosity increases
with stronger
intermolecular forces
and decreases with
higher temperature.
Surface Tension
Surface tension
results from the net
inward force
experienced by the
molecules on the
surface of a liquid.
Vapor Pressure
• At any temperature, some molecules in a
liquid have enough energy to escape.
• As the temperature rises, the fraction of
molecules that have enough energy to
escape increases.
Vapor Pressure
As more molecules
escape the liquid,
the pressure they
exert increases.
Vapor Pressure
The liquid and vapor
reach a state of
dynamic equilibrium:
liquid molecules
evaporate and vapor
molecules condense
at the same rate.
Vapor Pressure
• The boiling point of a
liquid is the
temperature at which
its vapor pressure
equals atmospheric
pressure.
• The normal boiling
point is the
temperature at which
its vapor pressure is
760 torr.
Intermolecular Forces in
Summary
Strength of Attraction
Molecular
(London dispersion)
Molecular
(dipole-dipole)
Molecular
(hydrogen bonding)
Ionic
1 to 50
Metallic
20 to 80
Network Covalent
5
15
50
100
Melting and Boiling Point
Molecular
(London dispersion)
Molecular
(dipole-dipole)
very low (nitrogen boils at
- 196 C) VARIABLE
low (H2S boils at -61 C)
Molecular
(hydrogen bonding)
medium
(H2O boils at +100 C)
Ionic
high
(NaCl boils at +1413 C)
variable (Hg @ +357 C,
W @ +5660 C)
Metallic
Network Covalent
very high
(SiO2 boils at 2600 C)
Properties of Solids
Molecular
(London dispersion)
Molecular
(dipole-dipole)
soft, waxy
Molecular
(hydrogen bonding)
Ionic
crystalline, brittle
Metallic
short-range crystalline,
ductile, malleable
long-range crystalline,
hard, brittle
Network Covalent
more rigid
long-range crystalline,
hard, brittle
Conductance of Heat and
Electricity
Molecular
(London dispersion)
Molecular
(dipole-dipole)
non-conductor
Molecular
(hydrogen bonding)
Ionic
non-conductor
Metallic
conductor in solid or liquid
phase
non-conductor in 3d form,
some conductance in 2d
Network Covalent
non-conductor
conductor in liquid or
dissolved phase
Solubility in H2O
Molecular
(London dispersion)
Molecular
(dipole-dipole)
Not soluble
Molecular
(hydrogen bonding)
Ionic
Soluble
Metallic
Not soluble
Network Covalent
Not soluble
Soluble
Soluble

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