[MUSIC]
Lighthouse Scientific
Education presents a lecture
in the Solutions
and Water series.
The topic;
Intermolecular Forces.
Material in this lecture relies
on an understanding of the
previous lectures Fundamental
Properties of Matter,
Bonding: An Overview,
Covalent Bonding and
Ions and Isotopes.
The lecture begins with
a brief description of
intermolecular forces to gives
some perspective
of the subject.
It includes bonding
types already familiar to
the student and a helpful look
at how to evaluate trends seen
with intermolecular bonding.
Before deepening the
investigation there will be
a review of the properties
of matter that underline
intermolecular forces.
An extensive review
is given due to the
advanced nature of the topic.
Students who have a strong
foundation in the subject
matter of ion formation and
polarity should skip ahead to
the ranking of
intermolecular forces.
The review includes a
look at atomic forces
Since the same
forces are at work at
intramolecular bonding
that are found in
intermolecular bonding it
will be helpful to look at
the role of electronegativity
and polarity.
Building on the review is
a progressive discussion of
intermolecular forces.
Moving from
strongest intermolecular
forces to the weakest there is
ion-ion, ion-dipole,
hydrogen bonds, dipole-dipole
ion or dipole-induced
dipoles and finally,
London or dispersion forces.
To give some perspective
to intermolecular forces we
can ask "What type of
bonds have we come across
so far in our studies?"
Mostly, we have seen
intramolecular bonding with
the term intra meaning within.
Intramolecular bonds hold
atoms together WITHIN a
molecule or compound.
To a lesser degree we have seen
intermolecular bonding
with inter meaning among or
between. Intermolecular
bonds are between
compounds or elements.
For the most part
intermolecular
bonds are stronger.
It then follows that
intramolecular bonding
is relatively weak.
Covalent, metallic and ionic are
types of intramolecular bonds.
The carbon-hydrogen bond
in methane is an example
of a covalent bond.
The delocalized electrons
around a core of metallic
nuclei with their inner shell
electrons is a representation
of metallic bonds.
And the crystal lattice of
solid sodium chloride
serves as an example of
ionic intramolecular bonding.
The ionic bonding type
actually finds itself
straddling both the definitions
of intra and intermolecular
bonds. That is,
depending on the context
of the discussion it can be
considered in either camp.
When an ionic bond is
spoken of as a compound than
the interactions between ions
is within and intramolecular.
But an ion itself can be
considered an entity on
its own and then an ion-ion
interaction is between
particles and intermolecular.
As we see, the strength of
an intermolecular bond is,
for the most part, based
on the charge, permanent
or temporary, of the
participating species.
If we look at the cast of
characters that are involved
in intermolecular bonding
the one that makes the
strongest bonds is the ion.
Ions have a large permanent
charge which readily interacts
with its neighboring particles.
Other players include
the permanent dipole.
Although it has a weaker
charge than an ion it to
readily interacts with
neighboring particles.
The only really
new character in
this lecture is
the induced dipole.
It is a short lived
dispersion of charge and it is
found in almost all
matter-matter interactions.
All three of these
species will be discussed
in detail along with
the interactions found
by pairing these 3 in
different combinations.
Some terminology clarification
is useful at this point.
Starting with
intermolecular forces.
They are the electrostatic
attraction or repulsion of
charged particles. Charge
being positive or negative.
The charges can be full
charges as seen on ions or
partial charges of dipoles.
As such the strength of
an intermolecular force
is dependent on the
sizes of the charges.
Larger the
charge-stronger the force.
Weaker the
charge-weaker the force.
And then there is the
intermolecular bond.
These are the electrostatic
interactions between different
atomic scale particles.
That includes elements,
ions, and compounds.
When dealing with
two particles, intermolecular
bond, intermolecular force
and intermolecular
interaction are generally
talking about the same thing.
To give some perspective on
the role intermolecular bonds
in chemistry we will look at
the trend these bonds produce
in phase transition.
More specifically,
the effects of intermolecular
bonding can be tracked by
following the temperature
of phase transitions.
Water will stand in as an
example molecule and its 3
familiar phases are
solid, liquid and gas.
A phase transition is moving
from one of the phases to
another and can be viewed
as the making or breaking of
intermolecular bonds.
The molecule that
makes up ice, liquid
water and steam is H20
and is not itself changed
in moving between states.
The states are based on
the degree to which the
water molecules
interact with each other.
All things being equal, with
the addition of heat there is
a characteristic temperature
at which ice melts.
It is called the melting point.
There is a
characteristic temperature
at which liquid water boils.
It is called the boiling point.
There's also transition
that occur with
the removal of heat.
In the 'Fundamental Properties
of Matter' lecture it was
argued that adding heat
adds motion to the molecules
and transitioning from
solid to liquid occurs
when the energy of motion
exceeds, to a certain degree,
the energy of
particle-particle interactions;
intermolecular bonds.
When enough of
the intermolecular bonds
have been broken the particle
begin to slide past each other.
That is liquid.
When enough additional heat
is added to break all of the
intermolecular bonds the
particles fly apart. That is gas.
Since the melting and boiling
points reflect the balance
between the intermolecular
forces between the
particles and the energy
of motion due to the heat
it reasonably follows
that the stronger, or more,
intermolecular bonds
will push melting points
and boiling points higher.
Stronger intermolecular
interactions require more
energy or heat to break
the bonds that are necessary
for a phase
transition to occur.
In short, substances with
strong intermolecular forces
melt and boil at higher
temperatures than substances
with fewer such forces.
As always there
are other factors
that can come into play.
Geometry can most certainly
affect melting and boiling
points of a substance.
Phase transition is
useful in giving context to
intermolecular bonding.
Now, we are going to step back
and build up the
foundation of intermolecular
forces starting with a review
of how and why atomic scale
particles interact. That
means going all the way back
to the structure of the atom.
Consider two atoms:
Atom A and Atom B.
Drawn here are just the
positively charge nuclei
of the atoms. But atoms
also have negatively charged
electrons zipping about in
shells around the nucleus.
Often this is shown
as an electron cloud.
Now, electrostatic interactions
results from attractive or
repulsive forces between
atoms, ions or compounds.
That is to say that the
negatively charged electrons
in atom A are ATRRACTED to
the positively charged nucleus
in atom B. Also, the
negatively charged electrons
in atom B are ATRRACTED to the
positively charged
nucleus in atom A.
Conversely, negatively
charged electrons in atom
A and B REPULSE each other
and positively charged
nuclei in atoms A and
B REPULSE each other.
Opposite charges attract.
Like changes repel.
Ultimately all atomic scale
interactions are going to come
down to a combination of these
4 electrostatic interactions.
As long as there is a net
attraction between particles
some form of
bonding will occur.
Furthering our
development of a base for
intermolecular forces is a
review of what we know about
intramolecular bonds. The
interaction between molecules
and compounds will
to a large degree
be determined by what type
of intramolecular bonds.
From the 'Bonding: an
Overview' lecture we
remember that intramolecular
bonds can be broadly
categorized as covalent,
polar covalent or ionic.
What distinguishes
the bond type
is how well they
share electrons?
Importantly, sharing means
time! How long does one atom
in the bond keep the
shared electrons compared
to the other atom?
Using a graphical
representation of sharing
between hypothetical
atoms A and B.
In a covalent bond
electrons spend
equal time or near equal
time around atom A and B.
In polar bonds the
sharing is less equal.
Electrons spend more
time around atom A than B.
This uneven distribution
gives a rise to a dipole which
can participate in
intermolecular bonding.
And for ionic bonding there
is no sharing. Electrons spend
ALL the time around atom A.
The transfer of the electron
from B to A causes
the formation of ions.
The charged ion is
a potent player in
intermolecular bonding.
How do we decide
which category to put
an intermolecular bond?
Quite simply it falls to
the electronegativity of the
atom in the bond.
Electronegativity, or EN,
is the capacity of an atom in
a molecule to attract shared
electrons from another atom.
Your textbook should
have a Periodic Table that
has laboratory determined EN
values for most of the atoms.
Electronegativity has been
covered several times before
but is of such importance that
we will quickly run through
how the categorization of a
polar or ionic bond is made.
delta EN is the difference
in electronegativity between
the atoms in the bond.
That value is compared to
a scale that indicates the
bond type. For instance take
the oxygen-hydrogen bond.
The difference in
electronegativity is found
by getting atomic specific
values from the table and
subtracting the smallest
from the largest.
Oxygen's value is 3.5
Hydrogen's is 2.1
3.5 minus 2.1 is 1.4.
Taking 1.4 to the scale shows
the bond to squarely be in
the polar covalent bond region.
While we are on the polar
bond let's further explore
what the uneven distribution
of electrons means in the
context of all the
bonds in the molecule.
Take water for instance.
It has 2 polar covalent bonds
with the larger electronegative
atom oxygen keeping
the electrons from the bond
for more than half of the time.
What that means from a net
charge perspective is that
oxygen will have more
negative charge than it started
with as just an atom. It will
adopt a partial negative charge
as denoted with the lower case
Greek delta symbol.
If the oxygen has more of the
negative charge then the
hydrogens will have less
negative charge than they
started with as an atom
and each will adopt a
partial positive charge.
To give some perspective,
a partial negative charge is
much less than the
charge on an electron
and a partial positive charge
is much less than the charge
on a proton. The partial
charges are significantly
weaker than the
charge on an ion.
Also keep in mind that while
there is a distribution of
charge there is no net
charge on the molecule.
It has the same number
of electrons as protons.
The amount of partial positive
charge is equal to the amount
of partial negative charge.
To put this in context of
intermolecular bonding we
need to bring in the term
electric dipole: it is the
separation of electric charge.
The first oxygen-hydrogen
is an electric dipole because
positive and negative
charge is separated in space.
As is the case for
the second dipole bond.
These are called permanent
dipoles because the
electronegativity of the
atoms calls for an uneven
distribution of electrons.
As was seen in the
'Covalent Bond' lecture,
these two polar bonds sum to
an a net dipole or
a molecular dipole.
Water is a polar molecule
because it has a net dipole.
In this lecture and in many
text books a molecular dipole
is given the cartoon
representation
of a flatten oval, with one
side carrying the partial
negative charge and the other
the partial positive charge.
Often a color scheme is applied
that has the partially
positive end colored blue
and the partially
negative end colored red.
In-between is considered
neural and is purple.
The color scheme allows
the distribution of charge
to be seen at a glance.
It really is a useful tool.
That sums up the foundation
component of the lecture
it is now time to look
at intermolecular forces
and bonding and we will do so
moving from strongest forces
and bonds to the weakest.
The ion produces the strongest
of the intermolecular forces.
The ionic bond is an
electrostatic interaction.
The positively charged cation
is attracted to the negatively
charged anion and vice versa.
The charges are full charges.
That means that they are
at least a +1 or a -1.
These charges are
much greater than
the partial charges associated
with polarity.
Because of the size of their
charge, ionic interactions
are considered the strongest
of the intermolecular forces.
We will put them at the top of
a list which ranks the
strength these forces.
Not all ionic interactions
have the same force.
A couple trends can be
established that allow
strength of interaction
within this ion-ion category.
A look at the left side of
the Periodic Table shows us
where some of the more
common cations come from.
Column 1 elements form
+1 cations and Column 2
elements form +2 cations.
The size of charge plays
a big role in the strength
of intermolecular
bonding for a substance.
This can be shown graphically
using phase transition.
The y-axis is the melting
point of an ionic compound
and the x-axis is the
size of the charge.
As seen with the Periodic
Table, ionic charge increases
going to the right. Since
the melting point reflects
the temperature at which
the energy of motion exceeds
some of the intermolecular
forces it is reasonable to
expect that the larger
the charge the stronger the
intermolecular bonds and
the higher the temperature
needed to melt the substance.
A second factor in the strength
of intermolecular forces in
ionic bonding is
the size of the ion.
The argument has it that the
larger the ion the more volume
it has to disperse its charge.
Larger ions of the same
charge make weaker ionic bonds.
On the Periodic Table ionic
radius or size of the ion
increases going down the Table.
Plotting size of ion (for any
one particular size charge)
on the x-axis should give
the opposite trend since the
bigger ions have smaller
intermolecular forces
to overcome and less
heat is needed to do so.
To summarize the
trends: higher charge-
-higher melting point.
Bigger ion
(for the same charge)
-lower the melting point.
Same set of arguments can
be used with boiling points.
Moving down in strength
is the ion-permanent dipole
intermolecular interaction.
It is the attraction of the
full charge of an ion to the
partial charge of a dipole.
Since a dipole has both a
partial positive and partial
negative charge it can
form favorable interactions
with either a
cation or a anion.
Strength of the
interaction is dependent on:
the magnitude of the charges
on the ion and the dipole and
the size of the ion (larger
the ion, smaller the force).
Both of these factors were
used to describe the strength
of an ion-ion intermolecular
bond. Since one of the
components in this type
of intermolecular bond
is an ion it stands to
reason the same set of factors
are still relevant.
The ion-permanent dipole
intermolecular bond will rank
second on our list because of
the strength of the
charge on the components.
A second consideration with
this type of intermolecular
bond has to do with the
dipole having both partial
positive and partial negative
charges. The ion charges
determine the
orientation of the dipole.
The end of the dipole with the
opposite charge faces the ion.
This makes sense from
the perspective of
electrostatic interactions.
If a cation found itself in
the polar solvent water,
how would a water molecule
approach the positively
charged ion so that it
maximized electrostatic
attraction but minimized
electrostatic repulsion?
It would orient itself
in a manner that pointed
the partially negative
oxygen at the positive ion.
In this way it promotes
attraction of opposite charges
and allows water keeps its
partially positive end as
far for the positive cation
as it can. Other water
molecules would do the same
thing and cation would
be considered solvated.
The opposite happens to a
negatively changed
anion in water.
Here the water molecule
orients itself so that its
partially positive side is
pointing at the negatively
charged anion allowing an
electrostatic attraction.
Other water molecules
orient themselves around the
anion in the same manner.
Ion-dipole forces are
important in solubility.
The term dissociation is used
with ions dissolving in water.
The lectures 'Ionic
Bonding' and 'Solutions and
Solubility' also cover
dissociation. The next kind of
intermolecular bond is
a rather particular one.
It is a variety of the dipole
- dipole intermolecular bond.
Dipole-dipole interactions are
between 2 permanent dipoles.
Hydrogen bonding, or
H-bond, is a variety or
a special case of dipole
- dipole interactions.
It is always between an atom
with a high electronegativity,
like oxygen, nitrogen or
fluorine in one molecule,
and a hydrogen bonded
to an atom with a high
electronegativity
in another molecule.
These very strong dipoles
makes the hydrogen bond
one of the strongest dipole
- dipole interactions.
They are so strong that they
get their own place on the
ranking. Third on our list.
Since dipoles have both a
positive and negative end
there will a favorable
orientation of molecules that
maximizes attraction
and minimizes repulsion.
The partial negative end
of one dipole (atom with a
high electronegativity)
is oriented the toward
the partially positive
end on another dipole
(the one with the
hydrogen bonded to an atom
with a high electronegativity).
While water
is the most familiar
example there are others that
we have come across.
Like ammonia; NH3.
It fulfills the requirements
for be capable of forming
hydrogen bonds by
having an atom with a
high electronegativity,
nitrogen, in one molecule
and a hydrogen
bonded to an atom
with a high electronegativity
in another molecule.
Same type of orientation
as found in water.
Since hydrogen bonding is so
central to aqueous solution
chemistry and biological
chemistry we should further
investigate it attributes. Many
of water's unique properties
can be understood from the
perspective of the numerous
hydrogen bonds it makes.
The 'Water; Aqueous Solution'
lecture in this series covers
the topic more in depth but
we will look at one property.
That will be hydrogen bond's
role in the phase transition
of boiling. Boiling
point trends can be
viewed as a function
of the Periodic Table.
The effect is quite stark.
Consider 3 columns of the
Periodic Table headed by
elements carbon, nitrogen
and oxygen. For this
demonstration each element
in these columns will be
represented by its compound
containing a maximum
number of hydrogens.
For carbon and its column
that is 4 hydrogens.
For nitrogen and its column
it is 3 hydrogens and for
oxygen's column
it is 2 hydrogens.
Of all these compounds
only NH3 and H2O make
hydrogen bonds.
Here is a graph of
the boiling points
of the compounds
in the 3 columns.
Each column is
its own tend line.
Carbon's column shows that as
the central atom gets larger
(moving down the Table) the
boiling points get larger.
None of these compounds make
hydrogen bond and
the trend is linear.
For the column
headed by nitrogen
a similar trend is expected
but there is a notable
absence where
ammonia should be.
It has a noticeably higher
boiling point. It takes more
energy to overcome the
hydrogen bonds in ammonia
and that leads to a
higher boiling point.
A similar pattern is
seen with oxygen's column.
Water is expected to
boil at something like
-80 degrees Celsius.
The stronger intermolecular
forces of its hydrogen bonding
requires more heat, higher
temperatures, to be overcome.
Water boils closer to
100 degrees Celsius.
The take home lesson here
is that overcoming hydrogen
bonds requires substantial
amounts of energy.
As mentioned earlier hydrogen
bonding is a particular form of
dipole-dipole interaction.
A separate category on the
list of intermolecular forces
is given to the general term
dipole-dipole.
Dipole-dipole is the
interactions between
2 permanent dipoles.
While there is a distribution
of charge in these dipoles they
have a net charge of zero.
The dipole-dipole
intermolecular bond is
between neutral molecules.
This isn't strictly true but
we will limit our discussion to
neutral molecules.
Dipoles are depicted here as
ellipses for convenience.
As for orientation:
the partial negative end of
one dipole is pointed toward
the partially positive
end of another dipole.
An electrostatic attraction.
Furthermore, in a solution
or mixture of dipoles the
most favorable arrangement is
for the dipoles to
orient themselves
such that they
maximizes attraction and
minimize repulsion.
This is not unlike the
orientation found
with ionic compounds.
But dipoles are much weaker
than ions and have to be much
closer in order to form
intermolecular bonds.
Still they are strong enough
to make it to 4th place on
our list of intermolecular
forces. Another similarity
that dipoles have with ions
is that the strength of the
charges (partial with
dipoles) and the geometry of
the compound influence the
strength of the interaction
The take home trend for this
intermolecular interaction
is that the stronger the
dipole, the more energy that is
needed to break its
intermolecular bonds.
We saw that with the H-bond
and the large effects on
water and ammonia.
Water's larger partial
charges have more of an
effect than ammonia's.
So far we have been dealing
with particles that we have
some history with: the
ion and the dipoles.
They both have what is
termed permanent charges.
Now it is time to introduce
a new concept, the temporary
or induced dipole.
All elements and compounds
are susceptible to
having induced dipoles.
This phenomenon arises out of
the very nature of the atom.
With an atom, the distribution
of electrons around the
nucleus is symmetrical.
The distribution of electrons
around a covalent bond
(non-polar) is also symmetrical.
Consider this drawing of
a spherical electron cloud
around its nucleus.
The symmetrical distribution
of electrons means that the
atom, as a whole, is neutral.
According to the charge bar
the atom should be colored
a neutral purple.
A similar argument can be
made with the distribution of
electrons around a
non-polar covalent bond.
The two atoms in the bond are
also colored purple because
of the symmetrical
distribution of electrons
around their nuclei.
Now imagine what
would happen if a charged
particle is brought
near the atom or
the covalent bond.
Proximity to a charge,
be it ion or dipole,
distorts the symmetry
of the electron cloud.
Proximity to a charged
particle induces a dipole.
For instance, bringing an atom
near a positive charge would
induce a dipole because
the electrons in the atom
are attracted to the
positively charge particle.
The atom would lose
its electron symmetry
as the side of the atom nearest
the charge would become
'electron heavy' and
adopt a partial
negative charge.
The other side of the atom
would have fewer electrons
that it did prior to
the introduction of the
positive charge and therefore
adopt a partial positive charge.
Bringing the atom near
a negative charge would
also induce a dipole
in the atom because
the electrons in
the atom nearest the
negative charge particle
would be repelled and pushed
to the other side of the
atom making the far side
adopt a partial
negative charge.
The side of the atom nearest
the negatively charged
particle would be
'electron light' and
adopt a partial
positive charge.
Same set of arguments
work for bringing a
non-polar covalent
bond near charges.
As it neared a positive
charge, electrons
would be drawn to the charge
inducing a dipole.
As it neared a negative charge
electrons would be pushed
to the opposite end
of the bond. The
redistribution would cause
an induced dipole.
Induced dipoles
are temporary dipoles.
They will go away if the
charged particle is removed.
Induced dipoles are
weaker dipoles than
permanent dipoles.
Still, induce dipoles
and charged particles
form intermolecular bonds.
The partial negative
charge of the induced dipole
forms a weak
intermolecular bond with
the permanent positive charge.
The partial positive
charge of the induced dipole
forms a weak
intermolecular bond with the
permanent negative charge.
The weakness of the induced
dipole has this form of
intermolecular bond placed
towards the bottom of our list.
Although it is a weak, it
still requires energy to break.
The strength of bond is
related to the strength of the
induced dipole which itself
is related to the strength of
the permanent charge
on the particle.
Therefore ions, with their
full charge, will induce
larger dipoles than
will permanent dipoles.
The size of the induced
dipole is proportional
to the size of the
permanent or full charge.
Since a full negative
charge is much larger than a
permanent dipole, an anion
will induce a larger dipole.
Likewise, a full positive
charge is much larger than a
permanent dipole and a cation
will induce a larger dipole.
This can be
summarized in a trend;
the stronger the ion charge or
the charge of the
permanent dipole
the stronger the
induced dipole and
the stronger the
intermolecular force or bond.
And that brings us to the last
of the intermolecular forces
and it involves
instantaneous dipole moments.
We have seen the permanent
dipole of molecules
like water, the induced
dipoles that requires a
charged particle and
now we are introduced
to the dipole that springs
up of its own volition.
An explanation of such
a dipole begins with
the symmetrical distribution
of electrons around the nucleus
of an atom or covalent bond.
Just like the start of the
discussion on induced dipoles.
Here, though, we will think
of the electrons as rapidly
moving charged particles.
When, through normal motion,
one side of an atom or bond
has more than half
of the electrons
an instantaneous
dipole moment occurs.
As a demonstration of
this instantaneous dipole
consider the third electron
in the 3 electron atom.
It is drawn here pretty close
to a dotted line that marks the
halves of the atom.
The positively
charged nucleus is also
sectioned into halves.
We can argue that with
this distribution that the two
halves of the atom each have
the same amount of negative
and positive charge.
Both halves are neutral.
But electrons are always moving
and at some point in time
the third electron will find
itself on one side with
another electron. The 2 halves
will no longer have the
same amount of charge.
One side has more negative
charge at the expense
of the other side.
One side is now
partial negative and the
other partially positive.
This type of instantaneous
or spontaneous dipole occurs
in all atoms and all bonds.
It is the very
nature of matter.
The partial charges are very
small and very short lived.
For completion sake we
will look at how a bond
can find itself as an
instantaneous dipole.
Here is the representation
of hydrogen, H2,
with the shared electrons
dawn as stationary dots
between the two atoms.
But electrons are not stationary.
At some point in time the
electron pair will find
themselves exclusively
on one of the two atoms.
The once neutral molecule
will now have a charge
distribution and become
an instantaneous dipole.
Again, it is very
weak and short lived.
Almost immediately the
electron pair will migrate back
and the molecule will
return to its neutrality.
The question becomes
whether these instantaneous
dipole can participate
in intermolecular bonding.
The answer is a resounding
yes. It is present in all
matter-matter interactions at
the atom and compound level.
These forces are called
dispersion forces.
All atoms and bonds have
instantaneous dipole moments.
It is a product of the
motion of electrons.
Previously we saw that the
proximity to a charge particle
like and ion or dipole
induces a dipole.
How big of a leap is it to
suggest that the proximity
to an instantaneous
dipole induces a dipole.
Bringing back that H2
molecule in its instantaneous
dipole form and putting
it in the proximity
to a neutral H2 molecule.
The small partial charge
of the instantaneous dipole
can still influence the
distribution of the electrons
in the neutral molecule.
The mild electron repulsion
will have the effect
of generating a
small but relevant
induced dipole in the
once neutral H2 molecule.
This is the weakest of
all of the induced dipoles.
Still, an intermolecular
bond is formed between the
partial negative side of the
instantaneous dipole and
the partial positive side
of the induced dipole.
Being that the instantaneous
dipole is so short lived and
so weak the attraction only
works over very short distances.
The forces are
called London dispersion
or Van Deer Waals forces.
All though they are the
weakest of all intermolecular
forces they are present in
all matter-matter interactions.
They explains how
non-polar molecules have
attractive interactions.
If it wasn't for these forces
all non-polar compound
would be gases.
They could not form
liquids or solids.
Dispersion forces bring
up the rear on our list.
They finish our list off.
While we have the list up
we should ask what
is the overarching
principle behind the
ranking of strongest
to weakest forces.
It is charge.
The weaker the charger
the weaker the forces.
As weak as they are dispersion
forces do have an effect on
physical properties as can
be seen with a few trends.
One is that the larger
the atom the stronger
the induced dipole.
Larger atoms
often have stronger
dispersion forces.
The more electrons an atom
has the bigger its radius.
This allows for stronger
instantaneous dipoles and
stronger induced dipoles.
Stronger intermolecular
interactions require more
energy to break.
That is why there is a trend
with atom size
and boiling points.
Moving down the noble gas
column on the Periodic Table
serves as a good
example of this trend.
There is an upward trend for
the boiling points of the noble
gases following atom size.
Helium is the smallest atom
and has the lowest boiling
point and xenon and radon are
the largest atoms in the
column and have the strongest
dispersion forces
as evident from
their higher boiling points.
In a similar manner bigger
molecules have more electrons
and a longer distance
for them to travel.
This results in
bigger induced dipoles.
Long thin molecules
can develop bigger
temporary dipoles due
to electron movement than
short fat ones even if
they containing the same
numbers of electrons.
This also results in
higher boiling points
for larger molecules.
A look at boiling points
of hydrocarbon as a function
of the number of carbons
shows the single carbon
methane to have a lower
boiling point. Well, lower than
the 2 carbon ethane
which is lower than the
3 carbon propane. Butane and
pentane continue the trend.
And finally, it is easier for
longer and thinner molecules
to have induces dipoles
because more of the molecule
can be approached.
Geometries do have a role.
6 forms of
intermolecular forces
were covered in this lecture.
As a recap. Intramolecular
bonds hold atoms
together WITHIN a
molecule or compound.
They are strong bonds.
Intermolecular bonds
are between compounds or
elements are generally weaker.
All bonding is the result of
electrostatic interactions.
The attractive and repulsive
forces of an electrostatic
interaction can be between
atoms, ions or compounds.
Polar covalent bonds have
their shared electrons
spend more time around
1 of the atoms forming a
permanent dipole. Dipoles are
often represented as an oval
or an ellipse and given
a color scheme to show
the charge distribution.
Raking intermolecular bonds
or forces was done by strength.
Strength is related to the
size of the charge be it
permanent or induced.
Our ranking has the ionic
bond as the strongest.
The ions have full charges.
Cations attract anions.
As for trends, bigger ions
tend to generate weaker
intermolecular forces.
This is seen with
lower melting points. The
higher the charge the stronger
the intermolecular force
and the more energy needed
to separate the ions. That
means higher melting points.
The ion-dipole
attraction comes next.
It has the full charge of an
ion attracted to partial
charge of a dipole.
These are important in
solvation. Then there
is a specific kind of the
dipole-dipole interaction.
It is the hydrogen bond.
It gives water many
of unique properties.
There are other types of
dipole-dipole interactions:
The partially negative end
of one dipole is pointed at
partially positive
end on another dipole.
Trends with this
interaction include
stronger dipoles having
higher melting points.
New to this lecture
was the induced dipole.
Proximity to a
charged ion or dipole,
distorts an atom's electron
symmetry and induce a dipole
The weakest force are
the dispersion forces.
Here an instantaneous dipole
moment induces a dipole.
Trends are based on
the capacity of an
element or compound to
form induced dipoles.
Larger the atom or molecule
the greater the instantaneous
dipole and the higher
the melting and boiling points.
And that completes our lecture.
Remember, intermolecular
bonding is about charge!
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