Intermolecular forces: Background
Electrostatics
Up until now, we have just discussed attractions between molecules in the area of the
covalent bond. Here, atoms within a molecule are attracted to one another by the
sharing of electrons. This is called an intramolecular force.
We know how the atoms in a molecule are held together, but why do molecules in a
liquid or solid stick around each other? What makes the molecules attracted to one
another? These forces are called intermolecular forces, and are in general much
weaker than the intramolecular forces.
We have, however, already discussed a very strong type of force that is responsible
for much of chemistry - electrostatics. The attraction of a positive charge with a
negative charge is the force that allows for the structure of the atom, causes atoms to
stick together to form molecules; both ionic and covalent, and ultimately is
responsible for the formation of liquids, solids and solutions.
London dispersion forces
The forces that hold molecules together in the liquid, solid and solution phases are
quite weak. They are generally called London dispersion forces.
We already know that the electrons in the orbitals of molecules are free to move
around. As such, if you would compare a "snapshots" of a molecule at an instant in
time, you would see that there would be slightly different charge distributions caused
by the different positions of the electrons in the orbitals. Just how much difference
one sees as a function of time is based on the polarizability of the molecule, which is
a measure of how well electrons can move about in their orbitals. In general, the
polarizability increases as the size of the orbital increases; since the electrons are
further out from the nucleus they are less strongly bound and can move about the
molecule more easily.
Given that two molecules can come close together, these variations in charge can
create a situation where one end of a molecule might be slightly negative and the near
end of the other molecule could be slightly positive. This would result in a slight
attraction of the two molecules (until the charges moved around again) but is
responsible for the attractive London dispersion forces all molecules have.
However, these London dispersion forces are weak, the weakest of all the
intermolecular forces. Their strength increases with increasing total electrons.
Dipole-dipole attractions
What would happen if we had a beaker of polar molecules, like ClF,
In addition to the attractive London dispersion forces, we now have a situation where
the molecule is polar. We say that the molecule has a permanent dipole. Now, the
molecules line up. The positive ends end up near to another molecule's negative end:
Since this dipole is permanent, the attraction is stronger. However, we only see this
sort of attraction between molecules that are polar. It is usually referred to as dipole dipole interaction. The strength of this attraction increases with increasing total
number of electrons.
Hydrogen bond
Hydrogen is a special element. Because its nucleus is really just a proton, it turns out
that it can form a special type intermolecular interaction called the hydrogen bond. If
the hydrogen in a molecule is bonded to a highly electronegative atom in the second
period (N, O, or F only) with a lone pair of electrons, a hydrogen bond will be
formed.
In essence the three elements listed above will grab the electrons for itself, and leave
the hydrogen atom with virtually no electron density (since it had only the one). Now,
if another molecule comes along with a lone pair, the hydrogen will try to position
itself near that lone pair in order to get some electron density back. This ends up
forming a partial bond, which we describe as the hydrogen bond. The strength of this
interaction, while not quite as strong as a covalent bond, is the strongest of all the
intermolecular forces (except for the ionic bond).
Hydrogen bonding interactions in water :
Could the ClF molecule exhibit hydrogen bonding? The answer is no, since the
hydrogen must be bound to either N, O, or F. Just having one of those species in the
molecule is not enough.
Trends in the forces
While the intramolecular forces keep the atoms in a molecule together and are the
basis for the chemical properties, the intermolecular forces are those that keep the
molecules themselves together and are virtually responsible for all the physical
properties of a material. For small molecules, the intermolecular forces increase in
strength according to the following:
London dispersion < dipole-dipole < H-bonding < ion-ion
Now, as these things increase in strength it becomes harder to remove the molecules
from each other. Therefore, one would expect the melting and boiling points to be
higher for those substances which have strong intermolecular forces. We know that it
takes energy to go from a solid to a liquid to a gas. This energy is directly related to
the strength of attraction between molecules in the condensed phases. Since energy is
directly proportional to the temperature, the above trends ought to hold true.
In addition, there are energies associated with making these phase transitions:
Each of these processes are endothermic, and scale with the magnitude of the
intermolecular forces. Thus, as these intermolecular forces increase, so do the energies
requires to melt, vaporize, or sublime (go from solid to a gas) a species.
Every substance also has an associated vapor pressure with it. The vapor pressure is
defined to be the amount of gas of a compound that is in equilibrium with the liquid or
solid. If the intermolecular forces are weak, then molecules can break out of the solid
or liquid more easily into the gas phase. Consider two different liquids, one polar one
not, contained in two separate boxes. We would expect the molecules to more easily
break away from the bulk for the non-polar case. This would mean that,
proportionately, there are more molecules in the gas phase for the non-polar liquid.
This would increase the vapor pressure. Thus, unlike the physical properties listed
above, the vapor pressure of a substance decreases with increasing intermolecular
forces.
Name: __________________________________________________________
Pre-Laboratory Questions
1. Read about intermolecular forces in your textbook. Construct a table comparing the following
intermolecular forces: dipole-dipole interactions, hydrogen bonds and London Forces.
Force
Origins
Occurs between
(polar or non polar
Molecules)
2. Show the hydrogen bonding interactions in the following
a) ammonia
b) A mixture of water and methanol, CH3OH
Strength
Example
Procedure
Part 1
Assemble at you table a dropper and a little bit of water, alcohol and acetone in small test tubes (~ 1 cm
high). Place one drop of each liquid on the table, in a place where the drops won’t be disrupted, and
observe which evaporates first. (They will evaporate even faster if you place them in the hood.) Make
sure your drops are not in your way as you will proceed with part 2 while you wait for your drops to
evaporate. You may blow lightly over the tops of the drops to speed the process as long as you try to blow
equally across all three. Write down your results later on.
Part 2
1. Rub a penny clean with a paper towel to remove any oil that might be on the surface. After this, try not
to handle the penny with your bare hands as this will deposit additional oil on the surface.
2. Place the penny on a paper towel. Predict how many drops of water you will be able to place on the
penny without getting the paper towel wet. Record your prediction in the table below.
3. Now place drops of water on top of the penny, one drop at a time, using a disposable pipet. Note the
shape of the water on the penny as the water increases in volume. Continue adding and counting water
drops until the water spills onto the paper towel. Record the number of drops in the table. Describe the
shape of the water on top of the penny or draw a picture.
4. To repeat for a 2nd trial, what variables should you keep constant so that you could get reproducible
results?
5. Do the 2nd trial and record the drops. Find an average of the two values. (If they are fairly different,
you may want to do and record a third trial before averaging).
6. Repeat the procedure above with alcohol, C2H5OH, and then with acetone, CH3COCH3, after you have
made predictions. Make sure the drops of each are about the same size as the drops of water you used
before. Note the shapes of each. Record your results in the table.
7. Now add 2 drops of liquid soap from the cart to your test tube with water and stir. Determine how
many drops of the soap/water mixture can be piled on top of the penny before it spills on to the paper
towel. Record your results of at least 2 trials in the table.
Data Sheet
Part 1 Data: (record these observations after you complete part 2)
Observations:
Part 2 Data:
Prediction
# of drops
Water
H2O
Alcohol*
C2H5OH
Acetone
CH3COCH3
Soap water
(as
explained
in step 6)
*ethanol
Trial 1
# of drops
Trial 2
# of drops
Trial 3
# of drops
Average
# of drops
Post Lab Questions
Critical Thinking Skill: Constructing convincing arguments.
Argument: A train of reasoning in which claims and supported reasons are
linked to establish a position.
1. State your position first: Which liquid had the strongest intermolecular forces and which had the
weakest:
2. When you build your argument, what data (results from three liquids in the evaporation
experiment and results from three liquids in the penny experiment) supports your position?
What data does not support your position?
Supports-
Does not support –
3. Now construct your argument. Write a full paragraph conclusion in which you rank the strength
of the intermolecular forces of these three liquids and explain how your data supports that
ranking. Your argument will be evaluated on three criteria: Is it complete? Is it correct? Is it
clear?
4. If you put your three liquids into beakers with a thermometer and heated them on a hotplate, one
liquid would start to boil at 56°C, one would start to boil at 78°C and one would start to boil at
100°C. Considering this boiling point data in connection with your evaporation data: which
liquid is water? ethanol? acetone?