Honors Chemistry
Liquids and Solids
Reading 1: What are intermolecular forces?
If you observe the world around you, you will see that many different molecules surround you.
Some of these molecules are in gas form at room temperature, while others are liquids and still
others are solids. What determines if a molecule forms a gas, a liquid or a solid at room
temperature? The answer, in part, is intermolecular forces.
What are intermolecular forces? You know that an interstate is a road that connects one state
to another. Well, an intermolecular force connects one molecule to another (or many
molecules to each other.) Compare these to intramolecular forces, otherwise known as
covalent or ionic bonds. Intramolecular forces are those that hold together the atoms within
one molecule. For example, intramolecular forces are responsible for holding the hydrogen
atoms to the oxygen atom, forming a water molecule. Intermolecular forces are responsible for
holding one water molecule close to another water molecule, causing water to be a liquid at
room temperature.
In terms of energy, intermolecular forces have much lower energy than covalent or ionic bonds.
It takes much less energy to move two molecules away from each other, than it does to break
apart the molecule itself.
There are three types of intermolecular forces: dipole-dipole attractions, hydrogen bonding,
and London dispersion forces.
Dipole-dipole attractions are the result of polar molecules discussed a short while ago. You
may recall, when molecules are polar, they create dipoles, where electrons are not dispersed
equally around the molecule. Recall that the greater the difference in electronegativity, the
greater the dipole. Recall, also, that a polar molecule has a positive area and a negative area,
similar to a magnet. So dipole-dipole attractions are caused by the positive area of one
molecule being attracted to the negative area of another molecule (and vice-versa). These
forces only exist in polar molecules.
Hydrogen bonding is a particularly strong type of dipole-dipole attraction. Hydrogen bonding is
present in polar molecules where hydrogen is bonded to a highly electronegative atom. The
highly electronegative atoms we most encounter are nitrogen, oxygen and fluorine. Two things
make hydrogen bonds particularly strong. One is the considerable difference in
electronegativity. The other is the small size of the hydrogen atom. Because hydrogen is so
small, molecules containing hydrogen can get closer to each other, allowing them to hold on to
each other tighter. These forces only exist in molecules with either an N-H, an F-H, or an O-H
bond.
1
The last type of intermolecular forces is called London dispersion forces. The word dispersion
refers to the way the electrons are spread around the molecule. Normally, the electrons are
spread fairly evenly around the molecule. However, electron movement is random and it is
possible for a majority of the electrons to wind up on one side of the molecule. What would
this do to the molecule? To understand this, consider an analogy. Imagine loading many
people on a boat and telling them to randomly run around the deck. Most of the time, you
would probably have people all around the boat, but imagine that at some time all of the
people wound up on one side of the boat. This would cause the boat to tilt. Now go back to a
molecule – the electrons would not cause the molecule to tilt, but it would create a very
temporary negative charge on one side of the molecule (and as a result, a positive charge on
the other side). This temporary dipole would go away quickly as the electrons keep on moving.
This situation occurs over and over again with the dipole appearing and then disappearing. Just
as in the dipole to dipole attractions, opposite charges in the molecules cause them to stick
together. You can see that because of the temporary nature, London dispersion forces are very
weak attractions. These forces exist in all molecules.
Reading 2: Identifying which forces are present in a molecule:
Let’s say we have the molecules NH3, CH4, and CH3Cl. How can we tell which intramolecular
forces are present in each molecule?
The first thing to consider is the polarity of the molecule. Is the molecule polar or nonpolar?
This means you need to draw the structure of the molecule and determine the polarity. So,
let’s say that you draw the three structures and you find that NH3 and CH3Cl are polar, while
CH4 is not. That means that both NH3 and CH3Cl will both have dipole-dipole forces, while CH4
will not.
If a molecule has polarity, then it is possible that it also has hydrogen bonding. Recall that this
means it must have both hydrogen and a highly electronegative element. For two of the
examples, both NH3 and CH3Cl have hydrogen, but only NH3 also has a highly electronegative
element (nitrogen) bonded to an H. This means that only NH3 has hydrogen bonding present.
Finally, what about London dispersion forces? Do all molecules have electrons? Do all
molecules have electrons that move randomly? Yes. Then all molecules will have dispersion
forces.
To summarize:
NH3 – dipole-dipole, hydrogen bonding, London dispersion forces
CH3Cl – dipole-dipole, London dispersion forces
CH4 – London dispersion forces
We now know that hydrogen bonding is stronger than dipole-dipole, which is stronger than
London dispersion. But what if we are looking at two molecules within a given category? For
2
example, compare PCl3 and PF3. Both are polar, so both would have dipole-dipole attractions,
but which one would have stronger attractions? The answer lies in electronegativity. Using the
electronegativity chart, the difference between P and Cl is 0.9 (2.1-3.0), but the difference
between P and F is greater at 1.9 (2.1-4.0). A greater difference in electronegativity will result in
a stronger dipole moment, so even though they both have the same type of forces, PF 3 will
exhibit a stronger attraction.
In molecules that are non-polar, we don’t need to worry about differences in polarity, but we
do need to worry about the size of the molecule. Compare C2H4 to C16H32. Both are non-polar
and have only London dispersion forces, but C16H32 has a stronger attraction. Why? Larger
molecules mean more electrons, which results in a larger temporary dipole. The larger (more
mass) the molecule is (generally speaking), the larger the London dispersion forces will be.
Reading 3: Intermolecular Forces – Phase Change (particularly water)
The Phase Change Graph
Recall the connection between intermolecular forces and changes of state. For example,
evaporation involves breaking intermolecular forces during a change in phase from liquid to gas
(no change in matter occurs). Think of all the possible changes of state with water, and list
them in the margin of your reading now.
Consider the following graph. Recall that this is a phase change (or change of state) graph and
it represents what is happening energetically as a substance goes from solid to gas (or vice
versa). We need to understand the various parts of the graph and why it looks the way it does.
3
What is the graph showing? Look at the x-axis label. A substance is being heated (i.e. energy is
absorbed going left to right). An example may be water as a block of ice heated on the stove.
According to the graph, energy is constantly being put into the substance. As energy is being
constantly being absorbed, what happens to the temperature? Is it constantly increasing also?
No. (Make sure you understand where, on the graph, temperature is not changing.) So what is
happening to the additional absorbed energy, if temperature is not changing?
Recall from physics two types of energy: kinetic and potential. Molecules have both types of
energy. You should recall that kinetic energy is energy of motion. Temperature is the
measurement of the molecules’ motion. Specifically, temperature = average kinetic energy of
the molecules. For a molecule, the potential energy is the stored energy between one molecule
and another, and between the bonds within the molecule. The further two molecules are apart,
the greater the potential energy.
So let’s apply this to our graph. The three slanted lines on the graph refer to the three states of
matter: solid, liquid and gas. So from the beginning of line A to the end of line A, we have a
solid. From the beginning of line C to the end of line C, we have a liquid. So what is happening
on line B? Well, what has to happen for a substance to change from a solid to a liquid? It
melts. Melting occurs on line B. Notice that during this time, the kinetic energy is not changing.
All of the energy is consumed to break the intermolecular bonds between the molecules while
melting occurs. Now apply this to D and E. What state of matter is present on line E? Gas. What
phase change is occurring on line D? Vaporization (a fancy word for boiling).
Reading 4: Intermolecular Forces – Vapor Pressure, Evaporation & Boiling
Imagine a beaker which contains a liquid, like the ones shown in the picture below. What is
happening in the space above the liquid?
Notice, the molecules moving about above the surface. The molecules show that evaporation
is occurring. And how is the beaker at the left, different than the beaker at the right? You
should notice that the beaker on the right has a lid on it.
4
When will evaporation stop in the beaker on the left? Because this is an open system,
evaporation will occur until the beaker is empty. But what about the beaker on the right?
When a lid is present, we have a closed system. Evaporation will occur until we have reached
equilibrium. What is equilibrium? Equilibrium does not mean equal, as in equal numbers of
molecules in the vapor and liquid phases. We might have several million molecules in the liquid
phase and only a few hundred molecules in the vapor phase. What equilibrium means, is that
the overall number of molecules in each phase is not changing.
Evaporation and condensation are two sides of the same coin. Evaporation occurs as
molecules gain enough energy to break their intermolecular bonds and enter the gas phase.
Condensation occurs as molecules lose energy and reenter the liquid phase. In a closed
container, both changes are occurring at the same time. Molecules are evaporating, and other
molecules are condensing. When these two events happen at the same time, equilibrium has
been reached. See the picture below for details.
If the molecules at the top of beaker represent molecules slowing down and condensing, while
the molecules at the middle of the beaker represent those which are evaporating, the left
beaker is not at equilibrium. Why? Because there are more molecules condensing than there
are evaporating. What about the right beaker? It is at equilibrium. Three molecules in, and
three molecules out. The beaker on the right is exactly what is happening in a sealed container
of a diet coke (CO2).
Now all of this evaporating means that we have a gas phase above the liquid phase in the
beakers. This means that we have pressure. We call it in this chapter vapor pressure. Vapor
pressure can be measured and it is not the same for all substances. In fact, the amount of
vapor pressure is directly related to the intermolecular forces in the substance being
evaporated.
Which type of substances will evaporate more quickly? Ones with weak intermolecular forces,
right? So which types of substances are more likely to have more molecules in the vapor phase
at any time? Ones with weak intermolecular forces, right? Will this result in higher vapor
pressure or lower vapor pressure? Higher vapor pressure; more molecules means more vapor
pressure.
5
Reading 5: Intermolecular Forces – Phase Diagrams
Earlier we addressed the relationship between state of matter and temperature. We learned
that for most substances, an increase in temperature will eventually mean a change in state.
We also learned that pressure can also affect a change of state. It is important to note that
these two things act together, which determines what state of matter a substance is in.
Consider the following graph for water:
Look at the graph, and think about what it tells us. Notice that all three phases of matter are
represented on the graph. Notice that some important numbers, such as standard pressure,
standard temperature, and 100C are also noted. Notice the lines that run between each
phase, as these represent where the change of state occurs. First, examine 0C, which is the
freezing/melting point of water. Is 0C always the freezing/melting point? If it were, then the
line between solid and liquid would be a straight line running straight up the middle of the
graph. What about boiling? We would expect the line between liquid and gas to be at the
100C mark, but it appears that we can have boiling at temperatures well below 100C, even
below 0C. So what’s going on? Remember that both temperature and pressure control the
state of matter. 100C is the boiling point of water only at standard pressure (1 atm). 0C is the
freezing/melting point of water only at standard pressure. Changing the pressure changes the
temperature where water boils or melts. At very low pressures, water can boil at very low
temperatures. At very high pressures, we must heat water much higher than 100C to get it to
boil.
There are several terms you should be familiar with when using this graph. The first is triple
point. To find this, look for where all phases of matter are present at the same time (on this
graph you will see a dot). There is also the critical point. Notice the curve between the liquid
6
and gas phases. It ends at a point. This is called the critical point because it represents the
temperature and pressure at which the liquid phase can no longer exist. For water, the critical
point is 374C and 218 atm. No increase of pressure at that temperature will succeed in
changing to a liquid. At a temperature higher than 374C, water will always be a gas.
7