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Ph a se C h an g e In str u ct io n al Ca s e: A s er i e s o f stu d en t - c en t er ed s ci en ce l e s son s
Teacher Background
Phase transitions are broadly defined as the
change from one phase of matter to another.
We commonly associate them with the
transitions between solid, liquid, and gas but
apply to any change of phase, such as
diamond converting to graphite. We also
normally think of such transitions as occurring
through change in temperature, but they can
also be induced by a change in pressure, as
described below. Phase changes are
examples of physical change because the
fundamental particles (atoms or molecules)
are not rearranging to make new products.
To understand phase transitions, students
must understand what makes the phases of
matter different. Differences in the
properties of solid, liquid, and gas with
respect to shape, volume, density, flow,
compressibility, etc.., are all due to the
relative strengths of the intermolecular
attractions between the particles (atoms or
molecules). In a solid, the attractions
between particles are so strong that the
particles are essentially fixed in place and can
only vibrate (definite shape and volume). The
particles in a liquid are still attracted to one
another such that they remain in constant
contact with their neighbors but can freely
move around (shape of container, definite
volume). The attractions between particles in
a gas are negligible; they can be thought of as
entirely independent from one another
(shape and volume of container).
Changing from one phase of matter to
another can be induced by adding heat to the
system. The heat is distributed amongst the
particles as thermal energy, but not in a
uniform way. Instead, there is a distribution of
energy; some particles have a great deal
whereas others have much less. If we start
with a solid, all the particles will be vibrating
around their fixed points. Adding heat
increases the energy the particles have to
vibrate. We observe this as an increase in
temperature. At a particular temperature
(the solid’s melting point), a particle may have
enough energy to overcome some of its
attractions to its neighbors such that it can do
more than simply vibrate in place and move
around. This is the beginning of the phase
change from solid to liquid (melting). Adding
more energy into the system will continue to
break these intermolecular attractions until
the solid is fully melted. Only once the solid is
melted will additional heat go to increasing
the temperature of the system. This
phenomenon can be observed as plateaus on
a heating curve.
If we continue to heat the liquid, the
temperature will increase until the liquid’s
boiling point is reached. At that point, a
particle on the surface may have enough
energy to overcome its attractions to all of its
neighbors (surface particles having fewer
neighbors) and enter the gas phase as an
independent particle
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This is known as vaporization. The
temperature will remain constant until all the
liquid has been converted into a gas.
The above are examples of endothermic
processes, requiring the input of energy.
Different substances have different amounts
of energy necessary to lead to a phase
transition based on the strength of their
intermolecular forces. Heat of fusion
(melting) is the energy associated with
changing from a solid to a liquid. That same
amount of heat would instead be released
when the reverse, exothermic process,
freezing, occurs. Heat of fusion is generally
smaller than the heat of vaporization, the
energy associated with the transition from a
liquid to a gas. This is because vaporization
requires complete breaking of attractions
between molecules whereas the particles in a
liquid are still in contact with one another, so
the intermolecular attractions are only
partially overcome. Again, the amount of
heat released upon condensation is the same
as the heat of vaporization since it the
reverse, exothermic process.
Another common phase change is
sublimation, the direct transition from solid to
gas. In this case, a particle on the surface of a
solid gains enough energy to completely break
its attractions to its neighbors and go directly
into the gas phase. The opposite process,
deposition, occurs if a particle comes into
contact with a solid and does not have
sufficient energy to break away from the
attractions of its new neighbors. This is
readily observed with dry ice, but water ice
also sublimes and deposits, as you can
observe by seeing ice crystals inside sealed
storage bags.
Changes in pressure can also lead to a phase
transition. In this case, increasing pressure
leads to a transition to the more densely
packed state, whereas decreasing pressure
leads to the less dense state. The fact that
both temperature and pressure determine the
stability of a particular phase is readily seen
on a phase diagram of the substance. Phase
transitions do not occur just at discrete
temperatures (e.g. melting points), but
actually along continuous lines of
temperatures and pressures.