Lesson 10.3 Changes of State
Suggested Reading
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Zumdahl Chapter 10 Section10.8
Essential Questions
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How do substances transition from on phase to another?
Learning Objectives
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Define the following terms: vaporization, enthalpy of vaporization
(∆Hvap), condensation, sublimination, enthalpy of fusion (∆Hfus), melting
point and boiling point.
Identify the trends observed in heating curves.
Calculate the energy of phase changes from heating curves.
Define phase diagram, critical temperature, critical pressure, critical
point and triple point.
Interpret phase diagrams.
Introduction
In the last few lessons, we looked at the states of matter. Depending on the
conditions, substances may change from one state of matter to another. A
change of a substance from one sate to another is called a change of state
or phase transition. In this lesson we will look at the different kinds of phase
transitions and the conditions under which they occur. A phase refers to the
particular state of a system.
Basic Terms
The students often confuse or terms listed below. Please make sure you
understand each terms.
Vaporization (Evaporation): The change of a solid or liquid to the
vapor. ENDOTHERMIC since energy must be absorbed so that the liquid
molecules gain enough energy to escape the surface and thus overcome
the liquid’s IMFs.
Enthalpy of vaporization (∆Hvap,heat of vaporization): The heat needed to
vaporize one mole of a liquid.
Condensation: The changed of a gas to either the liquid or the solid
sate. Water’s heat of vaporization is 40.7 kJ/mol. This is huge! Water
makes life on this planet possible since it acts as a coolant. The reason
its ∆Hvap is so large has everything to do with hydrogen bonding. The IMFs
in water are strong, so a great deal of the sun’s energy is needed
to evaporate the rivers, lakes, oceans, etc. of Earth. Perspiration is a
coolant for animals possessing sweat glands. Energy from your hot body is
absorbed by the water and evaporates.
Sublimination: The change of a solid directly to the vapor.
Enthalpy of fusion (∆Hfus, heat of fusion): The heat needed to melt one
mole of a solid (Melting is also called fusion in chemistry, just to create
some confusion).
Melting point: The temperature at which a crystalline solid changes to a
liquid or melts.
Boiling point: The temperature at which the vapor pressure of a liquid
equals the pressure exerted on the liquid (usually by the atmosphere).
In general, each of the three states of matter can change into either of the
other two states:
Vapor Pressure
Liquids, and some solids, are continuously vaporizing. If a liquid is in a
closed container with space above it, a partial pressure of the vapor state
builds up in this space. At the same time the rate of condensation starts to
increase. Equilibrium vapor pressure (VP) is reached when the molecules
leave and enter the liquid phase at the same rate.
The VP of a liquid can be measured easily using a simple barometer. A
small sample of the liquid is introduced into the Hg column of a barometer.
Being less dense than Hg, the liquid rises to the top of the Hg and begins to
evaporate. The Hg column drops a distance h because of the VP.
Volatile liquids have high VPs, thus weak IMF's. These liquids evaporate
readily from open containers since they have so little attraction for each
other. This requires little energy; the heat energy absorbed from a warm
room is usually enough to make these substances evaporate quickly. If
you smell an odor to a substance when you open the bottle, it means the
molecules have been banging against the lid wanting out!
VP increases significantly with T! Heat ‘em up, speed ‘em up, move ‘em
out! Increasing the temperature increases the KE which facilitates escape
AND the speed of the escapees! As molecules at the surface gain sufficient
KE (through collisions with neighboring molecules) they escape the
liquid. More molecules can attain the energy needed to overcome the IMFs
in a liquid at a higher T since the KE increases.
In general, as MM ↑ VP ↓
As molecules increase in molar mass (MM), they also increase in the
number of electrons. As the number of electrons increase, the polarizability
of the molecule increases so more induced dipole-induced dipole
(dispersion forces) exist, causing stronger attractions to form
between molecules. This decreases the number of molecules that escape
and thus lowers the VP. Water is an exception. It has an incredibly low VP
for such a light molecule due to the strong hydrogen bonding IMFs.
Boiling Point and Melting Point
Now that we have looked at VP lets revisit the definition for boiling point.
Boiling point is defined as the temperature at which the VP of a liquid equals
the pressure exerted on a liquid (atmospheric unless in a closed container).
As the T of a liquid is raised, the VP increased. When the VP = Patm, stable
bubbles of vapor form within the body of the liquid. This process is called
boiling. Once boiling starts, the temperature of the liquid remains constant, so
long as sufficient heat is supplied.
Because the pressure exerted on a liquid can vary, the BP can vary. The
normal BP of water is at 1 atm. However, at high altitudes where the Patm is
less, the BP is also less. This is why you sometimes find high-altitude
cooking instructions of foods.
The T at which a pure liquid changes to a solid is called the freezing point
(FP). The T at which the solid changes to a liquid is called the melting
point (MP). FP & MP are identical because melting and freezing occur at
the T where the liquid and solid are in equilibrium. Unlike BP, MP and FP
are only affected by large pressure changes. Furthermore, both MP and FP
are characteristic properties of a substance that can be used to identify it.
Heating Curves
Any change of state involves the addition or removal of energy as heat to or
from the substance. A heating curve is a graph that shows the phase
changes that occur as heat is added to a system at a constant rate. You
can see from the curve on the right, that the substance transitions from solid,
to liquid, to gas in a series of positive slants and flats. This is characteristic of
all heating curves. The slants show the increase in T as the substance moves
from one phase to the next and the flats represent a phase transition. The T is
constant during a phase transition because the system is in equilibrium.
However, energy is continually added, so eventually the system's T begins to
increase again, and the system moves away from equilibrium and toward the
next phase change.
On the plateaus, ∆E = q = ΔH(vap or fus) × n, where n is moles
On the slants, ∆E = q = m × c(liquid, solid, or gas) × T
Notice that the lengths of the flat regions vary. Because heat is being added at
a constant rate, the length of each flat region is proportional to the heat of
phase transition.
Heat of fusion (∆Hfus, enthalpy of fusion): The heat needed to melt one
mole of solid. For ice the heat of fusion is
6.01 kJ/mol.
H2O(s) → H2O(l); ∆Hfus = 6.01 kJ/mol
Heat of vaporization (∆Hvap, enthalpy of vaporization: The heat needed to
vaporize one mole of liquid. For water,
H2O(l) → H2O(g); ∆Hvap = 40.7 kJ/mol
Not that much more energy is required for vaporization than form melting.
This is because the relative strengths of the IMFs are about the same for
liquid and solid water. Melting only needs enough energy for the molecules
to escape form their sites in the lattice. For vaporization, enough energy
must be supplied to break most of the IMFs.
Example: Calculating the heat required for a phase change of a given
mass of substance.
A particular refrigerator cools by evaporating liquified
dichlorodifluoromethane, CCl2F2, How many kg of this liquid must be
evaporated to freeze a tray of water at 0∘C to ice at 0∘C? The mass of the
water is 525 kg, the heat of fusion of ice is 6.01 kJ/mol, and the heat of
vaporization of dichlorodifluoromethane is 17.4 kJ/mol.
Solution:
We want to transition from a liquid to a solid, so the water must move down
the heating curve, which requires the removal of energy in order to cool the
liquid to freezing.
First, we must convert grams of water to mole.
The minus sign indicates that heat energy is taken away from the water.
This result show that CCl2F2 must absorb -175 kJ of heat for the water to
freeze. The mass of CCl2F2 that must be vaporized to do this is
1.22 kg must be evaporated.
Phase Diagrams
A phase diagram is a graphical way to summarize the conditions under
which the different states of a substance are stable in a closed
system. Phase diagrams exist in a variety if 2-D and 3-D forms. They are
used in numerous applications and help us to make predictions about how
a material will behave under different conditions. Here are the basics.
Melting-Point Curve
The phase diagram for water consists of three curves that divide the
diagram into regions representing the different states. In each region, the
indicated state is stable. Every point along the curve represents the
experimentally determined temperatures and pressures at which two states
are in equilibrium. Thus, the curve AB divided the solid and liquid regions,
and represents the conditions under which the solid & liquid are in
equilibrium. This point give the melting points of the solid at various
pressures, and is often called the melting-point curve.
Since MP is only slightly affected by P, the MP curve is almost vertical. If
the liquid is more dense than solid, as is the case for water, the MP
decreases with P (curve leans left). For most substances, the liquid state is
less dense than the solid, and the cure leans right (MP increases with P).
Vapor-Pressure Curves for the Liquid and Solid
The curve AC divides the liquid and gaseous regions, and gives the vapor
pressures of the liquid at various T & P. It also gives the boiling points of
the liquids for various T & P. The curve AD divides the solid and gaseous
regions, and gives the vapor pressure of the solid at various T & P. Point A,
where all curves meet, is called the triple point. The triple point is the point
on a phase diagram representing the temperature and pressure at which
three phases of a substance coexist in equilibrium. For water, the triple
point occurs at 0.01°C, 0.0063 atm, and the solid, liquidm and vapor
phases coexist.
Critical Temperature and Pressure
The temperature above which the liquid sate can no longer exist is the
critical temperature (Tc). On the phase diagram pictured right this is the
denoted by point B. Notice that the vapor-pressure curve disappears at this
point. The vapor-pressure at the critical temperature is called he critical
pressure (Tp). The critical point is where the critical temperatures and
pressures intersect.
Many important gases cannot be liquified at room temperature. Nitrogen,
for example, has a critical temperature of -147 °C. This means that it
cannot be liquefied until the temperature is below -147 °C.
Supercritical Fluid
Lets let Martyn tell you about supercritical fluids. He knows a lot more
about them than I do.
Watch this You Tube video:
https://www.youtube.com/watch?v=yBRdBrnIlTQ
Homework:
Book questions pg 476 questions 37-40, 91
Study Guide book questions pg 233 questions 1-10, 44-53, 65