Name:_______________________________________ Date:__________________ Period:______
Phase Change Notes
I. Phases of matter and phase changes
a. Recall the three main phases of matter:
Energy is involved during the
transition from one phase of matter
to another. You should learn how
energy is involved in each of these
phase transitions, and you should
be able to name each of these
phase transitions.
(1)
(2)
(3)
(4)
(5)
(6)
Melting: energy absorbed
Freezing: energy released
Condensation: energy released
Vaporization: energy absorbed
Deposition: energy released
Sublimation: energy absorbed
The heat content of a chemical system is called the enthalpy (symbol: H)
The enthalpy change ( H) is the amount of heat released or absorbed when a chemical reaction
occurs at constant pressure.
H = H(products) - H(reactants) = negative
Energy is Released = exothermic
H = H(products) - H(reactants) = positive
Energy is Absorbed = endothermic
Figure 1: A pictorial summary of phase changes. For our purposes, "enthalpy" (symbol = H) is added heat energy.
II. Heat and temperature
a. Heat is a form of energy that flows from hot objects to cold objects. Heat is measured in joules.
b. Temperature is a measure of the average kinetic energy of the particles in a substance. Temperature
is measured in °C or K.
c. Although it seems counterintuitive, a substance that is
heated will not undergo a continuous rise in
temperature. Likewise, if a substance is cooled, its
temperature will not decrease uniformly.
In fact, as a substance such as ice is heated, its
temperature will only increase until a phase change
occurs. During a phase change, the temperature of a
substance does not change.
When the temperature of a substance that is being
cooled or heated does not change, it is undergoing a
phase change.
HS_Solutions_Phase_Change_Notes.doc
adapted from Ken MacGillivray
d. In the diagram Heating/Cooling Curve for H20, one can see that as ice is heated from -10°C, its
temperature increases only until 0°C. From -10°C to 0°C, all of the heat absorbed by the ice goes into
increasing the kinetic energy of the ice molecules (they vibrate faster).
e. At 0°C, however, the particles of ice can not move any faster. Therefore, the temperature does not rise. At
this point, any absorbed energy goes into making bonds between water molecules break. These bonds
between water molecules are intermolecular. This period is called the Heat of Fusion.
i. The amount of energy needed to melt a substance at its melting point is called the heat of fusion of
that substance. For H2O, that amount of energy is 6.02 kJ/mol.
ii. The temperature at which melting happens is called the normal melting point of the substance.
Believe it or not the melting point is the same as the freezing point. They are the same
temperature.
iii. The amount of energy that must be released in order for a substance to freeze at it s freezing
point is also called the heat of fusion of that substance. For H2O, that amount of energy is 6.02
kJ/mol.
f. After the ice has entirely melted, the water that forms can be heated to a higher temperature. The
temperature of water (at standard pressure) can only reach 100°C.
i. The temperature at which a substance boils is called the boiling point of as substance.
ii. The temperature at which a substance boils when the pressure is 1 atm (standard pressure) is
called the normal boiling point of as substance. The boiling point and melting point of a
substance depend on vapor pressure as well as temperature.
iii. The amount of heat
needed to boil a
substance at its
Boiling Point
boiling point is called
Temperature
the heat of
vaporization of that
substance. For H2O,
that amount of
energy is 40.6
kJ/mol.
iv. The amount of heat
that must be
released by a gas Freezing
Melting
at its boiling point –
Point
so that it can
condense to a liquid
is also called the
heat of vaporization
of that substance.
For H2O, that
amount of energy is 40.6 kJ/mol. The boiling temperature for a liquid must be the same as the
condensation temperature for the gas phase of that substance.
III. Calculations involving the heating/cooling curve.
a. There are five different regions to the heating/cooling curve for water.
i. The heating/cooling curve for other substances are constructed the same for other substances,
but will have different slopes and runs
ii. Water is a very important substance, it is the essence of life that without which the world as
we know would not be. Therefore, knowing how to perform the calculations for water’s phase
changes is very important to a scientist.
q = heat
b. The heating/cooling of ice: Use q=mCΔT
c. The melting/freezing of H2O: Use 6.02 kJ/mol
m = mass
d. The heating./cooling of water: Use q=mCΔT
C = specific heat of the substance
e. The boiling/condensation of H2O: Use 40.6 kJ/mol
ΔT = change in temperature
f. The heating/cooling of steam: Use q=mCΔT
HS_Solutions_Phase_Change_Notes.doc
adapted from Ken MacGillivray
IV. Intermolecular forces of attraction
a. These are the forces that hold molecule to molecule, not atom to atom.
b. Molecule to molecule = intermolecular; atom to atom in a molecule = intramolecular. Remember
the three kinds of intramolecular bonds: ionic bonds (
), covalent(
) or metallic bonds (
).
c. We care about three main intermolecular forces of attraction: hydrogen bonding, dipole-dipole
interactions, and London dispersion forces.
i. VERY STRONG: Hydrogen bonding – the attraction between an unshared pair on an atom of
one molecule for an electropositive (“electronstarved”) hydrogen atom on another nearby
molecule.
ii. SOMEWHAT STRONG: Dipole-dipole interactions – the attraction between to polar molecules
iii. WEAKEST: London dispersion forces aka Van Der Waals– the attraction that one nonpolar
molecule has for another nonpolar molecule, due to the temporary uneven distribution of
electrons within molecules which causes them to be weakly and temporarily polar.
V. Vapor pressure, equilibrium, and evaporation
a. All liquids evaporate when left out “in the air.” The highest energy molecules in the liquid phase, when they
are on the surface of the liquid, escape to the gas phase.
b. Even contained liquids will evaporate, that is, they will proceed from the liquid to the gaseous state of
matter.
c. However, in a contained liquid, as the liquid particles evaporate, the gas formed by them (the vapor)
begins to exert a pressure just like any other gas in a container would.
d. As the pressure increases due to the growing number of gas particles, the likelihood of two gas particles
interacting and re-liquefying (condensation) increases. After a while, the number of particles entering the
vapor phase equals the number of particles re-entering the liquid
phase. At this point, we say that a state of equilibrium has been
established between the gas and liquid phases. Equilibrium is
when maximum concentrations possible for a substance in all
phases is achieved.
e. The pressure exerted by a gas that is in equilibrium with its own
liquid phase at a certain temperature is called the vapor pressure
of that substance.
f. Huge tables are published in huge books detailing the vapor
pressures of substances. Alternatively, there are mathematical
equations that can predict (or at least estimate) the vapor
pressure of a liquid from some of its physical properties.
VI. Phase Diagrams
a. A particular substance can exist in several different phases – sometimes simultaneously – depending on the
pressure and temperature at which the substance is held.
b. The phase state behavior of a substance can be summarized with a phase diagram.
c. Key features:
i. Triple point: the conditions of temperature and pressure at which all three phases of matter can
coexist.
ii. Vapor pressure: the pressure at which a contained liquid is in equilibrium with its gas (vapor)
phase. Use the G-L line two different ways:
1. The boundary line between the “gas” and “liquid” sections of the graph represents the vapor
pressure of the substance. As temperature goes up, the vapor pressure goes up. So, you can
determine the vapor pressure for a substance at a given temperature.
2. The boundary between G and L also represents the boiling temperature of the substance at a
given pressure.
iii. Normal boiling point: this is the boiling point at standard atmospheric pressure (1 atm). Notice that
water can be boiled at many temperatures, depending on the pressure. In Denver, Colorado, water
boils at 95°C; on Mount Everest, water boils at 70°C. What non-scientists call “boiling point” is what
we call “normal boiling point”.
HS_Solutions_Phase_Change_Notes.doc
adapted from Ken MacGillivray
iv. Normal freezing point: This is the freezing point at standard atmospheric pressure (1 atm). Notice
that water can be turned to ice at many temperatures, depending on the pressure
d. Different substances have different phase diagrams because they behave differently as the pressure and
temperature change.
Test your knowledge of phase diagrams by answering the following questions:
What is the triple point of H2O?
What is the triple point of CO2?
What is the boiling temperature of water at 1 atm (101.3 kPa) of pressure?
Given the phase diagram for water state what phase(s) of water is/are present at each of the following
temperature-pressure conditions:
a) at any point on curve BO
b) at any point on curve OA
c) at any point on curve OC
d) at point O
2. Based on the phase diagram from the previous question, what effect would each of the following changes
have on a sample of water at any point on curve OA:
a) increasing the temperature at constant pressure
b) decreasing the pressure at constant temperature
c) decreasing the temperature at constant pressure
d) increasing the pressure at constant temperature
3. Given a sample of water at any point on curve BC in the phase diagram given above, how could more liquid
water in the sample be converted into a solid at room temperature?
HS_Solutions_Phase_Change_Notes.doc
adapted from Ken MacGillivray
Practice Questions
I.
Label each of the 5 phase changes in the diagram
below with the letter of the correct response. Not all
of the letters get used.
a. Vaporization
b. Sublimation
c. Melting
d. Freezing
e. Deposition
II. Label each of the blanks below as either “energy required” or “energy released”.
A= ENERGY REQUIRED
B= ENERGY RELEASED
III. The diagram below shows three different flasks. The substance in the flasks is H O. Match
2
each description with the letter of the appropriate flask in the diagram.
10) Which of the flasks shows a liquid that has reached equilibrium with its vapor?
11) Which of the flasks depicts a liquid that will eventually reach equilibrium with its vapor?
12) Which of the pictures shows a flask in which the liquid WILL NOT reach equilibrium with its vapor?
HS_Solutions_Phase_Change_Notes.doc
adapted from Ken MacGillivray
IV. The atmospheric pressure in Denver, Colorado on a particular day is 658 mm Hg. The table
below lists the vapor pressure of water at various temperatures. Given that at 760 mmHg
water boils at 100 °C determine the approximate temperature at which water will boil in
Denver, CO on this day.
13)
a) 100 °C
b) 105 °C
c) 96 °C
d) 106 °C
e) 0 °C
14. Order the intermolecular forces (dipole-dipole, London dispersion, and hydrogen bonding) from
weakest to strongest.
[A] dipole-dipole, London dispersion, hydrogen bonding
[B] London dispersion, dipole-dipole, hydrogen bonding
[C] hydrogen bonding, dipole-dipole, London dispersion
[D] London dispersion, hydrogen bonding, dipole-dipole
15. The intermolecular forces called hydrogen bonding will not exist between molecules of
[A] NH [B] H [C] HF [D] H O [E] any of these
3
2
2
16. At 1 atm of pressure and a temperature of 0°C, which phase(s) of H2O can exist?
[A] ice and water vapor [B] water only [C] ice only [D] ice and water [E] water vapor only
17. The normal freezing point of water is
[A] 0°F [B] 32°C [C] 273 K [D] 373°C [E] none of these
18. The normal boiling point of water is
[A] 373 K [B] 0°F [C] 32°F [D] 273 K [E] none of these
19. Calculate the quantity of energy required to change 3.00 mol of liquid water to steam at 100°C.
The molar heat of vaporization of water is 40.6 kJ/mol.
[A] 300 kJ [B] 13.5 kJ [C] 122 kJ [D] 40.6 kJ [E] none of these
20. Calculate the quantity of energy required to change 26.5 g of liquid water to steam at 100°C. The
molar heat of vaporization of water is 40.6 kJ/mol.
3
[A] 1.08 × 10 kJ [B] 59.8 kJ [C] 1.53 kJ [D] 27.6 kJ [E] none of these
21. The specific heat capacity of liquid water is 4.18 J/g°C. Calculate the quantity of energy required
to heat 10.0 g of water from 26.5°C to 83.7°C.
3
[A] 572 J [B] 837 J [C] 239 J [D] 2.39 × 10 J [E] none of these
22. The molar heat of fusion of water is 6.02 kJ/mol. Calculate the energy required to melt 46.8g of
water.
[A] 6.02 kJ [B] 7.77 kJ [C] 282 kJ [D] 2.32 kJ [E] none of these
23. The freezing point of helium is approximately −270°C. The freezing point of xenon is −112°C.
Both of these are in the noble gas family. Which of the following statements is supported by these
data?
a) The London forces between the helium molecules are greater than the London forces
between the xenon molecules.
b) The London forces between the helium molecules are less than the London forces between
the xenon molecules.
[C] As the molar mass of the noble gas increases, the freezing point decreases.
[D] Helium and xenon form highly polar molecules.
[E] none of these
24. Choose the state of water in which the water molecules are farthest apart on average.
[A] ice (solid) [B] steam (vapor) [C] all the same [D] liquid
25. The process of evaporation happens when which of the following occurs?
[A] A solid becomes a gas.
[B] A liquid becomes a solid.
[C] A solid becomes a liquid.
[D] A gas becomes a liquid.
[E] A liquid becomes a gas.
26. Which of the following processes must exist in equilibrium with the evaporation process when a
measurement of vapor pressure is made?
[A] fusion [B] condensation [C] vaporization [D] boiling [E] sublimation
HS_Solutions_Phase_Change_Notes.doc
adapted from Ken MacGillivray
27. The vapor pressure for water at
100.0°C is
[A] 760 torr [B] More
information is needed.
[C] 85
torr
[D] 1 torr
[E] 175 torr
28. As the temperature of a liquid
increases, the vapor pressure of the
liquid generally
[A] decreases
[B] stays
the same
[C] increases
[D]
depends on the type of
intermolecular forces
29. The boiling temperature of water
is always 100°C. [A] True [B] False
30. As the atmospheric pressure around a liquid decreases, the boiling temperature of the liquid
[A] decreases
[B] stays the same
[C] increases
[D] may increase or decrease depending on the liquid [E] none of these
31. When a substance undergoes a phase change from a solid to a liquid, the temperature of that
substance
[A] decreases
[B] stays the same
[C] increases
[D] may increase or decrease depending on the liquid [E] none of these
32. When a substance undergoes a phase change from a gas to a liquid, the temperature of that
substance
[A] decreases
[B] stays the same
[C] increases
[D] may increase or decrease depending on the liquid [E] none of these
33. The heating/cooling curve of chloroform is shown below. What is the heat of fusion of
chloroform?
[A] -63.6 °C
[B] 61.7 °C
[C] 8.80 kJ/mol
[D] 31.4 kJ/mol
[E] none of these
HS_Solutions_Phase_Change_Notes.doc
adapted from Ken MacGillivray