XI – States of Matter
Chemistry Student Notes
Chemistry Unit XI – States of Matter
PRE-TEST QUESTIONS
1. Which two elements are liquid at room temperature?
2. True/False: Nonmetals exist only as solids and liquids at room temperature.
3. What is pressure?
4. What does STP stand for?
5. What are the conditions of STP?
I. Kinetic Theory of Matter
A. Kinetic Theory of Matter
1. The word “kinetic” means ___________________.
a. Kinetic energy, then, is the energy due to motion.
2. The kinetic theory of matter states that all matter consists of tiny particles that are in
constant ___________________.
a. Very useful in explaining properties of the states of matter.
b. Depending on the substance, the basic particles can be ___________________,
___________________ or ___________________.
B. States of Matter
1. To review:
Properties of States of Matter
Solid
Liquid
Gas
Shape?
Volume?
Compressibility?
Particle Proximity?
Flows?
Picture
C. Kinetic Energy and Temperature
1. Temperature is defined as a measure of the __________________ _________________
___________________of a sample of matter (NOT how hot or cold something is).
a. Because some particles are moving faster than others, temperature is a measure of the
___________________ kinetic energy.
2. As this relates to energy…
a. As a substance absorbs energy, some is absorbed as ___________________ energy,
which does not cause the substance to move.
b. The remaining energy that is absorbed ___________________ up the particles in the
matter (that is, increasing the ___________________ energy).
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XI – States of Matter
Chemistry Student Notes
c. This increase in kinetic energy results in an increase in ___________________.
3. At any given temperature, the particles of all substances, regardless of physical state,
have the same average kinetic energy.
a. For example: at room temperature, the ions in table salt (a solid), the molecules in
water (a liquid) and the atoms in helium (a gas) all have the SAME average kinetic
energy.
b. Note that at higher temperatures, a wider range of kinetic energies exist.
4. Changes in kinetic energy and temperature…
a. When a substance heats up, the particles move faster, increasing their average kinetic
energy, and thus, the temperature of the sample.
b. When a substance cools down, the particles slow down, decreasing their average
kinetic energy, and thus, the temperature of the sample.
II. The Nature of Solids
A. A Model for Solids
1. A solid is a state of matter that has a ___________________ shape and volume.
2. The particles in a solid are strongly attracted to one another, so they do not
___________________ around, but instead, ___________________ in place around
fixed positions.
a. This results in solids having a ___________________ shape.
b. Solids do ___________________ flow easily since their particles are stationary.
3. The particles in a solid are strongly attracted to one another, and thus, are very
___________________ together.
a. This causes solids to have a ___________________ volume.
b. Solid particles are not easily compressed because the particles are already close
together.
4. Solids do not ___________________ much when heated.
B. Melting
1. The melting point is the temperature at which a solid changes into a liquid.
a. When the particles gain enough kinetic energy, they overcome the
_________________________ that hold them in ___________________ positions
and can move around.
b. However, as a liquid, they still lack the energy to overcome the attractions that keep
the particles close together, though they are now mobile as a liquid.
2. Please note that both the melting point and the freezing point are the
___________________ temperature.
a. At that temperature, the liquid and solid phases are in equilibrium.
C. Crystal Structure and Unit Cells
1. In a crystal, a solid in which the atoms, ions, or molecules are arranged in an orderly,
repeating, three-dimensional pattern called a crystal lattice.
2. In general, ionic solids have ___________________ melting points and molecular solids
have ___________________ melting points.
3. Not all solids melt when heated.
a. Some, like wood and cane sugar, ______________________ when heated.
D. Allotropes
1. Allotropes are two or more different molecular forms of the same element in the same
physical state.
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XI – States of Matter
Chemistry Student Notes
2. Carbon is an excellent example…
a. In diamond, each carbon atom is bonded to four other carbon atoms.
i. Result: Very strong bonds (___________________ solid)
b. Graphite consists of layers of carbon atoms bonded to three other carbon atoms in a
hexagonal honeycomb layer, where each layer is attracted to other layers.
c. Fullerenes are a class of compounds composed entirely of carbon. One fullerene is
comprised of _____ carbon atoms arranged in a ball-like cage with hexagons and
pentagons alternating around its surface, like a soccer ball.
i. Named in honor of Buckminster Fuller, who designed geodesic domes, are called
buckminsterfullerene, or buckyballs, for short.
d. All three of these allotropes are very different, though all are comprised of entirely
carbon:
i. Diamond has a high density and is very hard.
ii. Graphite has a relatively low density and is soft and slippery (pencil ‘lead’)
iii. Fullerenes have hollow cages that grant them great strength and rigidity.
3. The only other elements that have allotropes include:
a. ___________________ (white [aka yellow], red, and even violet and black)
b. ___________________ (second only to carbon in the number of allotropes)
c. ___________________ (O2 and O3)
d. ___________________
e. ___________________
E. Amorphous Solids
1. Some solids have no regular geometric shape.
2. An amorphous solid lacks an ordered internal structure.
a. Rubber, plastic and asphalt are amorphous solids.
3. A glass is a transparent fusion product of inorganic substances that have cooled to a rigid
state without crystallizing.
a. Glass is sometimes called a supercooled ______________________.
4. When a crystalline solid is shattered, the fragments tend to have the same surface angles
as the original solid.
a. By contrast, when an amorphous solid, such as glass, is shattered, the fragments have
irregular angles and jagged edges.
III. The Nature of Liquids
A. A Model for Liquids
1. A liquid is a state of matter that has a definite volume, but an indefinite shape.
2. The particles in a liquid never stay in one location, instead ______________________
past one another, in constant motion.
a. This results in liquids having an indefinite shape. They take the shape of their
container.
b. Liquids flow easily because the particles are mobile, moving past one another.
i. Thus, they are classified as a ______________________.
3. The particles in a liquid are attracted to one another pretty strongly, so while they may
slide past one another, they still remain close together.
a. This causes liquids to have a definite volume.
b. Liquid particles are not easily compressed because the particles are close together.
4. Liquids do not expand much when heated.
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XI – States of Matter
Chemistry Student Notes
B. Evaporation
1. Vaporization is any process whereby a liquid turns into a gas.
a. Evaporation is the process that occurs at a temperature ______________________
the boiling point, whereby a liquid turns into a gas.
b. A ______________________, then, is a substance in the gaseous state that is
normally a liquid under those specific conditions.
2. During evaporation, only those molecules with a certain minimum kinetic energy can
escape from the surface of the liquid.
3. Liquids evaporate more quickly if heated.
a. This is because their kinetic energy is ______________________.
4. As particles evaporate, the highest kinetic energy particles are removed _____________.
a. The particles remaining have ______________________ kinetic energy.
b. ______________________ kinetic energy = ______________________ temperature
c. Thus, evaporation is a ______________________ process!
IV. The Nature of Gases
A. A Model for Gases
1. A gas is a state of matter that will take both the shape and volume of its container.
2. The particles in a gas are in constant, ______________________ motion.
a. This results in gases having an indefinite shape. They take the shape of their
container.
b. Gases flow easily because the particles are moving.
3. Because gas particles are moving so quickly, their particles are generally far apart,
overcoming forces of attraction for one another.
a. This causes gases to have an indefinite volume. They will expand or compress to fit
their container. e.g., pumping up a tire.
b. Gas particles are easily compressed because their particles are far apart.
4. Gases expand when heated.
B. Ideal Gases
1. Gases are extremely variable in their properties.
2. As a result, scientists have defined what is considered an “___________________” gas.
a. The molecules in the gas can be considered small hard ______________________,
each with an insignificant volume.
b. All collisions between gas molecules are ______________________ and all motion is
frictionless (no energy is lost in collisions or in motion).
c. The distance between molecules, on average, is much ______________________
than the size of the molecules.
d. The gas molecules are constantly moving in ______________________ directions
with a distribution of speeds.
e. There are no attractive or repulsive ______________________ between the
molecules or the surroundings.
3. An ______________________ collision is one in which kinetic energy is
______________________ without ______________________ from one particle to
another, and the total kinetic energy remains constant.
C. Gases have many properties that will be discussed in greater detail later in this unit.
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XI – States of Matter
Chemistry Student Notes
V. Vapor Pressure and Boiling Point
A. Vapor Pressure
1. Vapor pressure is a measure of the ______________________ exerted by a gas above a
liquid in a sealed container; a dynamic physical equilibrium exists between the vapor and
the liquid.
a. This sealed container would be called a ______________________ system.
2. As particles leave the liquid, they encounter the walls of the container and exert a
pressure.
3. Condensation is the process whereby the vapor particles go from the gaseous state back
to the ______________________ state.
4. Eventually, the number of particles condensing will ___________________ the number
of particles vaporizing.
a. At this point, the vapor pressure will remain ______________________.
5. Note this is a dynamic physical equilibrium. Particles have not stopped evaporating and
condensing; they are simply doing both at the same ______________________.
B. Boiling Point
1. When a liquid is heated to a temperature at which particles throughout the liquid have
enough kinetic energy to vaporize, the liquid begins to boil.
2. The temperature at which the vapor pressure of the liquid is just equal to the external
pressure on the liquid is the boiling point.
3. Because liquids boil when its vapor pressure is equal to the external pressure, liquids
don’t always boil at the same temperature.
a. Atmospheric pressure is lower at high altitudes, so boiling points decrease at higher
altitudes.
4. Pressure cookers take advantage of this.
a. In a pressure cooker, the vapor cannot escape, so the vapor pressure increases. As a
result, the boiling point of water increases and food can cook more
___________________.
5. The temperature of a boiling liquid never rises above its boiling point.
6. Because boiling points can vary, you must specify the pressure when the boiling point
was measured.
a. Normal boiling point is measured at a pressure of ___________________ or
___________________.
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XI – States of Matter
Chemistry Student Notes
VI. Changes in States of Matter
A. Phase changes
1. Draw the diagram of states of matter.
2. Be sure to label the arrows as endothermic or exothermic, as well as labeling what each
change is called.
B. Sublimation
1. Sublimation is the process in which a solid changes to a gas or vapor without passing
through the ___________________ state.
a. Dry ice is solid carbon dioxide that transitions directly into a gas.
b. Ice shrinking in the freezer.
2. Sublimation occurs in solids with vapor pressures that exceed atmospheric pressure at or
near room temperature.
3. Dry ice has a low temperature of ___________________ and is used to package cold
goods because it never turns wet to damage them.
C. Heating and Cooling Curves
1. When a substance changes state, the temperature of the sample is absorbed merely as
potential energy and only when all particles change to the new state, does kinetic energy
(and temperature increase again.
2. Draw the heating curve for water:
3. Be sure to label your diagram appropriately.
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XI – States of Matter
Chemistry Student Notes
D. Phase Diagrams
1. A phase diagram is a graph that
shows the conditions of
temperature and pressure where a
substance exists as either a solid,
liquid or gas.
2. The conditions where two phases
of matter exist in equilibrium are
indicated by the lines separating
the phases.
3. The ___________________ point
represents the only conditions
where all three phases can exist in
equilibrium with one another at
once.
4. The ___________________ point
is the temperature above which
the substance will exist only as a
gas.
5. The slope of the line between solid and liquid is important, as well.
a. Note water’s line slopes to the ___________________. This indicates that water is
___________________ dense as a solid than it is as a liquid (ice floats in liquid
water!)
b. Almost every other phase diagram will lean to the ___________________, meaning
the solid phase is ___________________ than the liquid phase.
VII. Properties of gases
A. Pressure
1. Gas pressure is the result of simultaneous collisions of billions of rapidly moving
particles in a gas with an object.
a. If there are no particles, there are no ___________________.
b. If there are no collisions, there is no ___________________.
c. A ___________________ is an empty space with no particles and thus, no pressure.
2. The atmosphere exerts a pressure on us.
a. Atmospheric pressure results from the ___________________ of atoms and
molecules in air with objects.
b. Atmospheric pressure decreases as you climb a mountain because the density of
earth’s atmosphere decreases as the elevation increases.
3. Chemists use barometers to measure atmospheric pressure. Mercury was used because it
is extremely ___________________.
a. Mercury was filled in a tube and inverted into a pool of mercury.
b. Gravity pulled the mercury in the tube down at the same time the atmosphere is
pushing down on the mercury pool and trying to keep the mercury in the tube.
c. The mercury sank until gravity and atmospheric pressure balanced.
d. This height was found to be ___________________ mm high.
4. The SI unit for pressure is the Pascal (Pa).
a. Very small units, so usually measured in ___________________ (________ Pascals).
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XI – States of Matter
Chemistry Student Notes
b. Millimeters of mercury (mm Hg) are also used because of the above mentioned
barometer. This is also called the ___________________.
c. Atmospheres (atm) are also used.
d. One standard atmosphere (atm) is the pressure required to support 760 mm of
mercury in a mercury barometer at ___________________.
e. A bar (often mentioned in weather forecasts) is equal to ___________________, so
almost equal to atmospheric pressure.
5. The relationship: _____ atm = ____________ kPa = _______ mm Hg = _________ torr
6. Recall: STP is 0°C and 1 atm. Any of the above measurements, which are equal to 1 atm,
are acceptable.
EXAMPLE XI-01: Pressure Conversions
1. Convert 1.25 atmospheres into kilopascals.
2. Convert 800. mm Hg into kilopascals.
B. Compressibility
1. Seatbelts help save a person’s life in a car crash by making them a part of the car and
slowing them down to stop with the car.
a. The compression of the gas particles helps to absorb the ___________________ of
the impact.
2. Compressibility is a measure of how much the volume of matter decreases under
pressure.
a. Unlike other states of matter, gases can be squeezed into smaller volumes.
3. Gases are easily compressed because of the ___________________ between the
particles.
a. At room temperature, the distance between the particles in an enclosed gas is about
_____ times the diameter of a particle.
C. Factors Affecting Gas Pressure
1. The amount of gas, volume, and temperature are all factors that affect gas pressure.
D. Amount of Gas (n)
1. The number of ___________________ of gas is represented by ______.
2. The more gas particles added increases the number of collisions and thus, the pressure.
3. If the volume and temperature of the container do not change, an increase in the number
of moles increases the pressure as well.
a. But, once the pressure exceeds the strength of the container, it will
___________________! >.<
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XI – States of Matter
Chemistry Student Notes
4. Pressure will always flow to ___________________ itself, from high pressure to low
pressure.
a. When an aerosol can is pressed, or a straw is used, the pressure on the other side is
lowered so that the substances flow from higher pressure to lower pressure.
E. Volume (V)
1. You can raise the pressure exerted by a contained gas by _______________________ its
volume.
2. The smaller the volume of the gas, the more often the particles collide with the container
walls and thus, the ___________________ pressure that is produced.
3. Increasing the volume of a container gives the gas particles more room and thus, gas
particles exert ___________________ pressure.
F. Temperature (T)
1. An increase in the temperature of an enclosed gas causes an ___________________ in its
pressure.
2. As a gas is heated, the temperature increases and the average ___________________
energy of the particles in the gas increases.
a. Faster-moving particles impact the walls of their container with more energy.
VIII. The Gas Laws
A. Boyle’s Law
1. In 1662, Robert Boyle suggested a relationship between
pressure and volume.
2. Boyle’s Law states that for any given mass of gas, at
constant ___________________, the volume of the gas
varies inversely with pressure.
a. Mathematically, this is expressed as
3. Graphs of Boyle’s Law look like this:
B. Charles’s Law
1. In 1787, the French physicist Jacques Charles studied the
effect of temperature on the volume of a gas at constant
pressure.
2. Charles’s Law states that for any given mass of a gas, at
constant ___________________, the volume of the gas is
directly proportional to its Kelvin temperature.
a. Mathematically, this is expressed as
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XI – States of Matter
Chemistry Student Notes
3. Graphs of Charles’s Law look like this:
C. Absolute Zero
1. When graphed, data from Charles’s law can be extended to find the temperature at which
the volume of a gas would occupy ______ ___________________.
2. This theoretical temperature is called absolute zero and is equal to ________________.
3. The Kelvin temperature scale is based off absolute zero.
a. Because it is an absolute scale (no negative numbers), there are no ° symbols.
b. Absolute zero is ___________________, or ___________________
c. Celsius and Kelvin have the ___________________ increment scale.
d. So, a ___________________ of 1°C = 1 K.
4. Absolute zero is defined as the theoretical zero point on the Kelvin temperature scale,
equivalent to -273.15°C, where matter would ________________ occupy any volume.
a. This can never happen, though, because real gases ___________________ to liquids
or solids before reaching that temperature.
D. Gay-Lussac’s Law
1. Joseph Gay-Lussac (1778-1850), a French chemist, discovered the relationship between
the pressure and temperature of a gas in 1802.
a. Please note: there is some contention over whether or not this law should be named
after Gay-Lussac, but for simplicity, we will refer to this relationship as such.
2. Gay-Lussac’s Law states that for any given mass of a gas, at constant ______________,
the pressure of a gas is directly proportional to the Kelvin temperature.
a. Pressure cookers work under this principle.
b. Mathematically, this is expressed as
c. Graphs of Gay-Lussac’s Law look like:
E. The Combined Gas Law
1. The combined gas law describes the relationship among the ___________________,
___________________, and ___________________ of an enclosed gas.
2. The combined gas law allows you to do calculations for situations in which only the
amount (number of ___________________) of gas is ___________________.
3. The formula for the combined gas law is:
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UNIT XI Notes, Page 10
XI – States of Matter
Chemistry Student Notes
EXAMPLE XI-02: Combined Gas Law: Solve for volume
1. A helium-filled balloon has a volume of 50.0 L at 25°C and 820. mm Hg. What volume
will it occupy at 650. mm Hg and 10.°C?
EXAMPLE XI-03: Combined Gas Law: Solve for temperature
1. A 0.500-L balloon at 273 K has a pressure of 125 kPa. At what temperature will the
balloon have a volume of 0.750 L and a pressure of 130. kPa?
EXAMPLE XI-04: Boyle’s Law (using Combined Gas Law)
1. A sample of neon gas occupies a space of 3.5 L at a pressure of 553 torr. If the
temperature is held constant, and the volume is increased to 5.4 L, what pressure would
it be under?
IX. Ideal Gases
A. Ideal Gas Assumptions
1. Gases are comprised of large numbers of tiny particles, visualized as tiny spheres.
a. Particles are very far ___________________
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XI – States of Matter
Chemistry Student Notes
b. Thus, most of a gas is ___________________ space.
2. Gas particles are in constant motion, moving rapidly in straight lines in
all directions.
a. Thus, they have ___________________ energy
b. This kinetic energy overcomes the ___________________ forces
between them, except near temperatures where they condense.
3. There are no forces of ___________________ or ___________________
between gas particles.
a. As in pool, when the particles collide, they do not stick together.
4. Collisions between gas particles or between gas particles and container walls are
___________________.
a. No net loss of energy.
5. The average kinetic energy of a gas is ___________________ proportional to the
temperature (measured in Kelvin) of the gas.
a. KE = ½mv² where m is mass and v is the speed of the particle.
b. As long as they’re at the same temperature, all particles have the same kinetic energy.
c. Thus, heavier particles move ___________________ in order to have the same
kinetic energy.
B. The Ideal Gas Law
1. The combined gas law assumes the amount of gas does not vary.
a. It also compares two sets of conditions.
2. The ideal gas law is represented by the mathematical relationship:
a. P = _________________________________________________________
b. V = _________________________________________________________
c. n = _________________________________________________________
d. R = ideal gas constant (look it up depending on units of pressure)
e. T = _________________________________________________________
3. The ideal gas constant, R, has several values…
a. Copy them down (with units!)
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UNIT XI Notes, Page 12
XI – States of Matter
Chemistry Student Notes
EXAMPLE XI-05: Ideal Gas Law
1. What is the volume in liters occupied by .250 moles of oxygen at 20.0°C and 740 mm
Hg pressure?
C. Ideal Gases Vs Real Gases
1. An ideal gas is one that follows the gas laws at all conditions of ___________________
and ___________________.
a. Its particles would have ______ volume.
b. The particles are _____ attracted to one another.
2. Ideal gases do not exist in nature, however.
a. However, under many conditions of temperature and pressure, real gases _________
behave like an ideal gas.
3. Real gases behave like ideal gases at:
a. ___________________ temperatures (particles are moving ___________________
and overcome attractive forces between one another)
b. ___________________ pressures (particles are __________ ___________________,
helps reduce attractive forces as well as allowing them to have no volume)
X. Gases: Mixtures and Movement
A. Dalton’s Law of Partial Pressures
1. Gas pressure results from the collisions of gas particles with container walls, etc.
2. Most gases are mixtures of several gases. In this case, particles in a mixture of gases at
the same temperature have the same average kinetic energy.
a. As a result, the kind of particles is not important as far as its impact on pressure.
3. Composition of Dry Air
Composition of Dry Air
Component
Volume (%)
Partial Pressure (kPa)
Nitrogen
Oxygen
Carbon Dioxide
Argon and Others
Total
a. Notice, this is dry air, so it does not contain water vapor.
4. The contribution each gas in a mixture makes to the total pressure is called the partial
pressure is exerted by that gas.
5. Dalton’s law of partial pressures states that, at ___________________ volume and
temperature, the total pressure exerted by a mixture of gases is equal to the __________
of the partial pressures of the component gases.
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UNIT XI Notes, Page 13
XI – States of Matter
Chemistry Student Notes
a. Mathematically, represented by:
6. If the percent composition of a mixture of gases does not change, the fraction of the
pressure exerted by a gas does not change as the total pressure changes.
a. At the top of Mt. Everest, the total atmospheric pressure is 33.73 kPa, about _______
the value of atmospheric pressure at sea level.
b. The partial pressure of oxygen, then, is also reduced by _________, to 7.06 kPa.
EXAMPLE XI-06: Dalton’s Law of Partial Pressures
1. Three gases are mixed in a container. The pressure of the oxygen gas was 240. mm Hg,
the pressure of the nitrogen gas was 375 mm Hg and the pressure of the hydrogen gas
was 225 mm Hg. What was the total pressure of the mixture of gases?
7. In lab, many gases are collected in a process called ___________________ displacement,
where an inverted tube is filled with water and then the gas is piped in, displacing the
water from the container and thus, being collected.
a. The water has a partial pressure, as well, and can be determined by knowing the
___________________ of the water.
EXAMPLE XI-07: Dalton’s Law of Partial Pressures via Water Displacement
1. 2KClO3 (s) 2KCl (s) + 3O2 (g) : The oxygen was collected from this decomposition
by water displacement. The barometric pressure during the experiment was 731.0 mm
Hg and the temperature was 20.0°C. What was the partial pressure of the oxygen
collected?
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XI – States of Matter
Chemistry Student Notes
B. Graham’s Law
1. A perfume bottle opened in one corner of a room will eventually be smelled in the far
corner from where it was opened.
a. Molecules in the perfume evaporate and diffuse, or spread out, through the air in the
room.
2. Diffusion is the tendency of molecules to move toward areas of lower concentration until
the concentration is uniform throughout.
a. The prefix dis- means “___________________.”
3. During effusion, a gas escapes through a ___________________ ___________________
in its container.
4. Graham’s Law of Effusion states that the rate of effusion of a gas is
___________________ proportional to the ___________________ ________________
of the gas’s molar mass.
a. The relationship KE = ½mv² shows that for kinetic energy to remain constant, mass
and velocity are inversely proportional to one another.
b. That is, lower mass particles move ___________________.
5. Graham’s Law is expressed mathematically as follows:
EXAMPLE XI-08: Rate of Effusion
1. Compare the rate of effusion of H2 gas vs. O2 gas.
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UNIT XI Notes, Page 15