Transcript Chapter 08

Fundamentals of General, Organic,
and Biological Chemistry
5th Edition
Chapter Eight
Gases, Liquids, and Solids
James E. Mayhugh
Oklahoma City University
2007 Prentice Hall, Inc.
Outline
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8.1 States of Matter and Their Changes
8.2 Gases and the Kinetic–Molecular Theory
8.3 Pressure
8.4 Boyle’s Law: The Relation Between Volume and Pressure
8.5 Charles’s Law: The Relation Between Volume and Temperature
8.6 Gay-Lussac’s Law: The Relation Between Pressure and Temperature
8.7 The Combined Gas Law
8.8 Avogadro’s Law: The Relation Between Volume and Molar Amount
8.9 The Ideal Gas Law
8.10 Partial Pressure and Dalton’s Law
8.11 Intermolecular Forces
8.12 Liquids
8.13 Water: A Unique Liquid
8.14 Solids
8.15 Changes of State
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8.1 States of Matter and Their
Changes
► Matter exists in any of three phases, or states—solid,
liquid, and gas, depending on the attractive forces
between particles, and temperature, and pressure.
► Review Ch 5.9, polarity
► In a gas, the attractive forces between particles are very
weak compared to their kinetic energy, so the particles
move about freely, are far apart, and have almost no
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Hall © 2007on one another.
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Chapter Eight
influence
►In a liquid, the attractive forces between particles are
stronger, pulling the particles close together but still
allowing them considerable freedom to move about.
►In a solid, the attractive forces are much stronger than
the kinetic energy of the particles, so the atoms,
molecules, or ions are held in a specific arrangement and
can only wiggle around in place.
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►Phase change or change of state: The
transformation of a substance from one state to
another.
►Melting point (mp): The temperature at which solid
and liquid are in equilibrium.
►Boiling point (bp): The temperature at which liquid
and gas are in equilibrium.
►Sublimation: A process in which a solid changes
directly to a gas.
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Melting, boiling, and sublimation all have H > 0, and
S > 0. This means they are nonspontaneous,
below and spontaneous above a certain
temperature.
Matter (s)
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+ heat
- heat
Matter (l)
Chapter Eight
+ heat
- heat
Matter (g)
6
When isopropyl alcohol, (CH3)2CHOH vaporizes, what
are the signs of H and S for the vaporization
process?
1.
2.
3.
4.
H = – and S = –
H = – and S = +
H = + and S = –
H = + and S = +
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When isopropyl alcohol, (CH3)2CHOH vaporizes, what
are the signs of H and S for the vaporization
process?
1.
2.
3.
4.
H = – and S = –
H = – and S = +
H = + and S = –
H = + and S = +
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Chapter Eight
The figures below depict mercury in three different states. Identify
the change of state occurring when mercury is taken from the state
in figure (c) to the state in figure (b), and give the name of the
quantity of energy involved.
1.
2.
3.
4.
Boiling, heat of vaporization
Dissolution, heat of solution
Melting, heat of fusion
Sublimation, heatChapter
of sublimation
Seven
The figures below depict mercury in three different states. Identify
the change of state occurring when mercury is taken from the state
in figure (c) to the state in figure (b), and give the name of the
quantity of energy involved.
1.
2.
3.
4.
Boiling, heat of vaporization
Dissolution, heat of solution
Melting, heat of fusion
Sublimation, heatChapter
of sublimation
Seven
Chapter Seven
Chapter Seven
8.2 Gases and the Kinetic-Molecular
Theory
► The behavior of gases can be explained by a group
of assumptions known as the kinetic–molecular
theory of gases. The following assumptions account
for the observable properties of gases:
► A gas consists of many particles, either atoms or
molecules, moving about at random with no
attractive forces between them. Because of this
random motion, different gases mix together
quickly.
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► The amount of space occupied by the gas particles
themselves is much smaller than the amount of
space between particles. Most of the volume taken
up by gases is empty space, accounting for the ease
of compression and low densities of gases.
► The average kinetic energy of gas particles is
proportional to the Kelvin temperature. Thus, gas
particles have more kinetic energy and move faster
as the temperature increases. (In fact, gas particles
move much faster than you might suspect. The
average speed of a helium atom at room
temperature and atmospheric pressure is
approximately 1.36 km/s, or 3000 mi/hr, nearly that
of a rifle bullet, sound travels 760 mi/hr)
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► Collisions of gas particles, either with other
particles or with the wall of their container, are
elastic; that is, the total kinetic energy of the
particles is constant. The pressure of a gas against
the walls of its container is the result of collisions of
the gas particles with the walls. The number and
force of collisions determines the pressure.
► A gas that obeys all the assumptions of the kinetic–
molecular theory is called an ideal gas. All gases
behave somewhat differently than predicted by the
kinetic–molecular theory at very high pressures or
very low temperatures. Most real gases display
nearly ideal behavior under normal conditions.
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8.3 Pressure
► Pressure (P) is defined as a force (F) per unit area
(A) pushing against a surface; P = F/A.
► A barometer measures pressure as the height of a
mercury column. Atmospheric pressure presses
down on mercury in a dish and pushes it up a tube.
► Pressure units:
1 atm = 760 mm Hg = 14.7 psi = 101,325 Pa
1 mm Hg = 1 torr = 133.32 Pa
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A column of air
weighing 14.7 lb
presses down on each
square inch of the
Earth’s surface at sea
level, resulting in what
we call atmospheric
pressure.
Gas pressure inside a container is often measured
using an open-end manometer, a simple instrument
similar in principle to the mercury barometer.
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Chapter Seven
Chapter Seven
Oxygen gas is contained in the apparatus shown
below. What is the pressure of the oxygen in
the apparatus?
1.
2.
3.
4.
500 mm Hg
725 mm Hg
775 mm Hg
1000 mm Hg
Oxygen gas is contained in the apparatus shown
below. What is the pressure of the oxygen in
the apparatus?
1.
2.
3.
4.
500 mm Hg
725 mm Hg
775 mm Hg
1000 mm Hg
8.4 Boyle’s Law: The Relation Between Volume
and Pressure
► Boyle’s law: The volume of a gas is inversely
proportional to its pressure for a fixed amount of
gas at a constant temperature. That is, P times V is
constant when the amount of gas n and the
temperature T are kept constant.
► V  1/P or PV = k if n and T are constant
► If: P1V1 = k and P2V2 = k
► Then: P1V1 = P2V2
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The volume of a gas decreases proportionately as its
pressure increases. If the pressure of a gas sample is
doubled, the volume is halved.
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Learning Check
A sample of helium gas
has a volume of 6.4 L at
a pressure of 0.70 atm.
What is the new volume
when the pressure is
increased to 1.40 atm
(T constant)?
A) 3.2 L
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B) 6.4 L
C) 12.8 L
Solution
A) 3.2 L
Solve for V2: P1V1 = P2V2
V2 = V1P1
P2
V2
= 6.4 L x 0.70 atm = 3.2 L
1.40 atm
Volume decreases when there is an increase in the
pressure (Temperature is constant).
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8.5 Charles’ Law: The Relation Between
Volume and Temperature
► Charles’s law: The volume of a gas is directly
proportional to its Kelvin temperature for a fixed
amount of gas at a constant pressure. That is, V
divided by T is constant when n and P are held
constant.
► V  T or V/T = k if n and P are constant
► If: V1/T1 = k and V2/T2 = k
► Then: V1/T1 = V2/T2
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If the Kelvin temperature of a gas is doubled, its
volume doubles.
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Assume you have a
sample of gas at 350K in
a sealed container, as
represented in part (a).
Which drawing (b) - (d)
represents the gas after
the temperature is
lowered to 150K?
Using Charles’s law you should note that the volume
would decrease with decreasing temperature
Learning Check
Solve Charles’ Law expression for T2.
V1 = V2
T1
T2
T2 = V2T1
V1
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Learning Check
A balloon has a volume of 785 mL at 21°C. If
The temperature drops to 0°C, what is the new
volume of the balloon (P constant)?
1. Set up data table:
Conditions 1
Conditions 2
V1 = 785 mL
V2 = ?
T1 = 21°C = 294 K
T2 = 0°C = 273 K
Be sure that you always use the Kelvin (K)
temperature in gas calculations.
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Learning Check
2. Solve Charles’ law for V2
V1 = V2
T1
T2
V2 = V1 T2
T1
V2 = 785 mL x 273 K = 729 mL
294 K
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8.6 Gay-Lussac’s Law: The Relation Between
Pressure and Temperature
► Gay-Lussac’s law: The pressure of a gas is directly
proportional to its Kelvin temperature for a fixed
amount of gas at a constant volume. That is, P
divided by T is constant when n and V are held
constant.
► P  T or P/T = k if n and V are constant
► If: P1/T1 = k and P2/T2 = k
► Then: P1/T1 = P2/T2
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As the temperature goes up, the pressure also goes up.
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Calculation with Gay-Lussac’s Law
A gas has a pressure at 2.0 atm at 18°C. What
is the new pressure when the temperature is
62°C? (V and n constant)
1. Set up a data table.
Conditions 1
Conditions 2
P1 = 2.0 atm
P2
= ?
T1 = 18°C + 273
T2
= 62°C + 273
= 291 K
= 335 K
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Calculation with Gay-Lussac’s Law
(continued)
2. Solve Gay-Lussac’s Law for P2
P1 = P2
T1
T2
P2 = P1 T2
T1
P2 = 2.0 atm x 335 K = 2.3 atm
291 K
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Learning Check
Use the gas laws to complete with
1) Increases 2) Decreases
A. Pressure Increases
_______ when V decreases.
decreases
B. When T decreases, V _______.
C. Pressure decreases
_______ when V changes
from 12.0 L to 24.0 L.
D. Volume increases
_______ when T changes from
15.0 °C to 45.0°C.
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8.7 The Combined Gas Law
► Since PV, V/T, and P/T all have constant values for a
fixed amount of gas, these relationships can be
merged into a combined gas law for a fixed amount
of gas.
► Combined gas law: PV/T = k if n constant
► P1V1/T1 = P2V2/T2
► If any five of the six quantities in this equation are
known, the sixth can be calculated.
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Learning Check
A sample of helium gas has a volume of 0.180 L,
a pressure of 0.800 atm and a temperature of
29°C. At what temperature (°C) will the helium
have a volume of 90.0 mL and a pressure of 3.20
atm (n constant)?
1. Set up Data Table
Conditions 1
Conditions 2
P1 = 0.800 atm
P2 = 3.20 atm
V1 = 0.180 L (180 mL)
V2 = 90.0 mL
T1 = 29°C + 273 = 302 K
T2 = ??
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Learning Check
2. Solve for T2
P1 V1
T1
=
P2 V2
T2
T2 = T1 P2V2
P1V1
T2 = 302 K x 3.20 atm x 90.0 mL = 604 K
0.800 atm 180.0 mL
T2 = 604 K – 273 = 331 °C
45
Learning Check
A gas has a volume of 675 mL at 35°C and 0.850
atm pressure. What is the volume(mL) of the
gas at –95°C and a pressure of 802 mm Hg (n
constant)?
46
Learning Check
Data Table
T1 = 308 K
T2 = -95°C + 273 = 178K
V1 = 675 mL
V2 = ???
P1 = 646 mm Hg
P2 = 802 mm Hg
Solve for T2
V2 = V1 P1 T2
P2T1
V2 = 675 mL x 646 mm Hg x 178K = 314 mL
802 mm Hg x 308K
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8.8 Avogadro’s Law: The Relation
Between Volume and Molar Amount
► Avogadro’s law: The volume of a gas is directly
proportional to its molar amount at a constant
pressure and temperature. That is, V divided by n is
constant when P and T are held constant.
► V  n or V/n = k if P and T are constant
► If: V1/n1 = k and V2/n2 = k
► Then: V1/n1 = V2/n2
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► The molar amounts of any two gases with the same
volume are the same at a given T and P.
► Standard temperature and pressure:
(STP) = 0C (273.15 K) and 1 atm (760 mm Hg)
► Standard molar volume of a gas at STP = 22.4 L/mol
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► Avogadro’s law. Each of these 22.4 L bulbs contains 1.00 mol of
gas at 0 degrees C and 1 atm pressure.
► Equal volumes of gases contain equal numbers of moles,
regardless of the identity of the gas. Under standard temperature
andHall
pressure,
1.00 mole of Chapter
any gas
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© 2007
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Eight will occupy 22.4 liters.
8.9 The Ideal Gas Law
► Ideal gas law: The relationships among the four
variables P, V, T, and n for gases can be combined
into a single expression called the ideal gas law.
► PV/nT = R (A constant value) or PV = nRT
► If the values of three of the four variables in the
ideal gas law are known, the fourth can be
calculated.
► Values of the gas constant R:
For P in atm:
R = 0.0821 L·atm/mol·K
For P in mm Hg: R = 62.4 L·mm Hg/mol·K
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8.10 Partial Pressure and Dalton’s law
► Dalton’s law: The total pressure exerted by a gas
mixture of (Ptotal) is the sum of the partial pressures
of the components in the mixture.
► Dalton’s law Ptotal = Pgas1 + Pgas2 + Pgas3 + …
► Partial pressure: The contribution of a given gas in
a mixture to the total pressure.
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Chapter Seven
8.11 Intermolecular Forces
► Intermolecular force: A force that acts between
molecules and holds molecules close to one
another. There are three major types of
intermolecular forces.
► Dipole–dipole forces are weak, with strengths on
the order of 1 kcal/mol
► London dispersion forces are weak, in the range
0.5–2.5 kcal/mol. They increase with molecular
weight and molecular surface area.
► Hydrogen bonds can be quite strong, with energies
up to 10 kcal/mol. Chapter Eight
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Dipole–dipole forces: The positive and negative ends
of polar molecules are attracted to one another by
dipole–dipole forces. As a result, polar molecules have
higher boiling points than nonpolar molecules of
similar size.
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►Only polar molecules experience dipole–dipole
forces, but all molecules, regardless of structure,
experience London dispersion forces.
►(a) On average, the electron distribution in a
nonpolar molecule is symmetrical. (b) At any instant,
it may be unsymmetrical, resulting in a temporary
polarity that can attract neighboring molecules.
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A hydrogen bond is an attractive interaction between
an unshared electron pair on an electronegative O, N,
or F atom and a positively polarized hydrogen atom
bonded to another electronegative O, N, or F.
Hydrogen bonds occur in both water and ammonia.
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The boiling points of NH3, H2O, and HF are much higher
than the boiling points of their second row neighbor CH4
and of related third-row compounds due to hydrogen
bonding.
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Chapter Seven
Shown below is a model of liquid water. What is the
major intermolecular force of attraction responsible
for water being a liquid?
1.
2.
3.
4.
Covalent bonding
Dipole-dipole
Hydrogen bonding
London dispersion
Shown below is a model of liquid water. What is the
major intermolecular force of attraction responsible
for water being a liquid?
1.
2.
3.
4.
Covalent bonding
Dipole-dipole
Hydrogen bonding
London dispersion
Which of the following compounds form hydrogen bonds?
Methyl alcohol and methylamine both have hydrogen
attached to a very polar atom (oxygen, nitrogen) so they
may form hydrogen bonds.
8.12 Liquids
► Molecules are in constant motion in the liquid state.
If a molecule happens to be near the surface of a
liquid, and if it has enough energy, it can break free
of the liquid and escape into a state called vapor.
► Once molecules have escaped from the liquid into
the gas state, they are subject to all the gas laws.
The gas molecules make their own contribution to
the total pressure of the gas above the liquid
according to Dalton’s law. We call this contribution
the vapor pressure of the liquid.
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► Vapor pressure rises with increasing temperature
until ultimately it becomes equal to the pressure of
the atmosphere. At this point, bubbles of vapor
form under the surface and force their way to the
top; this is called boiling.
► At a pressure of exactly 760 mm Hg, boiling occurs
at what is called the normal boiling point.
► If atmospheric pressure is higher or lower than
normal, the boiling point of a liquid changes
accordingly. At high altitudes, for example,
atmospheric pressure is lower than at sea level, and
boiling points are also lower.
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At a liquid’s boiling point, its vapor pressure is equal
to atmospheric pressure. Commonly reported boiling
points are those at 760 mm Hg.
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► Surface tension: the resistance of a liquid to spread
out and increase its surface area. The beading-up of
water on a newly waxed car is due to surface
tension.
► Surface tension is caused by the difference between
the forces experienced by molecules at the surface
and those experienced by molecules in the interior.
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8.13 Water: A Unique Liquid
► Water covers nearly 71% of the Earth’s surface, it
accounts for 66% of the mass of an adult human
body, and it is needed by all living things.
► Water has the highest specific heat of any liquid,
giving it the capacity to absorb a large quantity of
heat while changing only slightly in temperature.
► As a result, large lakes and other bodies of water
tend to moderate the air temperature and the
human body is better able to maintain a steady
internal temperature under changing outside
conditions.
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► Water has an unusually high heat of vaporization
(540 cal/g), it carries away a large amount of heat
when it evaporates.
► Your body relies on the cooling effect of water
evaporation.
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Most substances are more dense as solids than as
liquids because molecules are more closely
packed in the solid than in the liquid. Water,
however, is different. Liquid water has a
maximum density of 1.000 g/mL at 3.98°C but
then becomes less dense as it cools. When it
freezes, its density decreases still further to 0.917
g/mL. Ice floats on liquid water, and lakes and
rivers freeze from the top down. If the reverse
were true, fish would be killed in winter.
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The large spaces
that form when
these crystals
form cause ice to
be less dense
than liquid water.
Therefore ice
freezes on the
surface of lakes
in cold climates,
insulating the
water below and
protecting
aquatic life from
extremely cold
temperatures.
8.14 Solids
► There are many different kinds of solids. The most
fundamental distinction between solids is that some
are crystalline and some are amorphous.
► Crystalline solid: A solid whose atoms, molecules,
or ions are rigidly held in an ordered arrangement.
Crystalline solids can be further categorized as ionic,
molecular, covalent network, or metallic.
► Amorphous solid: A solid whose particles do not
have an orderly arrangement.
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A summary of the different types of solids and their
characteristics is given below.
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8.15 Changes of State
► When a substance changes state, energy added is
used to overcome attractive forces instead of
increasing kinetic energy so temperature does not
change.
► Heat of fusion: The quantity of heat required to
completely melt a substance once it has reached its
melting point.
► Heat of vaporization: The quantity of heat required to
completely vaporize a substance once it has reached
its boiling point.
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A heating curve for water, showing the temperature
and state changes that occur when heat is added.
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What is the boiling point of the substance having
the heating curve shown below?
1.
2.
3.
4.
–50oC
12oC
75oC
125oC
What is the boiling point of the substance having
the heating curve shown below?
1.
2.
3.
4.
–50oC
12oC
75oC
125oC
Chapter Seven
Chapter Summary
►According to the kinetic–molecular theory of gases,
the behavior of gases can be explained by assuming
that they consist of particles moving rapidly at
random, separated from other particles by great
distances, and colliding without loss of energy.
►Boyle’s law says that the volume of a fixed amount
of gas at constant temperature is inversely
proportional to its pressure.
►Charles’s law says that the volume of a fixed amount
of gas at constant pressure is directly proportional to
its Kelvin temperature.
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Chapter Summary Cont.
►Gay-Lussac’s law says that the pressure of a fixed
amount of gas at constant volume is directly
proportional to its Kelvin temperature.
►Avogadro’s law says that equal volumes of gases at
the same temperature and pressure contain the
same number of moles.
►The four gas laws together give the ideal gas law,
PV = nRT, which relates the effects of temperature,
pressure, volume, and molar amount.
►At 0°C and 1 atm pressure, called standard
temperature and pressure (STP), 1 mol of any gas
occupies a volume of 22.4 L.
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Chapter Summary Cont.
►The pressure exerted by an individual gas in a
mixture is called the partial pressure. Dalton’s law:
the total pressure exerted by a mixture is equal to
the sum of the partial pressures of the individual
gases.
►There are three major types of intermolecular forces,
which act to hold molecules near one another in
solids and liquids. Dipole–dipole forces occur
between polar molecules. London dispersion forces
occur between all molecules as a result of temporary
molecular polarities. Hydrogen bonding, the
strongest of the three forces, occurs between a
hydrogen atom bonded to O, N, or F and a nearby O,
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Chapter Eight
N, or F atom.
Chapter Summary Cont.
►Crystalline solids are those whose constituent
particles have an ordered arrangement; amorphous
solids lack internal order. There are several kinds of
crystalline solids, ionic solids, molecular solids,
covalent network solids, and metallic solids,.
►The amount of heat necessary to melt a given
amount of solid at its melting point is its heat of
fusion. Molecules escape from the surface of a liquid
resulting in a vapor pressure of the liquid. At a
liquid’s boiling point, its vapor pressure equals
atmospheric pressure. The amount of heat necessary
to vaporize a given amount of liquid at its boiling
point is called its heat of vaporization.
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End of Chapter 8
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