Transcript Fundamentals of General, Organic and Biological Chemistry 6th Edition Chapter Eight Gases, Liquids, and Solids James E.

Fundamentals of General,
Organic and Biological
Chemistry
6th Edition
Chapter Eight
Gases, Liquids, and Solids
James E. Mayhugh
Copyright © 2010 Pearson Education, 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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Goals
► 1. How do scientists explain the behavior of gases? Be able
to state the assumptions of the kinetic–molecular theory of gases
and use these assumptions to explain the behavior of gases.
► 2. How do gases respond to changes in temperature,
pressure, and volume? Be able to use Boyle’s law, Charles’s
law, Gay-Lussac’s law, and Avogadro’s law to explain the effect
on gases of a change in pressure, volume, or temperature.
► 3. What is the ideal gas law? Be able to use the ideal gas law
to find the pressure, volume, temperature, or molar amount of a
gas sample.
► 4. What is partial pressure? Be able to define partial pressure
and use Dalton’s law of partial pressures.
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Goals Contd.
►5. What are the major intermolecular forces, and
how do they affect the states of matter? Be able to
explain dipole–dipole forces, London dispersion
forces, and hydrogen bonding, and recognize which
of these forces affect a given molecule.
►6. What are the various kinds of solids, and how
do they differ? Be able to recognize the different
kinds of solids and describe their characteristics.
►7. What factors affect a change of state? Be able
to apply the concepts of heat change, equilibrium, and
vapor pressure to 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, temperature and pressure.
► 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 influence on one another.
► In a liquid, the attractive forces between particles are
stronger, pulling the particles close together but still
allowing them considerable freedom to move about.
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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.
►Melting, boiling, and sublimation all have H>0, and
S>0. This means they are nonspontaneous below
and spontaneous above a certain temperature.
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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.)
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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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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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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 & 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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Graph (a) demonstrates the decrease in volume as
pressure increases, whereas graph (b) shows the
linear relationship between V and 1/P.
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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 & 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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As the temperature goes up, the volume also goes up.
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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 & 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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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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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 & 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 & 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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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 P total = Pgas1 + Pgas2 + Pgas3 + …
► Partial pressure: The contribution of a given gas in
a mixture to the total pressure.
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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.
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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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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.
► 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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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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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 Contd.
►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 Contd.
►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, N, or F atom.
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Chapter Summary Contd.
►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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Key Words
►Amorphous solid
►Avogadro’s law
►Boiling point (bp)
►Boyle’s law
►Change of state
►Charles’s law
►Combined gas law
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►Crystalline solid
►Dalton’s law
►Dipole–dipole force
►Gas constant (R)
►Gas laws
►Gay-Lussac’s law
►Heat of fusion
►Heat of vaporization
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Key Words Contd.
►Hydrogen bond
►Ideal gas
►Ideal gas law
►Intermolecular force
►Kinetic–molecular
theory of gases (KMT)
►London dispersion
force
►Melting point (mp)
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►Normal boiling point
►Partial pressure
►Pressure (P)
►Standard temperature
and pressure (STP)
►Standard molar
volume
►Vapor
►Vapor pressure
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End of Chapter 8
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