Gases, Liquids, and Solids
Chapter 5
Educational Goals
1.  Define, compare, contrast the terms specific heat, heat of
fusion, and heat of vaporization. Know the equations that
involve these concepts and be able to use them in
calculations.
2.  Describe the meaning of the terms enthalpy change,
entropy change, and free energy change. Explain if a
process is spontaneous or not based on the free energy
change.
3.  Convert between pressure units of atm, torr, and psi.
4.  List the variables that describe the condition of a gas and
give the equations for the various gas laws.
Educational Goals (cont.)
5.  Explain Dalton’s law of partial pressures.
6.  Define the terms density and specific gravity. Given the
density, and either the mass or volume, be able to
determine the volume or mass (respectively).
7.  Know that a liquid will boil when its vapor pressure is
equal to the atmospheric pressure.
8.  Describe, compare, and contrast amorphous solids and
crystalline solids.
9.  Describe the makeup of the four classes of crystalline
solids.
Phases of Matter and Energy
Example: Three phases of water
solid
liquid
gas
ice
water
steam
H2O(s)
H2O(l)
H2O(g)
Why are some molecular compounds solid
while others are gaseous and others are
liquid at room temperature?
Competing Powers
•  Intermolecular forces working to hold
particles together as liquids or solids
•  Kinetic Energy = Motion = Temperature,
work to separate particles
Kinetic Energy = Temperature
One major factor that is responsible for
the varied behavior of solids, liquids, and
gases is the nature of the interaction that
attracts one particle (atom, ion, or
molecule) to another.
What forces hold matter together to
make liquids and solids?
What forces hold matter together to
make liquids and solids?
The attractive forces that hold molecules
together are called Intermolecular forces.
3 Types of Intermolecular Forces
1) Dipole-Dipole Forces
2) Hydrogen Bonding
3) London Forces
–  also called Induced Dipole Forces
Other Noncovalent Interactions
Noncovalent interactions are interactions
that do not involve the sharing of valence
electrons (covalent bonding).
Other noncovalent interactions due to the attraction
of permanent charges.
•  1) Salt bridges
•  2) Ion-dipole interactions
A salt bridge is another name for ionic bond.
Ion-dipole interactions occur between ions with
a full charge and atoms with a partial charge.
Energy meets Matter
Adding energy to liquids will overcome the
forces holding the molecules together– boiling
Adding energy to solids will overcome the
forces holding the molecules together– melting
Phase Changes Language
To reduce a fever, rubbing alcohol (2propanol) can be applied to the skin.
As the alcohol evaporates (liquid
becomes a gas), the skin cools.
Explain the changes in heat energy as
this process takes place. Note: 2-propanol
vapors are flammable, so care must be taken
when using this technique.
Units of Energy
•  One calorie is the amount of energy needed
to raise the temperature of one gram of
water by 1°C
•  joule
–  4.184 J = 1 cal
•  In nutrition, calories are capitalized
–  1 Cal = 1 kcal = 1000 cal
Converting Between Calories and Joules
Example: Convert 60.1 cal to joules
Equivalence statement: 1 cal = 4.184 J
1 cal
4.184 J
60.1 cal
4.184 J
1 cal
Conversion
Factors
4.184 J
1 cal
=
251 J
Calculations Involving
Heat Energy
•  One of two things will happen if energy is added or
removed from matter (assuming no chemical
change takes place).
–  1) Change the phase of the substance
–  Example: melt, freeze, vaporize (boil)
–  2) Change the temperature of the substance
•  You can only do one of these at a time!!!
–  See graph on the next slide
1) Phase Changes
Energy calculations for phase changes may be carried out
using the tabulated values for:
•  heat of fusion (symbol = Hfus) for a substance (Table 5.2).
•  Energy required to melt one gram of a solid
•  Change sign to negative for freezing (liquid to solid)
1) Phase Changes
Energy calculations for phase changes may be carried out
using the tabulated values for:
§  heat of vaporization (symbol Hvap) of a substance
•  Energy required to vaporize one gram of a liquid
•  Change sign to negative for gas going to liquid
Calculations Involving
Changing the Phase
Energy Change = (mass) x (heat of fusion or vaporization)
Delta (∆) means
“change in ___”
ΔE = m x (Hfus or vap)
Get from a table
Calculations Involving
Heat Energy
Example: Determine the amount of energy needed
to melt 155 g of ice at 0°C, we use the heat of
fusion of water (79.7 cal/g) as a conversion factor.
ΔE = (mass) x (heat of fusion)
155 g x
79.7 cal
= 1.24 x 104 cal
g
Note: No Temperature Change!
Ice (0oC) → Water (0oC)
Group Work
A patient with a fever is sponged with 50.0 g
of 2-propanol. How much energy is drawn
from the patient when 2-propanol vaporizes?
(heat of vaporization for 2-propanol is 159 cal/g)
2) Changing Temperature of Matter
•  The amount the temperature of an object increases
depends on the amount of energy added (Q).
–  If you double the added heat energy the temperature
will increase twice as much.
•  The amount the temperature of an object increases
depends on its mass.
–  If you double the mass it will take twice as much heat
energy to raise the temperature the same amount.
Calculations Involving
Changing the Temperature
•  Energy calculations may be carried out using the
values for the specific heat of a substance.
•  Specific heat is the amount of energy required to raise
the temperature of one gram of a substance by one
Celsius degree.
Energy required = Specific Heat x Mass x Temperature Change
Q = S
x
m
x
(ΔT)
∆ is always:
(final) – (initial)
(∆T) = Tfinal-Tinitial
J
By definition, the specific heat of water is 4.184
g °C
Example:
Calculate the amount of heat energy (in joules)
needed to raise the temperature of 7.40 g of
water from 29.0°C to 46.0°C
Q = S x m x ΔT
Specific heat of water:
Mass = 7.40 g
4.184 J
g oC
Temperature Change (ΔT) = 46.0°C – 29.0°C = 17.0°C
Q = 4.184 J
g oC
(7.40 g) (17.0 oC) =526 J
Group Work
How much energy needs to be removed from 175 g
of water to lower the temperature from 23.0oC to
15.0oC ?
How much energy is required to convert 25g of
ice at -7.0oC to water at 50.0oC
Ice
-7.0oC
Step 1
Ice
0.0oC
Temperature Change
Q1=Sicexmx(∆T)
Step 2
Water
0.0oC
Water
50.0oC
Step 3
Phase Change
Temperature Change
∆E2=mxHfus
Q3=Swaterxmx(∆T)
∆Energy Total = Q1 + ∆E2 + Q3
New Topic:
Will a change occur?
Spontaneous vs. Nonspontaneous Changes
•  An important question to ask is why some changes
are:
–  spontaneous (continue to occur once they are started)
OR
–  nonspontaneous (will not run by themselves unless
something keeps them going).
•  Energy is the key factor in determining this.
Spontaneous vs. Nonspontaneous Changes
Spontaneous
vs. Nonspontaneous
Energy vs. Free Energy
The energy (E) of matter depends on the position
(potential energy) and velocity (kinetic energy) of
every molecule in the system.
E = Epotential + Ekinetic
This is not practical to measure in the lab or to
model in calculations!
When working at constant temperature and
pressure, it is mathematically convenient and
experimentally practical to look at the:
Free Energy (G)
Energy vs. Free Energy
Just like the energy (E), in nature, given the
chance, everything proceeds to the lowest
possible free energy (G)!
Free Energy (G)
The “free energy” (ΔG) of a process can be
thought of as the potential for change….
∆G = Gf - Gi
A spontaneous process has a negative ∆G and a
nonspontaneous process has a positive ∆G.
Gases and Pressures
Properties of Gases
How to visualize a gas:
Gas molecules or atoms are very far apart from
one another.
• different from liquids and solids!!
• Gas particles move in a straight line until they
collide with another particle or the container wall.
Properties of Gases
Gases Have Low Density
Because of the relatively large distances
between gas particles, most of the volume
occupied by a gas is empty space.
Properties of Gases
Gases completely fill their container.
Except for a few very heavy gases, most
gasses will completely fill their container.
Properties of Gases
Gases Are Highly Compressible
Gases are
compressible
Liquids and Solids
are not
Properties of Gases
Gases Are Highly Compressible
Compressibility is the ability to make the
space a substance takes up become smaller.
Properties of Gases
Gases can diffuse.
• Gaseous molecules travel at high speeds in all
directions and mix quickly with molecules of gases
in the air in a process called diffusion.
• Diffusion is the movement of one
substance within another substance until
it is evenly distributed.
Properties of Gases
Examples of diffusion.
• Ammonia
• Skunk in da house
Gas Pressure
•  Pressure = total force applied to a certain area
–  larger force = larger pressure
•  Gas pressure is caused by gas molecules
colliding with container walls or surfaces.
Air Pressure
•  Constantly present when air present
•  Decreases with altitude
–  less air
Air Pressure
•  Measured using a barometer
–  Column of mercury supported by air pressure
–  Force of the air on the surface of the mercury balanced by the
pull of gravity on the column of mercury
Various Units for Gas Pressure
•
•
•
•
•
1) atmosphere (atm)
2) height of a column of mercury (mm Hg, in Hg)
3) Torr
4) Pascal (Pa)
6) pounds per square inch (psi, lbs./in2)
Units we will use for pressure:
•  Atmospheres (atm)
•  Pounds per square inch (psi)
•  Millimeters of mercury (mm Hg)
–  also called torr (1mm Hg = 1 Torr)
760. mm Hg
1 atm
760. Torr
1 atm
14.7 psi
1 atm
Relationships:
1 atm = 760. mmHg
1 atm = 760. Torr
1 atm = 14.7 psi
1 atm
760. mm Hg
1 atm
760. Torr
1 atm
14.7 psi
Pressure Unit Conversions
A pressure of 690. Torr is how many
atmospheres?
1 atm = 760 Torr
690. Torr
1 atm
760 Torr
= .908 atm
Group Work
A pressure of 35.0 psi is how many atm?
1 atm = 14.7 psi
A pressure of 812 mm Hg is how many
atmospheres?
1 atm = 760. mm Hg
Gas Laws
Instructional Goals
Understand and be able to use the following gas
laws in calculations:
•  Boyles Law (relationship between pressure and volume)
•  Charles’ Law (relationship between volume and
temperature)
•  Gay-Lussac’s Law (relationship between pressure and
temperature)
•  Avogadro’s Law (relationship between moles and volume)
•  Combined Gas Law (relationship between pressure,
volume and temperature)
•  Ideal Gas Law (relationship between pressure, volume,
number of moles, and temperature)
The Gas Laws
•  The gas laws are the mathematical
equations that show the relationship
between volume, temperature, pressure,
and amount of gas.
•  As with all laws, they were discovered
by experiments.
Boyle’s Law
•  Boyle studied the relationship between volume and
pressure.
• The inverse relationship between pressure and
volume is known as Boyle’s law.
Boyle’s Law
•  When the volume decreases, the
pressure increases
Pressure Gage
low
high
Boyle’s Law
•  When the volume increases, the pressure
decreases
Pressure Gage
low
high
Boyle’s Law
•  Boyle also noticed that when the pressure and/
or volume of a gas is changed the product of
the pressure and volume remains the same.
•  P V = Constant
x
Initial
pressure
Initial
volume
Final
Final
pressure volume
P1 V1 = P2 V2
Boyle’s Law
•  Remember that when using Boyle’s
Law, that the temperature is never
changing.
•  Only the pressure and volume change.
Example
The initial volume of the gas in the piston below is 3.00 liters
and the initial pressure is 1.00 atm.
The piston compressed (at constant temperature) to a new final
volume of 1.00 L. What is the final pressure?
Pressure Gage
low
high
Solution
P1 V1 = P2 V2
P1 V1 = P2 V2
V2
V2
P1 = 1.00 atm P2 = ?
V1 = 3.00 L
P2 =
P1 V1
V2
=
V2 = 1.00 L
(1.00 atm) (3.00 L)
(1.00 L)
= 3.00 atm
Group Work
If the syringe shown has an initial volume of 0.50
mL and the gas in the syringe is at a pressure of
1.0 atm, what is the pressure inside the syringe if
your finger is placed over the opening and the
plunger is pulled back to give a final volume of
3.0 mL?
Charles’ Law
•  Charles observed that as the temperature increases, the
volume increases and vice versa.
Jacques Charles (1746-1823 )
• The direct relationship between temperature and volume
is known as Charles’ law.
Charles’ Law
When the temperature increases, the volume increases
Charles’ Law
When the temperature decreases, the volume decreases
Charles’ Law
•  Charles also noticed that ratio of volume to
temperature of a gas is always the same.
V = Constant
T
Charles’ Law
Charles’ Law
Initial
volume
Initial
temperature
V1
T1
=
V2
T2
Final
volume
Final
temperature
•  Remember that when using Charles’ Law,
that the pressure is never changing.
–  Only the temperature and volume change.
•  Temperature must be Kelvin (K).
–  Kelvin temperature scale is always positive
–  K = oC + 273.15
Example
The initial volume of the gas in the piston below is 1.35 liters
and the initial pressure is 1.00 atm.
The temperature is lowered from 373 K to 250. K (at constant
pressure). What is the final volume?
Pressure Gage
low
high
Solution
V2
V1
=
T2
T1
V2 =
T2 V1
T1
=
V1 = 1.35 L
V2 = ?
T1 = 373 K
T2 = 250. K
T2 V1
T1
=
(250.K) (1.35 L)
(373K)
V2 T2
T2
= 0.905 L
Group Work
A balloon is inflated to 665 mL volume at 27°C.
It is immersed in a dry-ice bath at −79°C. What is
its volume, assuming the pressure remains
constant?
V2
V1
=
T2
T1
Remember to convert
to Kelvin (K)
K = oC + 273.15
Gay-Lussac’s Law
Gay-Lussac’s observed that as the temperature
increases, the pressure increases and vice versa.
Joseph Gay-Lussac (1778–1850)
• The direct relationship between temperature
and pressure is known as Gay-Lussac’s Law.
Gay-Lussac’s Law
•  When the temperature decreases, the pressure
decreases.
Pressure Gage
low
high
Gay-Lussac’s Law
When the temperature increases, the pressure
increases.
Pressure Gage
low
high
Gay-Lussac’s Law
•  Gay-Lussac also noticed that ratio of pressure
to temperature of a gas is always the same.
P = Constant
T
Gay-Lussac’s Law
Initial
pressure
Initial
temperature
P1
T1
=
P2
T2
Final
pressure
Final
temperature
•  Remember that when using Gay-Lussac’s
Law, that the volume is never changing.
–  Only the temperature and pressure change.
•  Temperature must be Kelvin (K).
Example
The initial pressure of the gas in the container below is .870 torr
and the initial temperature is. 300.K.
The temperature is raised from 300. K to 1250. K (at constant
volume). What is the final pressure?
Pressure Gage
low
high
Solution
P1
T1
=
P2
T2 P1
T2
T1
=
P2 T2
T2
P1 = 0.870 torr P2 = ?
T1 = 300. K
P2 =
T2 P1
T1
=
T2 = 1250. K
(1250. K)(0.870 tor)
(300.K)
= 3.63 torr
Group Work
An aerosol can containing gas at 25 atm and 22°C
is heated to 55°C. Calculate the pressure in the
heated can.
P1
T1
=
P2
T2
Remember to convert
to Kelvin (K)
K = oC + 273.15
The Combined Gas Law
•  Boyles’s, Charles’s, and Gay-Lussac’s Laws can be
combined mathematically.
• The relationship between temperature, volume, and pressure
is known as the Combined Gas Law.
The Combined Gas Law
P1V1 P2V2
=
T1
T2
Example
At an ocean depth of 33 ft, where the pressure is 2.0
atm and the temperature is 285K, a scuba diver
releases a bubble of air with a volume of 6.0 mL.
What is the volume of the air bubble when it reaches
the surface, where the pressure is 1.0 atm and the
temperature is 298 K ?
Solution
P1 V1 P2 V2
=
T2
T1
P1 = 2.0 atm
P2 = 1.0 atm
V1 = 6.0 mL
V2 = ?
T1 = 285 K
T2 = 298 K
T2P1V1 (298 K)(2.0 atm)(6.0 mL)
V2 =
=
=
13
mL
P2T1
(1.0 atm) (285K)
Avogadro’s Law
Avogadro’s observed that the volume of a gas is
directly proportional to the number of gas molecules.
Amedeo Avogadro (1776–1856)
• The direct relationship between moles of gas
molecules and volume is known as Avogadro’s Law.
Avogadro’s Law
When the number of moles of gas decreases,
the volume decreases.
Pressure Gage
low
high
Avogadro’s Law
When the number of moles of gas
increases, the volume increases.
Pressure Gage
low
high
Avogadro’s Law
•  Avogadro noticed that ratio of volume to
the number of moles of a gas is always
the same.
V = Constant
n
Avogadro’s Law
Initial
volume
V1
n
1
Initial # moles
=
V2
n2
Final
volume
Final # moles
•  Remember that when using
Avogadro’s Law, that the pressure
and temperature are never changing.
– Only the number of particles and
volume change.
Example
The initial volume of the 3.5 moles of gas in a
container is 1.5 L.
Amadeo adds 2.0 moles of gas. (at
constant temperature and pressure). What
is the final volume?
Solution
V1
n1
=
V2
n2 V1
n1
n2
=
V1 = 1.5 L
V2 = ?
n1 = 3.5 mol
n2 = 5.5 mol
n2 V1
(5.5 mol) (1.5 L)
V2 =
n1
=
(3.5 mol)
V2 n2
n2
= 2.4 L
Group Work
A balloon has a volume of 2.4 L and contains
0.12 moles of air. A child blows more air into the
balloon until it has a final volume of 3.5 L. How
many moles of gas are in the balloon?
Gas Law Summary
V2
V1
n1 = n2
Combined Gas Law
P1 V1
T1
P2 V2
=
T2
The Ideal Gas Law
No gas perfectly obeys all four of these laws
under all conditions.
These assumptions work well for most
gases and most conditions.
One way to model a gas’s behavior is to
assume that the gas is an ideal gas that
perfectly follows these laws.
The Ideal Gas Law
If we combine all these equations,
we get the Ideal Gas Law.
P V = Cb
x
V
= Ca
n
PV
= R
nT
V
= Cc
T
Gas Constant
P
= Cg
T
The Ideal Gas Law
The gas constant (R) is a mathematical
combination of all the individual gas law constants
(Cb, Cc, Cg, Ca)
PV
= R
nT
The Ideal Gas Law is more commonly written
as:
PV = nRT
The Ideal Gas Law
The previous gas laws we studied involved
a change in either P, V, T, or n.
P1
T1
=
P2
T2
V1 V2
= n
n1
2
P1 V1 = P2 V2
V1
T1
=
V2
T2
The Ideal Gas Law
• The ideal gas law is used for any gas
system, any time.
• No changes are involved in the equation
PV = nRT
The Ideal Gas Law
The value of R is: .0821 L.atm
K.mol
PV = nRT
• When using this equation you must have
the following units:
• Pressure = atm
• Volume = liters
• Temperature = K
The Ideal Gas Law
There are 4 variables in this equation:
PV = nRT
Pressure Volume
# Moles
Temperature
In problems, we will always be given 3 of
the 4 variables, then solve for the unknown
variable.
Example: The Ideal Gas Law
How many moles of gas are contained in 11.2 liters at
1.00 atm and 0.0°C?
PV = n RT
atm
P = 1.00
_____
???
n = _____
L
V = 11.2
_____
T = 273.2
_____K
Partial Pressure
•  Dalton’s law of partial pressure states
that the total pressure of a mixture of
gases is the sum of the partial pressures of
its components.
– The partial pressure of a gas in a mixture is
the pressure that the gas would exert if alone.
Partial pressure
of gas B
Total
pressure
P T = PA + PB + PC
Partial pressure
of gas A
Partial pressure
of gas C
When two gases are present, the total pressure
is the sum of the partial pressures of the
gases.
Partial Pressures
The partial pressure of each gas in a mixture
can be calculated using the Ideal Gas Law
for gases A and B in a mixture
nA x R x T
nB x R x T
PA =
PB =
V
V
the temperature and volume of everything
in the mixture are the same
ntotal = nA + nB
ntotal x R x T
Ptotal = PA + PB =
V
Let’s Try It!
A 1.00L flask contains 5.00 x 10-2 mol of neon and 5.00 x
10-3 mol of argon. At 30.0 °C, what is the partial pressure of
each gas in atmospheres and what is the total pressure?
Liquids
Viscosity is the resistance to flow.
-
It is related to the strength of the noncovalent interactions
between the molecules that make up the liquid - the
stronger the attractions, the thicker the liquid.
-
Temperature has an effect on viscosity.
-
As temperature rises, the increase in the kinetic energy
of the molecules in the liquid helps the molecules pull
away from one another - higher temperature produces
lower viscosity.
Glycerol is able to form more hydrogen
bonds than 2-propanol. That is why glycerol
is thicker (more viscous) than 2-propanol.
Density of a liquid (or any other substance)
is the amount of mass contained in a
given volume.
m
d=
V
Density is the relationship between the
mass of a substance and its volume.
Vapor Pressure
Due to collisions that take place between particles
(atoms or molecules) that make up a liquid,
particles at the surface are continually evaporating being “bounced” off into the gas phase. At the
same time gas phase molecules are being trapped
and converted to liquid.
Vapor Pressure
•  The boiling point of a liquid is the temperature
at which the vapor pressure of the liquid equals
the atmospheric pressure.
•  Liquids boil when their vapor pressure equals
the pressure of the air above them.
Solids
•  The atoms, ions, or molecules that make
up a solid are held close to one another
and have a limited ability to move around.
•  Solids can be classified based on whether
or not the arrangement of these particles is
ordered (in crystalline solids) or not (in
amorphous solids).
Crystalline Solids
•  Ionic
–  consist of oppositely charged ions held to one another
by ionic bonds
•  Molecular
–  consist of an ordered arrangement of molecules
attracted to one another by noncovalent interactions
•  Covalent Networks
–  atoms are held to one another by an arrangement of
covalent bonds that extends through the solids.
•  Metallic
–  An array of metal cations immersed in a cloud of
electrons that spans the entire crystalline structure.
Metallic
Bonding
Metallic Solids
•  The valence electrons
in metals are free to
move about the entire
crystal of metal nuclei
and core electrons.
•  We can imagine it like a “sea of electrons” that
are bonding the positive nuclei together
Properties of Metallic Substances
Metallic substances are solid at room temperature.
Except for :_______________________
Metallic substances are malleable (they can be
hammered or beaten in thin sheets)
Metallic substances are ductile (they can be drawn,
pulled, or extruded through a small opening to
produce wire.
Metallic substances are good conductors of
electricity.
A few substances exist as:
Covalent Networks
•  Atoms are covalently bonded as if it was a huge
molecule
•  Not too many covalent network substances exist
•  Examples:
Diamond (carbon)
Amorphous Solids- no regular repeating pattern
of ions or molecules.
Example: rubber
Solids: Summary
a)  Ionic
b)  Covalent
c)  Molecular
d)  Metallic
e)  Amorphous