Chapter 14
Liquids and Solids
Review
Solid - Has a definite (fixed) shape and
volume (cannot flow).
Liquid - Definite volume but takes the
shape of its container (flows).
Gas – Has neither fixed shape nor fixed
volume (flows).
H2O(s)  H2O(l) H = ~6kJ/mol
H2O(l)  H2O(g) H = ~41kJ/mol
This tells us that liquids more closely
resemble solids than gases.
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Figure 14.1: Representations of the
gas, liquid, and solid states.
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Figure 14.7: The heating/cooling curve for water heated or
cooled at a constant rate.
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Figure 14.2: Intermolecular forces exist between molecules.
Bonds(intramolecular forces exist within molecules.
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Dipole-Dipole Attractions
All polar molecules have a dipole moment:
one end of the molecule has a partial
positive charge and the opposite end has a
partial negative charge.
When molecules with dipole moments are
put together, they orient themselves such
that opposite ends attract each other (like
ends repel).
This is called a dipole-dipole attraction,
these are only about 1% as strong as a
covalent or ionic bond.
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(a) Interaction of
two polar molecules.
(b) Interaction of
many dipoles in a
liquid.
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Hydrogen Bonding
Hydrogen bonding is just a really strong
form of dipole-dipole interaction that only
occurs between molecules in which H is
bound to a highly electronegative atom
(N, O, or F only).
Keep in mind hydrogen bonding refers to
the IMF’s, not the actual N-H, O-H, or
F-H bond within the molecule.
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Figure 14.4: Hydrogen bonding among water
molecules.
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Hydrogen Bonding
Hydrogen bonding has a major effect on
certain physical properties such as melting
and boiling points.
Think of IMF’s as molecules being “sticky”
Velcro® is sticky; the bigger the pieces
are, the stickier they are.
The same is true with IMF’s; the bigger
the dipole moment, the stickier the
molecules will be to each other.
The stickier they are, the higher the
melting and boiling points.
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London Dispersion Forces (LDF’s)
All molecules, even non-polar, exert
forces on each other.
Any substance can be cooled to a point
where it turns into a liquid or a solid; this
means the atoms or molecules have forces
between them that are holding them
together.
These forces are not important in polar
molecules where other IMF’s are stronger,
but are very important in nonpolar
molecules.
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Section 14.2: Water and Its Phase
Changes
Water covers about 70% of the Earth’s
surface and about 97% of that is located
in the oceans.
Pure water is a colorless, tasteless
substance that at 1 atm freezes to form
a solid at 0 C and vaporizes completely to
form a gas at 100 C. It is a liquid
between 0 and 100 C
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Water and Its Phase Changes
When liquid water is heated, the molecules
speed up and the temperature increases
until it hits 100 C (water’s normal boiling
point at 1 atm).
Once it hits 100 C, bubbles start to form
and the water begins to boil and change
from a liquid to a gas.
Once the phase change begins, the
temperature remains constant (100 C)
until all liquid has been converted to gas.
When the phase change is complete, the
temperature will begin to rise again.
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Water and Its Phase Changes
When liquid water is cooled, the molecules
slow down and the temperature decreases
until it hits 0 C (water’s normal freezing
point at 1 atm).
Once it hits 0 C, ice crystals start to
form and the water begins to freeze and
change from a liquid to a solid.
Once the phase change begins, the
temperature remains constant (0 C) until
all liquid has been converted to solid.
When the phase change is complete, the
temperature will begin to fall again.
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Figure 14.7: The heating/cooling curve for water heated or
cooled at a constant rate.
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Water and Its Phase Changes
Most compounds are the most dense as
solids and least as gases.
Water is the exception to the rule: water
is more dense as a liquid than as a solid.
The density of water varies with
temperature even within the same phase.
Water is the most dense around 4 C
(1.00 g/mL).
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Section 14.3: Energy Requirements for
the Changes of State
Remember: phase changes are physical
(NOT chemical) changes; no bonds are
broken. Only IMF’s must be overcome.
Also remember the definitions of
exothermic and endothermic.
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State Change Energy Requirements
Changing from solid  liquid  gas is an
endothermic process (energy in). Particles
must be sped up to overcome IMF’s.
Changing from gas  liquid  solid is an
exothermic process (energy out). Particles
must be slowed down to allow IMF’s to
take hold.
The energy required to melt 1 mol of a
substance is called the molar heat of
fusion; for water ice it is 6.02 kJ/mol.
To freeze it we need -6.02 kJ/mol.
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State Change Energy Requirements
The energy required to change 1 mol of a
liquid substance to its vapor is called the
molar heat of vaporization; for liquid
water it is 40.6 kJ/mol.
To condense it back to a liquid we need
-40.6 kJ/mol.
Note that it takes about seven times as
much energy to vaporize a mole of water
than to melt it.
This is because to vaporize it must
overcome more in terms of IMF’s.
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State Change Energy Requirements
 There are two types of energy changes:
 Temperature changes: Q = SmT
(note there is no phase change).
 Phase changes: Hfus or Hvap (note
there is no temperature change).
 A typical problem might involve changing
25.0 g water at 25 C to steam, 2
steps:
1) Heat from 25 C to 100 C: (4.18 J/gC)
(25 g)(75 C) = 7.8 kJ
2) Vaporize at 100 C: (40.6 kJ/mol)(1.4
mol) = 56 kJ  56 +7.8 = 64 kJ total
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Figure 14.7: The heating/cooling curve for water heated or
cooled at a constant rate.
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Energy Practice Problems
1) Calculate the energy released when 15.5
g of ice freezes at 0 C. The molar
heat of fusion of ice is 6.02 kJ/mol.
2) Calculate the energy required to vaporize
35.0 g of water at 100 C. The molar
heat of vaporization of water is 40.6
kJ/mol.
3) Calculate the energy required to heat
22.5 g of liquid water at 0 C and change
it to steam at 100 C. S = 4.18 J/gC
and Hvap is listed above.
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