Ch 13: Liquids and Solids
General properties of matter:
Solids:
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particles close together and highly ordered
very little motion of particles
cannot be compressed
maintains its shape
Liquids: -- particles stay close together but not as ordered
-- a lot of individual motion of particles but still in contact
-- takes shape of container
Gases:
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particles are far apart
Particles behave independently
can be compressed
fills volume of container
13.1 Water and Phase Changes
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Water is the most common liquid on earth
97% of all water found in the oceans
Colorless and tasteless
At 1.00 atm has freezing point (fp) = 0oC and boiling point (bp) = 100 oC
What happens when liquid water is heated?
-- As with gases, as temp is increased the motion of the molecules increases.
-- As the water boils the temperature stays at 100 oC until all the water is changed
to vapor. Only then does the addition of more energy increase the temperature of the
water vapor.
-- This is depicted in the following heating-cooling curve for water.
-- Also, see Fig 13.2 in text
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Temperature oC
Energy added
Phase changes occur at BC (solid to liquid) and DE (liquid to gas).
As mentioned, water temperature does not change during a phase change.
-- As ice melts the temperature of the ice-liquid water mixture stays at 0 oC until
all the ice has melted (see plateau at 0 oC).
-- As water boils the temperature of the liquid water-water vapor mixture stays at
100 oC until all the water turns to vapor (see plateau at 100 oC).
Generally, substances are denser as solids than in their liquid forms. This makes sense
because as liquids their particles (atoms or molecules) have more energy and are moving
around more. The space between particles is greater than when in solid form.
Water is an exception. As ice its molecules are locked in position and held slightly further
apart than in the liquid phase when there is more motion of the water molecules yet they
are closer together.
The density of liquid water is 1.00 g/1.00 mL = 1.00 g/ml
The density of water ice is 1.00 g/1.09 mL = 0.917 g/mL
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13.2 Energy Requirements for the Change of State
Remember, changes of state (solid to liquid, liquid to gas) are physical, not chemical,
changes.
No chemical bonds are broken. All physical states for a molecular substance consist of
the same molecules.
Intramolecular (within the molecule) forces are the bonding forces that hold the atoms
of molecules together. These bonding forces can be covalent or ionic (see Ch 11).
Intermolecular (between molecules) forces occur between molecules that cause them to
aggregate to form a solid or a liquid. (See fig 13.3, 13.4)
Intermolecular
forces
Intramolecular forces (bonds)
It takes energy to melt ice and vaporize liquid water because intermolecular forces must
be overcome. Once water has been vaporized there is little or no intermolecular force
between water molecules.
This applies to phase changes for all substances, not just water.
The energy required to melt one mole of a substance is called the molar heat of fusion.
For ice, the molar heat of fusion is 6.02 kJ/mole.
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The energy required to vaporize one mole of a liquid is called the molar heat of
vaporization. For water, the molar heat of vaporization is 40.6 kJ/mole.
Notice, in Fig 13.2 in your text or in the heating/cooling for water above, that the plateau
for vaporization is much longer than for melting. In liquid water, the molecules are still
fairly close together, so most of the intermolecular forces are still present. However, to
separate the molecules to form a gas, virtually all of the intermolecular forces must be
overcome, and this requires much more energy.
Example 1: How much energy is needed to melt 2.00 moles of ice at 0 oC?
Example 2: How much energy is needed to melt 75.0 grams of ice at 0 oC?
Example 3: How much energy (in kJ) is needed to heat 25.0 g of liquid water at
50.0 oC to water vapor at 125 oC?
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13.3 Intermolecular Forces
How do intermolecular forces arise? There are several mechanisms
1. Dipole-dipole attraction.
-- When molecules with dipole moments (i.e., centers of positive and negative
charge) are put together, they orient themselves so that the positive end of one molecule
is lined up with the negative end of another. See Fig 13.5 in text.
-- Dipoles find the best compromise between attraction and repulsion.
-- A dashed line is used to indicate the dipole-dipole attraction
-- Examples:
-- Dipole-dipole forces are typically only about 1% as strong as covalent or ionic
bonds, and become weaker as the distance between dipoles increases. In the gas phase
these forces are, therefore, relatively unimportant.
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2. Hydrogen bonding.
-- These are particularly strong dipole-dipole forces between molecules in which H is
bonded to a very electronegative atom, such as O, N or F. (See above for HCl and H2O.)
-- Hydrogen bonding has important effects on physical properties, such as boiling
points and vapor pressure.
Bond
O-H
N-H
S-H
∆EN
1.4
0.9
0.4
Molecule
water
ammonia
H2 S
BP
100 oC
-33.3 oC
-60 oC
3. London Dispersion forces
-- Even molecules without dipole moments, and even individual atoms, must exert
forces on each other. This is known because even substances such as noble gases can
exist in the liquid and solid phases at very low temperatures (see Table 13.2). Therefore,
there must be forces that hold them together. These forces are known as London
dispersion forces.
-- These forces occur because atoms can develop a temporary dipolar arrangement
of charge as electrons move around the nucleus.
-- This instantaneous dipole can induce a similar dipole in a neighboring atom or
nonpolar molecule
-- This interatomic or intermolecular attraction is weak and short-lived, but can serve
to hold these atoms or molecules closer together when temperatures are every low.
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