Liquids and Solids
Ch 10.2,10.3 & 10.4
Pg. 353 # 4, 5, 7, 8, 10-14,17, 20-22, 27,
28, 33
Liquids exist in the smallest temperature range, so
liquids are the least common state of matter . . .
 Kinetic Theory Description of the Liquid State

According to the kinetic theory, motion of liquid
particles can be described as . . .
Properties of Liquids and the Particles Model
– define each property Properties of Fluids
 Definite Volume
 Ability to Diffuse
 Fluidity
 Surface Tension
 Relative High Density
 Tendency to Evaporate
and Boil
 Incompressible
 Tendency to Solidify
 Dissolving Ability
13.1 Questions
 Why are liquids more dense than gases?
 Molecules are closer together so more molecules
in a given area
 Why are liquids harder to compress than gases?
 Same as above – molecules are closer
 Why do liquids diffuse slower than gases?
 Particles are not moving as fast as gases
 Can a liquid boil without increasing the
temperature? How?
 Yes – lower the atmospheric pressure
10.3 Solids
 “Solid as a rock, “ is the description of solid –
something that is hard, unyielding, with a definite
shape and volume. Many things other than rocks are
solids. In fact, solids are more common than liquids.
This diagram shows the particles of a gas, liquid and solid.
Kinetic-Theory Description of the Solid State
According to the kinetic theory, the motion of solid particles
can be described as….
 Lower kinetic energy, less motion, more packed
particles, and higher intermolecular forces (IMF)
 Properties of Solids and the Particle Model –
define each property: Properties of Solids






Definite shape and volume
Non-fluid
Definite melting point
High Density
Incompressible
Slow Diffusion
Crystalline Solids
 Classification of crystals by arrangement
and shape
 Crystal Lattice (define) - The total 3-D array of points
that describe the arrangement of the particles – a
collection of unit cells.
 The smallest portion of the crystal lattice that reveals
the 3-D pattern of the entire lattices is the unit cell.
Binding Forces in Crystals
Simple
Body-centered
(ex. Li, K, Cr)
Types of Crystals
Face-centered (ex.
Cu, Ag, Au)
Hexogonal (like oranges
in
a grocery store); (ex. Zn)
Binding forces in crystals
Binding Force
Lattice consists of
Formed When /
Binding Force
(+) and (-) ion as
arranged in regular
patterns
Group 1/2 metals
combine with Group
7/8 nonmetals
Covalent network
crystals
Lattice sites contain
single atoms
Atoms bond to
neighbors,
extending through a
network, large
chains form
Metallic crystals
(+) ions of the metal
surrounded by a
cloud of valence
electrons
Each e- and the (+)
metallic ions attract
electrostatically
Covalently bonded
molecules held
through IMF
For nonpolar
molecules, London
Forces; For polar
molecules, DipoleDipole.
Ionic crystals
Covalent molecular
crystals
Amorphous Solids
 Rubber, glass, plastics and synthetic fibers
are called amorphous solids.
“Amorphous,” comes from the Greek
for “without a shape.”
 Unlike crystals, amorphous solids do not
have a regular, natural shape, but instead
take on whatever shape imposed on them.
 Particle arrangement is not uniform; they are
arranged randomly, like particles of a liquid.
 Examples of amorphous solids – glass used
in fiberoptics (optical fibers transmit telephone
conversations by means of light waves.
Amorphous solids are prepared by
rapid cooling of thin film materials.
 Molecular examples
Crystalline vs. Amorphous
10.4 Changes of State
Possible Changes of State
Change of State
Name
Example
Solid -> Liquid
melting
ice -> water
Solid -> Gas
sublimation
dry ice -> CO2 gas
Liquid -> Solid
freezing
water to ice
Liquid -> Gas
vaporization
Br(l) -> Br(g)
Gas -> Liquid
condensation
water vapor -> water
Equilibrium
 What does equilibrium mean?
 It is a dynamic condition in which two
opposing changes occur at equal rates in a
closed system.
 What is a closed system?
 H2O in an open beaker
 H2O in a closed beaker
 When a liquid changes to a vapor, as in
evaporation, it absorbs heat energy and can be
shown as:


Open system evaporation – liquid + heat  vapor
Closed system evaporation – liquid + heat  vapor
 When a vapor condenses, as in condensation, it
gives off heat energy and can be shown as:

And condensation –
vapor  liquid + heat
 The liquid vapor equilibrium can be rewritten as:
 liquid + heat ↔ vapor
 “The double yields sign represents a reaction at
equilibrium”
Le Chatelier’s Principle
 What is it? LeChatelier

When a system at equilibrium is disturbed by
the application of stress, the system reacts to
minimize the stress.
 Is temperature an example of stress?

Yes.
 What happens when you increase the
temperature of a system? Equ. shift from heat

↓ liquid + increased heat ---> ↑ vapor
Le Chatelier’s Principle
 What happens when you decrease the
temperature of a system?

↓ vapor
---> ↑ liquid + decreased heat
 What factor is controlling the decrease and
increase of vapor and liquid?

the temperature (heat)
Equilibrium Vapor Pressure of a Liquid
 What is it?

At equilibrium, the molecules of a vapor exert
a specific pressure on its corresponding
liquid.
When equilibrium vapor pressure of water is
graphed, (draw figure 14 below):
 The strength of attractive forces is
independent of temperature. Higher
temperatures with resultant higher kinetic
energies make these forces less effective.
 Liquid water can exist in equilibrium with
water vapor only up to a temperature of
374.1ºC. Later you will learn that neither
liquid water nor water vapor can exist at
temperatures above 374.1ºC.
Water
Alcohol
Cooking Oil
At
80° C
355 torr
760 torr
10 torr
At
50° C
92 torr
400 torr
4 torr
At
20° C
20 torr
90 torr
1 torr
 What is equilibrium called when liquid
molecules enter into the gaseous state?

Vaporization
 Where does this occur?

On the surface of the liquid = evaporation,
throughout liquid = boiling
 Equilibrium vapor pressure depends on:


a) temperature and pressure
b) boiling point of a liquid (the type of liquid)
 If a liquid has high intermolecular forces,
then what happens to that liquid’s vapor
pressure? Why?

vapor pressure ↓
high IMFs =
increase hold on the molecules
Boiling. Freezing. Melting
 What is boiling?
 The conversion of a liquid to a vapor, within
the liquid as well as its surface when the
equilibrium vapor pressure of the liquid is
equal to the atmospheric pressure.
 What is the boiling point?
 The temperature at which the equilibrium
vapor pressure of the liquid is equal to the
atmospheric pressure (760 torr).
 Boiling happens throughout the
liquid…evaporation happens on the surface.
What is the molar heat of
vaporization?

The amount of heat energy required to
vaporize one mole of liquid at its boiling
point.
 How does a pressure cooker work?

It elevates pressure to raise boiling point and
shorten cooking time.
Freezing and melting
 What is the freezing?

The physical change of a liquid to a solid.
 What is melting?

The physical change of a solid to liquid.
 What is the molar heat of fusion?

The amount of heat energy required to melt
one mole of solid at its melting point.
solid + heat  Liquid
liquid
solid + heat
re-write the equation:
solid + heat ↔ liquid
heat of fusion
Are the freezing points and melting
points the same temperature?
 Yes
 at 0°C H2O with 6kJ is a liquid
 at 0°C H2O without 6kJ is a solid
Chapter 10 Calculations – not in book
 Molar heat of Vaporization
 The amount of heat energy required to
vaporize one mole of liquid at its boiling point.
 Joules are the standard unit to measure heat
energy.
 Molar heat of vaporization for water is 40.79
kJ/mole.
2.2 – Heat and Temperature – there
is a difference
• Heat transfers between objects – flows from hot
to cold - Law of Conservation of Energy
•
• Ex1:ice cube in a thermos of hot water - ice
melts, water cools - same amount of heat
• SI unit of heat - Joule (J)
calorie is also used
frequently
• Calorie - the amount of energy required to raise
the temperature of 1 g of water by 1 oC
• (Calories – capital letter – really means
kilocalories – used in food energy measurement)
Specific Heat Problems
• For water, Cp = 1.000 cal/g oC
or 4.184 J/g oC for
water
• Ex1: How many calories does it take to heat 20. g of
water from 10.0 to 40.0 oC? Also how many J?
• Ex2: How much heat is required to heat 75 g of Iron (Cp
= 0.444 J/gCo) from 15.5
to 57.0 oC?
Specific Heat Problems
 Ex3: What is the specific heat of an object if 250 calories
will heat 55 g of it from 25 to 100.0 oC?
 Ex4: - If a 100.0 g sample of silver (Cp = .237 J/g oC) at
80.0 Co loses 50. calories, what will its final temperature
be?
Not In Book
•
•
•
NIB: It also takes energy to melt or boil any substance. The amount
of energy required to melt or boil a substance can be expressed by
the following equations:
ΔH = nΔHfusion
ΔH = change in energy (J)
n = number of moles
ΔH = nΔHvaporization
ΔHfusion = the molar heat of fusion
(kJ/mol)
ΔHvaporization = the molar heat of
vaporization (kJ/mol)
•
ΔHfusion and ΔHfusion are constants and correspond to the amount of
energy it takes to freeze (fuse) or boil (vaporize) one mole of a
substance.
•
When doing heat calculations that involve both a change of state
and a change in temperature, make sure the answers for both
calculations are written in the same units before adding them
together!
Ex1: How much heat energy would be required to
vaporize 5.00 moles of H2O
 q = ΔHvap·(mol)
= 40.79 kJ/mol · 5.00 mol
= 204 kJ or 204,000 J
 Ex2: to vaporize 45.0g of H2O
 q = ΔHvap·(mol)
45.0g ·1mol
= 40.79 kJ/mol (2.50 mol)
1
18.0g
= 102 kJ or 102,000 J
 when....a liquid evaporates, it absorbs
energy. Energy is used to overcome
attractive forces. The energy doesn’t
increase the average energy of the particles,
so the temperature doesn’t change.
 when...a liquid evaporates, it takes energy
from its surroundings that’s why alcohol feels
cool to the skin.
 it’s also why we get cold when getting out of
the shower
Heat of vaporization - Hvap - energy needed to
vaporize a unit of substance (mass or moles)
 Formula: q = (H vap ) x ( unit )
unit = gram or mole
 Ex3 - How much heat does it take to
vaporize 50.0 g of water at 100.0 °C
50.0g · 1mol = 2.78 mol
1
18.0g

q = (H vap ) x ( unit )
= 40.79 (2.78)
= 113 kJ
Molar Heat of Fusion
 The amount of heat energy required to melt
one mole of a solid at its melting point.
 The molar heat of fusion of water is 6.008
kJ/mole.
 Ex1: How much energy would be required to
melt 12.75 moles of ice?

q = ΔHfus·(mol)
= 6.008 kJ/mol ·(12.75 mol)
= 76.60 kJ
Ex2: to melt 6.48 x 1020 kg of ice?
 6.48x1020kg · 1000 g = 6.48x1023 g
1
1kg
6.48x1023 g · 1mol = 3.6x1022 mol
1
18.0g
6.008kJ/mol(3.6x1022 mol) = 2.16x1023 kJ
Heat of Fusion - Hfus = heat of fusion - heat
required to change a unit of substance from solid
to liquid
 same formula: q = (Hfus) x (unit)
unit = g or mole
 Ex3: - How much ice can be melted by 2.9 x 104 J?
 2.9x104J · 1kJ = 29kJ
1
1000J
q = (Hfus) x (mol)
29 kJ = 6.008 kJ/mol x (mol)
= 4.8 mol ice
Temperature and Phase Changes