COMPARISON OF GASES, LIQUIDS AND SOLIDS
 Gases are compressible fluids. Their molecules are widely separated.
 Liquids are relatively incompressible fluids. Their molecules are more tightly
packed.
 Solids are nearly incompressible and rigid. Their molecules or ions are in
close contact and do not move.
 Vapors term customarily used for the gasesous state of a substance that
exists naturally as a solid or liquid at 25 C and 1 atmosphere
States of Matter
ENERGY REQUIREMENTS FOR PHASE CHANGES
Water (solid)
Water (liquid)
water (liquid)
water (gas)
= 6 kJ/mol
= 41 kJ/mol
What does this mean????
Takes more energy to convert water from a liquid to
a gas then to convert water from a solid to a liquid.
WHY??
Solid and liquid states of water are more similar
then the liquid and gas states
SECTION 13.1 WATER AND ITS PHASE CHANGES
FYI!!!
WAT E R I S U N I Q U E I N T H AT A S I T
C O O L S I T E X PA N D S
WHY???
E X P L A I N W H Y I C E F LOAT S ?
W H Y P I P E S B U R ST I N T H E W I N T E R ?
H OW P OT H O L E S FO R M ?
PHASE TRANSITIONS
Melting: change of a solid to a liquid.
Freezing: change a liquid to a solid.
H2O(s)  H2O(l)
H2O(l)  H2O(s)
Vaporization: change of a solid or liquid to H2O(l)  H2O(g)
H2O(s)  H2O(g)
a gas. Change of solid to vapor often
called sublimation.
Condensation: change of a gas to a liquid H O(g)  H O(l)
2
2
or solid. Change of a gas to a solid
H2O(g)  H2O(s)
often called deposition.
ENERGY OF HEAT AND PHASE CHANGE
Temperature does not change during the change from one phase to another.
notice: there is either a temperature change OR a phase change.
You cannot have BOTH at the same time
Increase in temperature = increase in kinetic energy (after all the definition of
temperature is average kinetic energy). There is no change in the potential
energy
Increase in heat = increase in potential energy. There is no change in the
kinetic energy because the temperature does not change!!!!
VAPORIZATION VS BOILING VS EVAPORATION
Boiling For water to boil, it must be
100 Celsuis. Boiling creates an
actual gas. The substance changes
phase.
Evaporation also involves liquids
become gaseous. However, the
body of liquid does not need to be
at the boiling temperature. It occurs
because the molecules of a liquid
are not tightly bound together, and
so some escape with time.
Vaporization is a blanket term
referring to both boiling and
evaporation. In the broadest sense,
it is liquid becoming gas.
BOILING POINT VS VAPOR PRESSURE
Boiling point the
temperature at which the
vapor pressure of a liquid is
equal to the pressure of the
external atmosphere.
Normal boiling point the
temperature at which the
vapor pressure of a liquid is
equal to atmospheric
pressure (1 atm).
Vapor Pressure the
pressure of the vapor over a
liquid at equilibrium in a
closed container
VAPOR PRESSURE
If a liquid is placed in a nonclosed container, some of the
molecules will escape the
surface and evaporate.
In a sealed container, some of
a liquid still evaporates but
cannot “leave” the container.
The molecules move back and
forth between liquid and gas
phases until they establish an
equilibrium and thus a
pressure in the vapor phase.
Vapor pressure: partial
pressure of the vapor over the
liquid measured at equilibrium
and at some temperature.
TEMPERATURE DEPENDENCE OF VAPOR PRESSURES
The vapor pressure above the
liquid varies exponentially with
changes in the temperature.
What does this graph indicate
about the relationship between
vapor pressure and
temperature??
There is a _______ relationship
between vapor pressure and
temperature. In other words as
vapor goes UP the temperature
goes _______ or as vapor
pressure goes DOWN the
temperature goes _______
PHASE DIAGRAMS
Graph of pressure-temperature
relationship; describes when 1,2,3 or
more phases are present and/or in
equilibrium with each other.
Lines indicate equilibrium state two
phases.
Triple point- Temp. and press. where
all three phases co-exist in
equilibrium.
Critical temp.- Temp. where substance
must always be gas, no matter what
pressure.
• Critical pressure- vapor pressure at critical temp.
• Critical point- point where system is at its critical pressure and
temp.
PHASE DIAGRAM
Phase Diagram of Water
TEMPERATURE, ENERGY AND HEAT
Temperature = average kinetic energy
Energy – the ability to do work
Heat – the transfer of energy
Energy and Reactions
Energy must be added to break bonds.
 Many forms of energy can be used to break bonds:
 heat
- electricity
 sound
- light
Forming bonds releases energy.
 Example: When gasoline burns, energy in the form of heat and light is
released as the products of the isooctane-oxygen reaction and other gasoline
reactions form.
Energy is conserved in chemical reactions.
 Chemical energy is the energy released when a chemical compound reacts
to produce new compounds.
 The total energy that exists before the reaction is equal to the total energy of
the products and their surroundings.
EXOTHERMIC VS ENDOTHERMIC REACTIONS
An exothermic reaction is a
chemical reaction in which
heat is released to the
surroundings.
An endothermic reaction is
a chemical reaction that
absorbs heat.
The graphs to the right
represent the changes in
chemical energy for an
exothermic reaction and an
endothermic reaction.
CATALYSTS
Catalysts provide alternative pathways for
a reaction, usually with a lower activation
energy. With this lower energy threshold,
more collisions will have enough energy to
result in a reaction. An enzyme is a large
organic molecule that folds into a unique
shape by forming intermolecular bonds
with itself. The enzyme’s shape allows it to
hold a substrate molecule in the proper
orientation to result in an effective
collision. The rate of a chemical reaction is
the change in the amount of reactants or
products in a specific period of time.
Increasing the probability or effectiveness
of the collisions between the particles
increases the rate of the reaction.
CALCULATING ENERGY CHANGES
Specific Heat Capacity: the
amount of energy required to
change the temperature of 1
gram of a substance 1 celsius
degree (in j/g ◦C)
Heat of vaporization: heat
(energy) needed for the
vaporization of a 1 mol of a
liquid.
H2O(l) H2O(g) DH =
40.6 kJ/mol
Heat of fusion: heat (energy)
needed for the melting of a I
mol of a solid substance.
H2O(s) H2O(l) DH =
6.02 kJ/mol
EQUATION
Q = sm∆t
Q= energy required (joules)
s = specific heat capacity
(given in table page 70 in
book)
M = mass of water (grams)
∆t = change in temperature
(Celcius)
MOLAR HEAT OF FUSION AND VAPORIZATION
Example: calculate the energy (in kJ) required to heat 25g of
water from 25C to 100C and change it to steam at 100C.
The specific heat capacity of water is 4.18J/g and the molar
heat of vaporization of water is 40.6kJ/mol
Step 1: Q = s
m
∆t
= 4.18J
x
25g
x
(100C – 25C)
g
= 7.8x103J
Convert to kJ
7.8kJ
Step 2: now use the molar heat of vaporazation to calculate heat
energy required to vaporize 25g of water at 100C. Heat of
vaporation is given in mols to we must convert 25g water to mols of
water
25g H2O x 1 mol H2O = 1.4 mol H2O
18 g H2O
Now calcualte energy need to vaporize the water
40.6 kJ x 1.4 mol H2O = 57kJ
mol
Step 3: TOTAL THE ENEGY NEEDED: 7.8kJ + 57 kJ = 65kJ
FYI
One calorie is the amount of heat (or
energy) needed to raise the
temperature of 1 gram of water by
1°C.
4.184 joules = one calorie.
1,000 calories = 1 Calorie
INTRA VS. INTER
MOLECULAR FORCES
Intramolecular Forces:
• The attractive forces between atoms and ions
within a molecule
• Ex: ionic bonds, covalent bonds
• STRONG
Intermolecular Forces
• The attractive forces between molecules
• Ex: London dispersion forces , dipole-dipole forces and hydrogen
bonding
• WEAK
I.e. much less energy to melt H2O (inter)
than for it to decompose into H2 and O2
(intra)
COMPARISON OF ENERGIES FOR INTERMOLECULAR FORCES
Interaction Forces
Approximate Energy within Bond
or Attraction between molecules
Intermolecular
London
1 – 10 kJ
Dipole-dipole
3 – 4 kJ
Hydrogen bonding
10– 40 kJ
Chemical bonding
Ionic
100 – 1000 kJ
Covalent
100 – 1000 kJ
LONDON DISPERSION FORCES
•Aka: van der waals forces
•Attractive forces between
ALL molecules
•temporary attractive force
that results when the
electrons in two adjacent
atoms occupy positions
that make the atoms form
temporary dipoles
•Exist with nobles gases
and nonpolar molecules,
H2, N2, I2
Intermolecular Forces
London Dispersion Forces
-
Induced Dipole – Induced Dipole
Weakest of all intermolecular forces.
It is possible for two adjacent nonpolar molecules to
affect each other.
The nucleus of one molecule (or atom) attracts the
electrons of the adjacent molecule (or atom).
This attraction causes the electron clouds become
distorted.
In that instant a polar molecule (dipole) is formed
(called an instantaneous dipole).
DIPOLE-DIPOLE
• Forces of attraction between oppositely
charge ends of polar molecules
• Line up + and - ends
HYDROGEN BONDING
• Strong/Special type of dipole-dipole force
• IT IS NOT A BOND.
• Between the positive H atom attached to an
N, O or F and the negative N, O, or F of
another molecule
RELATIVE STRENGTH
London Dispersion < Dipole-Dipole < H-bonds
DETERMINING INTERMOLECULAR FORCES
STEP ONE
Is the compound
polar or
nonpolar???
a.Nonpolar =
ONLY london
dispersion forces
b.Polar = go to
step 2
STEP T WO
Does the
compound contain
N-H, O-H, or F-H
Bonds???
a.NO = dipoledipole and london
dispersion forces
b.YES = hydrogen
bonding, dipoledipole and london
dispersion forces
EXAMPLES: WHAT TYPE(S) OF INTERMOLECULAR
FORCES EXIST BETWEEN EACH OF THE FOLLOWING ?
HBr???
HBr is a polar molecule: dipole-dipole
forces. There are also london dispersion
forces between HBr molecules.
CH4???
CH4 is nonpolar: london dispersion forces.
SO2???
SO2 is a polar molecule: dipole-dipole
forces. There are also london dispersion
forces between SO2 molecules.
DETERMINING RELATIVE BOILING POINTS
STEP ONE
ST E P T WO
Determine the types of
intermolecular forces
a. Look at the size of
the compounds
a.Different forces =
H > D> L
Stronger force =
higher boiling point
b. Same forces = go to
step 2
Larger molecules =
higher boiling point