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Molecules of a cold sample of liquid have
lower kinetic energy than those in a
warmer sample
If a particle near the surface has enough
kinetic energy, it can overcome the
attractive forces in a liquid and escape
into the gaseous state
Known as a phase change
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Viscosity:
The friction or resistance to motion that exists
between the molecules of a liquid when they
move past one another
 The stronger the attraction between the
molecules in a liquid, the greater the
resistance to flow
 Liquids with large intermolecular forces tend
to be highly viscous
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Surface Tension:
 The resistance of a liquid to an increase in its surface
area
 Which liquids will have high surface tensions and
why?
Because of decreased volume and increased molecular
interaction, liquids expand and contract only very
slightly with temperature change
Boiling Point:
 The point at which the liquid’s vapor pressure is
equal to the atmospheric pressure
 Rapidly converting from liquid to the vapor phase
within the liquid as well as at the surface
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The attraction of the surface of a liquid to the surface of
a solid
Liquids will rise very high in a narrow tube if a strong
attraction exists between the liquid molecules and the
molecules that make up the tubing
Pulls liquid up against force of gravity
Concave meniscus
Polar liquids exhibit capillary action
The spontaneous rising of a liquid in a narrow tube,
due to:
 Cohesive forces – the intermolecular forces among the
molecules of the liquid
 Adhesive forces – the forces between the liquid and its
container
 Which of these are stronger for water?
 Adhesive
 Evaporation (vaporization) – a process by
which the molecules of a liquid can escape the
liquid’s surface and form a gas
 Endothermic process
 Heat of vaporization (enthalpy of
vaporization) – energy required to vaporize
one mole of a liquid at a pressure of 1 atm
 Symbol: Δhvap

Condensation – process by which vapor
molecules re-form a liquid
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Eventually, enough
vapor molecules are
present so that the
rate of condensation
equals the rate of
evaporation
The system is said to
be at equilibrium
The pressure of the
vapor present at
equilibrium is called
vapor pressure
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What will happen if the temperature is
increased?
The number of liquid molecules will be
reduced
The number of gaseous molecules will be
increased
The rates of evaporation and condensation will
become equal again
This illustrates what is known as ;
Le Châtelier’s Principle
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Reversible reactions – conversion of reactants to products and vice
versa occur simultaneously
Change in conditions is imposed on a system at equilibrium, the
equilibrium will shift in the direction that tends to reduce that change
in conditions.
CHANGES IN CONCENTRATION:
 Substance is added  reaction consumes added substance
 Substance is removed reaction shifts to produce more
2NO2 (g)  N2O4 (g)
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Which direction will the above reaction shift if we add NO2?
Which direction will the above reaction shift if we remove N2O4?
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CHANGES IN PRESSURE:
Increase: shift in direction that produces fewer molecules (moles)
of gas.
Decrease: shifts in direction that produces more molecules of gas.
In the reaction below, if we increase the pressure which
direction will the reaction shift?
NH4Cl (s)  NH3 (g) + HCl (g)
CHANGES IN TEMPERATURE:
EXOTHERMIC: Reaction gives off heat (product)
ENDOTHERMIC: Reaction absorbs heat (reactant)
Consider heat as a component of the reaction.
H2 (g) + I2 (g)  2 HI (g) + Heat
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If we raise the temperature (add heat) which way will the
reaction shift?
If we want the reaction to go to the right do we add or remove
heat?
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During WWII, Allied forces blocked the
Germans from acquiring sodium nitrates used
for explosives from mines in Chile in hopes of
shortening the war.
Fritz Haber used Le Chatelier’s Principle to
come up with a new process of making
ammonia.
 N2(g) + 3H2(g) ―› 2NH3(g) + Heat
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Concentration: Remove the products (NH3)
Pressure: Increased the pressure
Temperature: Kind of complicated; reducing
heat meant less pressure, but increasing heat
would shift toward reactants therefore kept
temperature moderate (500C)
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1.
2.
3.
Can be classified into very broad categories:
Crystalline solids – highly regular arrangement
of components
Amorphous solids – have considerable
disorder in their structure
Polycrystalline solid – an aggregate of a large
number of small crystals in which the structure
is regular but the crystals are arranged in
random fashion
Crystalline Solid
Amorphous Solid
Polycrystalline Solid
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Lattice structure – a 3D system of points
designating the positions of the components
Unit cell – the smallest portion of a crystal
lattice that is repeated throughout the crystal
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Atomic solid containing strong directional covalent
bonds
Allotropes – forms of the same element that differ in
crystalline structure
• Differ in properties because of differences in
structure
• Example: Diamond is one allotrope of carbon in
which each carbon is covalently bonded to four
other carbon atoms in a tetrahedral direction.
• Graphite is another allotrope of carbon,
covalently bonded to form hexagonal sheets
• What is a buckyball?
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Melting point – the temperature at which atomic
or molecular vibrations of a solid become so great
that the particles break free from their fixed
positions and start to slide past each other in a
liquid state
Heating curve – a plot of temperature versus time
for a substance where energy is added at a
constant rate
Sublimation – when a solid goes directly to a
gaseous state without passing through the liquid
phase
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Heat of fusion(ΔHfus) – the amount of energy
required at the melting point temperature to
cause the change of phase to occur
Heat of vaporization (ΔHvap) – the amount of
heat needed to vaporize 1 gram of a liquid at
constant temperature and pressure
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Way to represent the phases of a substance as a
function of temperature and pressure
Triple point – the point at which all three states of
a substance are present
Critical temperature – the temperature above
which the vapor cannot be liquefied no matter
what pressure is applied
Critical pressure – pressure required to produce
liquefication at the critical temperature
Together, the critical temperature and critical
pressure define the critical point
http://www.teamonslaught.fsnet.co.uk/c
o2%20phase%20diagram.GIF
Both solids and liquids are condensed states of
matter
 Relatively weak forces which occur between
molecules
 Both dipole-dipole and London dispersion forces
are known as Van der Waals forces
*It is important to recognize that when a substance
such as water changes from solid to liquid to gas,
the molecules remain intact. The changes in state
are due to changes in the forces among the
molecules rather than within the molecules*
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Dipole-dipole Forces
•The
attractive force resulting
when polar molecules line up
so that the positive and
negative ends are close to each
other
•Try
to maximize the (+)----(-)
interactions
•In
the gas phase, these forces
are unimportant
•Weaker than ionic or covalent
bonds
London Dispersion Forces
•Forces
which exist among
all covalent molecules but
is the only force for
nonpolar molecules.
•Weak
attractive forces
between molecules
resulting from the small,
instantaneous dipoles that
occur because of the
varying positions of the
electron during their
motion about nuclei.
Hydrogen Bonding
•Unusually
strong dipoledipole attractions that occur
among polar molecules in
which hydrogen is bonded
to a highly electronegative
atom such as O-H, N-H, F-H
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Nonpolar tetrahedral
hydrides show a steady
increase in boiling point
Polar tetrahedral hydrides,
the lightest member has an
unexpectedly high boiling
point
This is due to hydrogen
bonding that exist among
the smallest molecule with
the most polar X—H bond.
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Bonds are WAY stronger than forces
Ionic>Ionic/Dipole>H>Dipole/Dipole>
Dipole/Induced> Induced/Induced
The stronger the intermolecular forces the
higher the melting and boiling points
Solids have highest intermolecular forces
followed by liquids and gases.
1st
Law of Thermodynamics (AKA Law of
conservation of Energy): Energy cannot be created
or destroyed. It remains constant in the universe.
E = q + w
E = change in system’s internal energy
q = heat
w = work
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Heat: energy that flows into or out of a system
because of difference in temperature between the
system and its surrounding
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Thermodynamics: Science of the relationships
between heat and other forms of energy
Thermochemistry: Study of heat absorbed or
given off by chemical reactions.
Energy: Ability to do work (measured in
Joules)
Work = Force (N) x Distance (m)
Types of energy:
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Kinetic =1/2 mv2
Potential = mgh
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Heat of Reaction (q)
 Before any reaction the system and its
surrounding are at the same temperature.
 When the reaction starts, the temperature
changes.
The value of q needed to return the system to the
given temperature at the completion of the reaction
is known as the heat of reaction.
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Exothermic Reaction (-q): Heat is given off.
Products contain less energy than reactants.
Endothermic Reaction(+q): Heat is absorbed.
Reactants have less energy than products.
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Measured in joules and calories
Heat capacity – the amount of heat needed to
raise the temperature of an object exactly 1˚ C
q=C∆t
Depends on mass and chemical make-up
Specific heat (s)– the amount of heat it takes to
raise the temperature of 1g of the substance 1˚C
q = s x m x ∆t
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Can be measured using a calorimeter
The heat released by the system is equal to the
heat absorbed by the surrounding and vice
versa
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Enthalpy (H): Extensive (meaning it depends
on the amount of substance) property of a
substance that can be used to obtain the heat
absorbed or evolved in a chemical reaction.
All chemical reactions absorb or give off heat.
This change in energy is known as the change
in enthalpy (heat content) of a system.
ΔH = H (products) – H (reactants)
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The entropy (S) of the universe increases
for any spontaneous process
ΔS universe = ΔS system + ΔS surroundings
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Entropy: A measure of the degree of
disorder
Reactions are driven by the need for a
greater degree of disorder and the drive
towards the lowest heat content
Reactions with negative ΔH’s are
exothermic and those with positive ΔS’s
are proceeding to greater disorder
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When a gas is formed from a solid
CaCO3(s)  CaO(s) + CO2(g)
When a gas is evolved from a solution
Zn(s) + 2H+  H2(g) + Zn2+(aq)
When the number of moles of gaseous product
exceeds the moles of gaseous reactant
2C2H6(g) + 7O2  4CO2(g) + 6H2O(g)
When crystals dissolve in water
NaCl(s)  Na+(aq) + Cl-(aq)
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ΔH for an endothermic reaction is positive
ΔH for an exothermic reaction is negative
Changes in enthalpy are independent of
the path taken to change a system from
the initial to final state
Heat absorbed or given off varies with the
temperature
Standard enthalpies of formation are given
at 25°C and 1 atm pressure (ΔH0)
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Standard enthalpy of formation – the
change in enthalpy that accompanies
the formation of 1 mole of a
compound from its elements with all
substances in their standard states at
25°C
These values are known
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Example:
How much heat is liberated when 10.0
grams of CH4 (g) reacts with excess O2(g)?
CH4(g) + 2O2(g) → CO2 (g) + 2 H2O(l): ∆H = -890.3 kJ
Convert grams CH4 → moles of CH4 → kilojoules
of heat
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Free energy – energy available to do
work
The Gibb’s Free Energy Equation:
ΔG = ΔH –TΔS
The sign of ΔG can be used to predict
the spontaneity of a reaction
ΔH
ΔS
ΔG
Will it
Comment
happen
Exothermic
(-)
+
Always
negative
Yes
Exothermic
(-)
-
At lower
Probably
temperatures
At low
temperature
Endothermic
(+)
+
At higher
Probably
temperatures
At high
temperatures
Endothermic
(+)
-
Never
No
exceptions
No
No
exceptions
If a series of reactions are added together,
the enthalpy for the total reaction is the
sum of the enthalpy changes for the
individual steps
What is the ∆Hf of the following reaction?
2C(graphite) + O2(g)→ 2CO(g)

2C(graphite) + 2 O2(g)→ 2 CO2(g)
2 CO2(g)→ 2CO (g) + O2(g)
2C(graphite) + 2 O2(g)→ 2 CO2(g)
2(0) + 2(0)
2(-393.5 kJ)
(-787.0kJ) – 0 = -787.0 kJ
2 CO2(g)→ 2CO (g) + O2(g)
2 (-393.5kJ)
2 (-110.5) + 0
(-221.0kJ) – (-787.0kJ) = 566.0 kJ
∆Hf = -787.0 kJ + 566.0 kJ = -221.0 kJ