Energy and
Phases
 Potential
Energy - stored energy
(stored in bonds, height)
 Kinetic
Energy - energy of motion,
associated with heat
Forms of Energy
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Light - light waves, electromagnetic radiation
Electrical
Chemical
Heat
Mechanical - moving parts, machines
Atomic/Nuclear - changes in mass of atom
Conservation of Energy
 Energy can be converted from one form
to another but never destroyed.
 The total amount of energy is always
constant.
Exothermic Reactions
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Energy (heat) exits
Releases energy (heat) when new products
are formed
Potential Energy of the reactants is greater
than the potential energy of the products
Surroundings feel warm because heat was
released
Heat is a product
Ex: AB  A + B + heat
Endothermic Reactions
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Endothermic – energy (heat) goes in
Reaction absorbs heat from the
environment
Potential energy of the reactants is less
than the potential energy of the products
Surroundings feel cold because heat was
absorbed from the surrounding
Heat is a reactant
Ex: A + B + heat  AB
Activation Energy
• The energy needed to break the bonds (to
get the reaction started)
• Energy to get over the “hill”, difference
between your starting point and the top of
the hill
• All reactions need activation energy
Heat of Reaction (H)
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Heat absorbed/released by a reaction
H = Hp - HR
If H is negative, Exothermic
If H is positive, Endothermic
Table I
1. Is the reaction, C(s) + O2(g) → CO2(g),
exothermic or endothermic?
2. What is the value of ΔH for the reaction,
CO2(g) → C(s) + O2(g)
Catalyst
• Lowers the activation energy
• Speeds up the reaction
Heat and Temperature
• Heat and temperature are not the same
– Heat – measure of the energy transferred from one
substance to another
– Temperature – measure of the average kinetic
energy of a substance’s particles
• The faster the particles move (more KE), the
higher the temperature
• Heat flows from high to lower until an equilibrium
is established (Hot  Cold)
• Calorimeter – measures the heat given off by a
reaction
Heat/Temp Examples
1.
If two systems at different temperatures have contact
with each other, heat will flow from the system at
a.
b.
c.
d.
20oC to a system at 303K
30oC to a system at 313K
40oC to a system at 293K
50oC to a system at 333K
2. Which is not a form of energy?
a. Light
c. Heat
b. Temperature
d. Motion
3. Which has the most kinetic energy?
a. 10.0 g of H2O at 70oC
c. 25.0 g of H2O at 60oC
b. 10.0 g of H2O at 5oC
d. 25.0 g of H2O at 10oC
Solids
• Definite shape
• Definite volume
• Very high attractive forces between
molecules
• Neighboring particles are very
close together
• Crystalline structure
• Kinetic Energy – solids have KE
– Particles are constantly vibrating
(around their positions in the
crystal)
– Positions do not change in
relation to the other particles in
the crystal
– At absolute zero (O K) all
movement stops (theoretically)
Melting Point
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
As energy is added to the solid, KE of the
particles increases until they have sufficient
energy to overcome the forces holding
them in the crystal, the substance begins to
melt
Temperature where the solid and liquid
phase exist in equilibrium
Heat of Fusion
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Heat needed to melt
Amount of heat needed to convert a unit mass
of a substance from a solid to liquid at a
constant temperature
q = mHf
q = heat (J)
m = mass (g)
Hf = heat of fusion (J/g)
Hf for water = 334 J/g
Hf Examples
1. How much energy is needed to change
75g of ice at OoC to water at the same
temperature?
2. 11,000J of heat are released as a sample
of water at OoC freezes. Calculate the
mass of the sample.
Melting/Freezing
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Melting is an endothermic process, the H2O
absorbs 334J/g
Freezing is the reverse process, so it is
exothermic, the H2O releases 334J/g
Sublimation
• Solid phase  gas phase (skip liquid)
• Endothermic
Example:
dry ice
CO2(s)  CO2(g)
Iodine
I2(s)  I2(g)
Naphthalene (moth balls)
Liquids
• Definite Volume
• Take the shape of the
container
• High attractive forces
between molecules
(but not as high as
those found in solids)
Evaporation (Vaporization, Boiling)
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Change from liquid to gas, endothermic
As temperature increase, rate of
evaporation increases
Increased temperature, more KE, easier
to overcome intermolecular forces and to
break them
Condensation - change from gas to
liquid, exothermic
Vapor Pressure
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Pressure of a gas on a liquid
In a closed system (sealed container) the
vapor evaporating from the liquid exerts
pressure on the liquid
Vapor Pressure increases as the
temperature of the liquid increases
It has specific values for each substance
at any given temperature
Table H
Boiling Point
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Temperature where vapor pressure =
atmospheric pressure
Water boils at 100oC at 101.3kPa (1atm)
– Pressures below 101.3kPa (high elevations),
water boils below 100oC
– Pressures greater than 101.3kPa (below sea
level), water boils above 100oC
Table H Examples
1. What is the vapor pressure of water at
105oC?
2. If the pressure is 30kPa, what
temperature will water boil at?
3. What pressure is needed for ethanol to
boil at 50oC?
4. Which liquid on Table H has the
strongest intermolecular forces?
Heat of Vaporization
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Heat to vaporize
Amount of heat needed to convert a unit
mass of a substance from liquid to vapor
at a constant temperature
Does not increase KE, so temperature
does not change
Heat of Vaporization
•
q = mHv
q = heat (J)
m = mass (g)
Hv = heat of vaporization
Hv for water = 2260 J/g
Examples:
1. How much heat must be supplied to evaporate
50.g of H2O at 100oC?
2. 12,750 J of heat are used to boil a sample of
water. Calculate the mass of the sample.
Boiling and Condensation
• Boiling is an endothermic process, the
H2O absorbs 2260J/g
• Condensation is the reverse process, so it
is exothermic, the H2O releases 2260J/g
Deposition
• Reverse of sublimation
• Direct change from the gas to the solid
phase, skip liquid
• Exothermic
• Example: Frost
Phase Diagrams / Heating and
Cooling Curves
• When a sample of matter is heated its
temperature usually increases.
 Increase in KE causes an increase in
temp.
• Sometimes matter can gain or lose heat
without changing temperature.
 What is going on at this point?
Points to remember:
• When heat is used to increase the speed of
particles, the temperature increases – at this
point the KE is changing
• When heat does not cause a change in
temperature, it is being used to change phases
(phase changes occur at the flat parts of the
graph) – at this point PE is changing
• Melting Point - point when the solid begins to
melt, both the solid and liquid phases are
present
• Boiling Point - point when the liquid begins to
boil, both the liquid and gas phases are present
Rate of Heating
• Total amount of heat absorbed = time x
rate of heating
q= t x rate
Example:
A solid is heated at a constant rate of 150
J/min for 3 minutes. How much heat is
absorbed?
Heating Curve
Cooling Curve
Change in Temperature Calcs
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q = mcT
q = heat (Joules)
m = mass
C = specific heat capacity
(for water; c = 4.18 J/goC)
• T = change in temperature
Examples
1. How many joules does it take to raise the
temperature of a 5.0g sample of water by
20.oC?
2. A 1000. gram mass of water in a
calorimeter has its temperature raised
5.0oC. How much heat energy was
transferred to the water?