States of Matter
 Solid




Lowest energy/heat
Molecules barely moving
Definite, uniform shape
Example: ice
States of Matter
 Liquid




Medium energy/heat
Molecules slowly moving
Shape of container
Example: water
States of Matter
 Gas




High energy
Molecules moving rapidly and bouncing
Fills its container
Example: steam
http://ds9.ssl.berkeley.edu/LWS_GEMS/2/part2.htm
© UC Regents
Phase Changes
• Phase Change – when a
during phase change!
substance changes from
• Examples:
one state to another
• Change of energy in
“system.”
• Solid liquid
• Liquid  gas
• When energy is added, it • Solid  gas
is used to heat the
substance, then it is used • Gas  liquid
to change the substance
• Liquid  solid
into a different state.
• Temp doesn’t change
• Gas  solid
Phase Changes
• Solid liquid
melting
• Liquid  gas
boiling, evaporation
• Solid  gas
sublimation
• Gas  liquid
condensation
• Liquid  solid
freezing
• Gas  solid
deposition
Phase Changes
GAS
Condensation
Sublimation
Deposition
Evaporation
Melting
SOLID
Freezing
LIQUID
When a solid is heated…
 The temperature increases
UNTIL
 It reaches it’s melting point
THEN it turns to a liquid at that temperature
The temperature will not
change unless all matter is in
the same state
 Ice will warm, then melt, the liquid will warm
 It takes energy for those molecules to change state!
Heating/Cooling Curve
Heat of Vaporization
Heat of Fusion/Crystallization
Energy is added as Heat
Heating Curves
 Energy is added to the system at a constant rate
 Temperature increases at a constant rate UNLESS it is
changing phase
Phases
 The phase of a substance depends on:
 TEMPERATURE
 PRESSURE
Phase diagram
If you know
T and P, you
can figure out
the state!
The phase diagram is graph of pressure vs. temperature that shows
conditions under which the phases of a substance exist
http://itl.chem.ufl.edu/4411/2041/lec_f.html
 Triple Point: Unique temperature and
pressure where all three phases exist…AT
THE SAME TIME!!!
Phase Diagrams for Water and CO2
*Atmospheric Pressure is 1.0 atm
Phase Diagram Worksheet
 Try answering the questions about the mystery
substance!
Phase Changes
 A heating curve shows a substance’s change in
temperature while adding heat energy
Heating Curve ws
Changes in Matter and Energy
 Matter cannot be created or destroyed.
 But it can be changed and when it does, that is how
we get energy!
 Energy - capacity to do work or produce heat
 Energy is always involved in physical and chemical changes.
 Energy can take several forms: Heat, light, (sound, chemical,
electrical)
 Measured in calories, Calories (kcal), and joules
 Law of Conservation of Energy: energy can be absorbed or
released, but it cannot be created or destroyed through
ordinary chemical reactions.
 Energy can be transferred.
Kinetic and Potential Energy
Kinetic energy: is the energy of motion.
Potential Energy: energy of Position
 Stored energy (chemical bonds)
 Since energy is constant and cannot be created or
destroyed ….
Total Energy = KE + PE
 Temperature: kinetic energy of all particles within matter.
 There are times during phase changes when temperature
does not change, but stays constant while the energy works
to change the phase (ie: the heating curve of water)
Heating/Cooling Curve
Heat of Vaporization
Heat of Fusion
Energy is added as Heat
Exothermic
 energy is released by the substance into the surroundings
 less PE, more KE, so temperature rises
 Ex: a match burning
Endothermic
 energy is absorbed by the substance from the surroundings
 more PE, less KE, so temperature drops
 Ex: water freezing
Energy
Calorie (cal): the amount of energy (heat) required to raise
the temperature of one gram of water by one Celcius
degree
 Standard unit for energy is the joule (J)
 1 cal = 4.184 J
60.1 cal x 4.184 J = 251 J
1 cal
Specific Heat (s) : amount of energy required to change the
temperature of one gram of a substance 1 oC
 Varies from one substance to another
 Heat always travels from high concentration to low concentration!!
 Heat lost = Heat gained
 Water has a specific heat = 1 cal/goC or 4.184 J/goC
 Water has the second highest specific heat capacity of all known
substances. So it requires high amounts of heat energy to raise
water temperature.
 water also has a high energy/heat requirement for evaporation
 SIRON = 0.449 J/goC
 Which would heat up faster, 5.00 grams of iron or 5.00 grams of
water?
 Which would cool down faster, 5.00 grams of iron or 5.00 grams
of water?
 Which is a better thermal conductor?
 Which is a better insulator?
Q = s x m x DT
Q = energy (heat) required (J) or (cal)
s = specific heat capacity (J/goC) or (cal/goC)
m = mass of the sample in grams
DT = change in temperature in oC
 A 2.8 g sample of a pure metal requires 10.1 J of energy
to change its temperature from 21 oC to 36 oC. What is
the specific heat of the metal?
s= Q
m x DT
=
10.1 J = 0.24 J/goC
(2.8 g x 15oC)