States of MatterSolid
Lowest energy/heatMolecules barely movingDefinite, uniform shapeExample: ice
States of MatterLiquid
Medium energy/heat Molecules slowly movingShape of containerExample: water
States of MatterGas
High energyMolecules moving rapidly and bouncingFills its containerExample: steam
© UC Regents http://ds9.ssl.berkeley.edu/LWS_GEMS/2/part2.htm
Phase Changes
Phase Changes
• Solid liquid
• Liquid gas
• Solid gas
• Gas liquid
• Liquid solid
• Gas solid
melting
boiling, evaporation
sublimation
condensation
freezing
deposition
Phase Changes
Condensation
Evaporation
Melting
Freezing
SublimationDeposition
GAS
LIQUIDSOLID
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 stateIce 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 CurvesEnergy is added to the system at a constant rate
Temperature increases at a constant rate UNLESS it is changing phase
PhasesThe phase of a substance depends on:
TEMPERATURE PRESSURE
Phase diagram
The phase diagram is graph of pressure vs. temperature that shows conditions under which the phases of a substance exist
If you know T and P, you can figure out the state!
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 WorksheetTry answering the questions about the mystery
substance!
Phase ChangesA heating curve shows a substance’s change in
temperature while adding heat energy
Heating Curve ws
Changes in Matter and EnergyMatter 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 heatEnergy 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 EnergyKinetic 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
Exothermicenergy is released by the substance into the surroundings less PE, more KE, so temperature risesEx: a match burning
Endothermic energy is absorbed by the substance from the surroundingsmore PE, less KE, so temperature dropsEx: water freezing
EnergyCalorie (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/goCWater 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/goCWhich 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 TQ = energy (heat) required (J) or (cal)
s = specific heat capacity (J/goC) or (cal/goC)
m = mass of the sample in grams
T = 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 = 10.1 J = 0.24 J/goC m x T (2.8 g x 15oC)
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