1. States of Matter
Chapter 3 pg. 68 – 97
Chapter

2. States of Matter Solid Particles are tightly packed
Stuck to each other in a pattern
Definite Shape
Definite volume

3. States of Matter Liquid Particles are tightly packed
Able to slide past each other
Can flow
Variable Shape
Definite volume

4. States of Matter Gas Particles are spread out
Flying all over the place
Can flow
Variable volume
Variable Shape

5. Kinetic Theory
Kinetic energy – energy an object has due to its motion (Kinetic means motion)
Three main parts of the theory
All matter is made of tiny particles
These particles are in constant motion and the higher the temperature, the faster they move
At the same temperature, heavier particles move slower.

6. Motion in Gases
There are forces of attraction among particles in all matter.
The constant motion of particles in a gas allows a gas to fill a container of any shape or size.

7. Behavior of Liquids
A liquid takes the shape of its container because its particles in a liquid can flow to new locations.
The volume of a liquid is constant because the forces of attraction are strong enough to keep the particles together.

8. Behavior of Solids
Solids have definite volume and shape because the particles in a solid vibrate around fixed locations.
Strong forces of attraction
Vibration is repetitive back and forth Particles do not change places with neighboring atoms.

9. Gas Laws Pressure is F/A (N/m2) Force distributed over an area
Simplified to Pa (Pascals)
1 kPa = 1000 Pa
Air Pressure lb/in kPa = 1 atm = 760 mm Hg = 760 torrs
Collision Theory
Collisions between the particles of a gas and the walls of the container cause the pressure in a closed container of gas.
More collisions more pressure

10. Factors that affect Gas Pressure
Temperature (Kinetic energy of gas particles)  higher temperature, more KE, more speed leads to more collisions
Volume (Space between particles)  Lower volume, less space, more collisions can occurs
Number of Particles  More particles, more collisions can occur
Each assumes the other two are constant

11. Boyle’s Law
Volume of a gas is inversely proportional to its pressure if the temperature and number of particles are constant.
P1V1 = P2V2

12. Examples
1.00 L of a gas at standard pressure is compressed to 473 mL. What is the new pressure of the gas? (Standard P = 1 atm)
Synthetic diamonds can be manufactured at pressures of 6.00 x 104 atm. If we took 2.00 liters of gas at 1.00 atm and compressed it to a pressure of 6.00 x 104 atm, what would the volume of that gas be?

13. The volume of the lungs is measured by the volume of air inhaled or exhaled.  If the volume of the lungs is L during exhalation and the pressure is KPa, and the pressure during inhalation is KPa, what is the volume of the lungs during inhalation? 

14. Temperature Conversions
Pg. 20
For combined gas law temperature must in Kelvin (K)
No negative Kelvin temperatures
K = oC oC = K - 273

15. Charles’s Law
The volume of a gas is directly proportional to its temperature in Kelvins if the pressure and number of particles of the gas are constant.
V1 = V2
T1 T2
Absolute Zero  “No Volume of Gas”

16. Examples
A man heats a balloon in the oven. If the balloon initially has a volume of 0.4 liters and a temperature of 20 0C, what will the volume of the balloon be after he heats it to a temperature of 250 0C?
How hot will a 2.3 L balloon have to get to expand to a volume of 400 L? Assume that the initial temperature of the balloon is 25 0C.

17. Combined Gas Law
Relationship of Effect on Gas from Temperature, Volume, and Pressure
P1V1 = P2V2
T1 T2
Standard Temperature and Pressure – T - 0oC, P = 1 atm

18. Energy and Phase Changes
Energy - the ability change or move matter
Measured in Joules (J)
Energy is either absorbed or released during a phase change.
Endothermic Change –absorbs energy
Melting, Boiling, Sublimation
Exothermic Change –released energy
Freezing, Condensation, Deposition

19. Phase Changes
Reversible physical Change that occurs when a substance changes from one state to another.

21. Phases Changes Molecules and atoms don’t change during a phase change
The mass doesn’t change
The volume does change
During the phase change the temperature doesn’t change.
Only the arrangement and attraction to molecules changes

22. Heating Curve

23. Work and Heat
Heat flows spontaneously from hot objects to cold objects
Temperature is the measure of how hot or cold something is to a reference point
Based on Kinetic Energy (how much motion is occurring randomly by the particles inside of a substance)

24. Thermal Energy
Thermal Energy depends on mass, temperature, and phase (solid, liquid, gas) of an object.
More mass  More thermal Energy (More total kinetic energy)

25. Thermal Expansion
Thermal expansion occurs because particles of matter tend to move farther apart as temperature increases.
How thermometers work
Doors in summertime
Bimetallic strips (Thermostats)

26. Specific Heat
Specific Heat  amount of heat needed to raise temperature of one gram of a material by one degree Celsius
Lower Specific Heat  More its temperature increases when heat is added
Material (100 kPa)
Specific Heat (J/GoC)
Water
4.18
Plastic
1.84 – 2.09
Air
1.01
Iron
0.449

27. Specific Heat Q = m x c x ∆ T Mass  grams
C  specific heat (value of amount of heat needed to raise one gram by one degree Celsius)
∆ T  temperature change in Celsius

28. Measuring Heat Changes
Calorimeter – instrument used to measure changes in thermal energy
Heat flows from a hotter objects to a colder object until both reach the same temperature.