1. The Physical Behavior of Matter

2. What is Matter? Anything that takes up space & has mass
Substances – variety of matter that has the same composition and properties throughout.
Two types are elements (Na) and compounds (NaCl).
Mixtures – two or more substances mixed together (not united).
Two types are homogeneous
(salt water) and heterogeneous (sand and sugar).

3. Substances Elements Compounds
Cannot be decomposed (broken down)
Ex: hydrogen (H2), Oxygen (O2) and nitrogen (N2)
Compounds
Can be decomposed by a chemical change
Chemically united
Definite proportions
Different properties
Ex: H2O, NaCl, CO2
A Binary Compound is a compound that only has two elements (NaCl)

4. Differences between Mixtures
Heterogeneous
Uneven mixture of two different things.
Example: milk
Homogeneous
Solutions that are considered one thing.
Example: White vinegar

5. Ways To Separate Mixtures (1)
Distillation- a mixture of liquids can be separated by their boiling points.
Examples: distillation of salt water.
distillation of petroleum liquids.

6. Ways to Separate Mixtures (2)
Filtration- separates the solid and liquid parts of a mixture.
Example: coffee filter which separates the coffee grounds from the brewed coffee.

7. Ways of Separating Mixtures (3)
Chromatography- way of separating different molecules in a mixture.
Example: separating components of chlorophyll.

8. Ways of Separating Mixtures (4)
Centrifuge- a spinning machine that pushes the most dense particles to the bottom of the tube.
Example: separate isotopes such as separating uranium hexafluoride, and uranium-235.

9. Ways to Separate Mixtures (5)

10. Properties Chemical Properties Physical Properties
Are found by making a substance react, and form a new substance.
Ex. Burning, reaction w/ water or acid, changing to a new substance.
Physical Properties
Can be found without changing the substance to something else.
Ex. Color, hardness, phase, solubility, odder, density, mass, volume

11. Energy The ability to do work.
Exothermic- energy given off in a chemical reaction.
Endothermic- energy absorbed in a chemical reaction.
You measure energy in joules (J).

12. Table T Heat q= mCΔT q=heat m= mass
q= mHf C= specific heat capacity(table B)
ΔT= change in temperature
q= mHv Hf= heat of fusion
Hv= heat of vaporization

13. Sample Problem:
How much heat energy in joules if absorbed by 100g of water when it is heated from 20ºC to 30ºC?
q= m· C · ΔT
q= 100g x 4.18 Joules/gºC x (30 – 20) ºC
q= 100g x 4.18 J/gºC x 10ºC
q= 4,180 joules

14. Heating Curves
Heat of Fusion- is the amount of heat needed to change a solid into a liquid at a constant temperature.
Heat of Vaporization- is the amount of heat needed to change a liquid into a gas at a constant temperature.
Table B
Physical Constants for Water
Heat of Fusion J/g
Heat of Vaporization J/g
Specific heat capacity of H2O(l) J/gºC

15. Temperature
The measure of the average kinetic energy of the molecules.
The higher the temperature the more kinetic energy it has.
Heat flows from a higher temperature to a lower temperature until they are the same temperature.
Measured with a thermometer.

16. Temperature continued
Boiling Point- when the vapor pressure equals the atmospheric pressure.
Freezing Point- the temperature at which a liquid solidifies under a specified pressure.
Absolute Zero= -273oC or 0K
Kelvin-
K= Kelvin ºC = degrees Celsius
K = oC + 273

17. Changes Physical Change
-change in appearance, but no new substance is produced
Ex- tearing a piece of paper, heating ice
Chemical Change
- Produces a new substance with different properties
Ex- burning magnesium

18. Solids Melting point- temperature when a solid changes into a liquid.
Definite shape, definite volume, and crystalline structure, geometric pattern.
Closely packed particles that vibrate but don’t change position.
Melting point- temperature when a solid changes into a liquid.
Sublimation- the change from a solid directly to a gas. Ex. Dry ice (CO2) & Iodine

19. Liquids Evaporation- when liquid changes into a gas. Ex: water vapor.
Definite volume, takes shape of its container. Particles are close together and move :water.
Evaporation- when liquid changes into a gas. Ex: water vapor.
Vapor Pressure- the pressure that the vapor exerts on the sides of the container.

21. Gases No definite volume or shape.
Particles are far apart, and can expand anywhere.
* When there is a phase change from solid to liquid to gas entropy increases

22. As pressure increases, volume decreases at constant temperature
Gas Laws
Boyle’s Law
As pressure increases, volume decreases at constant temperature

23. As volume increases, temperature increases at constant pressure
Gas Laws
Charles’ Law
As volume increases, temperature increases at constant pressure

24. What is STP? Standard Temperature and Pressure
When you have a combined gas law at STP, use 273 K as your temperature and kPa as the pressure.

25. Kinetic Molecular Theory (Ideal Gas Law)
A model that tells how gasses should behave
Ideal gas- perfect gas that agrees with John Daltons 5 assumptions
Tiny particles
Elastic collisions
Gases are in constant motion
No force of attraction
Temperature is related to speed

26. Ideal Gas Real Gas Particles have no volume. No attractive forces
Examples: H2, He
Real Gas
Particles have volume
Attractive
Examples: Cl2, H2O(g)

27. How do the Gas Laws relate to the Kinetic Molecular Theory?
Boyle’s Law
Boyle’s law states that pressure and volume are inversely proportional.
If you have a million molecules in a container and you decrease that container the molecules will hit twice as often, therefore twice the pressure.
Charles’ Law
Charles’ Law states that as temperature increases, volume increases.
If you heat the air in a balloon, there will be more pressure on the sides. This makes the balloon bigger in volume.

28. Combined Gas Law
P1 x V = P2 x V2
T T2

29. THE END