1. Properties of Matter
GPS 3

2. Phases of Matter (Station #1)
Simulation
SOLID
Definite Shape
Definite Volume
LIQUID
Shape varies depending on container
Definite Volume
GAS
Takes on the shape and volume of the container
KINETIC ENERGY INCREASES

3. Phases of Matter (Station #1)
Solid – definite shape, definite volume
Liquid – definite volume, no definite shape
Gas – no definite shape or volume
Plasma – A state of matter that is formed when a gas is heated to a temperature at which its electrons dissociate from the nuclei
usually hotter than 100,000 ºC
ENERGY
INCREASES
The universe is about 99% plasma (stars and outerspace mostly plasma). Not seen as often on earth because earth is cooler. Volcanoes, plasma tv screens

4. Phase Changes (Station #2)
Melting – solid to a liquid
Freezing – liquid to a solid
Evaporating – liquid to a gas
Condensing – gas to a liquid
Sublimating – solid to a gas
Energy released or absorbed?

5. Heat vs Temperature Heat Temperature
The transfer of energy (from hotter objects to colder objects)
Temperature
The degree or intensity of heat present in a substance or object
As the temperature increases, the average kinetic energy for that sample increases

6. Kinetic-Molecular Theory (Station #3)
The Nature of Gases:
A gas has mass
Gas particles are spread far apart
Particles are on constant, rapid, random motion
Collisions are elastic
Energy depends on temperature

7. Kinetic-Molecular Theory (Station #3)
What did you observe happening to the food dye in the cold water verses the hot water?
How would you try to explain your observations?

8. Kinetic-Molecular Theory (Station #3)
-15°C °C °C
The diagrams above depict water as a solid, liquid, and gas at different temperatures. Which phase has the greatest average kinetic energy?

9. Kinetic-Molecular Theory (Station #3)
Nitrogen at 20°C Chlorine at 100°C Oxygen at 195°C
The diagrams above show different gases at different temperatures. Which substance has the least average kinetic energy?

10. Heating Curve for Water (Station #4)
Simulation
Simulation
“Horizontals”: potential energy is increasing (particles are spread farther apart to change phases) so there is no increase in temperature

11. What phase or phases are present at D?
What letter corresponds to only gases being present?
What is happening at B?

12. What is the melting point of this substance?
About how many minutes did it take for this substance to begin to boil?

13. What is the boiling point of this substance?

14. General Phase Diagram
Beyond the critical point, the liquid and gas phase become indistinguishable (kind of a hybrid form); forms what we call a supercritical fluid
Add in melting, freezing, vaporization, condensation, sublimation

15. Phase Diagram for Water
Hydrogen bonding causes unusually high boiling point
In most phase diagrams, the line between liquid and solid is slanted to the right. For water it is slanted to the left. For most liquids, their solid forms are more dense. However, solid water (ice) is actually less dense than liquid water. Ice is less dense than liquid water because when ice freezes, the hydrogen bonds in the crystaline structure cause ice to become more spread out when it freezes. So, ice is more dense and takes up more volume that liquid water. At high pressures, molecules of water are pushed as close to each other as possible. When molecules of water are closest, they are in the liquid form because the liquid form takes up less volume.

16. What phase is present at B?
What take place if a substance changes from B to A?
What take place if a substance changes from C to A?

17. What phase or phases are present at F?
What phase change takes place from A to G?
Will energy be gained or released?

18. Label the phase changes taking place.

19. Definitions (Station #5)
Matter – anything that has mass and takes up space
Ex: A quarter, a chair, air
Element – a pure substance that is made up of only one specific type of atom
Ex: gold, silver, oxygen, nitrogen
Atoms – the smallest particle of an element that has all the properties of that element

20. Definitions (Station #5)
Compound – a substance containing more than one type of element, which are chemically bonded
Ex: Table salt (NaCl: sodium chloride)
Molecule – a group of two or more atoms bonded together
Ex: A molecule of water
Substance – has the same composition throughout. Made of either elements or compounds.
Ex: A block of gold, a glass of water

21. Definitions (Station #5)
Mixture – a physical blend of at least two substances (homogeneous mixture and heterogeneous mixture)
Homogeneous Mixture (a.k.a. Solution) – the same composition and properties throughout
Ex: sugar water, salt water
Heterogeneous Mixture – not the same composition and properties throughout
Ex: rock salt and sugar crystals mixed together

22. Determining the type of Mixture (Station #6)
To determine if a mixture is homogeneous or heterogeneous:
Take several random samples of the mixture
Compare the composition (ratio of components in the mixture)
If the ratio of the components is the same in all of the samples, then it is homogeneous.
Ex 1: all samples 50:50 salt water
Ex 2: all samples 30:70 salt water
Ex 3: sample 1 = 20:80 salt water, sample 2 = 20:80 salt water, sample 3 = 10:90 salt water

23. Determining the type of Mixture (Station #6)
Sample
Homogeneous or Heterogeneous mixture?
Salt, iron filings, and sand
Salt water
Sand and water
Oil and water
Rubbing alcohol and water

24. Separating Mixtures Common Method: Common Method:
Filtering (with filter paper):
Used to separate a solid from a liquid
Common Method:
Distillation:
Used to separate a liquid from a liquid OR a solid and liquid

25. Separating Mixtures Paper Chromatography:
Another method used to separate the liquid components of a mixture
- Use a polar solvent (water or rubbing alcohol)– more polar components go farther than nonpolar/less polar components

26. Separating Mixtures Common Method:
• Decanting (with a stirring rod held flat across the top of the beaker)
Used to pour off a liquid when it has separated from a solid that has settled to the bottom of the beaker

27. Physical and Chemical Changes (Station #7)
Physical Change
- The same substance is present before and after the change
- Ex: ripping a piece of paper
Chemical Change
- A different substance is present after the change (in other words, a reaction takes place)
- Ex: burning a piece of paper
Ask about other examples: boiling water (boiling point), pounding metal (malleability), baking a cake (chemical change), melting iron (melting point)
Other physical changes: conductivity, ductility (ability of a material to be deformed)

28. Physical and Chemical Changes (Station #7)
Is it a physical change or a chemical change?
Baking a cake
Boiling water
An iron nail rusting
Pounding a piece of copper
Sodium reacting with chlorine

29. Physical and Chemical Properties (Station #7)
Physical Properties
- A property that can be observed without changing the substance
- Ex: Iron is magnetic
Chemical Properties
- A property of a substance when it is involved in a chemical change
- Ex: Iron reacts with air and water to form rust

30. Density: A Physical Property (Station #8)
Simulation
d = m/v
Density = mass/volume
Density is the measure of mass per unit volume of a substance ( given in g/mL)
Is the density of ice greater than or less than the density of water?
What would happen to fish in a lake during the winter if ice were more dense than water?
Explaining density: people say lead is heavier than feathers. But a truckload of feathers is heavier than a single piece of lead. Mass and volume must be accounted for. The density of lead is greater than the density of feathers
Density Demo (regular coke vs diet coke)
Show how ice crystals freeze – they are less dense because the volume increases as ice forms
Show how water circulates when a lake is freezing – 4 deg C ice starts to form on surface (traps 4 deg. Water below so fish survive even as the surface decreases to zero deg)
*Density Lab- design their own lab – density of various materials
*Density Lab – Instruct students how to properly use balances (paper, etc.)

31. Density
Ex: Find the density of an object that has a mass of 1.25 g and a volume of 0.55 mL.
d = 2.3 g/mL

32. Density
Find the mass of an object with a volume of 25.0 mL and a density of 2.1 g/mL.
m = 53 g

33. Density
Find the volume of an object with a mass of 35.8 grams and a density of 25.9 g/ml.
1.38 ml

34. Pepsi Station Notes (Station #9)
What was your hypothesis? Why?
What did you observe?
Try to explain what you observed.

35. Spheres Station Notes (Station #10)
What was your hypothesis? Why?
What did you observe? Explain why you think the spheres “behaved” the way they did.

36. Graduated Cylinder Station Notes (Station #11)
What did you observe?
What do you think keeps these liquids from mixing together?
What do you think these liquids could be made of?

37. “Lava” Lamp Station Notes (Station #12)
What did you observe?
How would you try to explain your observations?

38. Resources

39. SOL covered in these lessons
CH 2 h
CH 5 c