1. Behavior of Gases Low Density Compression and Expansion
Gases have low density because particles have a small mass in a large volume.
Compression and Expansion
A large amount of empty space exists between gas particles which allow gases to be compressed into a smaller volume
When compression stops, the gas expands to its original shape (like a sponge)

2. Behavior of Gases Diffusion Effusion
When 2 or more gases mix together until distributed evenly.
Depends on mass (lighter particles diffuse more quickly)
Effusion
Gas escapes through a tiny opening
The heavier the particle, the slower the effusion rate

3. Behavior of Gases Graham’s Law of Effusion
Rate of Effusion α 1/√molar mass
Rate of Diffusion A = √(molar mass B/molar mass A)
Rate of Diffusion B

4. Force / Area (force divided by area)
Pressure
Force / Area (force divided by area)
Force: a push or pull in a certain direction
Pounds (lbs), Newtons (N)
Area: Length x Width
in2, ft2, cm2, m2
The air in a tire exerts a force of 120 lbs. The surface area of the tire is 40 in2. What is the pressure in the tire?
120 lbs = 3 lbs/ in2 40

5. Gas Pressure Scientific measure of pressure is called the Pascal (Pa).
It is a very small unit, so we use kilopascals (kPa = Pa).
1 Pa = 1 Newton m2
A box with dimensions 2 m length and 3 m width exerts a force of Newtons on the floor. What pressure is the box exerting in kPa.
Area = 2 x 3 = 6 m2
N = Pa or 16.6 kPa
6 m2

6. Gas Pressure Barometer Atmospheric Pressure
Pressure that air particles exert on everything
Increase altitude, decrease air pressure
Increase depth, increase air pressure
Barometer
Instrument used to measure air pressure
Seen in weather reports

7. Units of Gas Pressure Unit Compared with 1atm Compared with 1kPa
Kilopascal (kPa)
1atm = 101.3kPa
mm Hg
1atm = 760 mmHg
1kPa=7.501 mm Hg
Pounds per square inch (psi)
1atm = 14.7 psi
1kPa = .145 psi
Atmosphere (atm)
1kPa = atm

8. Liquids Have no definite shape Have a fixed volume
Particles are able to move, but still have some attractive forces between them.
Molecules in a liquid are closer together than in a gas.

9. Properties of Liquids Density and Compression
Liquids are more dense than gases because of he intermolecular forces
Liquids are harder to compress than gases
Fluidity – the ability to flow
Gases and liquids both flow, but gases flow more easily than liquids

10. Properties of Liquids Viscosity – resistance to flow Surface Tension
Intermolecular forces determine liquid viscosity (stronger forces, higher viscosity)
As temp goes up, particles move faster, lowers viscosity, liquid flows more easily
Surface Tension
E required to increase the surface area of a liquid by a given amount (measure of the inward pull by particles below liquid’s surface)

11. Properties of Liquids Surfactants Capillary Action
Compounds that lower surface tension of water (break H bonds)
Capillary Action
Sum of adhesion and cohesion to draw a liquid up a narrow tube by molecular motion
Adhesion – attraction between different molecules
Cohesion – attraction between same molecules

12. Solids Solids are more dense than liquids and gases
Solid particles are closer together
Solid particles have stronger intermolecular forces between them
Crystalline solids – solid whose atoms, molecules, or ions are arranged in an orderly, geometric, 3D structure

13. Types of Solids Atomic Solids Molecular Solids
Noble gases in solid form
Dispersion forces hold them together
Soft, low melting point, poor conductors
Molecular Solids
Usually liquid or solid at room temp
Dispersion and dipole forces hold them together
Soft, low to moderate melting point, poor conductors
Water, Iodine crystals, sucrose

14. Types of Solids Covalent Network Solids Ionic Solids
Atoms that form multiple covalent bonds
Stronger than a regular covalent bond
Very hard, very high MP, poor conductors
Diamond (C) and quartz (SiO2)
Ionic Solids
Ionic bonds between oppositely charged ions
Crystalline, brittle, high MP, poor conductors
NaCl, KBr

15. Types of Solids Metallic Solids Amorphous Solids
+ metal ions surrounded by sea of mobile e-
Ductile, malleable, excellent conductors
Soft to hard, low to high MP
All metallic elements
Amorphous Solids
Can change shape
No crystals
Wax, rubber, plastic

16. Phase Changes Requiring E
Melting
Particles absorb heat E, speed up, and break intermolec forces between them to change to (l) phase
Melting Point
Temp at which solids turn to liquids
Vaporization
Particles absorb heat E, speed up, and break intermolec forces between them to change to a (g) phase

17. Phase Changes Requiring E
Evaporation
Only the surface of a (l) changes to (g).
How your body controls temperature (sweat)
Vapor Pressure
Pressure exerted by a (g) over a (l)
Boiling Point
Temp at which a (l) changes to a (g), or when VP of (l) = atmospheric pressure
Gas bubbles push out of the (l) against air pressure

18. Phase Changes Requiring E
Sublimation
A solid absorbs enough heat E to change directly to a (g) without becoming liquid
Results from particles speeding up rapidly with the increased heat
Examples are dry ice (frozen carbon dioxide), moth balls (solid naphthalene), snow, ice cubes.

19. Phase Changes Releasing E
Condensation
Process by which a gas (vapor) loses heat E and changes to a liquid
Particles slow down
Deposition
Gas loses heat E and changes to a (s) without becoming a (l)
Frost on a window

20. Phase Changes Releasing E
Freezing
Process by which a (l) loses heat E and changes to a (s)
Particles slow down
Freezing Point
Temp at which a (l) turns to a (s)

21. Phase Diagrams Phase Diagram Triple Point Critical Point
Graph of varying pressure v. temp that shows which phases a substance can exist.
Triple Point
Temp and pressure where a substance can be in all 3 phases (6 phase changes) at once
Critical Point
Critical temp and pressure above which only (g) exists