1. Chapter 13 & 14 YOU NEED TO READ!!!!
States of Matter
Chapter 13 & 14
YOU NEED TO READ!!!!

2. Properties of Gases
Low Density- gases are much less dense than liquids or solids
Compressible
Will diffuse / effuse- gas particles flow easily past each other. Random motion causes them to mix until evenly distributed
Diffusion-movement of one material through another.
High concentration--> low concentration
Effusion – movement of gas particles through a tiny opening (like a pinhole)
The rate of diffusion and the rate of effusion depend on the molar mass of the gas; heavier gases will effuse more slowly than lighter gases at the same temperature

3. Gas Pressure SI unit is the Pascal (Pa) Standard pressure:
1.00atm = 760mmHg = 101.3kPa
(be able to convert)
Dalton’s Law of partial pressure: total pressure of a mixture of gases is = sum of the pressure of all the gases.
(P total = P1+ P2+P3…..Pn)
** also, Dalton’s law states that the mole fraction of a gas (χ) times the total pressure equals the partial pressure of that gas

4. The Ideal Gas Law
The ideal gas law is a mathematical equation that relates gas pressure (P) to the volume (V), temperature (T), and number of moles (n) present in a sample of gas.
PV=nRT
where R is the universal gas constant ( L∙atm/mol∙K)

5. Other Gas Laws
Boyle’s law says that at constant temperature, the pressure and volume of a sample of gas are inversely proportional
Charles’s Law says that at constant pressure, the volume and temperature of a sample of gas are directly proportional
Gay-Lussac’s law says that if volume is held constant, temperature and pressure of a sample of gas are directly proportional
Avogadro’s Law says that if pressure and temperature are constant, the number of moles of gas and volume are directly proportional

6. Kinetic Molecular Theory (a model for explaining ideal gases)
Describes the behavior of ideal gases in terms of particles in motion, making several assumptions about size, movement, energy and attraction.
The theory applies when a gas is behaving ideally
REAL gases do not behave ideally at low temperature and/or high pressure

7. Kinetic Molecular Theory points
Particles are very small, so volume is negligible with regard to distance between them. (each particle is assumed to have zero volume)
Particles are in random straight-line motion constantly until collisions occur/ collisions cause pressure.
Collisions are elastic, no energy lost or gained because there is no attraction or repulsion between the particles.
The average kinetic energy of the particles is directly proportional to Kelvin temperature.
K= °C + 273
* At 0 K (-273°C, absolute zero) all molecular motion theoretically stops

8. Real Gases Real gas particles do have volume
Real gas particles do have forces of attraction between them (intermolecular forces)
SO…Real gases will condense to liquids and solids when the conditions are right

9. Forces of Attraction
Intermolecular Forces- the forces between 2 or more molecules.
Intramolecular forces – the forces within one molecule (i.e., bonds)
Intermolecular forces are weaker than intramolecular forces.

10. Types of Intermolecular forces
Dispersion Forces: weak forces that result from temporary shifts in density of electron clouds.
Dipole-Dipole Forces: relatively strong attraction between oppositely charged regions of polar molecules.
Hydrogen bonds: very strong attraction between the hydrogen of one molecule and nitrogen or oxygen or fluorine of another molecule

11. Properties of Liquids
Somewhat compressible; fixed volume that takes shape of its container
Density- liquids are denser than gases due to intermolecular forces holding particles together.
Fluidity- ability to flow
*Liquids have a slower rate of diffusion than gases at the same temperature due to intermolecular forces

12. Properties of Liquids (cont.)
Viscosity- resistance to flow; determined by type of intermolecular forces, shape of particles and temperature.
Surface Tension- measure of inward pull by particles in the interior

13. Properties of Liquids (cont.)
Vapor Pressure- pressure caused by the vapor over the surface of a liquid. In a closed container, vapor gets trapped and eventually the space above the surface of the liquids gets “saturated” with vapor then equilibrium happens (evap. at the same rate as condensation)
Vapor pressure is temperature dependent. Increase Temp. , Increase VP.
Example- puddles evaporate on a warm day
(be able to read and answer questions from a vapor pressure diagram)

14. Properties of Liquids
When VP over a liquid= atmospheric pressure, the liquid boils
Examples: covering a spaghetti pot will cause water come to a boil faster or water will boil at higher temperature than 100°C when altitude is high
Normal boiling pt. = boiling temp. when P= 1.00 atm
NOTE: Boiling (vaporization) is different from evaporation

15. Solids
Particles have no freedom of movement but vibrate around fixed points
Solids tend to be dense and incompressible
In most cases particles are tightly packed in a highly organized pattern
Heating a solid breaks down the intermolecular forces- eventually the solid melts.

16. Types of solids
Ionic Solids – have high melting pts because ionic bonds must be overcome. They are crystalline- particles arranged in an orderly repeated pattern called a crystal lattice; the smallest group of particles within a crystal that retain the geometric shape is a unit cell.
Molecular solids- have lower melting pts because intermolecular forces are weaker than ionic bonds. They are often amorphous- lack an ordered structure of particles.
Metallic solids – generally have high melting pts because the atoms of the metal are held together by a “sea” of electrons (this also explains why they are electrically conductive)
Covalent network solids – have extremely high melting pts because they contain a network of covalent bonds (ex: diamonds, quartz)

17. Phase Changes Endothermic (taking in energy)phase changes:
Melting- solid becoming liquid
Boiling- liquid becoming gas (at the boiling point)
Evaporation- liquid becoming gas at a temperature other than the boiling point
Sublimation- solid becoming gas without going through the liquid phase

18. Phase Changes Exothermic( releasing energy):
Freezing- liquid becoming solid
Condensing- gas becoming liquid
Deposition- gas becoming solid (without going through the liquid phase)

19. Heating (cooling curve)
Be able to analyze and perform calculations associated with a heating or cooling curve such as the one shown below.

20. Phase Diagram
A graph of P vs T that shows which phase a substance exists under various conditions (be able to analyze and answer questions from a phase diagram, such as the one shown below)