1. States of Matter
High points

2. Gases Kinetic theory of matter: all particles are in constant motion
The particles in gases are considered to be small spheres with insignificant volume, in rapid, constant, and random motion
All collisions are perfectly elastic…no loss of energy…this results in gas pressure: the simultaneous collisions of multitudes of particles in an object

3. Gas Pressure
Atmospheric pressure: the result of all the particles in air pushing against everything around them
Barometer: a device for measuring atmos. pressure, originally a column of mercury, pushed 760 mm high.
Pascal: SI unit,1atm=760mmHg=101.3kPa
at 25C.

4. KE and Temperature
Realize that we are not looking at an individual particle but the average KE of a large number of particles
See 388 for graph showing relative temps.
Temperature is simple a measurement that uses a scale to show the relative KE of the particles of a given substance

5. Liquids Similar to gases, except:
The particles in a liquid are attracted to each other…these intermolecular attractions keep the particles closer together, which determines the physical properties of the liquids. This will be especially clear in terms of viscosity, adhesion/cohesion, surface tension, etc.

6. Liquids: evaporation Vaporization: phase change at boiling pt.
Evaporation: molecules going from liquid to gas at a temperature below the BP.
This requires a higher KE than the average shown by the molecules of the liquid.
This can be speeded up with heat!

7. Liquid: vapor pressure
VP: the pressure/force exerted by a gas above a liquid. IF the vapor pressure is high, it means that the liquid is really trying to move the equilibrium to the gas phase!
Note table 13.1, page 392 for discussion; see next slide
This pressure can be measured with a manometer, which allows for the pressure difference to be measured (13.7, pg. 393)

8. Vapor Pressure Curves

9. Boiling Point
The point at which the vapor pressure is equal to the atmospheric pressure (still at equilibrium, those particles liqgas and gas liq).
Any time more energy goes into the liquid, the liquid starts boiling off (more  gas).
If the atmospheric pressure is lessened, the boiling point is lowered! Normal BP is defined at 1 atm, kPa.

10. Solids
Movement between molecules is even more restricted, as the orderly arrangement is now called a crystal lattice, where the ions, atoms, or molecules are packed tightly together.
When a solid is heated, the particles vibrate more rapidly, to the point where it eventually overcomes the attractions holding the particles in place…we call this melting and the temperature is the melting point.

11. Crystal Structures
The shape of a crystal reflects the internal arrangement of the particles within the solid.
A crystal has faces, outer boundaries of the lattice of a given substance, which can form in the proper environment.
There are seven groupings/systems of how crystal lattices form.

12. Crystal Systems
Each system is defined by the lengths of the crystal axes and the angles at which these axes intersect.
See fig , page 397 for a color picture of these.
The outward shape is dependent upon the geometric shape of a unit cell, the smallest grouping of particles for the given substance.
See fig , 398 for the 3 possible for cubic system.

13. Allotropes Two or more different molecular forms of the same element.
Only a few elements: C, P, S, O, B, Sb.
Carbon: diamond, graphite, fullerene.

14. Non-Crystalline solids
Amorphous: lacking an ordered internal structure
Examples: rubber, plastic, asphalt

15. Changes of state Phase change triangle
Sublimation: occurs in solids when VP exceed atm. pressure at or near room temp.
Phase change diagram: fig , 403
See handout