1. States of Matter
Chapter 13

2. Temperature Temperature: average kinetic energy of particles
Standard Temperature: 0oC
Units: Celsius, Kelvin, Farenheit
Absolute Zero: KE = 0; all motion of particles stops
Celcius  Kelvin: TK = TC + 273

4. Pressure
Gas Pressure: the force exerted by a gas per unit surface area of an object.
 Created by the collision
of gas particles
with a surface.
Vacuum: empty space; no particles; no pressure

5. Units of Pressure SI Unit = pascal (Pa)
Other units: atmospheres (atm), millimeters of mercury (mmHg), torr.
Standard Pressure: kPa = 1 atm = 760 mmHg = 760 torr = 14.7 psi

6. Practice with Pressure Conversions
7.31 psi = _______ mmHg
7.31 psi 760 mmHg =
14.7 psi
1140 torr = _______ kPa
1140 torr kPa =
760 torr
19.0 psi = ________ kPa
202 kPa = ________ psi
377.9 mmHg
kPa
130.9
29.3

7. Pressure
Atmospheric Pressure: created by the gases that make up Earth’s atmosphere.
Atmospheric pressure decreases
as elevation increases
(lower density of gases).
Measured using a Barometer.

8. States of Matter State of Matter Arrangement of Particles
Motion of Particles
Definite Shape
Definite Volume
Relative Forces of Attraction
Solid
Close
Vibrate 
 yes
 highest
Liquid
 Close
 Flow
 no
 high
Gas
 Far Apart
Random crazy motion No
none

9. States of Matter

10. Phase Changes Phase changes indicate EQUILIBRIUM.
Example: During vaporization, you have just as many particles going from:
liquid  gas
Gas liquid

11. Phase Changes Freezing: liquid solid Melting: solid liquid
Condensation: gas  liquid
Vaporization: liquid  gas
Boiling – usually at high temps, bubbles within liquid
Evaporation – room temp, surface of liquid
Sublimation: solid  gas
Deposition: gas  solid

12. Phase Diagrams
Phase Diagrams – Show the pressure and temperature for the three states of matter for a particular substance.

13. Phase Diagram

14. Phase Diagram Terms
Triple Point: the temperature and pressure at which three phases of a substance can coexist.
Lines: indicate T and P at which two phases exist in equilibrium
Critical Point: critical temperature above which no amount of pressure can change the vapor into a liquid.
Normal Pressure = Standard Pressure (1 atm or 760 torr)

15. Phase diagrams: H2O vs. CO2
Water
Carbon Dioxide
At standard pressure, CO2 is a solid at very cold temperatures. But it turns to a gas when heated at 1.0 atm, instead of a liquid  SUBLIMATION

16. Kinetic Molecular Theory
Kinetic Theory – The particles that compose matter are in constant motion.

17. Gases GASES: no definite shape or volume; easily compressible
Particles are very far apart from one another.
Between particles is empty space.
No attractive or repulsive forces between particles. (They move independently.)
Move randomly at high speeds
Travel in straight paths, but can change directions due to a collision.
Collide elastically
no kinetic energy lost
Ideal Gas

18. Vapor Pressure
Vapor Pressure – the pressure exerted by a vapor over a liquid in a closed container.
Vapor pressure (of a contained liquid) increases as temperature increases.
Volatile – A substance that evaporates easily (has weak bonds)
A volatile substance has a high vapor pressure.

19. Boiling Point
Boiling point: the temperature at which the vapor pressure equals the atmospheric pressure exerted on the surface of the liquid.
Normal boiling point – temperature where liquid  gas at standard pressure.
Boiling point of a substance increases when the pressure on the liquid increases.
Why must foods cooked in boiling water at high altitudes be cooked longer?
- Lower pressure lowers the boiling point

20. Vapor Pressure Diagramb a c

21. Vapor Pressure Diagram
1) Identify the normal
boiling point for each
substance.
2) Which substance is the
most volatile? Least
volatile?
3) Which substance has
the strongest forces of
attraction? Weakest?
4) How does vapor pressure change as temperature increases? (What kind of relationship is that?)