1. Chapter 13 States of Matter
13.1 Gases
The Kinetic-Molecular Theory
Explaining the Behavior of Gases
Gas Pressure

2. The Kinetic Molecular Theory
Basic Assumptions
Particle Size
Gas particles have no volume (pin point particles)
The space between particles is extremely large compared to the volume of the particles. Due to this distance, there is no significant attractive or repulsive force acting on the particles.

3. The Kinetic Molecular Theory
Basic Assumption
Particle Motion
Gas particles are in constant random motion.
Collisions between particles are elastic (Energy can be transferred from one particle to another during a collision, but no energy is lost when particles collide)

4. Basic Assumptions Particle Energy
The mass and velocity of a particle determine the kinetic energy of a particle
Temperature is a measure of the average kinetic energy of particles in a sample.

5. The Kinetic Molecular Theory
Mass/Velocity Relationship Questions
Condition #1:
Two particles (one heavy and one light) traveling at the same velocity.
Which exhibits the greatest kinetic energy?
Condition #2:
Two particles of the same size traveling at different velocities (fast and slow).

6. Explaining the Behavior of Gases
Properties
Low Density (pinpoint mass/volume of empty space)
Random Motion
Behaviors
Compression
Gases can be compressed due to the large space that exists between particles
Gases expand to fill their containers due to constant random motion

7. Explaining the Behavior of Gases
Properties
No attractive or repulsive forces acting on particles
Particles exhibit constant random motion
Behaviors
Particles can flow easily past each other in a process called diffusion. The rate of diffusion is dependent on the mass of the particles.
Question: Based on this equation which particles diffuse faster, heavy or light particles?

8. Explaining the Behavior of Gases
Property
Particles exhibit constant random motion
Behavior
Effusion (similar to diffusion, where particles escape through a tiny opening)
Graham’s Law of Effusion
Rate of effusion

9. Questions/Problems
What assumption of the Kinetic-Molecular Theory explains why a gas can expand to fill a container?
How does the mass of a particle affect its rate of effusion?
Calculate the ratio of diffusion rates for CO and CO2.
What is the rate of effusion for a gas that has a molar mass twice that of a gas that effuses at a rate of 3.6mol/min?

10. The force that a gas exerts per unit area.
Gas Pressure
The force that a gas exerts per unit area.
Measuring Air Pressure Using a Barometer
Invented by Evangelista Torricelli
A rise in air pressure will cause the height of mercury to rise.
Two forces affect the height of the mercury column
Gravity and Atmospheric Pressure
A decrease in air pressure will cause the height of mercury to fall

11. Pressure and Altitude Question
Which condition would cause the level of mercury in a barometer to fall below 760 mm?
Being below sea level or being on the top of a mountain.

12. Units of Pressure The SI unit for pressure is the pascal (Pa)
At sea level and 0oC conditions (STP-standard temperature and pressure) a barometer will read 760mm Hg.
Equivalent pressure units
760 mm Hg = kPa = 1atm = 760 torr = 14.7 psi

13. Factor-Label is in the Air
Equivalent pressure units
760 mm Hg = kPa = 1atm = 760 torr = 14.7 psi
Convert 362 torr to kPa
Convert 35.4 psi to torr
Convert 48.9 kPa to psi

14. Dalton’s Law of Partial Pressures
“The total pressure of a mixture of gases is equal to the sum of the pressures of all the gases in a a mixture”
The partial pressure of a gas is dependent on the number of moles of gas, the size of the container and the temperature of the mixture.

15. Partial Pressure Problems
1. What is the partial pressure of hydrogen gas in a mixture of hydrogen and helium if the total pressure is 600 mm Hg and the partial pressure of helium is 439 mm Hg?
2. Find the total pressure in kPa for a mixture that contains three gases with partial pressures of 122 kPa, 35 psi, and 722 torr.

16. Chapter 13 States of Matter
13.2 Attractive Forces
Dispersion Forces
Dipole-dipole Forces
Hydrogen Bonds

17. Intramolecular Forces
Forces that occur between atoms, ions or molecules within a molecule.
Bonding
Attractive Parties
Ionic
Cations and anions
Molecular
Positive nuclei and shared electrons
Metallic
Metal cations and mobile electrons

18. Intermolecular Forces
Forces that occur between molecules to hold them together
Dispersion
Dipole-Dipole
Hydrogen Bonding

19. Dispersion Forces
Weak forces that occur between non-polar molecules that result from a temporary shift in the density of electrons in electron clouds.
(Butane)
The electrons of two non-polar repulse one another which causes the temporary shift.

20. Dispersion Forces
With increasing atomic number, the number of electrons in a molecule increases which results in a greater dispersion force. This explains why Cl2 is a gas, Br2 is a liquid and I2 is a solid.
Greater Force = smaller distance between molecules

21. Dipole-dipole Forces
Attractive force that occurs between molecules that have a permanent dipole.
Stronger than dispersion forces as long as the two molecules have about the same mass.

22. Hydrogen Bonding
For a hydrogen bond to form, hydrogen must be bonded to oxygen, fluorine or nitrogen.

23. Attractive Force Review
Why are dipole-dipole forces typically stronger than dispersion forces?
Which molecules listed below can form hydrogen bonds? Which ones could only experience dispersion forces (how would you know)?
H2, NH3, HCl, HF
3. Predict the relative boiling points of the noble gases.

24. 13.3 Liquids and Solids

25. Liquid Behavior Density- much denser than gases, not compressible
Fluidity- diffuse slower than gases, still “flow”
Viscosity- measure of resistance to flow
Effect of temperature- higher temp, lower viscosity

26. Liquid Behavior Surface Tension Capillary Action
Causes drops and meniscus
Capillary Action
Water can climb narrow tubes

27. Solid Behavior
Density- more dense than gases and liquids and incompressible
Crystalline Solids
Unit Cells – smallest particle of a crystal that has the same shape as the crystal
Crystal Structure

28. Types of Solids
Molecular Solids-dispersion, dipole, or H-bonds (ex: sugar)
Covalent Network Solids- covalent bonds with self (ex: diamond, graphite)
Ionic Solids- ionic attraction (ex: salt)
Metallic Solids- mobile electrons (ex: copper)
Amorphous Solids- irregular pattern (ex: glass)

29. Phase Changes

30. Heating Curve

31. Endothermic Phase Changes
Melting- solid absorbs energy until particles have enough speed to break free of IM forces holding them in place
Vaporization-liquid absorbs energy until particles have enough speed to break free of IM forces holding them close together

32. Liquid to Gas Evaporation- occurs at surface
Boiling- occurs throughout when vapor pressure equals atmospheric pressure

33. Sublimation-solid to gas

34. Exothermic Phase Changes
Condensation
Deposition
Freezing

35. Phase Diagrams
Triple Point