1. States of Matter
Chapter 12

2. Gases
Kinetic molecular theory = describes how gases behave, has 3 major assumptions:
Particle size = particles are small compared to volume of empty space, no forces between particles
Particle motion = particles are in constant motion, move in a straight line until they collide with other particles or container walls, elastic collisions (no loss of kinetic energy)

3. 12.1…
3. Particle energy = determined by mass and velocity of a particle (KE=1/2 mv2)
In a sample of a gas, all particles have same mass but different velocity (so KE of particles varies)
Temperature = measure of average KE of particles in a sample (so at a given temp., all particles have same average KE)

4. 12.1 – Gas Behavior Low density = lots of space between gas particles
Compression and expansion = empty space in a gas container can easily be pushed into a smaller volume (ex. Squishing memory foam)
Diffusion and effusion = random motion of gas particles causes them to mix until they are evenly distributed

5. 12.1 - Diffusion and Effusion
Diffusion = movement of one material through another from area of high concentration to area of low concentration (ex. Spraying perfume)
Effusion = gas particles escape through a tiny opening (ex. Popping a balloon)
At the same temp., heavier particles effuse/diffuse slower than lighter ones (Graham’s law of effusion: rate is inversely proportional to molar mass.

6. 12.1 – Gas Pressure
Pressure = force exerted by gas particles colliding with the walls of the container
Barometer = measures atmospheric pressure, height of mercury determined by gravity exerting downward force on mercury and upward force of air pressing down on surface of mercury (increase in air pressure = mercury rises, decrease in air pressure = mercury falls), air pressure depends on temp. and humidity

7. 12.1 – Units of Pressure Pascal (Pa) = SI unit of pressure
Average air pressure at sea level =
760 mm Hg
760 torr
101.3 kilopascals (kPa)
14.7 psi
One atmosphere (1 atm)

8. Pressure Demos!
How to get an egg inside a milk bottle (vacuum inside the bottle)
Extreme Shaving Cream
Monster Marshmallows

9. 12.1 – Dalton’s Law of Partial Pressures
Dalton’s Law of Partial Pressures = total pressure of a mixture of gases is equal to the sum of the pressures of all gases in the mixture
Portion of the total pressure contributed by one gas is called its partial pressure
Partial pressure of a gas depends on the number of moles of gas, size of the container, and temp.

10. 12.2 – Forces of Attraction
Intramolecular forces = attractive forces that hold particles together (bonds)
Intermolecular forces = attractive forces between particles

11. 12.2 – Intermolecular Forces
Dispersion forces = weak force between two non-polar molecules (those with evenly distributed electrons)
Electrons always moving, electron clouds of non-polar molecules repel each other
Electron density around the nucleus shifts so one side of the molecule is more negative than the other
Leads to weak attraction between these temporary dipoles

12. 12.2 – Dispersion Forces

13. 12.2 – Intermolecular Forces…
Dipole-Dipole forces = forces between oppositely charged regions of polar molecules (permanent dipoles)
Ex. HCl
Partially positive H in one molecule of HCl lines up with partially negative Cl in another molecule

14. 13.2 – Intermolecular Forces…
Hydrogen bonds = a type of dipole-dipole attraction between molecules that contain hydrogen and a highly electronegative atom with at least one lone pair
Explains why water is a liquid at room temperature and not a gas  life on Earth!
Hydrogen atom has large partial positive charge and is attracted to oxygen’s large partial negative charge on another molecule

15. 13.2 – Hydrogen Bonds

16. 12.3 – Liquids and Solids

17. 12.3 – Liquids and Solids…
Liquids are much denser than gases due to intermolecular forces that hold particles together
Liquids cannot be compressed as much as gases because particles are already closely packed together
Molecules in a liquid

18. 12.3 – Liquids and Solids…
Fluidity = ability to flow, both gases and liquids are considered fluids because they can diffuse through one another
Liquids are less fluid than gases due to intermolecular forces that slow down their movement
Ex. Water pipe leak vs. natural gas leak

19. 12.3 – Liquids and Solids… Viscosity = resistance of a liquid to flow
Stronger intermolecular forces  higher viscosity
Bigger molecules  higher viscosity
Lower temperature  higher viscosity
Ex. Motor oil = need less viscous oil in winter to make sure it keeps flowing at low temperatures, need more viscous oil in summer to make sure it stays thick enough to lubricate the engine at very high temperatures

20. 12.3 – Liquids and Solids
Surface tension = energy required to increase surface area of a liquid by a given amount
Stronger attractions between particles  higher surface tension
Water has a very high surface tension due to hydrogen bonding, allows small insects to walk on water, why we need water AND soap to clean dishes/clothes
Surface tension of water
Magic milk

21. 12.3 – Liquids and Solids…
Capillary action = occurs when adhesion to walls of container is greater than cohesion in the liquid
Adhesion = force of attraction between molecules that are different
Cohesion = force of attraction between identical molecules
Ex. Meniscus in graduated cylinder, absorbance of paper towels and diapers

22. 12.3 – Liquids and Solids…
Solids = not considered to be fluids, very ordered, more dense than liquids because particles are more closely packed together
Exception = water, solid ice is less dense than liquid water due to three dimensional structure of water molecules (ice cubes and glaciers float, keeps aquatic life alive in the winter)
Why ice floats...

23. 12.3 – Liquids and Solids…
Crystalline solids = atoms, ions, or molecules arranged in an orderly, geometric, 3-D structure
Individual pieces are called crystals
Unit cell = smallest part of a crystal retaining the crystal shape
Molecules in solids

24. 12.3 – Liquids and Solids 4 types of crystalline solids:
Molecular solids = soft, poor conductors of heat and electricity, low to moderate melting points
Allotropes = 2 different solid forms of the same element (ex. Carbon as graphite or diamond), very hard, poor conductors, high melting points
Ionic solids = hard, brittle, poor conductors, high melting points
Metallic solids = soft to hard, excellent conductors, low to high melting points

25. 12.3 – Liquids and Solids…
Amorphous solids = do not contain crystals, no repeating pattern, form when a molten material cools too quickly
Ex. Glass, rubber, plastics

26. 12.4 – Phase Changes
Intro to phase changes

27. 12.4 – Phase Changes Why do substances melt?
As molecules are heated, they gain enough KE to overcome attractive forces that hold them together as a solid
Melting point = temperature at which forces holding crystal lattice together are broken and solid becomes liquid
REQUIRES energy

28. 12.4 – Phase Changes Why do substances freeze?
As heat is removed, molecules lose KE and slow down
Freezing point = temperature at which a liquid is converted into a crystalline solid
RELEASES energy

29. 12.4 – Phase Changes Why do liquids become gases?
When energy is added to a liquid, the temperature increases and molecules gain KE and escape the liquid
Vaporization = process by which a liquid changes to a gas or vapor
Evaporation = vaporization that occurs on the surface of a liquid that is not boiling
REQUIRES energy

30. 12.4 – Phase Changes…
Vapor pressure = pressure exerted by a vapor over a liquid
Boiling point = temperature at which vapor pressure of a liquid is equal to atmospheric pressure, molecules throughout the liquid have enough energy to vaporize and bubbles of vapor rise to the surface
Water boiling in a vacuum Demo!

31. 12.4 – Phase Changes… Why do gases become liquids?
When a vapor molecule loses energy, it slows down and is more likely to collide with another molecule and bond, releasing energy and becoming a liquid
Condensation = Process by which a gas or vapor becomes a liquid
RELEASES energy
Evaporation and Condensation

32. 12.4 – Phase Changes… What happens when we sweat?
Water molecules in sweat gain energy from heat coming off the body
Some molecules gain enough energy to evaporate, taking energy with them
Molecules left behind have less energy and therefore, a lower temperature

33. 12.4 – Phase Changes…
Sublimation = solid changes directly to a gas without passing through the liquid phase
Ex. Dry ice
REQUIRES energy
Deposition = gas or vapor changes directly to a solid without passing through the liquid phase

34. 12.4 – Phase Changes
Phase diagram = Pressure vs. Temperature, shows relationship between pressure/temperature/phase of a substance
Triple point = point on the diagram that shows the temp. and pressure at which all three phases can exist in equilibrium with eachother
Critical point = point on the diagram that shows the temp. and pressure at which water can no longer exist as a liquid

35. Phase Diagram