Download Presentation
Chapter 13 States of Matter

Loading in 2 Seconds...

1 / 19

# Chapter 13 States of Matter - PowerPoint PPT Presentation

Chapter 13 States of Matter. Read pgs. 384 - 409. Kinetic Molecular Theory. The kinetic molecular theory describes the behavior of gases in terms of particles in motion. A. Gases consist of point-masses. No volume B. Gas particles are in constant motion.

I am the owner, or an agent authorized to act on behalf of the owner, of the copyrighted work described.
Download Presentation

## Chapter 13 States of Matter

An Image/Link below is provided (as is) to download presentation

Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author.While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server.

- - - - - - - - - - - - - - - - - - - - - - - - - - E N D - - - - - - - - - - - - - - - - - - - - - - - - - -
Presentation Transcript

### Chapter 13States of Matter

Read pgs. 384 - 409

Kinetic Molecular Theory
• The kinetic molecular theory describes the behavior of gases in terms of particles in motion.
• A. Gases consist of point-masses. No volume
• B. Gas particles are in constant motion.
• C. Collisions between gas particles are elastic.
• D. There are no repulsive or attractive forces acting between gas particles.
• E. Temperature is a measure of the average speed of gas particles.
Behavior of Gases
• Gases have low densities. Gases have lots of empty space between particles.
• Gases have no definite volume. They can be compressed or expanded.
• Gases diffuse evenly through out a space based upon their temperature and molar mass.
• Kinetic Energy = ½ m v2 where:
• m = mass and v = velocity, speed
Graham’s Law of Effusion
• At the same temperature, two gases will effuse at a rate based upon their molar mass.
• The heavier the gas the slower it moves.
• The lighter the gas the faster it moves.
• See board for equation.
Gas Pressure
• Because gas has mass, it creates a pressure. Pressure = Force / Area
• The higher up in the atmosphere you go, the less air, the less air pressure.
• Air pressure is the weight of air over an object.
• Units used to measure pressure: Atmospheres, (Atm), Kilopascals, (KPa), millimeters of mercury, (mmHg)

Standard Pressure is pressure at sea-level at 0oC.

• 1.0 atm = 101.3 kPa = 760 mm Hg
• Dalton’s Law of Partial Pressures

Ptotal = P1 + P2 + P3 + . . . + Pn

The total pressure in a closed container is equal to the sum of all pressures in the container.

Intermolecular Forces
• Intermolecular forces are forces that hold particles together. They are not the same as a chemical bond.
• Dispersion Forces – weak forces caused by the movement of electrons.
• Dipole Forces – attractive forces between polar molecules.
• Hydrogen Bonds – special type of dipole force between hydrogen and an unshared pair of electrons.
Properties of Liquids
• The kinetic molecular theory can also be applied to liquids and solids.
• Liquids are more dense than gases because the particles are closer together.
• Liquids are only slightly compressible.
• Liquids have no definite form so they can flow. Liquids are fluid.
• Liquids have a viscosity. Viscosity is a measurement of the resistance to flow.

Viscosity is temperature dependent.

• The higher the temperature, the more a liquid flows.
• Think pancake syrup.
• Cold syrup – pours slowly. Warm syrup – pours very quickly.
• Surface Tension – is a measurement of the inward pull by particles on the surface of a liquid.
• The more surface tension a liquid has, the more it will ball up.
• Think about water on a freshly waxed car.
• Raindrops are round because of this.
Properties of Solids
• Solids vibrate about fixed points.
• Solids tend to be the densest phase of matter for most substances.
• Two types of solids: Crystalline and Amorphous
• Crystalline solids have a definite arrangement of atoms.
• Crystals come in definite shapes
• Crystalline solids have definite melting points

Amorphous solids has no regular repeating pattern.

• Amorphous solids have no definite melting point.
• You can also classify solids by the type of bonds that they have.
• Molecular solids – covalent bonds
• Ionic solids – ionic bonds
• Metallic solids – metallic bonds
Phase Changes
• All phase changes are reversible.
• All phase changes comes in pairs.
• Freezing – Melting solids and liquids
• Vaporization – Condensation liquids and gases
• Sublimation – Deposition solids and gases
• All phase changes involve a change in energy. One is exothermic and one is endothermic.
• For a phase change to occur, particles must have some minimum kinetic energy.
Melting - Freezing
• Melting points and freezing points are always the same.
• When particles have enough energy to overcome Intermolecular forces, phase change happens.
• Temperature remains constant during phase change.
Vaporization / Evaporation
• Vaporization is the change of a liquid to a gas.
• Vaporization or Boiling take place at a specific point. Evaporation can take place at any temperature.
• Boiling points vary, but take place through out the whole liquid.
• Evaporation only happens at the surface of a liquid.
Evaporation
• Evaporation is a slow change from a liquid to a gas.
• Evaporation can be increased by several ways.
• Increase the room temperature.
• Increase air currents over liquid.
• Increase surface area of liquid.

Evaporation can even take place at temperatures below freezing.

Boiling Points
• For boiling to take place, particles have to have a certain amount of energy.
• Boiling point is also affected by atmospheric pressure.
• Normal boiling point is the boiling point at standard pressure.
• Water can boil at any temperature.
Sublimation / Deposition
• Sublimation is the direct change from the solid phase to the gaseous phase.
• Deposition is the opposite of sublimation.
• Some examples of sublimation:
• Dry ice, frozen CO2
• Freezer burn
Phase Diagrams
• Phase diagrams show the phases of a substance at different temperatures and pressures.
• Each substances phase diagram is a little different.
The triple point is the exact temperature and pressure where all three phases exist at the same time.
• The critical point is where you can no longer change a gas back to a liquid.