Chapter 13 States of Matter

1. Chapter 13States of Matter Read pgs. 384 - 409

2. 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.

3. 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

4. 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.

5. 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)

6. 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.

7. 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.

8. 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.

9. 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.

10. 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

11. 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

12. 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.

13. 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.

14. 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.

15. 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.

16. 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.

17. 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

18. Phase Diagrams • Phase diagrams show the phases of a substance at different temperatures and pressures. • Each substances phase diagram is a little different.

19. 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.