1. Liquids The particles that make up liquids are in constant motion too. Liquid particles are free to slide past one another. This allows liquids, as well as gases, to “flow.”

2. Liquids and the Kinetic- Molecular Theory A liquid can be described as a form of matter that has a definite volume and takes the shape of its container. The attractive forces between particles in a liquid are more effective than those between particles in a gas. This attraction between liquid particles is caused by the intermolecular forces: dipole-dipole forces(hydrogen bonding) London dispersion forces Induced Dipole

3. Liquids But, the particles in liquids are close enough to be attracted to each other. These attractive forces are intermolecular forces. The particles are in motion, BUT their average kinetic energy isn’t high enough to allow the particles to break away from the liquid and become gas particles. Hydrogen bonds Surface Tension

4. Liquids However, liquid particles can escape from the surface of the liquid and become a gas. You can heat a liquid and boil it and convert it from a liquid to a gas. When a liquid converts to a gas, and it is not boiling, we call that process evaporation.

5. Liquids Both evaporation and boiling cooling are actually “cooling” processes. -- Because the particles with the highest energy tend to escape into the gas phase first, the particles which are left behind have a lower energy. –Because those particles have lower energy (the ones who were left behind), the average temperature decreases. The particles are the surface of a liquid are constantly breaking away to form a gas above the surface of the liquid.

6. Liquids Those particles form a gas above the surface of the liquid. Since that gas a “pressure” on the surface of the liquid, we call that the “vapor pressure” of the liquid. At the same time as the gas particles are forming, some of the gas particles condense to reform a liquid. We call the process is which a gas becomes a liquid “condensation,” which is essentially the opposite of evaporation.

7. Liquids Over time, the number of particles escaping the liquid to become a gas = the number of gas particles condensing to become a liquid. When this happens, we say we have reached a state of equilibrium, where rate of evaporation = rate of condensation Boiling point is defined as the point at which the vapor pressure of the liquid equals the atmospheric pressure. Because of this, The boiling point of a substance can change as the atmospheric pressure goes down. Boiling Lake in the Dominican Republic

8. The Other State of Matter

9. Solids The particles in solids are still in constant motion, but they don’t have as much freedom to move, as they do in liquids or gases.The particles in solids are still in constant motion, but they don’t have as much freedom to move, as they do in liquids or gases. Therefore they tend to vibrate in place as opposed to sliding or moving from place to place.Therefore they tend to vibrate in place as opposed to sliding or moving from place to place.

10. TRV POLKA GASES LIQUIDS SOLIDS

11. Solids As you heat a solid, the particles begin to vibrate more rapidly as the kinetic energy increases. The organization of the particles within the solid begins to break down. The solid eventually melts and becomes a liquid. The melting point is the temperature at which the forces holding the solid together are overcome by the kinetic energy of the vibrating particles.

12. Solids When a liquid turns back into a solid, we call this process freezing. Most solids are crystalline in nature and we call them crystals. Look at page 339 for a discussion of various crystal structures. I will not ask about crystal structures on the test, however, for those who are curious, take a look.

13. Solids allotropes Some solids can exist in more than one form. We call these different forms of the same element allotropes. Carbon is an example:  Diamond, Graphite, Buckyballs

14. Amorphous Solids Solids which do not exist in crystalline form are called “amorphous” solids, because they lack an ordered internal Rubber, plastic and structure. Rubber, plastic and asphalt asphalt are all examples of amorphous solids. GlassGlass is another example of an amorphous solid. Glasses are supercooled” liquids sometimes called “supercooled” liquids.

15. Section 4 Phase Changes

16. Changes of Phase and Equilibrium phase A phase is any part of a system that has uniform composition and properties. (Solid,Liquid or Vapor) Phase change A Phase change occurs when the dynamic equilibrium of a system is stressed and modified. Equilibrium Equilibrium is a dynamic condition in which two opposing changes occur at equal rates in a closed system.

17. Equilibrium

18. Possible Changes of State

19. Phase Diagrams The relationships among the solid, liquid and vapor states is defined by the heat absorbed or released by the substance. All three Phases can be represented in a graph called a “phase diagram.”

20. Phase Diagrams Pressure vs.Temperature

21. Phase Diagrams Each region represents a phase. The lines represent the conditions at which those two phases are in equilibrium. The graph has temperature on the independent (x) axis and pressure on the dependent (y)axis. all the points The Curve represents all the points in which a phase change can occur.

22. There is a phase diagram for every pure substance. When you cross a line in the graph you are going through a phase change!

23. Melting & Boiling Points normal melting pointIf you look at the graph at standard pressure (101.3 KPa or 1atm), the point on the line where the solid and liquid are in equilibrium is known as the “normal melting point.” normal boiling point.Continuing across the graph at standard pressure you will find the point on the line where the liquid and vapor are in equilibrium. This point is known as the “normal boiling point.”

24. Phase Diagrams critical pointAt the top right of the diagram you see something known as the critical point, which for water is 225 atm and 374K. critical pointDo some research and discover the definition of the critical point. (Pg 347) critical temperature While you are at it define critical temperature.

25. Triple Point triple point A unique feature of a phase diagram and matter is the “triple point,” This point is the temperature and pressure at which all three phases of matter are present and in equilibrium.