1. Chapter 10 Solids and Liquids

2. Intermolecular Forces These are considered to be “weak” forces… That is not to say that they do not serve an important function.

3. Types of Intermolecular Forces 1.VanDerWaals 2.Dipole-Dipole Interactions 3.Hydrogen Bonds 4.Dispersion (London) Forces

4. VanDerWaal Forces These are electrostatic forces between molecules. VanDerWaals are responsible for many of the physical properties of substances. These are relatively weak.

5. Dipole-Dipole Interactions Polar molecules have regions of unevenly distributed electrical energy. The stronger the polarity, the stronger the force. They are similar to forces in crystal structures, but much weaker.

6. Hydrogen Bonds This is a bond that includes a hydrogen atom. The other member of the bond must be a highly electronegative atom. This can be a very strong bond. The best example of this is a water molecule.

7. London Forces This force involves nonpolar molecule like gasoline and wax. Electrons as they move sometimes momentarily concentrate on one end of the molecule. This can cause a VERY weak polar moment.

8. Solids Particles in solids vibrate, but not at great distances. Due to the relative close proximity of particles in solids, they are considered to be dense.

9. Solids Some solids have a natural orderly shape. They form three-dimensional patterns with distinct edges and sharp angles. Even when broken, the smaller pieces retain these edges and angles. These solids are called Crystalline Solids.

10. Solids Solids that have no preferred shape are called Amorphous Solids. The difference between the two is the way in which the particles (atom, ions, and or molecules) are stacked.

11. Melting and Freezing Melting and freezing are changes between the solid and liquid state. Every substance can melt or freeze!!!

12. Sensible Heat Heat, that when applied, results in a temperature change in the substance.

13. Latent Heat The heat energy that results in a phase change.

14. The Melting of Water Ice can be heated until it reaches 0 degrees. Any heat added at this point, does not raise the temperature. It serves to begin to break the bonds. Not until the ice completely melts, does the temperature begin to rise.

15. Evaporation Molecules in a liquid are not all moving at the same rate. The speeds are similar, but random motion rules out uniformity. Molecules with above average speeds can sometimes “break” away. This process is called evaporation. This can happen above or below boiling point!

16. Evaporation Evaporation is actually a cooling process. The molecules with the most kinetic energy leave. As a result, the ones left behind draw heat from the surrounding environment. This heat does not raise the temperature of the molecules left behind, it merely replaces the energy that the others took.

17. Heat of Vaporization The heat of vaporization is the amount of heat required to convert one gram of a liquid at its boiling point to its vapor at the same temperature. This differs for different liquids.

18. Condensation Fluids can easily reenter the liquid state. Condensation is the reverse of vaporization. Dynamic Equilibrium occurs when the process of condensation and evaporation oppose each other so that no net effect can be seen.

19. Vapor Pressure The pressure exerted by evaporated molecules is called the vapor pressure. As more molecules evaporate, the more molecules present to press on the walls of the container. As temperature increases, so does vapor pressure.

20. Boiling Boiling is the rapid change of state between a liquid and a gas. Vapor escapes from a liquid’s surface and also may form internally and collect in bubbles. Boiling occurs when the vapor pressure of a liquid equals the atmospheric pressure.

21. Boiling Boiling point is defined as the temperature at which the vapor pressure equals the applied pressure. Normally this is considered pressure of 760 torr.