1	Chapter 7 Gases, Liquids, and Solids
2	Gases, Liquids, and Solids
3	Hydro-power from potential energy Fig. 7.1 The water in the lake behind the dam has potential energy as a result of its position.
4	Kinetic Energy and Collisions
5	Properties of Gases, Liquids, and Solids Table 7.1
6	Arrangement of Particles in three States Fig. 7.3 (a) In a solid, the particles are close together. (b) In a liquid, the particles slide freely over one another. (c) In a gas, the particles are in random motion.
7	Gases in random motion Fig. 7.4 Gas molecules can be compared to billiard balls in random motion, bouncing off one another.
8	Popping into different states
9	Compressibility of gases Fig. 7.5 When a gas is compressed, the amount of empty space in the container is decreased.
10	Barometer: Measuring the Pressure of gases Fig. 7.6 The essential components of a mercury barometer are a graduated glass tube, a glass dish, and liquid mercury.
11	Boyle and the gas Laws Fig. 7.7 Robert Boyle was self-taught. Through his efforts, the true value of experimental investigation was first recognized.
12	Boyles’ Law: Fig. 7.8 Data illustrating the inverse proportionality associated with Boyle’s law.
13	Charles Law Fig. 7.9 When the volume of a gas at constant temperature decreases by half, the average number of times a molecule hits the container walls is doubled.
14	Application of Boyle’s Law Fig. 7.10 Filling a syringe with a liquid is an application of Boyle’s law.
15	Charles and gas laws Fig. 7.11 Jacques Charles in the process of working with hot-air balloons made the observations that led to the formulation of what is now known as Charles’s law.
16	Charles Law Fig. 7.12 Data illustrating the direct proportionality associated with Charles’s law.
17	Dalton and the Partial Pressure Fig. 7.13 John Dalton had an interest in the study of weather.
18	Dalton’s law of partial pressures Fig. 7.14 A set of four containers can be used to illustrate Dalton’s law of partial pressures. The pressure in the fourth container equals the sum of the first three.
19	Blood gases and their solubilty
20	Summary of gas laws
21	Changes of states Fig. 7.15 There are six changes of state possible for substances.
22	Sublimation of Iodine The beaker contains iodine crystals.
23	Vapor pressure and equilibrium Fig. 7.17 (a) the liquid level drops for a time, (b) then becomes constant. At that point a state of equilibrium has been reached in which (c) the rate of evaporation equals the rate of condensations.
24	Boiling point of liquids Fig. 7.18 Bubbles of vapor form within a liquid when the temperature of the liquid reaches the liquid’s boiling point.
25	How you cook? Low or high pressure Fig. 7.19 The converse of the pressure cooker “phenomenon” is that food cooks more slowly at reduced pressure.
26	Intermolecular forces: Dipole-dipole Fig. 7.20 There are many dipole-dipole interactions possible between randomly arranged CIF molecules.
27	Vapor pressure and the temperature Table 7.2
28	Pressure cooker: how does it work? Table 7.4
29	Boiling point and Eelvation Table 7.3
30	Strongest intermolecular force: The hydrogen bonding Fig. 7.21 Depiction of hydrogen bonding among the water molecules.
31	Three types of hydrogen bonding Fig. 7.22 Diagrams of hydrogen bonding involving selected simple molecules.
32	Hydrogen bonding in ice
33	How does hydrogen bonding affect properties Fig. 7.23 If there were no hydrogen bonding between water molecules, the boiling point of water would be approximately -80C.
34	London Dispersion Forces: Fig. 7.24 Nonpolar molecules can develop instantaneous dipoles and induced dipoles.
35	Summary of intermolecular forces