To ensure that you understand the material in this chapter, you should review the meanings of the following bold terms in the following summary and ask yourself how they relate to the topics in the chapter.

A phase is a certain form of matter that has the same physical properties throughout. Three phases are common: the solid, the liquid, and the gas phase. The strength of the intermolecular interactions determines whether a substance is a solid, liquid, or gas under any particular conditions. Covalent network bonding, exemplified by diamond, is a very strong form of intermolecular interaction.  Ionic interactions, the forces of attraction due to oppositely charged ions, are also relatively strong. Covalent bonds are another type of interaction within molecules, but if the bonds are polar covalent bonds, then the unequal sharing of electrons can cause charge imbalances within molecules that cause interactions between molecules. These molecules are described as polar, and these interactions are called dipole-dipole interactions. A certain rather strong type of dipole-dipole interaction, involving a hydrogen atom, is called hydrogen bonding. On the other hand, equal sharing of electrons forms nonpolar covalent bonds, and the interactions between different molecules is less because the molecules are nonpolar. All substances have very weak dispersion forces (also called London forces) caused by the movement of electrons within the bonds themselves.

In the solid phase, intermolecular interactions are so strong that they hold the individual atoms or molecules in place. In many solids, the regular three-dimensional arrangement of particles makes a crystal. In other solids, the irregular arrangement of particles makes an amorphous solid. In liquids, the intermolecular interactions are strong enough to keep the particles of substance together but not in place. Thus, the particles are free to move over each other but still remain in contact.

In gases, the intermolecular interactions are weak enough that the individual particles are separated from each other in space. The kinetic theory of gases is a collection of statements that describe the fundamental behavior of all gases. Among other properties, gases exert a pressure on their container. Pressure is measured using units like pascal, bar, atmosphere, or mmHg (also called a torr).

Gas laws quantify simple relationships between the variables used to describe a quantity of gas. Boyle’s law relates the pressure and volume of a gas, while Charles’s law relates the volume and absolute temperature of a gas. The combined gas law relates the volume, pressure, and absolute temperature of a gas sample. All of these gas laws allow us to understand the changing conditions of a gas. The ideal gas law relates the pressure, volume, amount, and absolute temperature of a gas under any conditions. These four variables are related to the ideal gas law constant, which is the proportionality constant used to calculate the conditions of a gas. Because the conditions of a gas can change, a set of benchmark conditions called standard temperature and pressure (STP) is defined. Standard temperature is 0ºC, and standard pressure is 1.00 atm.