Vaporization is the escape of a molecule from a liquid. The energy of a individual liquid molecule constantly changes as it collides with other molecules of higher and lower energy. If an individual liquid molecule has a large enough amount of kinetic energy and finds itself at the surface of the liquid it can overcome the attractive forces holding it in the liquid, escape the surface, and enter the gas phase.

When both the liquid and gaseous phases of a substance are present simultaneously, the gaseous part of the sample is called a vapor. At room temperature neither nitrogen or oxygen can exist as a liquid so the nitrogen and oxygen gases making up the atmosphere are technically not vapors. Substances like gasoline or water do have stable liquid phases at room temperature, however and the gaseous molecules evaporating from these liquids are correctly called vapors.

The vapor pressure of a gaseous substance above its liquid phase increases with increasing temperature. As the temperature of a liquid is increased by adding heat (energy), the heat energy is stored in the molecules of the liquid as kinetic energy (as well as rotational and vibrational energy). Increasing kinetic energy (1/2 mv 2 ) mean increasing velocities and more molecules are able to escape the liquid’s surface per second.

If a liquid is not enclosed in an airtight container, molecules leaving its surface may never return. Sweat evaporates from your body. Puddles of rainwater disappear A pan of water boils dry. In an airtight container (closed system), however the molecules in the gas phase will eventually strike the surface of the liquid and be reabsorbed into the liquid phase or condense.

The rate of condensation of a gas back into the liquid phase depends upon the number of molecules in the gaseous state. The rate evaporation of liquid molecules depends upon the temperature. In a closed system, a liquid will evaporate until there are enough molecules in the gaseous state for the rate of condensation to equal the rate of evaporation.

When these two rates are balanced, the system is said to be in an equilibrium state. The vapor pressure under these conditions is called the equilibrium vapor pressure. The equilibrium vapor pressure of a liquid always increases as the temperature of the liquid increases. When two dynamic processes oppose each other and reach an equilibrium, it is called a dynamic equilibrium. The amount of heat absorbed during vaporization is exactly equal to the amount of heat liberated during condensation. A dynamic equilibrium is indicated by two arrows pointing in opposite directions.

In a closed system (bell jar) a dynamic equilibrium will always occur if enough liquid water is present at the start. In an open system (puddle of water) the rate of evaporation is generally greater than the rate of condensation (unless the relative humidity is 100%) and a dynamic equilibrium cannot occur. Whenever a molecule leaves the surface of a liquid, it carries with it kinetic energy. Since the liquid subsequently contains less energy, its temperature becomes lower. Evaporation is a cooling process. Nature uses it to reduce body temperature by the evaporation of sweat.

Starting with a liquid whose vapor pressure is less than atmospheric pressure, the liquid can be made to boil by: Raising the liquid’s temperature until the vapor pressure equals atmospheric pressure. Lowering atmospheric pressure until it equals the vapor pressure of the liquid at the original temperature. In the first case one might start with water at 70 o C at sea level and heat the water until it started to boil at 100 o C In the second case one might start with water at 70 o C at sea level and carry the water to the top of Mt. Everest (where atmospheric pressure is much less) where it would spontaneously start to boil at 70 o C.

Since the boiling point of a liquid depends upon atmospheric pressure, the normal boiling point of a liquid is defined as the temperature at which a liquid will boil when the atmospheric pressure is exactly 1 atm. The value of the normal boiling point of a liquid depends upon the strengths of the secondary forces holding the molecules of the liquid together.

The arrangement of atoms or molecules in a solid is in one of two forms: Crystalline: Regular repeating arrangement of atoms or molecules. Amorphous: Lacking a regular repeating arrangement of atoms of molecules.

Crystalline substances have solid crystals that have sharply defined melting points and exist in well defined shapes, such as prisms, octahedrons, or cubes.

There are three types of crystals: Covalent crystals: The atoms are held in place by a network of covalent bonds. Examples of covalent crystals are diamond and quartz. Covalent crystals have very high melting points. Ionic crystals: The atoms are held together by simple ionic bonds. Examples of ionic crystals are NaCl and CaCl 2. Ionic crystals have intermediated melting points. Molecular crystals: The molecules in molecular crystals are held together by much weaker secondary forces. Examples of molecular crystals are solid ethanol and ice.

Certain substances can evaporate from the solid state directly to the gaseous state without the intervening liquid state. This process is called sublimination. Examples of substances that sublime at atmospheric pressure are dry ice (solid CO 2 ), solid I 2, and p-dichlorobenzene, a clothes moth repellant. When a substance sublimes, the energy required for a molecule to enter the gas phase from the solid phase is the same as if it had gone through two separate transitions: melting and evaporation.

Chapter 6 Summary Molecules and the States of Matter Molecules retain their sizes and shapes whether in the solid, liquid, or gaseous state. Molecules in the solid state are highly organized and in close contact. Molecules in the liquid state also are in contact but are less well organized. In the gaseous state, the amount of space between molecules and the degree of disorder among them are at a maximum.

Chapter 6 Summary Attractive Forces Between Molecules Attractive forces between molecules are called secondary forces. The secondary forces that lead to the formation of liquids and solids are London forces, dipole–dipole forces, and hydrogen bonds. The relative strengths of these forces are in the order: hydrogen bonding dipole–dipole attractions London force. The relative strength of the secondary forces between molecules controls differences between melting and boiling points of different substances

Chapter 6 Summary Changes in the States of Matter Phase transitions take place when the balance between kinetic energy and secondary forces is changed by the addition or removal of heat. Melting, the conversion of solid into liquid, happens when only a part of the secondary forces is overcome. Vaporization, the conversion of liquid into gas, occurs when all secondary forces are overcome.

Chapter 6 Summary Vapor Pressure and Dynamic Equilibrium A gas in contact with its liquid form is called a vapor and exerts a pressure called the vapor pressure. Vapor pressure increases with increase in temperature. Vaporization and condensation of a liquid take place simultaneously. When the two processes proceed at the same rate, the amounts of liquid and gas (vapor) remain constant and the system is said to be in dynamic equilibrium.

Chapter 6 Summary Vaporization of Water and the Regulation of Body Temperature The body’s ability to maintain a constant temperature chiefly depends on the large value of water’s heat of vaporization.

Chapter 6 Summary Formation of Solutions If the secondary forces type as in another substance, mix to form a solution.

Chapter 6 Summary The Structure of Solids In contrast with the structural disorder found in gases and liquids, the atoms or molecules of solids are arranged in distinct and ordered patterns. The regular patterns of atoms within solids lead to their geometrically well defined shapes and fixed melting points.