Water is the only substance we routinely encounter as a solid, a liquid, and a gas. At low temperatures, it is a solid in which the individual molecules are locked into a rigid structure. As we raise the temperature, the average kinetic energy of the molecules increases, which increases the rate at which these molecules move.
There are three ways in which a water molecule move: (1) vibration, (2) rotation, and (3) translation. Water molecules vibrate when H--O bonds are stretched or bent. Rotation involves the motion of a molecule around its center of gravity. Translation literally means to change from one place to another. It therefore describes the motion of molecules through space.
To understand the effect of this motion, we need to differentiate between intramolecular and intermolecular bonds. The covalent bonds between the hydrogen and oxygen atoms in a water molecule are called intramolecular bonds. (The prefix intra- comes from the Latin stem meaning "within or inside." Thus, intramural sports match teams from the same institution.) The bonds between the neighboring water molecules in ice are called intermolecular bonds, from the Latin stem meaning "between." (This far more common prefix is used in words such as interface, intercollegiate, and international.)
The intramolecular bonds that hold the atoms in H2O molecules together are almost 25 times as strong as the intermolecular bonds between water molecules. (It takes 464 kJ/mol to break the H--O bonds within a water molecule and only 19 kJ/mol to break the bonds between water molecules.)
All three modes of motion disrupt the bonds between water molecules. As the system becomes warmer, the thermal energy of the water molecules eventually becomes too large to allow these molecules to be locked into the rigid structure of ice. At this point, the solid melts to form a liquid in which intermolecular bonds are constantly broken and reformed as the molecules move through the liquid. Eventually, the thermal energy of the water molecules becomes so large that they move too rapidly to form intermolecular bonds and the liquid boils to form a gas in which each particle moves more or less randomly through space.
The difference between solids and liquids, or liquids and gases, is therefore based on a competition between the strength of intermolecular bonds and the thermal energy of the system. At a given temperature, substances that contain strong intermolecular bonds are more likely to be solids. For a given intermolecular bond strength, the higher the temperature, the more likely the substance will be a gas.
The kinetic theory assumes that there is no force of attraction between the particles in a gas. If this assumption were correct, gases would never condense to form liquids and solids at low temperatures. In 1873 the Dutch physicist Johannes van der Waals derived an equation that not only included the force of attraction between gas particles but also corrected for the fact that the volume of these particles becomes a significant fraction of the total volume of the gas at high pressures.
The van der Waals equation is used today to give a better fit to the experimental data of real gases than can be obtained with the ideal gas equation. But that wasn't van der Waals's goal. He was trying to develop a model that would explain the behavior of liquids by including terms that reflected the size of the atoms or molecules in the liquid and the strength of the bonds between these atoms or molecules. The weak intermolecular bonds in liquids and solids are therefore often called van der Waals forces. These forces can be divided into three categories: (1) dipole-dipole, (2) dipole-induced dipole, and (3) induced dipole-induced dipole.
Many molecules contain bonds that fall between the extremes of ionic and covalent bonds. The difference between the electronegativities of the atoms in these molecules is large enough that the electrons aren't shared equally, and yet small enough that the electrons aren't drawn exclusively to one of the atoms to form positive and negative ions. The bonds in these molecules are said to be polar, because they have positive and negative ends, or poles, and the molecules are often said to have a dipole moment.
HCl molecules, for example, have a dipole moment because the hydrogen atom has a slight positive charge and the chlorine atom has a slight negative charge. Because of the force of attraction between oppositely charged particles, there is a small dipole-dipole force of attraction between adjacent HCl molecules.
The dipole-dipole interaction in HCl is relatively weak; only 3.3 kJ/mol. (The covalent bonds between the hydrogen and chlorine atoms in HCl are 130 times as strong.) The force of attraction between HCl molecules is so small that hydrogen chloride boils at -85.0oC.
Dipole-Induced Dipole Forces
What would happen if we mixed HCl with argon, which has no dipole moment? The electrons on an argon atom are distributed homogeneously around the nucleus of the atom. But these electrons are in constant motion. When an argon atom comes close to a polar HCl molecule, the electrons can shift to one side of the nucleus to produce a very small dipole moment that lasts for only an instant.
By distorting the distribution of electrons around the argon atom, the polar HCl molecule induces a small dipole moment on this atom, which creates a weak dipole-induced dipole force of attraction between the HCl molecule and the Ar atom. This force is very weak, with a bond energy of about 1 kJ/mol.
Induced Dipole-Induced Dipole Forces
Neither dipole-dipole nor dipole-induced forces can explain the fact that helium becomes a liquid at temperatures below 4.2 K. By itself, a helium atom is perfectly symmetrical. But movement of the electrons around the nuclei of a pair of neighboring helium atoms can become synchronized so that each atom simultaneously obtains an induced dipole moment.
These fluctuations in electron density occur constantly, creating an induced dipole-induced dipole force of attraction between pairs of atoms. As might be expected, this force is relatively weak in helium -- only 0.076 kJ/mol. But atoms or molecules become more polarizable as they become larger because there are more electrons to be polarized. It has been argued that the primary force of attraction between molecules in solid I2 and in frozen CCl4 is induced dipole-induced dipole attraction.