A phase is a certain form of matter that includes a specific set of physical properties. That is, the atoms, the molecules, or the ions that make up the phase do so in a consistent manner throughout the phase. Three of the phases are: the solid phase, in which individual particles can be thought of as in contact and fixed in place; the liquid phase, in which individual particles are in contact but moving past each other; and the gas phase, in which individual particles are separated from each other by relatively large distances. Not all substances will readily exhibit all phases. For example, carbon dioxide does not exhibit a liquid phase unless the pressure is greater than about six times normal atmospheric pressure. Other substances, especially complex organic molecules, may decompose at higher temperatures, rather than becoming a liquid or a gas.  Several other phases exist but will not be considered here.

Which phase a substance adopts depends on the pressure and the temperature it experiences. Of these two conditions, temperature variations are more obviously related to the phase of a substance. When it is very cold, H₂O exists in the solid form as ice. When it is warmer, the liquid phase of H₂O is present. At even higher temperatures, H₂O boils and becomes steam.

Pressure changes can also affect the presence of a particular phase, as indicated for carbon dioxide, but its effects are usually less obvious. We will focus on the temperature effects on phases. Chemical substances follow the same pattern of phases when going from a low temperature to a high temperature: the solid phase, then the liquid phase, and then the gas phase, although some compounds decompose without entering the more mobile phases. However, the temperatures at which these phases are present differ for all substances and can be rather extreme. Table 8.1 “Temperature Ranges for the Three Phases of Various Substances” shows the temperature ranges for solid, liquid, and gas phases for three substances. As you can see, there is extreme variability in the temperature ranges.

What accounts for this variability? Why do some substances become liquids at very low temperatures, while others require very high temperatures before they become liquids? It depends on the strength of the intermolecular interactions between the particles of substances. (Although ionic compounds are not composed of discrete molecules, we will still use the term intermolecular to include interactions between the ions in such compounds.) Substances that experience strong intermolecular interactions require higher temperatures to become liquids and, finally, gases because it takes more energy to overcome the intermolecular attractions and mobilize the particles to enter the liquid or gas phase.   Substances that experience weak intermolecular interactions do not need much energy to become liquids and gases and thus will exhibit these phases at lower temperatures.  The following paragraphs will consider substances with weakest to strongest intermolecular attraction, thus generally lowest to highest melting/boiling points.

All molecules  experience temporary charge separation caused by chance uneven distribution of electrons in the molecule, which also induces charge separation in adjacent molecules by attracting or repelling that molecule’s electrons.  While these charge separations are tiny and temporary, they never-the-less cause attractions between neighboring molecules.  These very weak intermolecular interactions are called London dispersion forces.  Molecules that experience no other type of intermolecular interaction will at least experience dispersion forces. Substances that experience only dispersion forces are typically soft in the solid phase and have relatively low melting points. Because dispersion forces are caused by the instantaneous distribution of electrons in a molecule, larger molecules with a large number of electrons can experience substantial dispersion forces. Examples include waxes, which are long hydrocarbon chains that are solids at room temperature because the molecules have so many electrons.

Weak London dispersion forces are the only attractive forces between nonpolar covalent molecules, resulting in generally low melting and boiling points.  Molecule may be nonpolar by having only nonpolar bonds or by having polar bonds that cancel each other.   Carbon dioxide (CO₂) and carbon tetrachloride (CCl₄) are examples of nonpolar molecules having polar bonds that cancel each other.  Under normal atmospheric pressure, carbon dioxide sublimes rather than melting and boiling.  Carbon tetrachloride melts at -22.9 ^oC and boils at 76.7 ^oC.  (Figure 8.1 “Nonpolar Molecules”).

Figure 8.1 Nonpolar Molecules. Although the individual bonds in both CO₂ and CCl₄ are polar, their effects cancel out because of the spatial orientation of the bonds in each molecule. As a result, weak London dispersion forces are the only attractions between molecules and melting and boiling points are generally low.

In addition to London dispersion forces, polar molecules experience dipole-dipole interactions.  As discussed in Sections 4.4 and 4.5, molecules are polar if they have polar bonds that do not cancel each other, resulting in distinct areas of the molecule having δ+ and δ- charges.  Figure 8.3 shows the dipole-dipole attractions between molecules of acetone.  Acetone melts at -95.4 ^oC and boils at 56.5 ^oC.  Despite acetone’s stronger dipole-dipole attractions, its melting and boiling points are lower than those for nonpolar carbon tetrachloride.  However, carbon tetrachloride is almost 3 times more massive than acetone, so its London dispersion forces  add up to a good deal of attraction between molecules.

When a polar molecular compound has a hydrogen atom bonded to an atom of one of most electronegative elements, fluorine, oxygen, or nitrogen, the δ+ on the H atom and the δ- on O, N or F atom form especially strong dipole-dipole attractions known as hydrogen bonds between molecules, again in addition to the weaker London forces and any other dipole-dipole attractions. A hydrogen bond is about 10% as strong as a covalent bond.  The physical properties of water, which has two O–H bonds, are strongly affected by the presence of hydrogen bonding between water molecules. Figure 8.3 “Hydrogen Bonding between Water Molecules” shows how molecules experiencing hydrogen bonding can interact.  The melting point of water is 0 ^oC, and its boiling point is 100 ^oC.  Note that water is a much smaller molecule than previous examples carbon tetrachloride and acetone, yet has higher melting and boiling point due to the relatively strong intermolecular hydrogen bonding experienced by water molecules.

Substances with the highest melting and boiling points have covalent network bonding.   In these substances, all the atoms in a sample are covalently bonded to other atoms; the entire sample is essentially one large molecule, which cannot truly melt or boil, but can decompose.  Many of these substances are solid over a large temperature range because it takes so much energy to disrupt all of the covalent bonds at once. One example of a substance that shows covalent network bonding is diamond, shown in Figure 8.5,  which is a form of pure carbon. At temperatures over 3,500°C, diamond finally vaporizes into gas-phase atoms.

The phase that a substance adopts at a given temperature and pressure depends on the type and magnitude of the intermolecular interactions the particles of a substance experience. If the intermolecular interactions are relatively strong, then a large amount of energy—in terms of temperature—is necessary for a substance to change phases. If the intermolecular interactions are weak, a low temperature is all that is necessary to move a substance out of the solid phase.  The strengths of the individual attractions fall in this order: ionic interaction >>hydrogen bonding > dipole-dipole interaction>>London dispersion forces.  In future chapters, this concept will be used to predict and explain relative melting and boiling points for organic compounds.