Sublimation is a phase transition where a substance changes directly from a solid state into a gaseous state, completely bypassing the intermediate liquid phase. This physical change is unique because most common substances, like water, move sequentially from solid to liquid, and then from liquid to gas when heated. The process is endothermic, meaning it requires the absorption of heat energy to occur.
Understanding the Phase Transition
Phase transitions depend on the energy of a substance’s molecules and the forces holding them together. Sublimation occurs when the energy supplied to the solid is sufficient to overcome the intermolecular forces, allowing the molecules to escape into the vapor phase. This direct jump contrasts with melting, where molecules gain enough energy only to move freely past one another as a liquid.
The energy required for this direct solid-to-gas change is defined as the enthalpy of sublimation (\(\Delta H_{\text{sub}}\)). This value represents the total heat needed to convert one mole of a solid directly into a gas at a constant temperature and pressure. According to Hess’s Law, the enthalpy of sublimation equals the sum of the enthalpy of fusion (solid-to-liquid) and the enthalpy of vaporization (liquid-to-gas).
The reverse transformation, where a substance transitions directly from a gas to a solid, is called deposition or desublimation. An example is the formation of frost, where water vapor converts directly into ice crystals on a cold surface. Sublimation and deposition represent an equilibrium point between the solid and gas phases.
The Role of Temperature and Pressure
Sublimation is governed by the conditions of temperature and pressure relative to a substance’s phase diagram. The existence of the liquid phase depends on the pressure being above a threshold known as the triple point. The triple point is the singular temperature and pressure where the solid, liquid, and gas phases of a substance can coexist in equilibrium.
If the surrounding pressure is kept below the triple point pressure, a substance cannot exist as a liquid, regardless of the temperature. Any heat added to the solid under these low-pressure conditions will cause it to sublime. For water, the triple point pressure is extremely low (about 0.006 atmospheres), which is why ice can sublime into vapor on a dry, cold day or in the vacuum of space.
For dry ice (solid carbon dioxide), the triple point pressure is 5.11 atm, which is significantly higher than normal atmospheric pressure (1 atm). Since the ambient pressure is below this threshold, solid carbon dioxide cannot melt into liquid carbon dioxide at standard atmospheric pressure. Instead, it sublimates into carbon dioxide gas at a temperature of -78.5 °C. The molecular mechanism involves the solid’s vapor pressure exceeding the surrounding ambient pressure, allowing molecules to escape the crystal lattice.
Common Examples and Practical Applications
A common example of sublimation is the use of dry ice for cooling or theatrical fog effects, where solid carbon dioxide converts into a fog of cold gas. Other substances that readily sublime include iodine and naphthalene, the chemical once used in mothballs that slowly disappears over time. The diminishing size of ice cubes left in a freezer over long periods is also an everyday instance of water ice subliming.
Sublimation is a valuable tool in chemistry and industry, particularly for purification and preservation. In the laboratory, chemists use it to separate volatile solids, such as camphor or ferrocene, from non-volatile impurities. The impure solid is heated under reduced pressure, causing the desired compound to sublime and then deposit as a pure solid on a cooled surface, called a cold finger.
Another application is freeze-drying, or lyophilization, a process used to preserve sensitive materials. The material is first frozen and then placed under a vacuum; the low pressure induces the frozen water (ice) to sublime out of the material, leaving behind a dry, structurally intact product. Dye-sublimation printing also uses heat to transform solid dyes into a gas, which then bonds with synthetic fabrics or specialized paper to create durable images.