11.11
Phase Transitions: Sublimation and Deposition
Some solids can transition directly into the gaseous state, bypassing the liquid state, via a process known as sublimation. At room temperature and st…
It is commonly observed that regular ice melts under ambient conditions, but dry ice does not; instead, dry ice transitions directly into the gas phase. This transition from solid to gas — without passing through the liquid phase — is known as sublimation.
Generally, compounds that sublimate exhibit weak intermolecular forces in the solid state. In dry ice, or solid carbon dioxide, weak dispersion forces exist between CO2 molecules.
At atmospheric pressure, dry ice remains a solid below negative 78.5 °C. However, at −78.5 °C, the surface molecules acquire enough thermal energy to completely overcome the attractive forces and transform directly into the vapor phase. This is the sublimation point of dry ice.
The amount of energy required to sublimate one mole of a solid is called its molar heat of sublimation or its molar enthalpy of sublimation. As sublimation is an endothermic process, its enthalpy value is always positive.
The reverse of sublimation — that is, the direct transition from vapor to solid — is called deposition. When gas molecules collide with cooler solid surfaces, they lose heat. Multiple collisions result in a significant loss of heat, and the molecules ultimately deposit.
Since deposition involves loss of energy, it is an exothermic phase change with a negative enthalpy value. Although the enthalpy of deposition is negative, its magnitude is the same as the enthalpy of sublimation.
When sublimation occurs in an open system, most sublimed molecules disperse in air and never come back. Consequently, the rate of sublimation is greater than the rate of deposition.
However, in a closed system, a solid–vapor equilibrium is established at the solid’s sublimation point.
The partial pressure exerted by the gas in dynamic equilibrium with its solid is called its vapor pressure. Solids that sublimate have high vapor pressures. Dry ice, for example, has a vapor pressure as high as 56.5 atm at 20 °C.
However, since most solids have low vapor pressures at easily accessible temperatures, sublimation is not common.
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Q1: What is sublimation and how does it differ from regular melting?
Sublimation is the direct transition of a solid to gas without passing through the liquid phase. Unlike melting, which converts solid to liquid, sublimation completely bypasses the liquid state. Dry ice exemplifies this process, transforming directly into carbon dioxide vapor at −78.5 °C. Sublimation occurs in solids with weak intermolecular forces that cannot maintain the solid structure when sufficient thermal energy is applied.
Q2: Why do solids with weak intermolecular forces sublimate more readily?
Solids exhibiting weak intermolecular forces require less energy to overcome attractive forces between molecules. In dry ice, weak dispersion forces between CO2 molecules allow surface molecules to escape directly into the vapor phase when they acquire sufficient thermal energy. Stronger intermolecular forces would require more energy, making the solid more stable and preventing sublimation under normal conditions.
Q3: What is the molar heat of sublimation and why is it always positive?
The molar heat of sublimation, or molar enthalpy of sublimation, is the energy required to convert one mole of solid directly into gas. This value is always positive because sublimation is an endothermic process requiring energy input to overcome intermolecular attractions. The energy must be supplied to break bonds holding molecules in the solid state and allow them to escape as gas.
Q4: What is deposition and how does it relate to sublimation?
Deposition is the reverse of sublimation—the direct transition of gas to solid without forming liquid. When gas molecules collide with cooler solid surfaces, they lose heat and deposit onto the surface. Frost formation exemplifies deposition. While sublimation is endothermic with positive enthalpy, deposition is exothermic with negative enthalpy of equal magnitude.
Q5: Why does sublimation occur more readily in open systems than closed systems?
In open systems, sublimed molecules disperse into air and rarely return, making sublimation rate exceed deposition rate. In closed systems, a solid-vapor equilibrium establishes at the sublimation point, where sublimation and deposition rates balance. This equilibrium creates a dynamic state where gas molecules continuously deposit back onto the solid surface.
Q6: What is vapor pressure and why do subliming solids have high vapor pressures?
Vapor pressure is the partial pressure exerted by gas in dynamic equilibrium with its solid at the sublimation point. Solids that sublimate have high vapor pressures because their weak intermolecular forces allow molecules to escape easily into the gas phase. Dry ice, for example, exhibits a vapor pressure of 56.5 atm at 20 °C, far exceeding most ordinary solids.
Q7: How can the enthalpy of sublimation be estimated using other phase transitions?
Sublimation can be modeled as sequential melting followed by vaporization using Hess's Law. The enthalpy of sublimation approximately equals the sum of the enthalpy of fusion and enthalpy of vaporization. This approach works because converting solid to gas requires completely overcoming intermolecular attractions, whereas melting only partially overcomes them.