11.7

Phase Transitions: Vaporization and Condensation

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In a liquid, when the thermal motions of molecules overcome the intermolecular forces holding them together, the molecules break free and enter the gas state. This transition from liquid to gas is known as vaporization and can occur under two conditions: non-boiling or boiling.

Under the non-boiling condition, vaporization occurs only at the surface and below the liquid’s boiling point. This is called surface evaporation and occurs without the formation of vapor bubbles in the bulk liquid.

In contrast, if vaporization occurs at the liquid’s boiling point, vapor bubbles form in the bulk liquid, and the process is called boiling. Boiling is not a surface phenomenon and occurs at all points inside the liquid.

Vaporization is a temperature-dependent endothermic process: the greater the heat supplied, the higher the vaporization rate.

The amount of energy required to vaporize one mole of a liquid is called its molar heat of vaporization or its molar enthalpy of vaporization. As vaporization is an endothermic process, its enthalpy value is always positive.

Intermolecular forces influence the molar enthalpy of vaporization. For instance, due to the strong network of hydrogen bonds between water molecules, one mole of water requires a substantial amount of heat energy — around 40.65 kilojoules — to transform into water vapor.

In comparison, the weaker dipole–dipole forces between acetone molecules can be overcome with merely 31.3 kilojoules per mole of heat energy.

The reverse of vaporization, that is, the transition from gas to liquid, is called condensation.

When gas molecules collide with cooler liquid or solid surfaces, they lose heat. Multiple collisions result in a significant loss of heat, and the molecules ultimately condense. Condensation is, therefore, an exothermic process.

Although the enthalpy of condensation is negative, its magnitude is the same as the enthalpy of vaporization.

When the opposing transitions — vaporization and condensation — occur in a closed system, the system reaches a state of dynamic equilibrium called the vapor–liquid equilibrium.

Overview

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The physical form of a substance changes on changing its temperature. For example, raising the temperature of a liquid causes the liquid to vaporize (convert into vapor). The process is called vaporization—a surface phenomenon. Vaporization occurs when the thermal motion of the molecules overcome the intermolecular forces, and the molecules (at the surface) escape into the gaseous state. When a liquid vaporizes in a closed container, gas molecules cannot escape. As these gas phase molecules move randomly about, they will occasionally collide with the surface of the condensed phase, and in some cases, these collisions will result in the molecules re-entering the condensed phase. The change from the gas phase to the liquid is called condensation.

Vaporization is an endothermic process. The cooling effect is evident after a swim or a shower. When the water on the skin evaporates, it removes heat from the skin and cools the skin. The energy change associated with the vaporization process is the enthalpy of vaporization, ΔH_vap. For example, the vaporization of water at standard temperature is represented by:

Chemical structure of morphine, C17H19NO3, in skeletal formula representation.

The reverse of an endothermic process is exothermic. And so, the condensation of a gas releases heat:

Static equilibrium, ΣFx=0, ΣFy=0, ΣFz=0, force diagram illustrating vector components.

Vaporization and condensation are opposing processes; consequently, their enthalpy values are identical with opposite signs. While the enthalpy of vaporization is positive, the enthalpy of condensation is negative.

Different substances vaporize to different extents (depending on the strengths of their IMFs) and hence display different enthalpy of vaporization values. Relatively strong intermolecular attractive forces between molecules result in higher enthalpy of vaporization values. Weak intermolecular attractions present less of a barrier to vaporization, yielding relatively low values of enthalpies of vaporization.

This text is adapted from Openstax, Chemistry 2e, Section 10.3: Phase Transitions.

Transcript

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In a liquid, when the thermal motions of molecules overcome the intermolecular forces holding them together, the molecules break free and enter the gas state. This transition from liquid to gas is known as vaporization and can occur under two conditions: non-boiling or boiling.

Under the non-boiling condition, vaporization occurs only at the surface and below the liquid’s boiling point. This is called surface evaporation and occurs without the formation of vapor bubbles in the bulk liquid.

In contrast, if vaporization occurs at the liquid’s boiling point, vapor bubbles form in the bulk liquid, and the process is called boiling. Boiling is not a surface phenomenon and occurs at all points inside the liquid.

Vaporization is a temperature-dependent endothermic process: the greater the heat supplied, the higher the vaporization rate.

The amount of energy required to vaporize one mole of a liquid is called its molar heat of vaporization or its molar enthalpy of vaporization. As vaporization is an endothermic process, its enthalpy value is always positive.

Intermolecular forces influence the molar enthalpy of vaporization. For instance, due to the strong network of hydrogen bonds between water molecules, one mole of water requires a substantial amount of heat energy — around 40.65 kilojoules — to transform into water vapor.

In comparison, the weaker dipole–dipole forces between acetone molecules can be overcome with merely 31.3 kilojoules per mole of heat energy.

The reverse of vaporization, that is, the transition from gas to liquid, is called condensation.

When gas molecules collide with cooler liquid or solid surfaces, they lose heat. Multiple collisions result in a significant loss of heat, and the molecules ultimately condense. Condensation is, therefore, an exothermic process.

Although the enthalpy of condensation is negative, its magnitude is the same as the enthalpy of vaporization.

When the opposing transitions — vaporization and condensation — occur in a closed system, the system reaches a state of dynamic equilibrium called the vapor–liquid equilibrium.

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