Heat of Vaporization: Definition & Equation

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• 0:05 What Is Heat of Vaporization?
• 0:47 Example
• 3:13 By Substance
• 3:43 Equation
• 5:02 Lesson Summary
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Lesson Transcript
Instructor: Jodie Stackhouse
In this lesson, you'll learn about the heat required to transform a liquid to its gas phase. Specifically, we'll explore the factors affecting heat of vaporization, including how to estimate it and how it changes between different substances.

What Is Heat of Vaporization?

Substances exist in three different phases: solid, liquid and gas. Heating and cooling a substance can transform it from one phase to another. Vaporization is the phase change that occurs when a liquid is transformed to a gas.

Substances require a specific amount of heat to undergo the physical changes necessary to switch phases. The energy or heat consumed per unit mass during the vaporization of a liquid is called heat of vaporization or enthalpy of vaporization. To condense water vapor to its liquid phase, energy must be removed from the gas. The energy per unit mass required to condense water vapor is equal to the heat of vaporization.

Heat of Vaporization Example

Water in a teapot undergoes an increase in temperature when heat is provided by the flame of your stove. Let's say we heat one kilogram of room temperature water to boiling in a teapot. In this graph, the temperature of one kilogram of water is plotted against the amount of heat absorbed.

The graph shows three distinct parts:

• Part I: First the temperature increases from room temperature to 100 degrees Celsius when boiling begins.

• Part II: At this point, energy continues to be absorbed, but the temperature remains at 100 degrees. This is the case when the physical characteristics of water change in order to be transformed to vapor. The consumed heat is utilized to drive those changes.

• Part III: After all the water is transformed to vapor (steam), the energy absorbed by the vapor is used to increase the temperature again.

Molecules must be rearranged and intermolecular forces such as ionic bonds, hydrogen bonds, and London dispersion forces must be broken for a substance to change phases. These changes are represented by the demand of energy. As you can see, the arrangement of molecules in solids is very different to that of liquids and gases. Molecules in a solid typically arrange in a periodic lattice, whereas in liquids, molecules do not have a periodic arrangement. Molecules in a gaseous substance move constantly, making their position hard to track.

Additionally, the strength of intermolecular forces between molecules is significantly higher in solids and liquids when compared to gases. Molecules in a liquid and a solid are close-packed due to strong intermolecular forces. In a gas, however, there are no intermolecular forces between molecules, and the average distance between molecules increases considerably.

The heat absorbed from a burning flame is used to break the intermolecular forces in water and increase the interatomic distance between water molecules. Because energy is being used to break intermolecular forces, there is no increase in temperature when water is boiling.

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