Energy changes accompany changes of states. The symbol ΔH represents the change in heat content that accompanies each phase transition, known as enthalpy.
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A positive ΔH means that heat is added to the substance and heat is absorbed from the surroundings. Changes with a positive ΔH are endothermic. Changes in states that have a negative ΔH means that heat is lost from the substance and released to the surroundings.
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These phase changes are exothermic.
|
Change of State
|
Energy Change
|
ΔH
|
|
Solid to liquid
(fusion, melting)
|
Heat absorbed, endothermic
|
+
|
|
Liquid to solid
(crystallization, freezing)
|
Heat released, exothermic
|
-
|
|
Liquid to gas
(vaporization, evaporation)
|
Heat absorbed, endothermic
|
+
|
|
Gas to liquid
(condensation)
|
Heat released, exothermic
|
-
|
|
Solid to gas
(sublimation)
|
Heat absorbed, endothermic
|
+
|
|
Gas to solid
(deposition)
|
Heat released, exothermic
|
-
|
Figure 3. Changes of state and accompanying energy changes.
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In order to change a liquid into a gas, energy must be added. This addition of energy is called the heat of vaporization, making it endothermic.
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The reverse process is called condensation and it releases energy into the environment, making it exothermic.
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This energy released is also known as the heat of condensation. The amount of heat transferred during the vaporization and condensation is exactly the same. The same is true for the heat of melting and freezing and the heat transfers in sublimation and deposition.
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Water requires a lot of energy to boil because water molecules have very strong forces of attraction. The hydrogen bones in water molecules cause water molecules to be highly attracted to one another.
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These “forces must be overcome before water molecules can escape the gas phase” and become independent of the other water molecules.
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This is also why a large amount of energy is released when steam condenses.
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Often steam burns are more serious than hot water burns.
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Gibbs Free Energy
Gibbs free energy (ΔG) describes the energy associated with a chemical reaction that can be used to do work. The free energy of a system is the sum of its enthalpy (ΔH) plus the product of the temperature (Kelvin) and the entropy (ΔS) of the system.
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Entropy (ΔS) is the measure of disorder in a system. Gas particles have random motion and therefore have high entropy values. Liquids have much lower entropies and solids even lower. Solutions have higher entropies than pure liquids because their particles are more separated and random.
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Gibbs free energy is a balance between enthalpy and entropy. That is, the enthalpy of melting is independent of temperature. The entropy is also independent of temperature, but since in Gibbs free energy equation ΔG = ΔH – TΔS, then the phase that has higher entropy will be the one that exists at higher temperature.
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Alternatively, the one with lower entropy will be the one that exists at lower temperatures. Processes that have a -ΔG are spontaneous and process that have a +ΔG are non-spontaneous.
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Matter will change spontaneously to minimize its free energy. One analogy for this is the fact that a ball on a rap will roll downward, thus decreasing its potential energy.
There is also a particular temperature where ΔH is equal to TΔS, and therefore there is no spontaneous tendency to go in either direction, and both phases would be in equilibrium indefinitely.
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Superheated and Supercooled Liquids
Superheated and supercooled liquids illustrate that water can exist at temperatures below 0
o
C and above 100
o
C. They are not the lowest energy state, but they still exist. They will spontaneously change phase when given the chance.
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