7.3 Phase Changes

A phase change is a physical process in which a substance goes from one phase to another. Usually the change occurs when adding or removing heat at a particular temperature, known as the melting point or the boiling point of the substance. The melting point is the temperature at which the substance goes from a solid to a liquid, and it is exactly the same temperature as the freezing point, the temperature at which a substance goes from a liquid to a solid). The boiling point is the temperature at which a substance goes from a liquid to a gas at normal atmospheric pressure. The nature of the phase change depends on the direction of the heat transfer. Heat going into a substance changes it from a solid to a liquid or a liquid to a gas. Removing heat from a substance changes a gas to a liquid or a liquid to a solid.  Adding energy to molecules makes them more faster and become more disorderly.  Removing energy from molecules makes them slow down and become more orderly.

Two key points are worth emphasizing. First, at a substance’s melting point or boiling point, two phases can exist simultaneously. Take water (H2O) as an example. On the Celsius scale, H2O has a melting point of 0°C and a boiling point of 100°C. At 0°C, both the solid and liquid phases of H2O can coexist. However, if heat is added, some of the solid H2O will melt and turn into liquid H2O. If heat is removed, the opposite happens: some of the liquid H2O turns into solid H2O. A similar process can occur at 100°C: adding heat increases the amount of gaseous H2O, while removing heat increases the amount of liquid H2O.

Note

Water is a good substance to use as an example because many people are already familiar with it. Other substances have melting points and boiling points as well.

Second, the temperature of a substance does not change as the substance goes from one phase to another. In other words, phase changes are isothermal, which means “constant temperature”). Again, consider H2O as an example. Solid water (ice) can exist at 0°C. If heat is added to ice at 0°C, some of the solid changes phase to make liquid, which is also at 0°C. Remember, the solid and liquid phases of H2O can coexist at 0°C. Only after all of the solid has melted into liquid does the addition of heat change the temperature of the substance.

For each phase change of a substance, there is a characteristic quantity of heat needed to perform the phase change per gram (or per mole) of material. The heat of fusion (ΔHfus) is the amount of heat per gram (or per mole) required for a phase change that occurs at the melting point. The heat of vaporization (ΔHvap) is the amount of heat per gram (or per mole) required for a phase change that occurs at the boiling point. If you know the total number of grams or moles of material, you can use the ΔHfus or the ΔHvap to determine the total heat being transferred for melting or solidification using these expressions:

heat = n × ΔHfus (where n is the number of moles)
or
heat = m × ΔHfus (where m is the mass in grams)

For the boiling or condensation, use these expressions:

heat = n × ΔHvap (where n is the number of moles)
or
heat = m × ΔHvap (where m is the mass in grams)

Remember that a phase change depends on the direction of the heat transfer. If heat transfers in, solids become liquids, and liquids become solids at the melting and boiling points, respectively. If heat transfers out, liquids solidify, and gases condense into liquids.

Example 4

How much heat is necessary to melt 55.8 g of ice (solid H2O) at 0°C? The heat of fusion of H2O is 79.9 cal/g.

Solution

We can use the relationship between heat and the heat of fusion to determine how many joules of heat are needed to melt this ice:

heat = m × ΔHfus
heat = (55.8 g)(79.9[latex]\frac{\text{cal}}{\text{g}}[/latex])=4,460 cal

Skill-Building Exercise

  1. How much heat is necessary to vaporize 685 g of H2O at 100°C? The heat of vaporization of H2O is 540 cal/g.

 

Table 7.4 “Heats of Fusion and Vaporization for Selected Substances” lists the heats of fusion and vaporization for some common substances. Note the units on these quantities; when you use these values in problem solving, make sure that the other variables in your calculation are expressed in units consistent with the units in the specific heats or the heats of fusion and vaporization.

Table 7.4 Heats of Fusion and Vaporization for Selected Substances
Substance ΔHfus (cal/g) ΔHvap (cal/g)
aluminum (Al) 94.0 2,602
gold (Au) 15.3 409
iron (Fe) 63.2 1,504
water (H2O) 79.9 540
sodium chloride (NaCl) 123.5 691
ethanol (C2H5OH) 45.2 200.3
benzene (C6H6) 30.4 94.1

Looking Closer: Sublimation and deposition

There is also a phase change where a solid goes directly to a gas:

solid → gas   (sublimation)

We encounter sublimation in several ways. You may already be familiar with dry ice, which is simply solid carbon dioxide (CO2). At −78.5°C, solid carbon dioxide sublimes, changing directly from the solid phase to the gas phase:

CO2(s)  →  CO2(g)  at −78.5°C

Solid carbon dioxide is called dry ice because it does not pass through the liquid phase but goes directly to the gas phase at normal atmospheric pressure. Dry ice has many practical uses, including the long-term preservation of medical samples.

Even plain old ice, solid H2O, slowly sublimes at temperatures below 0°C. For example, ice cubes in a freezer may get smaller over time. Although frozen, the solid water slowly sublimes, releasing water molecules into the air in the freezer. Water also sublimes from frozen foods, giving them an unattractive, mottled appearance called freezer burn. It is not really a “burn,” and the food has not necessarily gone bad, although it looks unappetizing. Freezer burn can be minimized by lowering a freezer’s temperature and by wrapping foods tightly so water does not have any space to sublime into.

Deposition is the opposite of sublimation in that molecules leave the gas phase and go directly to the solid phase when they hit a cold object.

gas → solid   (deposition)

The formation of frost on a car windshield is an example of deposition, as is the formation of snowflakes in clouds.  The lacy appearance of frost and snow flakes reflects their molecule-by-molecule deposition on a cold surface.  In contrast, ice looks smooth because a water sample went from liquid to solid in bulk, not molecule-by-molecule.

Concept Review Exercises

  1. Explain what happens when heat flows into or out of a substance at its melting point or boiling point.

  2. How does the amount of heat required for a phase change relate to the mass of the substance?

Answers

Key Takeaway

  • There is an energy change associated with any phase change.

Exercises

  1. How much energy is needed to melt 43.8 g of Au at its melting point of 1,064°C?

  2. How much energy is given off when 563.8 g of NaCl solidifies at its freezing point of 801°C?

  3. What mass of ice can be melted by 558 cal of energy?

  4. How much ethanol (C2H5OH) in grams can freeze at its freezing point if 1,225 cal of heat are removed?

  5. What is the heat of vaporization of a substance if 10,776 cal are required to vaporize 5.05 g? Express your final answer in joules per gram.

  6. If 1,650 cal of heat are required to vaporize a sample that has a heat of vaporization of 137 cal/g, what is the mass of the sample?

  7. What is the heat of fusion of water in calories per mole?

  8. What is the heat of vaporization of benzene (C6H6) in calories per mole?

  9. What is the heat of vaporization of gold in calories per mole?

  10. What is the heat of fusion of iron in calories per mole?

Answers