Updated April 25, 2017
By Calia Roberts
A phase change, or transition, occurs when a substance undergoes a change in state on a molecular level. In most substances, changes in temperature or pressure result in a substance phase change. There are several processes of phase changes, including fusion, solidification, vaporization, condensation, sublimation and physical vapor deposition.
Fusion occurs when a substance changes from a solid to a liquid. Prior to melting, strong intermolecular bonds or attractions hold the atoms, molecules or ions that comprise a solid substance tightly together in the solid form. Upon heating, the particles gain enough kinetic energy to overcome the bonds that are holding them together and become mobile. This results in the fusion of the substance.
Solidification occurs when a substance changes from a liquid to a solid. While in the liquid state, the particles in a substance possess enough kinetic energy to move around in close proximity to each other. When a drop in temperature occurs, the particles lose their kinetic energy and band together. Gradually, the particles settle into a fixed position, causing the substance to take shape and become a solid.
Vaporization occurs when a substance changes from a liquid to a gas. The molecules in a liquid are in constant motion while staying relatively close together due to intermolecular forces. When an increase in temperature occurs, the molecules' kinetic energy also increases. This increase in temperature allows the molecules to gain kinetic energy and overcome the intermolecular forces, resulting in the vaporization of the substance.
Condensation occurs when a substance changes from a vapor to a liquid. In a vapor, there are molecules with high and low kinetic energy that often collide with surfaces and each other. When molecules with low kinetic energy collide, intermolecular forces cause them to stick together. As temperature decreases, the kinetic energy of the molecules also decreases causing the molecules to stick together and resulting in condensation.
Sublimation occurs when a substance changes from a solid into a gas. Increases in temperature causes the kinetic energy of particles to also increase. This allows the particles to overcome the intermolecular forces and become mobile. Low pressure also increases the particles' kinetic energy. As the particles escape the solid and disperse as a gas, sublimation occurs.
Physical vapor deposition occurs when a substance changes from a gas into a solid. In low-pressure situations, thin films of vaporized materials develop on various surfaces due to plasma sputter bombardment or high-temperature vacuum evaporation.
Phase changes are the transformations from one state of matter to another due to thermodynamics. The processes of phase change between solid and liquid are called melting and freezing. Phase changes between liquid and gas are vaporization and condensation. Phase changes between gas and solid are deposition and sublimation. Phase changes can be spontaneous or non-spontaneous.
We're going to talk about phase changes, going from different forms of matter; for the solid, liquid gases and how they interact with each other and how they change from one phase to another. So let's use this as an examp- as a good diagram to show us how these things interact with each other. Alright, so if we're going to get from solid to a liquid we're actually call that melting which I know we've heard that word many times before and that actually requires energy, we need heat to melt something, so we're going to call that endothermic process meaning that it requires energy or requires heat for that reaction or that that phase that to occur. So if we're going to from liquid the opposite from liquid to solid, we're going to call that freezing which I know we've heard many times before too. That actually releases some sort of energy, it's going to be an exothermic process, meaning it releases energy. These guys are opposite of each other, melting and freezing, of the same thing. Let's go over to liquid and gases. If we're going from liquid to gas, we're going to call that vaporization; we're going to vaporize that particular liquid. That actually, requires energy as well. We need heat or some sort of energy to make that happen. So we we're going to call that endothermic process. The opposite will be cond- condensation; when we're are condensing something from a gas down to liquid and that's an exothermic process meaning that's going also to release some sort of energy.There are rare instances where substances will go straight from the gas phase to the solid phase. It doesn't happen as often as you probably know but they do happen with different substances so if we're going from the gas phase down to the solid phase, we're actually going to release that sort of energy because we know gas is in high has higher energy than solid phase so we're going to release that energy we're going to call that process deposition. The opposite would be sublimation going from solids to a gas we've seen this probably before when ice or solid CO2 maybe iodine crystals they go from this solid phase straight to the gas phase skipping over the liquid phase that actually it releases sort releases some sort of energy and that we call it endothermic process.Alright so let's actually look at this in a different way this you might see more often in class. This is actually a different a graph describing all those things that we just talked about. Alright so on the x ax- sorry on the y axis we have temperature on the x axis we got x axis we're going to have energy okay so we know in this case we're going to talk about water the phase change of water and we know that below 0 degree Celsius that is in solid phase okay? So as we increase energy, our temperature of that solid is going to increase until we hit 0 degree Celsius which we know it is melting and freezing point s the increase if we increase energy it's going to melt and if we're going from liquid to solid it's going to start freezing but notice the temperature it's not changing even though we're increasing temperature why is that? Well that energy that we're pumping into the into this solid molecule this substance is actually being used to break up those intermolecular forces that are holding it together in a solid so here's the picture water and these blue dots are the hydrogen bonds that are holding it together in a solid so because solids have more hydrogen bonds than liquids that energy is going to be used to break up some of those bonding some of those forces that are holding it together. Then as you go from 0 degree Celsius to 100 degree Celsius we're going to be in a liquid phase all that and all the energy is going to be used to increase the temperature of that particular liquid in this case water. And here we have the same thing we have this plateau and 100 degree Celsius we know that is it's vaporization point or it's con- condensation point again it's flat and again that energy is being used to break apart more of these hydrogen bonds once at 100 degree Celsius these bonds are going to be pretty wear because their energy is being used to break them apart and have them flow around allover the place and then up at higher temperatures above it's always going be in a gaseous phase. If we go straight from the solid to a gas which water doesn't do if it were to we would call that sublimation going to skipping this liquid phase completely, if we're going from gas to solid we're going to call it deposition we're going to complete again skipping that liquid phase so this actually cycle talks about the different phase changes that substances tends to undergo.
- Describe what happens during a phase change.
- Calculate the energy change needed for a phase change.
Substances can change phase — often because of a temperature change. At low temperatures, most substances are solid; as the temperature increases, they become liquid; at higher temperatures still, they become gaseous.
The process of a solid becoming a liquid is called melting [an older term that you may see sometimes is fusion]. The opposite process, a liquid becoming a solid, is called solidification. For any pure substance, the temperature at which melting occurs — known as the melting point — is a characteristic of that substance. It requires energy for a solid to melt into a liquid. Every pure substance has a certain amount of energy it needs to change from a solid to a liquid. This amount is called the enthalpy of fusion [or heat of fusion] of the substance, represented as ΔHfus. Some ΔHfus values are listed in Table 10.2 “Enthalpies of Fusion for Various Substances”; it is assumed that these values are for the melting point of the substance. Note that the unit of ΔHfus is kilojoules per mole, so we need to know the quantity of material to know how much energy is involved. The ΔHfus is always tabulated as a positive number. However, it can be used for both the melting and the solidification processes as long as you keep in mind that melting is always endothermic [so ΔH will be positive], while solidification is always exothermic [so ΔH will be negative].
| Water [0°C] | 6.01 |
| Aluminum [660°C] | 10.7 |
| Benzene [5.5°C] | 9.95 |
| Ethanol [−114.3°C] | 5.02 |
| Mercury [−38.8°C] | 2.29 |
What is the energy change when 45.7 g of H2O melt at 0°C?
Solution
The ΔHfus of H2O is 6.01 kJ/mol. However, our quantity is given in units of grams, not moles, so the first step is to convert grams to moles using the molar mass of H2O, which is 18.0 g/mol. Then we can use ΔHfus as a conversion factor. Because the substance is melting, the process is endothermic, so the energy change will have a positive sign.
Without a sign, the number is assumed to be positive.
Test Yourself
What is the energy change when 108 g of C6H6 freeze at 5.5°C?
Answer
−13.8 kJ
During melting, energy goes exclusively to changing the phase of a substance; it does not go into changing the temperature of a substance. Hence melting is an isothermal process because a substance stays at the same temperature. Only when all of a substance is melted does any additional energy go to changing its temperature.
What happens when a solid becomes a liquid? In a solid, individual particles are stuck in place because the intermolecular forces cannot be overcome by the energy of the particles. When more energy is supplied [e.g., by raising the temperature], there comes a point at which the particles have enough energy to move around but not enough energy to separate. This is the liquid phase: particles are still in contact but are able to move around each other. This explains why liquids can assume the shape of their containers: the particles move around and, under the influence of gravity, fill the lowest volume possible [unless the liquid is in a zero-gravity environment — see Figure 10.16 “Liquids and Gravity”].
The phase change between a liquid and a gas has some similarities to the phase change between a solid and a liquid. At a certain temperature, the particles in a liquid have enough energy to become a gas. The process of a liquid becoming a gas is called boiling [or vapourization], while the process of a gas becoming a liquid is called condensation. However, unlike the solid/liquid conversion process, the liquid/gas conversion process is noticeably affected by the surrounding pressure on the liquid because gases are strongly affected by pressure. This means that the temperature at which a liquid becomes a gas, the boiling point, can change with surrounding pressure. Therefore, we define the normal boiling point as the temperature at which a liquid changes to a gas when the surrounding pressure is exactly 1 atm, or 760 torr. Unless otherwise specified, it is assumed that a boiling point is for 1 atm of pressure.
Like the solid/liquid phase change, the liquid/gas phase change involves energy. The amount of energy required to convert a liquid to a gas is called the enthalpy of vaporization [or heat of vaporization], represented as ΔHvap. Some ΔHvap values are listed in Table 10.3 “Enthalpies of Vaporization for Various Substances”; it is assumed that these values are for the normal boiling point temperature of the substance, which is also given in the table. The unit for ΔHvap is also kilojoules per mole, so we need to know the quantity of material to know how much energy is involved. The ΔHvap is also always tabulated as a positive number. It can be used for both the boiling and the condensation processes as long as you keep in mind that boiling is always endothermic [so ΔH will be positive], while condensation is always exothermic [so ΔH will be negative].
| Water [100°C] | 40.68 |
| Bromine [59.5°C] | 15.4 |
| Benzene [80.1°C] | 30.8 |
| Ethanol [78.3°C] | 38.6 |
| Mercury [357°C] | 59.23 |
What is the energy change when 66.7 g of Br2[g] condense to a liquid at 59.5°C?
Solution
The ΔHvap of Br2 is 15.4 kJ/mol. Even though this is a condensation process, we can still use the numerical value of ΔHvap as long as we realize that we must take energy out, so the ΔH value will be negative. To determine the magnitude of the energy change, we must first convert the amount of Br2 to moles. Then we can use ΔHvap as a conversion factor.
Because the process is exothermic, the actual value will be negative: ΔH = −6.43 kJ.
Test Yourself
What is the energy change when 822 g of C2H5OH[ℓ] boil at its normal boiling point of 78.3°C?
Answer
689 kJ
As with melting, the energy in boiling goes exclusively to changing the phase of a substance; it does not go into changing the temperature of a substance. So boiling is also an isothermal process. Only when all of a substance has boiled does any additional energy go to changing its temperature.
What happens when a liquid becomes a gas? We have already established that a liquid is composed of particles in contact with each other. When a liquid becomes a gas, the particles separate from each other, with each particle going its own way in space. This is how gases tend to fill their containers. Indeed, in the gas phase most of the volume is empty space; only about one one-thousandth of the volume is actually taken up by matter [see Figure 10.17 “Liquids and Gases”]. It is this property of gases that explains why they can be compressed, a fact that is considered in Chapter 6 “Gases”.
Under some circumstances, the solid phase can transition directly to the gas phase without going through a liquid phase, and a gas can directly become a solid. The solid-to-gas change is called sublimation, while the reverse process is called deposition. Sublimation is isothermal, like the other phase changes. There is a measurable energy change during sublimation; this energy change is called the enthalpy of sublimation, represented as ΔHsub. The relationship between the ΔHsub and the other enthalpy changes is as follows:
ΔHsub = ΔHfus + ΔHvap
As such, ΔHsub is not always tabulated because it can be simply calculated from ΔHfus and ΔHvap.
There are several common examples of sublimation. A well-known product — dry ice — is actually solid CO2. Dry ice is dry because it sublimes, with the solid bypassing the liquid phase and going straight to the gas phase. The sublimation occurs at temperature of −77°C, so it must be handled with caution. If you have ever noticed that ice cubes in a freezer tend to get smaller over time, it is because the solid water is very slowly subliming. “Freezer burn” isn’t actually a burn; it occurs when certain foods, such as meats, slowly lose solid water content because of sublimation. The food is still good but looks unappetizing. Reducing the temperature of a freezer will slow the sublimation of solid water.
Chemical equations can be used to represent a phase change. In such cases, it is crucial to use phase labels on the substances. For example, the chemical equation for the melting of ice to make liquid water is as follows:
H2O[s] → H2O[ℓ]
No chemical change is taking place; however, a physical change is taking place.
Heating Curves
A plot of the temperature versus the amount of heat added is known as a heating curve [see Figure 10.18]. These are commonly used to visually show the relationship between phase changes and enthalpy for a given substance.
In Figure 10.18, the solid gains kinetic energy and consequently rises in temperature as heat is added. At the melting point, the heat added is used to break the attractive intermolecular forces of the solid instead of increasing kinetic energy, and therefore the temperature remains constant. After all the solid has melted, once again, the heat added goes to increasing the kinetic energy [and temperature] of the liquid molecules until the boiling point. At the boiling point, once again, the heat added is used to break the attractive intermolecular forces instead of supplying kinetic energy, and the temperature remains constant until all liquid has been turned to gas.
- Phase changes can occur between any two phases of matter.
- All phase changes occur with a simultaneous change in energy.
- All phase changes are isothermal.
Questions
- What is the difference between melting and solidification?
- What is the difference between boiling and condensation?
- Describe the molecular changes when a solid becomes a liquid.
- Describe the molecular changes when a liquid becomes a gas.
- What is the energy change when 78.0 g of Hg melt at −38.8°C?
- What is the energy change when 30.8 g of Al solidify at 660°C?
- What is the energy change when 111 g of Br2 boil at 59.5°C?
- What is the energy change when 98.6 g of H2O condense at 100°C?
- Each of the following statements is incorrect. Rewrite them so they are correct.
- Temperature changes during a phase change.
- The process of a liquid becoming a gas is called sublimation.
- Each of the following statements is incorrect. Rewrite them so they are correct.
- The volume of a gas contains only about 10% matter, with the rest being empty space.
- ΔHsub is equal to ΔHvap.
- Write the chemical equation for the melting of elemental sodium.
- Write the chemical equation for the solidification of benzene [C6H6].
- Write the chemical equation for the sublimation of CO2.
- Write the chemical equation for the boiling of propanol [C3H7OH].
- What is the ΔHsub of H2O? [Hint: see Table 10.2 “Enthalpies of Fusion for Various Substances” and Table 10.3 “Enthalpies of Vaporization for Various Substances”.]
- The ΔHsub of I2 is 60.46 kJ/mol, while its ΔHvap is 41.71 kJ/mol. What is the ΔHfus of I2?
- Melting is the phase change from a solid to a liquid, whereas solidification is the phase change from a liquid to a solid.
- The molecules have enough energy to move about each other but not enough to completely separate from each other.
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- Temperature does not change during a phase change.
- The process of a liquid becoming a gas is called boiling; the process of a solid becoming a gas is called sublimation.