Phase Changes

Phase Change Definition

Matter can exist in several forms called phases or states. These are:

1. Solids

2. Liquids

3. Gases

The amount of energy within a species determines its phase. Below is a diagram showing what each phase looks like:

Fig. 1 - Particles are spaced differently in each state of matter

Here's a breakdown of what each illustration means:

• Solid particles are held tightly in fixed positions and are in constant contact with one another. They do have enough energy to vibrate in place.
• Liquid particles are still close together and in constant contact, but have enough energy to switch positions with each other.
• Gas particles have enough energy to move freely. They only come in contact occasionally when they collide with each other.

When species either gain or lose enough energy, they can go through a phase change.

Since these points are where a phase change occurs, both phases (solid/liquid or liquid/gas) can exist at that temperature. We will talk about this in more detail later.

Sublimation Phase Change

While most phase changes involve a liquid at some point, there are a couple of phase changes that skip the liquid phase entirely. One of them is known as sublimation.

You've probably seen sublimation in action if you have ever seen "dry ice" (solid CO₂). At room temperature, dry ice sublimes from a solid to a gas, which is why it appears as smoke.

Phase Change Examples

Now that we've learned the basics of phase changes, let's look at some examples. There are 6 phase changes in total:

• Melting (solid \( \rightarrow \) liquid)

  • Ex: Snow melting in the sun

• Solidifying/Fusion (liquid \( \rightarrow \) solid)

  • Ex: Putting water in the freezer to make ice

• Boiling (liquid \( \rightarrow \) gas)

  • Using a humidifier to turn water into water vapor

• Condensing (gas \( \rightarrow \) liquid)

  • Water vapor condenses into clouds

• Sublimation (solid \( \rightarrow \) gas)

  • Dry ice becomes carbon dioxide

• Deposition (gas \( \rightarrow \) solid)

  • Water vapor turns to frost in the winter

Phase Change Diagram

Interestingly, when a phase change occurs, the temperature doesn't change. Instead, whatever heat is being added\subtracted is going into making the phase change happen.

Below is the phase change diagram for water:

The melting point of water is 0 °C, and you can see that once that temperature is reached, it flat-lines for a bit. The heat that is being added is used to move the ice particles apart until they become water (i.e., liquid phase).

The same process occurs at 100 °C, the boiling point of water. The heat energy being added is pulling apart the water molecules, so they can become steam (i.e., gaseous phase).

When the reverse is occurring (liquid to solid, gas to liquid), the temperature is still unchanged as the phase change h. Heat/energy is being released, so the particles end up closer together.

As mentioned previously, at the boiling point/melting point, both phases can exist at that temperature. Once you hit the needed temperature, a certain amount of heat needs to be added for a complete phase change. Before that happens, there exists either a solid+liquid or a liquid+gas, as one is being converted into the other.

The temperature at which something boils/melts is dependent on pressure. When we refer to the boiling/melting point of water, we are referring to these points at 1 atm (atmosphere at sea level). However, the boiling/melting point will change depending on the pressure.

Below is a phase diagram for CO₂ (carbon dioxide).

Fig. 3 - Phase changes for carbon dioxide.

The border between each phase represents where that phase change will occur. At standard pressure (1 atm), and room temperature (~20 °C), carbon dioxide is a gas/vapor, but at very low temperatures it is solid. This explains why dry ice sublimes instead of melts The graph above also shows us two important points:1. The triple point2. The critical point

Essentially, the triple point is a crossroads, where you can choose any direction (to get to any phase). Let's say we had a sample of dry ice at 5.2 atm and -100 °C. As the temperature increases and reaches -57 °C, the dry ice will start to convert into both liquid and gaseous CO₂.

If the temperature continues to increase, the dry ice will be completely converted into a gas. However, if the pressure increases, the dry ice would instead become a liquid.

The critical point is essentially the "endpoint" for the boiling point curve (border between liquid and gas). After this point, boiling cannot occur. The species has combined characteristics of a liquid and gas, so it is difficult to differentiate between the two. Converting from a liquid to a gas cannot happen, since the species isn't a liquid or a gas anymore.

Energy of Phase Changes

Each phase change has its own energy associated with it. For all molecules/elements, it takes a set amount of energy for a phase change to occur. The heat required for boiling to occur is called the heat of vaporization (ΔH_vap), while the heat required for melting is the heat of fusion (ΔH_fus).

While the magnitude is the same, the signs of each are dependent on the direction of the phase change. If heat is being added, ΔH_vap and ΔH_fus are positive, while they are negative if heat is being released.

These values are dependent on the interactions between molecules. If molecules have strong forces between them, it will take more energy to change their state.

Here is a table of some common heats of vaporization/fusion values.

The total heat required is based on the mass of the species. The formula for the heat change is:

$$ \text{heat}=m*\Delta H_{x} \quad or \quad \text{heat}=n*\Delta H_{x} $$

Where m is the mass of the species, n is the number of moles of the species and \(\Delta H_{x}\) is either the heat of vaporization or the heat of fusion

Let's try an example problem:

Let's do one more example using the other formula:

Endothermic Phase Changes

As we mentioned earlier, phase changes are caused by either a gain or release of heat. Phase changes caused by a gain of heat are called endothermic phase changes (ΔH > 0). These are:

1. Boiling (liquid --> gas)

2. Melting (solid --> liquid)

3. Sublimation (solid --> liquid)

Exothermic phase changes are the opposite. Heat is released during these phase changes (ΔH < 0). These are:

1. Condensing (gas --> liquid)

2. Solidifying (liquid --> gas)

3. Deposition (gas --> solid)