The Physics Philes, lesson 130: It’s Just a Phase
We’re at the end our our discussion of the thermal properties of matter. So far we’ve been talking about the simplest molecular viewpoint, ideal gases. In analyzing ideal gases, we can ignore the interactions between molecules. That certainly makes things simpler, but it’s these interactions that cause gases to condense into liquids and solids. Looking at the theoretical basis for these phase transitions gets messy and complicated, so we’re not going to do it right now. Instead let’s look at phase transitions in general.
To do that, let’s look at a pT phase diagram:
A phase is the physical state of matter, like solid, liquid, and gas. A phase change or a phase transition is the transition between phases (obviously). Each phase is stable at only certain ranges of temperature and pressure. These transitions usually occur when the substance is at phase equilibrium between two phases. For any given pressure, phase equilibrium occurs at one specific temperature. We can represent this on a phase diagram like the one above. As you can see from the graph, for a given pressure and temperature, only one phase can exist, except for points on the solid lines. At those point, two phases can exist in phase equilibrium.
Take a minute to examine the graph. The lines divide the diagram into three general areas: the solid phase, liquid phase, and gaseous or vapor phase. The red line is called the sublimation curve. Both solid and vapor phases exist there. The green line is the fusion curve, where both solid and liquid phases exist. The blue line represents the vaporization curve, where both liquid and vapor phases exist. Where these three curves meet is called the triple point. It’s the only place on the graph where all three phases can exist at the same time.
Let’s pretend that the phase diagram represents water. Let’s hold the water at some constant pressure below the triple point pressure. Now, let’s raise the temperature of the water. As you heat the water at this constant pressure, you’ll notice that eventually we cross straight into the vapor phase without passing through the liquid phase. We went straight from solid to vapor. This process is called sublimation. At any pressure less than the triple point pressure, a substance – in this case, water – isn’t allowed to have a liquid phase.
The triple point is only one of the important points on our pT phase diagram. Over on the far right side of the diagram is the critical point. As the name suggests, it’s an important point. But what is it? Well, it requires a little explanation.
A liquid-vapor phase transition only occurs at certain temperatures and pressures. Where this curve ends on a pressure-Volume diagram corresponds to this critical point on the pT phase diagram. A substance at a pressure and temperature above the critical point changes smoothly from a gas to a liquid or vice versa instead of going through a phase change.
So that’s weird, right? I mean, it has to go through a phase change, right? Nope. As a substance approaches the critical point, the differences in the physical properties – like, density and viscosity – between liquids and gases become smaller and smaller until at the critical point those differences become zero. The distinction between liquids and solids disappears.
If this is still hard to wrap your head around, you’re not alone. Don’t feel too bad about it. Most things we come into contact with in our daily lives don’t do this at normal pressures. Unless you’re peeking into high-pressure steam boilers on the reg, you’re unlikely to have a lot of real world experience with this behavior.
OK! That’s it for thermal properties of matter. But we’re not done with heat. Oh no, far from it. Next week we’re going to really sink our teeth into the first law of thermodynamics. Heat forever!
Featured image credit: Erik Fitzpatrick via Flickr