Chemistry Myth: Glass Is A Liquid

I've heard it before, and I keep coming across this idea: That glass is a liquid, it's just exceedingly slow-moving/viscous/lazy and unwilling to get off the couch. I don't blame people for believing it, I myself did for a long time. My science teacher told me this myth and I passed it on. It's just that there's something about glass, we take it for granted but whenever we actually start thinking about it, glass does seem like an amazing substance doesn't it? Granted it doesn't seem so amazing when you try walking straight into a glass door. So it seems credible that this incredibly smooth and ice-like substance might be some kind of exception to the rules of chemistry. Well I don't mean to shatter the mystery, but I promise to replace it with a much more interesting truth.

Glass, as you might already know is made from a mineral, silicon dioxide, which is the principal component of white sand. We heat the sand up, blow it into interesting shapes, and toss it into water to quench, or cool it rapidly (this is the key, as you will see in a minute.)

To explain glass, I need to explain the nature of liquid to solid transition. To do this, I'm going to show you a standard cooling curve for a container of water:

This curve looks similar for many substances, though some substances sublime under normal conditions, which means they turn directly from solid to gas and back. The reason I chose water here is because water is something we are all familiar with, and because I hope to get people to try experimenting at home. We all know water freezes at zero, but it won't go any lower until all of the water has turned to ice in a solution (we think of a mixture of water and ice as a solution in this example, even though they are the same compound.) However, you will notice a little dip in the curve that goes below zero, even though the water is still a liquid. This is liquid that is colder than the compound's melting point. A supercool liquid. How is this possible?

The answer lies in the simple fact that solids are solids because their molecules take on a rigid structure. Look at this rough representation of water sliding around in liquid form:

Now this diagram is somewhat problematic in many ways, but it will do for this explanation. Here we see our old friend the Mickey Mouse model of water. Now the reason hydrogen bonds to oxygen is that oxygen likes to donate an electron to hydrogen, and oxygen is more than happy to give one electron to each of two hydrogen (since hydrogen only has one positive charged proton to cancel out). This is how molecules try to equalize their charges. In nature, electromagnetic charges do their best to cancel each other out. However, oxygen is always hungry, it's so much more massive that hydrogen (about sixteen times as massive) and it has a slight negative charge even when its fully bonded to hydrogen. Hydrogen too, retains a slight positive charge.

Now in liquid water under normal conditions, these slight charges hold the water molecules together enough that most don't zing off into the air (cohesion). Though if you leave it lying around long enough it will evaporate, the water is warm enough that the molecules are bouncing around rather happily and don't stick in any rigid formation.

Now if you cool the water down slowly enough (and you may have had this experience with water, but it is common in alcoholic beverages.) the molecules will slow down very gradually, but they won't fall into any kind of order. They won't slide across each other as much, but they will kind of almost stand still in one place. This is the supercool liquid, there is no real structure to the molecules, they're not arranged in a way that would look like a solid. However, if you take this slowly cooled liquid (and once again, I encourage you to try this at home) and agitate it, it will freeze up and turn into ice almost instantaneously. This is because you have physically forced the molecules to turn, and they freeze up in an arranged structure, usually according to their partial charges, which looks a little like this:


As you can see, the slightly positively charged hydrogen ions are attracted to the slightly negatively charged oxygen ions. This is what a classic crystalline structure looks like.

"But Chemist," You ask, "What does this have to do with glass?"

Well, I'm glad you asked. Glass is like any other liquid when melted. As a mentioned before, it's silicon dioxide, which you may be familiar with in the form of sand, but also quartz. Yet you can't really see through most quartz crystals that well, they tend to let light in one way but not another. Glass doesn't have this problem. This is because unlike the ice in your refrigerator, when we freeze glass (that's the term we use for cooling anything into a solid) it's in an amorphous state, not crystalline. This means it's in that same disordered state I showed you when water is still liquid. This allows light to pass through it in all directions with minimal interference.

How do we get that to happen? It's the quenching process that I said was key earlier. While the glass is still in its plastic state having been melted, we rapidly cool it. This means that we get the temperature down low enough while its still a supercool liquid (which for glass is actually kind of hot) that it freezes instantaneously without rearranging its molecular structure much. Hence glass is not a liquid, it's a solid, albeit an amorphous solid. Its not super viscous, viscosity is a measure of how much molecules are attracted to each other in a liquid, the molecules in glass are locked in place, but they have no particular order.

The exact same thing can happen with water. Cool water rapidly enough and it will turn into amorphous ice, and will have similar glassy properties. However I don't recommend you try cooling water this rapidly at home. My guess is you would need a dangerously cold substance to cool water this rapidly, and I don't want anyone getting frostbite or worse. If you think you can pull it off with household materials you can try it (you might want to work with tiny amounts of water) just don't break out the liquid nitrogen, mmkay?

There you have it, a chemistry myth discredited. Questions? Comments?