Real vs Ideal - Do Students "Get It"?

The first time I taught an Honors chemistry class, I noticed there was one additional topic in our unit on gases that differed from our General chem and Concepts chem classes: ideal vs. real gas behavior. When I thought about what made this specific topic better suited for Honors, my first impression usually revolved around the ideas of the math being more difficult (Van der Waals equation) or it being too theoretical for the other classes to grasp. Given how I taught it at the time, it was reasonable to assume these things.

Within the next couple of years, as I began to better understand this topic, I started to realize that there was actually a beautiful opportunity lurking in the shadows that I thought more of my students should be given a chance to experience. That opportunity centered around new observations that would contradict our current models of gas behavior that we had worked so hard to produce and had consistently provided us with the fruits of our labor each time we deployed them--until now. 

Given the fact that most real gases tend to behave ideally under typical conditions, it is understandable how teaching ideal vs. real gas behavior can easily become a bit too theoretical. Sure, I can give students word problems involving extremely low temperatures or high pressures and ask them to predict or even calculate how the pressure of a gas can deviate depending on whether it is treated as an ideal or real gas, but it always felt like something was missing. For example, when working to develop the typical gas laws, we actually interacted with, observed, and measured the properties of a gas. Students could mess around with the volume of gas in a syringe or heat up a sample of gas in a test tube and see how these changes would affect other variables. For some engaging demos, I could crush a can, make an egg slide through a small opening, show them the fire syringe, or even inflate a balloon within a flask. But with the topic of ideal vs. real gases, how was I supposed to engage students with a similar kind of mystery that was not just some fictional scenario on a piece of paper?

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That's when I came across Tom Kuntzleman's ChemEdX blog post about The Case of the Misbehaving Balloon. In a series of several posts set up as chemical mysteries, one of his mysteries involves placing several balloons in liquid nitrogen. What separates Tom's approach to this demo from the numerous online videos of balloons placed in liquid nitrogen is the setup. Like any good magician, there is something he knows about the balloons that we don't. So, when the balloons are placed in the liquid nitrogen, our minds immediately anticipate the decrease in volume that occurs. But as this process plays out, we begin to notice that one of the balloons does not shrink nearly as much as the others; almost like it's misbehaving. Tom really sells this by verbalizing his apparent confusion and tries to resolve the issue by pouring more liquid nitrogen on the stubborn balloon--no luck. It simply won't budge! Then, at the end of the video he reels you in with a challenge by stating,

"Well, if you know your chemistry, you'll be able to explain to me what's going on here."

Instantly, I'm hooked. Given the fact that I'm hooked, I immediately begin to realize what an engaging learning opportunity this could be for my students. This was the achoring phenomenon I was looking for to make the concept of real vs. ideal gases less theoretical and more practical. 

Fortunately, Tom followed his original blog post with a fairly detailed solution to this mystery. In addition to informing us that the misbehaving balloon was actually full of helium while the others were full of air, what I love most about Tom's explanation of this demo is that he incorporates a quantitative approach by using the ideal gas law to predict what we should expect the volume of the balloon to shrink to. Students can easily perform this same calculation. Then, when several of the balloons shrink to a volume of nearly zero, for the first time students experience an instance where our trusty model of describing gas behavior actually fails. In other words, an opportunity naturally arises for a need to rethink our model and account for this new behavior--beautiful! 

As Tom points out, the funny thing about the demo is that even though the balloon that doesn't shrink nearly as much as the others is labeled as "misbehaving," it is actually the other balloons that are misbehaving. The balloons should have decreased to about 25% of their original volume but nearly all of them decreased to a volume that resembles something closer to a volume zero. Whereas the behavior of the balloon that seemingly refused to continuously shrink like the others more closely aligns with the ideal gas law prediction of a decrease in volume to about 25% of its original size. 

Since Tom did such a great job explaining the solution to the mystery in a clear and concise way, I won't simply repeat what he already elequently did and I would direct you to his previous blog post on this topic. That being said, I wanted to share a few things about how I have interweaved this demo into my classroom.

Use as an anchoring phenomenon

As mentioned before, students already had numerous experiences gathering overwhelming evidence suggesting that the ideal gas equation is a pretty good model for making quantiative predictions about gas behavior. Once established, introducing this demo has served as a great opportunity to show students that even the best scientific models may need some tweaking when confronted with new evidence. The demo can function as an anchor that drives all subsequent conversations back to explaining why we saw what we saw. I remind students of the assumptions we make about ideal gases. Specifically, the ones about gas particles having no attraction to each other and that gas particle volume is essentially zero. Before students have even heard of formal names for the attractive forces at play, I lead students along a path of reasoning that allows them to naturally arrive at the idea that in order for us to account for what we saw, attractions between gas particles must be present to some degree. Because this goes against a key assumption of an ideal gas, there is now a need for us to recognize that some gases do not behave ideally in all circumstances--let's call them real gases! Once agreed upon, we begin specualiting how we might tweak the ideal gas equation to account for these attractions. With proper guidance, we arrive at a consensus that some kind of correction factor must be included in the equation. In my upper-level class, this is where we begin to actually derive Van der Waals equation. By the end of our discussion, we have a new model that can account for real gas behavior and can therefore be applied when the ideal gas equation doesn't quite agree with reality.

Use as a formative assessment

If the concept of real vs. ideal gas behavior has already been established, this demo can easily serve as a quick way to evaluate whether students actually understand what it means to behave ideally. I give them the initial and final temperature conditions, they predict the expected decrease in volume, and they are tasked with identifying which balloon behaves more like an ideal gas. The benefit of this is that it can easily allow the teacher to gain a sense of how many students truly understand the concept compared to those whose understanding is typically camouflaged by a more algorithmic thinking that allows them to hide behind the fact that they happen to be good with making calculations. Moments like this are a good reminder that quantitative success does not always mean they actually understand the implications of what their calculations really mean. 

Embed a video of the demo and use as test question

Since I have been teaching in a hybrid model, this was the first time I chose to include a video of this demo within a summative unit test. Here is a screenshot of that very question from our test this year. 

Figure 1: Screenshot of question taken directly from my unit test on gases. Hyperlink takes students to Tom's original mystery video.

It is worth mentioning that not everyone readily has access to liquid nitrogen. However, it is worth looking into your area to for any local suppliers. For example, even though Tom lives in a different state than me, we both got our liquid nitrogen from Airgas, a nationwide distributor of industrial, medical, and specialty gases. If you would like to see if an Airgas store is within your area, click here for a map of their locations. I was surprised that I could get several liters of liquid nirogen for roughly $3.50/liter.

Video 1: Liquid N₂ - Real vs Ideal Gas, ChemEd X Vimeo Channel (5/20/21)

For a more informal overview of how I typically do this demo, I made a video of the entire process. View it above if you are looking for a bit more clarity.

Even if you can't physically do the demo in your classroom due to lack of supply, storage or safety concerns (if you have not been trained to use liquid nitrogen, you should not try this yourself), you can always utilize Tom's original mystery video and embed it within your lesson however you want. It is perfect for students since Tom doesn't supply any answers during the video and none of the Youtube comments over the past six years have provided a single answer. 

Whatever you decide, I hope this can serve as a useful supplement to a topic that can spark some great discussions!