It will be really sublime if my protein samples make it cross-country in time. I hope they reach their destination before all the dry ice is lost through sublimation! I’m preparing for my first time shipping protein samples. I spent a lot of time purifying and preparing these samples so I really want them to make it there safely. I’m not just worried about them getting lost in the mail, I’m also worried about them getting damaged – not from dropping – but from degradation. So no “fragile” sticker could help me here – instead I turned to dry ice, whose long-lasting super-cold-yet-dry-ness will hopefully keep it safe. And I did a little research to figure out what makes the solid form of carbon dioxide (CO₂) we affectionately call “dry ice” so magical…

If you Google dry ice you’re likely to find lots of little demo experiments showing dry ice used to make fog and a brief snippet of an explanation like “dry ice is the solid form of carbon dioxide (CO₂), which, through a process called sublimation, goes directly from a solid to a gas” – those videos are cool and all, but those little explanations don’t do it justice. So, I hope you like the video I made, but if you’re interested in learning a bit more, this post is for you!

The path we’re probably all familiar with is a solid melts into a liquid which can evaporate into a gas (and in the reverse direction, a gas condenses into a liquid which freezes/crystallizes into a solid). But turns out there’s a “shortcut” called sublimation in which a solid changes directly into a gas, and this is exemplified by the solid form of carbon dioxide, affectionately known as “dry ice” because it bypasses the whole wet part. 

In order to understand what’s going on it’s important to keep in mind that solid, liquid, and gas are all just “states of matter” of the same matter. Take water, for instance. Each molecule of water has 2 hydrogen atoms and 1 oxygen atom, giving it the formula H₂O. And that same molecule of water can be present in a solid (ice), a liquid (water), or a gas (water vapor). You can predict which form its likely to be at under given conditions by looking at a phase diagram, which plots temperature (on the x-axis) against pressure (on the y-axis).

Each molecule “wants” to be free (have maximum entropy (randomness/disorder) so they can move around and do their own thing). But they need sufficient energy (supplied by heat) and low enough pressure to escape. When those molecules don’t have much energy (i.e. they’re at a low temperature) they stay stuck in one place, only able to vibrate, and we call this a solid. When the temperature increases, they get some more energy and become able to slide around a bit – but they can’t get too far because they’re attracted to other molecules and don’t have enough energy to permanently break free from them. But give them some more energy and they *can* break free, and escape as a gas, a state where each molecule has lots of energy and can roam around without having to worry about getting sidetracked by other molecules (for the most part). 

What’s there to be sidetracked by? What are molecules anyway and what makes them stand out on the dating scene? Molecules are made up of atoms (think individual carbons, oxygens, hydrogens, etc.) which have positive atomic nuclei consisting of positively-charged protons (and neutral neutrons) surrounded by a cloud of negatively-charged electrons. Pairs of electrons are shared with other atoms to form the strong covalent bonds that hold molecules together – one pair for single bonds like you see in water, and 2 pairs for double bonds like you see in CO₂.

But that’s not the main difference between water and carbon dioxide – with carbon dioxide, instead of dealing with 1 oxygen attached to 2 hydrogens, you’re dealing with 2 oxygens attached to 1 carbon, so CO₂. And instead of being attached in a mickey-mouse-like configuration like you see with water, the molecules in CO₂ lie in a line. And even though oxygen is more electronegative (electron-hogging) than carbon, since they’re on opposite sides, the oxygens pull on the carbon’s electrons evenly in both directions, giving you a “non-polar” molecule. This is in stark contrast with water, where the oxygen pulls on the diagonally-placed hydrogens’ electrons to create a charge imbalance where the oxygen-side of the molecule is partly-negative and the “hydrogen side” is partly positive. Since opposite charges (even partial ones) attract, water is really sticky.

But CO₂ doesn’t have this stickiness situation. So it’s easier to escape other CO₂ molecules and thus doesn’t get “stuck” in the liquid state. It just needs high enough temperature and low enough pressure. I kinda glossed over the whole pressure issue before, but it’s super important, because pressure basically “pushes molecules together” so they need more energy to “push back” and escape. And pressure often also comes with increasing concentrations of molecules, so there’s more to bump into and get sidetracked by when you’re trying to escape. 

Carbon dioxide only gets stuck in the liquid state when it’s at really high pressures – like in fire distinguisher canisters (ever wonder why those canisters are so sturdy? It isn’t just to keep them fire-safe!)

A cool look at what’s going on can be found when you stick dry ice in liquid water – as the dry ice absorbs heat from the liquid, it sublimes (goes from solid directly into a gas). But it’s still surrounded by water – and the water molecules are in interconnected attractive networks with other water molecules, so it’s not like a few CO₂ molecules can just casually snake their way through the water to the water-air interface and out into the air. Instead, they have to team up in “bubbles.” As CO₂ gas forms, the gas molecules, trapped in a watery sea, form bubbles. As more and more CO₂ gas forms, there are more and more gas molecules banging on the bubble walls, pushing against the water and helping the bubbles rise to the surface and escape into the air.

So you can see bubbles form, rise, and then, wait, is that fog I see??? As in condensed water vapor? Yup. Where’d that come from???? NOT the air! Instead, the fog you see comes from the water that the ice was dropped into. We’ve been focusing on the CO₂ wanting to escape, but water does too, remember, so water molecules at the water/CO₂ bubble interface can, if they get sufficient energy, vaporize into the bubble and then hitch a ride out with the CO₂. But water’s stickier than CO₂, remember, so it needs more energy to stay a gas, so once it’s out in the open, it can easily condense back into liquid form, leading to the fog you see. 

So you have CO₂(s) -> CO₂(g) which form bubbles in the water

And then liquid water molecules at the water-bubble interface vaporize into the bubble: H₂O(l, in bulk water) -> H₂O(g, in bubble). 

But the bubble is cold, so once in there, it condenses into water vapor: H₂O(g, in bubble) -> H₂O(l, in bubble). 

And when those interface water molecules left the interface to enter the bubble, they freed up space for new water molecules to get a chance at escape. So those molecule vaporize in too. So you get increasing fogginess in your bubbles.

And then when the bubbles make it to the surface, you get more fogginess in the air!

Pretty cool, right? Well, theater and Halloween aficionados think so too. So dry ice is commonly used to produce an eerie fog effect. 

What’s not so cool is when dry ice is so cool (CO₂ sublimes at -78.5 °C (-109.3°F)) that it gives you freezer burn. So you don’t want to touch it. Instead, if you’re going to work with it, it’s important to wear proper personal protective equipment (PPE) like freezer gloves and safety glasses. And, to make things even safer, dry ice companies often cell dry ice in the form of pellets that look kinda like rabbit poop – their curved shape means there’s less surface on any “side” to get stuck to you. 

more on topics mentioned (& others) #365DaysOfScience All (with topics listed) 👉

Leave a Reply

Your email address will not be published.