Around Halloween you might see a bunch of spooky fog and bubbling (yet cold) cauldrons. And, if you’re like me, you might wonder how these effects are created. And, if you’re really like me, you might investigate further and then write a blog post to explain to people in more detail than they probably wanted to know. But there always are people who, like me, really love to know the “how” behind the “wow” and often have trouble finding that ideal level of detail. So, this post is for all of you fellow proud geeks. For everyone else, spoiler alert – it’s DRY ICE (solid carbon dioxide (CO₂)). But I do hope you’ll stick around for more of the post because there’s some cool stuff (and videos!). 

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). 

so, we have: 

with increasing temp and/or decreasing pressure: solid -> melts to liquid -> evaporates/vaporizes to gas

and, in reverse, 

with decreasing temp and/or increasing pressure: gas -> condenses to liquid -> freezes/crystallizes into a solid

But turns out there’s a “shortcut” called sublimation in which a solid changes directly into a gas, 

so, solid -> sublimates into gas

This strange phenomenon is exemplified by the solid form of carbon dioxide, which is 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. 

That might not seem like that big of a deal…but it is! You see, oxygen is more electronegative (electron-hogging) than carbon and hydrogen. So, when oxygen covalently-bonds bonds with one of those, it pulls the shared electrons towards itself, leaving the oxygen partly negative and the carbon or hydrogen slightly positive. These sorts of *bonds* which involve unequal electron sharing are called “polar covalent bonds” and whether they lead to a “polar *molecule*” (a molecule with unequal distribution of charge) depends on how these bonds are arranged and whether they “cancel out.” 

In water, where you have the Mickey Mouse situation, 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. This is called a “molecular dipole.” Since opposite charges (even partial ones) attract, water is really sticky.

But, in carbon dioxide, where you have a straight line, O=C=O, since the oxygens are on opposite sides, they pull on the carbon’s electrons evenly in both directions, giving you a “non-polar” molecule. So, CO₂ doesn’t have the stickiness situation. 

Since CO₂ is less sticky, it’s easier to escape other CO₂ molecules and thus it 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. 


Note: I’m going to introduce the notation we use to report states of matter – stick a little letter in parentheses after the chemical formula. So, (s) means solid, (g) means gas, (l) means liquid. You might also sometimes see (aq) which stands for “aqueous” which means “dissolved in water.”

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, they pop and 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. 

But what if you want those bubbles to stick around for a bit? Add soap! 

with soap

The reason why bubbles burst is surface tension, which is a sort of “cinching” effect like if you had a drawstring on the surface of something. Imagine a water drop. And remember that water’s really sticky because all those  partly ➖ O’s & partly ➕ H’s are attracted to one another. In a water drop, water molecules in the drop’s INTERIOR are surrounded by other water molecules on ALL SIDES, so they have lots of potential binding partners & are attracted in all directions. And because they’re in *all* directions, these attractive forces cancel out so there’s no NET force.

BUT water molecules on drop’s surface have air on one side instead of water, so they have fewer water binding partners. As a result, they form stronger bonds with the ones they do have. And since they’re not being pulled by water in the air-facing direction, they DO have a net force. This force is at a right angle (perpendicular) to the surface & it pulls the water towards the center of the sphere.

That was a water drop. But now imagine a bubble, like the the pockets of CO₂ gas forming when you drop dry ice into water. So here there’s water all around gas. Before the bubble comes to the surface, there’s gas on the inside and water all around (the opposite situation of our water drop). The gas molecules doing their best to escape by moving around randomly, and thereby putting pressure on that water, forming bubbles and propelling them towards the surface (they don’t “know” where the surface is, but they get there because it’s easier to get there because there’s less water weight pushing down on it as it goes toward the surface). 

When those bubbles reach the surface, they burst because water has too high of surface tension to form bubbles –  it’s kinda like trying to blow up a ballon made of cement. But when soap molecules get in there and break things up a bit, the water molecules are able to spread out more and the bubble can expand. 

And the soap does more than just that….

In a soap bubble what you get is a film where you have 2 layers of soap molecules sandwiching a thin layer of water. This strange arrangement comes about because soap molecules are amphiphilic – they have a hydrophilic part (a polar head) and a hydrophobic part (a nonpolar, fatty tail). They arrange themselves so that their tails face the air (inside or outside the bubble) in order to allow as many heads as possible to hang out with the water. And this kinda insulates the water so that it doesn’t evaporate and “burst your bubble” and it limits the stickiness between the water molecules so the bubble doesn’t just “implode” 

As a result, if you add dry ice to soapy water you can get bubbles that can survive for a time outside the water – until the surface tension gets overwhelmed by the vapor pressure inside the bubble caused by the gas molecules banging on the walls of the bubble as they try to escape. 

In addition to this spook-tacular stuff, we use dry ice for more run-of-the mill lab-y stuff. If you stick some dry ice in a styrofoam box it’ll keep the stuff inside of it cold. Really cold. We commonly use it when shipping things like proteins that we normally keep in ~80°C freezers.

This works because CO₂ sublimes at -78.5 °C (-109.3°F)), which is good for our proteins but also “good” for giving 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. 

This post is part of my weekly “broadcasts from the bench” for The International Union of Biochemistry and Molecular Biology. Be sure to follow them on all the social media places if you’re interested in biochemistry! They’re a really great international organization for biochemistry.⠀

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