What state do you live in? Why does it matter? I want to know how much freedom you have! And I want an excuse to talk about one of my favorite molecules – liquid nitrogen! Molecules want to be free -> universe tends towards increasing ENTROPY, or “randomness” (2nd Law of Thermodynamics). Molecules “want” to be gases bc In a gas, each molecule gets to move around freely & do it’s it own thing.
So molecules in a liquid, like a glass of water, try to escape into the gas phase. BUT they need enough energy to overcome the ESCAPE KINETIC ENERGY (the energy required to break free of the bonds “chaining them” to neighboring molecules)
Solid ice, liquid water, & water vapor (gas) are all made up of the same molecule, good ole H2O. The difference between these PHASES or STATES is how tied to each other they are which depends on how much energy they have.
All water molecules are inherently “sticky” bc they’re highly POLAR – their charge is unevenly distributed so that the oxygen (O) is partly negative & the hydrogens (H’s) are partially ➕. The partial charges of neighboring molecules attract each other – INTERMOLECULAR FORCES (IMFs between water molecules lead to the formation of hydrogen bonds (H-bonds). They’re weak individually, but have great strength in numbers!
In a SOLID, intermolecular forces (IMFs) are strong; the molecules are “glued together” &, although they want to escape, they can only vibrate
But add energy (heat) & molecules start vibrating more strongly – some bonds break & the molecules can briefly escape, they try to escape from their neighboring molecules, but the molecules are still close together & on the way out they bump into other molecules that they bind to -> so molecules in a LIQUID are moving around but forming & breaking bonds as they go (kinda like a swing dance) – and a really fast one! bonds are constantly breaking and reforming. When a molecule breaks free it normally forms a new bond (with that same water molecule or a different one) within 0.1ps (picoseconds (1 trillionth of a second!)
But add even more energy and the molecules are moving so fast they don’t have time to get “caught” before they leave -> they escape into the air as a GAS Once they’re in the air and out of the close vicinity they can move far away from other water molecules so they don’t have to worry about getting caught any more
Temperature is a measure of the average amount of energy the molecules have. So the more heat you add, the higher the energy of the molecules and the greater the chance they have to escape – this is why we can evaporate water by putting it on a stove
Above a solution’s boiling point, even molecules in the interior can gain enough energy to escape
But you don’t have boil to have evaporation (which is a good thing for crystallographers using evaporation to supersaturate protein solutions! I don’t think the protein would like that…)
In fact, molecules are always escaping the surface of water, but in a stable system others are getting “pulled in” (CONDENSING) at the same rate -> there’s a DYNAMIC EQUILIBRIUM where the rate of molecules coming & going cancel out, so there’s no net movement
But it’s harder to escape at a lower temperature, so below the boiling point, only molecules at the surface have a chance of escaping
- they have a shorter distance to travel – fewer other molecules to get caught by on their way out
- if the air’s warmer than the liquid they’ll get some energy from it
Water won’t go full-on boil at room temp because it’s above its boiling point. But liquid nitrogen (aka LN2) (LN2) (The abbreviation refers to the fact that the atoms of nitrogen (N) are present in pairs (diatomic)) – that’s a different story. It has a really low boiling point which means you don’t have to add much heat to get it to boil. In fact, the heat of room temp is way higher than it needs.
So when liquid nitrogen is exposed to room temp air, from its perspective its like sticking it on a stove. Except that the heat’s hitting it from the top first, so the boiling can start there without the bubbles having to make their way up. And it’s not just the air that feels scorching hot to the nitrogen – basically anything it contacts including the walls of the containers you put it in for short term use – so it’s bubbles all around. And pretty soon you don’t have any liquid nitrogen left.
You do have a cloud of foggy stuff and that is NOT nitrogen gas. You can’t see the nitrogen gas because it just slips in along with the other gas molecules – N2 is a component of “normal air” anyway. In fact, instead of thinking of nitrogen as a “weird” *liquid* you can think of it as a “weird” *gas.* It’s “usually” a gas (under our normal earthy conditions) so the thing that makes it unique is that we can capture & hold it in the liquid form at all! And water’s super weird because we can easily get it to be a solid, a liquid, & a gas.
And speaking of water -> the cloud you see isn’t nitrogen gas that’s escaped – it isn’t a gas at all – it’s water that’s “unescaped” from the air because the nitrogen takes its heat. And without that heat, the water molecules in the air don’t have enough energy to resist the pull of nearby molecules.
And the water in the air isn’t the only thing liquid nitrogen can take heat from. We commonly use it to “flash freeze” samples (like proteins we’ve purified that are currently dissolved in a liquid) – we take those dancing molecules in that liquid swing dance and turn that swing dance into a freeze dance – stop the “music” by nitrogen taking the heat. And this is freeze dance – not musical chairs – you don’t have time to go find an empty seat – you have to freeze in place. So this prevents water from “sitting” into crystals of ice (orderly lattices) that could mess with your proteins (think busting pipes but on the molecular scale).
So liquid nitrogen’s super useful to biochemists, but if it’s so evaporation-happy, how do we even keep liquid nitrogen liquid? In a super-duperly-well-insulated container called a dewar. Since LN2’s always ditching the L (and we use it a lot) we have to refill this little guy from a big tank & when it’s filling, it reminds me of some Seussian contraption.
This storage dewar looks like R2D2 & it was rescued by WD40 – it used to squeak SOOO badly. It’s really well insulated but it has a loose lid. You need this because nitrogen expands about 175 X when it turns into a gas. Reaching room temp increases its volume another few times – you end up with a gas volume ~650X the liquid volume you started with. So we need to keep liquid nitrogen containers covered enough that they’re kept cold but not airtight so they don’t explode
You definitely should use caution when working with it – insulated gloves, goggles, etc so you don’t get frostbitten. Trust my thumbs, those simple nitrile gloves aren’t very protective! But the reason we’re able to (and do) work with it all the time without catastrophe is that it’s not as dangerous as you might expect based on how super cold it is. This is because of something called the Leidenfrost effect – it boils so fast it’s cushioned by gas. And it can actually “ride this cushion” to “levitate” across surfaces as little beads of liquid (cleaning dust as they go) – it’s pretty *cool* stuff!
This post is part of my weekly “Bri-fings from the Bench” for The International Union of Biochemistry and Molecular Biology (@theIUBMB). Be sure to follow the IUBMB if you’re interested in biochemistry! They’re a really great international organization for biochemistry.