Ever wonder why we salt driveways to keep them from getting icy? Or why we use antifreeze to protect our pipes? It all has to do with a concept called freezing-point depression. Ice forms when water molecules stick to one another and stay stuck in place because they’re too cold to slide around like they’d be able to in a liquid or to move freely as they could do if they were a gas. If water molecules can’t find each other as easily, it’s harder for them to stick to one another and thus harder to freeze. When you dissolve salt in water you introduce a bunch of non-water molecules that “distract” the water molecules and make it harder for them to form strong water-water bonds. So water stays water and doesn’t freeze! Now, if you’ll indulge me in some details, please?

At any given temperature/pressure combo a substance prefers to be solid, liquid, or gas (we call these phases or states of matter). Real talk – if they had their way they’d always prefer to be a gas, because that’s where they’re most free and thus most happy. But being a gas requires having enough energy to break free of intermolecular (between molecule) bonds. The more pressure pushing molecules together, the harder it is to break free and stay free. And the lower the temperature, the less energy molecules have to resist one another’s lure. At low temp and or high pressure, molecules in a substance are stuck in place, and can only vibrate (think of couples doing that slow dance thing where they just stand and sway). We call such substances solids. At higher temps and/or lower pressures, the molecules start to be able to temporarily break free from one another and slide around, but they can’t get far without getting stuck to another molecule (a livelier dance where couples change dancing partners frequently). That’s a liquid for you. And at even higher temps and/or lower pressures, the molecules break free and get so far away from one another that they don’t have to worry about getting dragged back (I finally escape from the dance hall!). Now you have a gas. 

The act of a substance going solid->liquid (melting) or liquid->gas (vaporizing) or doing that in reverse (gas->liquid (condensing) or liquid->solid (freezing)) is referred to as a “phase change” and such phase changes happen at specific temperature/pressure combos. 

These “special” temperatures depend on what the substance is made of. For example *pure* water freezes at 0°C (32°F) and pure nitrogen freezes at (−210 °C; −346 °F), which is why we can use liquid nitrogen to keep samples really really cold in the lab. 

I asterisked pure because adding solute (dissolving things such as salts in the pure substance) changes the what the substance is made up of and thus changes those phase change points. Salt water makes a nice case in point…

Water is “sticky” (it has strong Intermolecular Forces (IMFs)) due to something called polarity whereby the O’s in H₂O are partly negative and the H’s are partly positive and opposites attract. Solutes break up some of the “stickiness” so water molecules are “freer.” When you hear “freedom” in a chemistry context, think entropy. Entropy is a measure of “randomness” – the more that molecules are free to move around, the harder it is to predict their location, and thus the higher the entropy of the liquid. “Nature likes entropy” (it’s one of the laws of thermodynamics), so solutes make liquid a “happier place.”

The MORE dissolved particles you have in a liquid solution, the happier it is as a liquid and thus the more it wants to remain a liquid. This has a couple consequences

  1. It’s harder to get it to vaporize (go liquid to gas). We call this vapor pressure depression and it leads to boiling point elevation. note: if you’re confused by the term “vapor pressure depression,” think about a closed container of vaporizing liquid. The vaporized molecules bang into the top of the liquid, putting pressure on it. The more vaporization that’s occurring, the more banging you have, and thus the higher the vapor pressure. If you have less vaporization happening, there’s less banging, and thus lower pressure, hence “vapor pressure depression” 
  2. it’s also harder to go the “other way” – you have to remove more heat (get to a lower temperature) in order to freeze (go liquid to solid).  We call this freezing point depression

So, if you want to prevent something from freezing, don’t get it cold… or ADD SOLUTE! We call solutes we add to prevent freezing cryoprotectants and they’re commonly used to keep water from freezing. Why water? Well, for one thing, it’s the liquid we deal with the most – and for another thing, water has this weird property that can damage your property…

What I’m referring to here is that water’s weird in that it *expands* when it freezes. Most liquids don’t do that, but water does because, when it freezes, it does so in an orderly crystal lattice in which, in order to achieve optimal binding angles, it has to spread out. If water freezes inside pipes, the pipes can burst and if it freezes inside our protein (where “rivers” of solvent flow through gaps in the protein’s 3D structure), they can “break.” So we add cryoprotectants to protect pipes in winter & proteins when we freeze them – these solute molecules get in the way of water forming that lattice. 

Another place we see solute used to prevent ice formation is the use of road salt. The idea here is that you spread an ice-water slurry on roads and then, when new snow falls it will meet a salty solution that discourages freezing. It only works if the salt is dissolved, because it wouldn’t be helpful to have a layer of salt and a layer of ice. Instead, the two have to play together for the method to work nice. If you just salt the road, you’re counting on the road heat to help melt the initial ice/snow. And if you just salt on top of ice/snow you’re counting on car tires, etc. to smash that salt in. The salt isn’t going to melt the snow/ice, but it will make it harder for ice to form in the first place or reform when it melts (even below freezing point you still have molecules temporarily melting and the same time you have molecules freezing – it’s a dynamic 2-way street). 

Adding salt thus lowers the freezing point of water – how much? Depends on how much salt! (or other molecule as we’ll see). With a 10% salt solution, you’re looking at a freezing point of 20°F (-6°C). Bump it up to 20% and you knock of another 18°F (so you freeze at 2°F or -16°C). http://bit.ly/3pehIzj  

I think typical street salting typically gives you a freezing point ~15°F http://bit.ly/2NsYqJe  

Why all the different temps? Freezing point depression is a “colligative property” (as are vapor pressure depression & boiling point elevation). When it comes to colligative properties, it’s the number of dissolved particles, not the identity of the particles that matters (think about more stuff interfering with water-water bonds).

And, when you add an electrolyte like NaCl (table salt) which dissociates into its component ions (charged particles) when you dissolve it (e.g. (Na⁺ & Cl⁻), you get more particles than you put in, so you get a sort of “2-for-1” with regards to colligative power. And it’s cheap. But you also get some downsides. For example, because it breaks into charged particles, it can have undesirable consequences like corroding metals (just ask the bottom of your car). It can also have consequences for wildlife. Road salt is usually NaCl, so your “normal” table salt, except you wouldn’t want to serve it on your table because it’s less pure. Impurities such as mineral contamination can seep into soils by roadsides and even run-off into water ways, harming plants and aquatic life. Therefore, it’s important not to over-salt, and alternative salts are also sometimes used. 

You might sometimes see calcium chloride (CaCl₂) used. The benefit of this salt is that it’s a “3-for-1” deal in terms of colligative properties – when it dissolves it dissociates into 1 Ca²⁺ and 2 Cl⁻) so you can get a lower freezing point for the same amount of salt. http://bit.ly/3tZwkXc  

You might sometimes see sand used – but don’t get confused! Sand is not working to lower the freezing point (to do that it would need to be water-soluble and oceans disprove that idea…). Instead, sand just provides some traction so cars don’t slip and slide. 

You definitely don’t want to salt your pipes because of that whole corrosion thing, and we don’t want to overreact our proteins when we freeze them because that could mess up the protein interactions, but thankfully there are alternatives. You want to add something that’s more “water-like” in terms of not disrupting the system (but less water-like in terms of “stickiness”). Basically, you want something kinda “blah” – and SMALL so we can pack lots of it in there (colligative properties, remember). Small polyols (chemicals with multiple -OH groups) like glycerol & ethylene glycol (aka “antifreeze”) fit the bill and thus are commonly used as cryoprotectants.

more on how we freeze protein samples: http://bit.ly/flashfreezingdance 

more on topics mentioned (& others) #365DaysOfScience All (with topics listed) 👉 http://bit.ly/2OllAB0

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