Salt
Hi folks, sorry for the hiatus. It’s amazing how much life can get in the way of one’s volunteer blogging, especially when there are great juicy science topics as of late. How could I miss the USA 193 shoot-down, the lunar eclipse, and the Shuttle landing, which all happened on one day? Missed opportunities, I guess.
Anyway, onto the rant of the day. Salt.
The Albany Times-Union, my hometown paper, tried to actually publish a little something on science. Good! In Tuesday’s paper, writer Stephanie Earls wonders why salt both melts ice and seasons food.
Generally, this is a decent article, but I wanted to clear up a few finer points of what really happens with salt.
“But what is it about this ‘rock salt,’ or sodium chloride (that’s NaCl, chemically speaking) that makes it good at battling what foul weather hath wrought, on the roads we drive, or our own icy front steps? And does it differ from the stuff we sprinkle on our fries? …
Because of its unique chemical makeup, sodium chloride lowers the freezing temperature of water from 32 degrees [Fahrenheit] to about 15 degrees.”
Its “unique chemical makeup”? Not really. Any salt has what are called “colligative properties”. A salt’s colligative properties describe things that happen when the material is put into a solution.
In general, a salt is any two kinds of atoms joined by an ionic bond. An ionic bond occurs when two atoms physically trade an electron. When you go to the grocery store to buy salt, however, you’re getting a specific salt — “table salt”, or sodium chloride (NaCl). As Earls points out, chemically “table salt” and “rock salt” are completely identical.
One of the key rules in chemistry is that eight electrons in the outer shell is a “magic number” of sorts. The sodium atom (Na) in table salt is very happy to give up the single electron in its outer shell — this would mean that the next lower shell, containing eight electrons, would be exposed to the world. This makes sodium very stable and very pleased with its lot in life! On the other hand, chlorine (Cl) has seven electrons in its outer shell, so when it picks up sodium’s extra electron, it’s in hog heaven. The two go together like cold cuts and mustard, and before you know it you have two happy atoms.
When salt is dissolved in water, making a solution, the electrons are able to stay permanently exchanged. We call this process “dissociation” — it’s exactly what it sounds like. The sodium and chlorine ions stop associating with each other and go their own way. To differentiate, we now say that the sodium atom (Na) is now a sodium ion (Na+) and the chlorine atom is now a chloride ion (Cl-).
There’s nothing unique about NaCl’s chemical properties though. In fact, this behavior only occurs because sodium has one valence electron and chlorine has seven. Other elements in group 1 of the periodic table, lithium, potassium, rubidium, cesium, and the very rare francium, will be more than happy to step into sodium’s role. Similarly for chlorine — any group 17 atom will do. This includes fluorine, bromine, iodine, and astatine. Cesium bromide (CsBr) would have essentially a similar effect as sodium chloride.
We’re not even limited to groups 1 and 17; in fact, another common salt used in refrigeration is calcium chloride (CaCl2). Calcium is an alkaline earth metal and sits in group 2 of the periodic table. In the case of calcium chloride, calcium’s two valence electrons divide up when it dissociates in solution. Each chlorine atom gets an electron to become a chloride ion, and they go off happy as clams. Chemically speaking, we write that CaCl2 → Ca2+ + 2Cl-.
Why then do we use NaCl over CsBr? Well, sodium chloride is used because it is by far the most widely available. It occurs in mines, left over from evaporated seawater. CaCl2 is also popular because it can be produced from limestone. Both occur in nature readily; none of the more obscure compounds can be easily found in nature.
“At minus 6 degrees, however, the great salt debate is settled. ‘At 6 degrees (below zero), it becomes inert, just like gravel,’ [Salt Institute President Richard] Hanneman said.”
Well, technically true, but salt doesn’t just “shut off” at -6°F. NaCl is only capable of lowering the freezing point of water down to -6°F. The salt doesn’t become inert; this implies that if the temperature gets below -6°F the salt no longer works and the salt-water mixture is the same as if there were no salt to begin with.
Salt won’t completely melt ice all the way down to absolute zero, for the same reason that your kitchen doesn’t heat up to 212°F when you boil water. There’s only so much of an effect that a small amount of salt will have in a large amount of water. If you take a gallon jug of water and add a grain of salt, the salt has virtually no effect on the freezing point of the water. There’s way too much water and too little salt. This is why you see the road crews sprinkling salt on the roads very liberally; rather have too much than not enough.
I did learn something, however!
“The particle size in table salt is not sufficient. It doesn’t have enough weight or mass to bore through, so it tends to melt only the surface, whereas a large particle of rock salt will, in fact, melt through the ice down to the pavement and spread out on the pavement, destroying the bond between the pavement and the ice,” Hanneman said.
I would have thought that finely-ground table salt would be a much more effective melting agent. The smaller the particles, the greater the exposed surface area, and so the greater the “melting power” of the salt. (For the same reason that fine-grained sugar dissolves quickly in hot tea, while old fashioned sugar cubes take a long time.)
However, I’m not going against Hanneman on this one. Especially since salt is applied to roads before and during snowstorms, it makes sense that the heavier the salt the deeper it will penetrate into a layer of snow. It will just take longer for the melting point of the salt solution to decrease. So while the road crews are sacrificing a little bit of time, the effectiveness of the salt is greater when it can prevent ice from forming on the roadway. To counter this, road crews salt early, even before storms begin when possible. Chemically speaking, there’s no difference between salting snow that’s already there or having falling snow encounter salt sitting on a roadway.
Perhaps someday a traffic engineering podcaster will take us through winter snow removal techniques and engineering roadways for winter.
Note to the Times-Union: Not bad! B+. Generally good, but you could use a little touching up on the finer points. But, we appreciate the effort, and keep up the good work!
Tags: chemistry, earls, media, msm, physics, salt, science, timesunion
Physics Is Phun | Jim | 
February 21, 2008 5:26 pm
7 Comments
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Salting the Roadway : Talking Traffic — February 21, 2008 @ 8:32 pm
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By Annie, February 21, 2008 @ 6:27 pm
I knew about the table salt and rock salt difference in real world ice melting effectiveness. I tried this as a kid either just for the heck of it - what else to do on a snow day besides waste a container of salt - or as part of a science class (or possibly both - learned it in class, had to go home and replicate).
By Bill Ruhsam, February 21, 2008 @ 8:23 pm
Nice throw. I’ll put this on the list, although I suppose I better move it to the top to keep it relevant to the season.
Not having been involved much in maintenance, I’m not that familiar with the mixture that is spread on the roads to deter icing. I’ll look at that.
By Amy, February 22, 2008 @ 8:42 am
Around here they use brine before a storm, so no actual salt crystals, which changes the size of the crystal argument. It is very easy to tell where they’ve been as there are parallel, white salt stripes down the lanes.
Once it starts snowing I’m not sure what they are using to continue the salt effect.
By Annie, February 22, 2008 @ 2:25 pm
@Amy
Around here they’ll put down the pre-storm liquid stuff. I always thought it was some higher potency chemical stuff. But, yeah, you can always see where they’ve applied it and do the road importance assessment (meaning, how important is this road that it remain clear). Hint, I live on the very unimportant end of a rather unimportant road.
By Kindergarten teacher, May 11, 2009 @ 6:00 pm
Why does table salt (NaCl) form different shaped crystals than epsom salt (MgSO4)? My guess is that it has to do with the different makeup but need a more scientific answer for my daughters science fair project. A fourth grade answer would be great. She has not yet learned about electrons and such so if there is a easy scientific answer that would be appreciated :0)
By Jim, May 11, 2009 @ 9:04 pm
@Kindergarten teacher:
Hmm… a simple answer, huh? I’ll do the best I can.
There are a couple things that make MgSO4 different from NaCl. Firstly, MgSO4 is made up of both ionic and covalent bonds, whereas NaCl only uses ionic bonds. Atoms joined by ionic bonds tend to want to hold onto each other more tightly than covalent bonds.
Epsom salt is MgSO4, certainly, but there’s a little something missing. Epsom salt is a heptahydrated compound. This means that for every molecule of MgSO4 in the crystal, there are seven molecules of water. This does some crazy things to the crystal structure.
So, the really basic answer: NaCl has a simple crystal structure because it’s only got two different kinds of atoms, and those atoms can only bond one way. MgSO4 ċ 7H2O has four different types of atoms and several different types of bonds. The added complexity of Epsom salt means a different crystal structure.
Does this help at all?