Happy Mole Day!

For all you geeks out there, it's National Mole Day! As I took two different chem classes in high school (Chem 1 and AP Chem), I still remember what a mole is. For those of you who don't know, a mole is the amount of any substance you have when you have exactly Avogadro's number's worth of atoms or molecules of that substance. Now, Avogadro's number is expressed as:

6.02 x 10²³

When you have this many atoms or molecules of something (water, for example), you have one mole of that substance. This usually isn't a big deal for most applications where other units to measure the amount of something are better, like using ounces or cups in cooking. For chemical applications, however, moles are highly significant; a mole is a standard amount of any substance, since the number of atoms/molecules of the substance are always the same. This allows any substance to be measured in a comparable way to any other substance.

For example, there are two measures used chemically to determine concentrations of chemical solutions: molality, defined as moles of solute per kilogram of solvent and signified with a lower-case m, and the much more common molarity, defined as moles of solute per liter of solution and signified with a capital M. The higher the molality or molarity of solution, the more concentrated it is. So a solution of hydrochloric acid that measures 0.01 molarity (or properly, 1/100th of a mole of hydrochloride molecules [HCl] per liter of solution) might not do much more to you than provide an unpleasant burn, but crank up the molarity to 2.0 or 3.0, and you're looking at losing some substantial flesh (as well as a great deal of pain). For chemical purposes, molality and molarity are much more precise than something like "parts per million," especially considering the possible disastrous effects of certain chemical interactions.

Not that you'd want to be fooling around with hydrochloric acid, but that's another story. So anyway, enjoy National Mole Day. Thanks for reading along.