This week, Sam Kean takes a look at some ridiculously precise standards -- the meter, the second, and other international standard units -- and the role that elements have played in defining, redefining, and re-redefining them over the ages.
My high school teacher had a Velcro rodent she would dismember to help us visualize it—a "mole" that detached in different places and could be stuck back together to illustrate a half a mole, a quarter of a mole, or whatever. I'm not sure it helped.
Basically, a mole measures the amount of a substance, but measures it in a clever way. Let's say you wanted to manufacture calcium sulfide, CaS, and you worked in a very competitive industry where you couldn't waste any calcium or sulfur. That means you need the exact same amount of each to mix together. But defining amount gets tricky here, because a sulfur atom has fewer neutrons and protons and therefore weighs less than a calcium atom. So if you have ten kilos of both, you actually have far more sulfur atoms than you need. The mole solves this problem: It provides a way to convert from kilograms (or whatever) into the amount of X that will react with Y. In this case, you'd want to mix one mole of each element to get a perfect yield.
The international definition of a mole has been based on common elements like oxygen and hydrogen in the past, but ever since 1960, scientists have defined one mole as exactly the number of atoms in 12.0000... grams of carbon-12. But really, this definition papers over some predicaments—it fudges things.
You might remember a number, Avogadro's number, associated with a mole—a mole always has 6.022141793... x 1023 particles of whatever. That's an absolutely ginormous number. Counting one atom per second, with thirty million or so seconds in your average year, it would take twenty million billion years to count that high, over a million times the age of the universe. So while you might know you have exactly one mole of carbon twelve, you only have a vague idea of how many atoms that is: Because after the ellipsis in 6.022141793..., it's anyone's guess—and there are a lot of decimal places to go.
What's more, if you've had a sneaking suspicion this whole time that the mole sounds a little redundant—since the "amount of a substance" is an awful lot like the "mass of a substance"—you're onto something. In fact, issues related to enumerating atoms have led to even bigger problems with defining the last standard we'll look at, the kilogram.
Tune in tomorrow for the final installment of our exploration of the standards that make science tick. The series is written by Sam Kean, author of The Disappearing Spoon—a collection of funny and peculiar stories hidden throughout the periodic table.
I like these little blurbs on the industry standards. Although the units mentioned are commonly known, the history behind them is less known and I've found it's rather fascinating. Thanks for the insight!
I'm studying Moles right now at school, so this is quite relevant, and factual, as far as I can tell.
I'm wondering if Americans are ever going to adapt international standard units as the norm?
"I'm studying Moles right now at school, so this is quite relevant, and factual, as far as I can tell."
Right now at school, Moles are studying, as far as I can tell, so I'm quite relevant and factual.
I found a wiki that cleared this story up qute a bit more.
I searched on unit of measurement mole. ( I don't want to violate whatever link rule they have here).
Anyway, it also had the following little tidbit:
"October 23 is called Mole Day. It is an informal holiday in honor of the unit among chemists in North America. The date is derived from Avogadro's number, which is approximately 6.02×1023. It officially starts at 6:02 A.M. and ends at 6:02 P.M."
I don't see why it should be so hard for school children to understand what a mole is. It's just a numerical amount of something(atoms, molecules, formulaic units), with a particularly convenient choice of number.
It is chosen such that a mole of a substance weighs in grams exactly the same numerical value as the individual constituent molecules weigh measured in molecular mass.
The unified atomic mass unit, being defined as 1/12th of a carbon-12 atom, is the clever part. There a slight deviation due to binding energy, but it is accurate to within 1%, the isotope mass is equal to the sum of the number of protons and neutrons.
U-238 should have an isotopic mass of about 238 u(actual: 238.051 u). Fe-56 should have an isotopic mass of about 56 u(actual 55.934 u). Hydrogen-1 should have an isotopic mass around 1(actual 1.0078 u).
It would have been perfectly possible to do without moles. To simply deal with unitless numerical amounts of atoms and to have table values for the mass of the different elements in kilograms. This is inconvenient because atoms weigh so little and visible amounts of atoms are great big thundering hoards; it's not easily compatible with slide-rules.