# Happy Mole Day! (What's a mole?)

This is National Chemistry Week. It's always chosen to coincide with whichever calendar week includes October 23 (or 10/23), since October 23 is "Mole Day".

"Huh? Why would chemists celebrate a furry critter that burrows underground?"

Not that mole. The mole chemists celebrate is a unit.

1 mole = 6.02 x 1023 of whatever it is you want a mole of.

You can think of a mole as being sort of the chemists' equivalent of a dozen -- it's a convenient sized bundle for working with the kind of stuff chemists work with, namely, atoms and molecules. If they were working with eggs, or shoes, or some other stuff, they would probably use a different unit than the mole.

Remember that atoms are really, really small. They are so small that you've never seen one with your naked eye. (Depending on how you score things like diamonds, you may have seen a molecule with your naked eye, but even most molecules are way beyond the limits of your visual acuity.) If passing intro chem depended on doing experiments where you had to weigh out single atoms of reagents, there would be no chemists.

Weighing out 6.02 x 1023 atoms, though, is do-able. 6.02 x 1023 atoms of carbon weigh 12.01 grams. 6.02 x 1023 atoms of copper weigh 63.55 grams. 6.02 x 1023 atoms of helium weigh 4.003 grams.

I don't have any of this memorized. I just turned to my handy, dandy periodic table. Under each element's symbol is listed an atomic weight. Carbon's atomic weight is 12.01. That means a single atom of carbon has a mass of 12.01 atomic mass units (amu). However, since we've already noted that you hardly ever want to deal with one atom at a time (at least if you're a chemist -- maybe physicists would rather deal with single atoms), the more important fact that you can glean from the table is that 6.02 x 1023 atoms of carbon have a mass of 12.01 grams. Or, simply, 1 mole of carbon is 12.01 g.

Once you recognize that atomic masses tell you how many grams in a mole of an element, you can measure your desired amount of multi-element molecules, too. Table salt is NaCl. Na has atomic mass of 23, Cl has atomic mass of 35.5, so a mole of NaCl (or 6.02 x 1023 molecules of NaCl) is 58.5 g. To find the mass of a mole of H2O, you have 2 moles of H (atomic weight 1) and one of O (atomic weight 16) -- 18 g of H2O holding all 6.02 x 1023 molecules.

From this, it's a good bet that mixing an equal number of moles of table salt and water will not yield a saline solution so much as a slightly damp pile of salt.

In addition to being the number of atoms or molecules (or shoes, or eggs, or pencils, etc.) in a mole, 6.02 x 1023 has a name that sticks in people's heads: Avogadro's number. Avogadro also has a law:

Equal volumes of gases at the same temperature and pressure contain the same number of molecules.

If we're dealing with "ideal gases (gases whose molecules are not too enormous, and which don't have significant attractions between the molecules), at standard temperature and pressure (0 o C and 1 atm), 22.4 liters of gas will contain 6.02 x 1023 molecules (or 1 mole). So, at standard temperature of pressure, balloons filled with 22.4 L of H2, N2, O2, Cl2, He, Ne, and Ar will each contain a mole of gas molecules. However, the masses of gasses in these balloons will be different (2 g, 28 g, 32 g, 71 g, 4 g, 20 g, and 40 g, respectively). This means a mole of argon is denser -- by a lot -- than a mole of helium.

That Avogadro's number of [anything] equals a mole of [that stuff] has also given rise to clever marketing, like this:

The label clarifies that there are not actually 6.02 x 1023 avocados in each package of guacamole (it's closer to 5 or 6). But it's a fine store bought guacamole, and the chemistry joke on the package just makes it more appealing.

Happy Mole Day!

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Oh, mole. The English word. Lo siento.

Just remember, Mole Day only goes from 6:02am until 6:02pm!
6:02 10/23

The honors chemistry society here is having a pub night to celebrate Mole Day. Mind you, this would just so happen to be the night before I take the GRE, so I can't celebrate chemistry too much...

I was hoping that this might have something to do with my favorite Mexican ingredient.

That means a single atom of carbon has a mass of 12.01 atomic mass units (amu).

Minor quibble: No single atom has an atomic mass of 12.01 amu. The atomic weights listed in the periodic table are averages for natural abundances.

That is, there are three isotopes of carbon, carbon-12 at 12 amu per atom, carbon-13 at ~13 amu per atom, and carbon-14 at ~14 amu per atom. For a typical chunk of carbon, they're will be about 98.89% 12C, 1.109% 13C, and a trace amount of 14C, so on average an atom of carbon weighs 12*0.9889 + 13*0.01109 + 14*trace = 12.011 amu. But no single carbon atom weighs exactly 12.011 amu. (But 6.02x1023 atoms of carbon will weigh 12.011 g)

One more High School Science song.
Mole day Song
(to McCartney and Lennon's Let It Be)

When I find myself in times of trouble
Avogodro's number comes to me
Six point oh two two, e twenty three.

The mole is a counting number
And you'll find few who disagree
Six point oh two two, e twenty three.

Twenty three, twenty three
Twenty three, ya twenty three.
A mole is six point oh e twenty three
(that's sig figs kids)

The molar mass of an atom is on
The periodic table you will read
The number of atoms is six point oh e twenty three

To convert from grams to atoms
It's the molar mass that you will need
Times by six point oh two e twenty three

Twenty three twenty three
Twenty three, ya twenty three
A mole is six point oh e twenty three

It's always great fun to ask students in their second or third semesters of Chemistry why there are 6E23 in a mole. Why that number and not, say, 7E23?

Not one student has ever given even a semi-decent answer. After a bit of pushing, the odd person does stumble towards the light. 6E23 is the reciprocal of 1.66E124. And 1.66E-24 is the mass of 1 hydrogen atom (in grams) or 1u. so, if one atom has a mass of XXXu, then 6E23 of them have a mass of XXXg. It's the same number, only the units change. If Avogadro's constant were different, we would have two numbers to deal with, and chemists would keep getting them mixed up.

One of the interesting questions on that subject: why did SI decide to allow the mol to be effectively described in terms of grams rather than kg, since the latter is the SI unit of mass? It would have been, for example, just as easy to define the "kmol" as the fundamental unit of (chemical) quantity as the mass of 12 kg of 12C.

By Taxorgian (not verified) on 24 Oct 2008 #permalink

im a student and dont realy understand chem. just yet.

By ashley m. (not verified) on 23 Oct 2009 #permalink