How is a mole (the unit) different from the number 6.02214179e23?

A mole of anything is 6.02214179e23 of it. A mole of air molecules is 6.02214179e23 air molecules, and so forth. So, how is the mole different from just a number? It got elevated to the status of a unit in the SI, and I understand that this was a controversial move, detractors complaining that it is just a number. I understand this view. What I don’t understand is, what is the opposing argument? Apparently it was successful. Why and how is this unit different from just a dimensionless number?

Because saying "I have a mole of (whatever) is easier than saying “I have 6.02214179e23 (whatever).”

A dozen is just twelve, a gross is just 144, but they are not just numbers, they are a count. A count of *something *and not just an abstract number.

One mole of anything is that thing’s molecular mass in grams. So one mole of carbon atoms is 12 grams. You could consider a mole to be a derived unit, then – “the number of molecules of anything such that its mass in grams is equal to the molecular mass.”

Nitpick–a mole of carbon is 12 grams only if the sample is pure carbon-12. But there will be some amount of carbon-13 and carbon-14 in there, depending on the source. For elements with multiple stable isotopes there will be a typical mixture.

Carbon-12 is the only stable isotope, but the other isotopes are created via various processes. And this is how carbon dating works, a sample that contains lots of carbon-14 is new because it incorporates newly created carbon-14, a sample that contains little carbon-14 is old because the carbon-14 it used to contain has decayed.

And if you were counting twelve-gram standard masses, one mole of them would be quantity one? :slight_smile:

(The reason for this all made sense to me in high-school chemistry, but it’s been a loong time…)

It’s easier than chemistry texts make it. Avogadro’s number is just the number of protons you need to add up to 1 gram. Well, pretty much, anyway.

A mole is a unit,
Or haven’t you heard?
It’s 6 times 10 to the…23rd.

I think this is incorrect. A mole of Argon atoms is a count of something, and not an abstract number. But that is because of the words “of Argon atoms”, not because of the moles.

To quote the BIPM’s brochure on the SI, “When the mole is used, the elementary entities must be specified and may be atoms, molecules, ions, electrons, other particles, or specified groups of such particles.”

Nitpick–Carbon-13 is also a stable isotope. It’s about 1/100th of all carbon.
I don’t really understand why it wouldn’t be a unit. 6.02x10^23 would make calculations very cumbersome. I suppose you might argue that they could have gone with 1x10^23 to make it sort of metric. I have no idea what the prefix for 10^23 in metric is. Maybe they could have done that. There’s no going back now though. All of our masses and calculations are based on it.

And thanks for the mole day reminder!

For most uses, the “of atoms” is not needed. I don’t weigh out a mole of sodium carbonate molecules, and moles of molecules per liter is not a concentration unit anyone is familiar with.

It’s a unit of measurement of something specific, viz., molecules (including monoatomic ‘molecules’). You would not have a mole of moles, the burrowing animals, consisting of 6.02*10[sup]23[/sup]*weight of a typical mole. You would not say that there is 0.0167 mole of stars in the Virgo cluster of galaxies.

What makes it a useful quantity is that it denotes a chemically equal amount of any substance. One mole of hydrogen fluoride will react evenly with one mole of lead tetraiodate, even though the masses are vastly dissimilar – they contain exactly the same number of molecules.

It’s not an abstract ratio or constant universally applicable, like pi or e. It happens to be the number which defines the number of molecules of any substance that are found in a sample with a mass of the substance’s molecular mass number times 0.001 the weight of a particular platinum-iridium cylinder sitting in Paris.

While the idea that there are an equal number of molecules in any measurement system – the long ton or the old Japanese momme, for example, or whatever Star Trek Vucans or wookiees use for units of mass – Avogadro’s number is specific to the SI. For any molecule with molecular mass units equal to X, it defines the number of molecules in X grams. You could moderately easily calculate what the equivalent number would be for whatever mesaurement of mass you choose.

Near as I can tell, the opposing argument was that chemists were uncomfortable with fifth-grade math, and considered scientific notation scary, so they gave a name to a common big number so they could avoid the scary scientific notation. Yes, yes, you can say that one mole of oxygen reacts with two moles of hydrogen, and people say that that makes things easier, but really, it’s simple enough without using moles, too.

I recognize you are not on the side of the opposing argument you are summarizing. However, I still don’t quite understand what is going on here beyond having a convenient name for a large number that comes up often. [It seems recognizing it as a “unit” means something more than this, although I may be wrong about that]

Can’t you also say one atom of oxygen reacts with two atoms of hydrogen? Why does arbitrarily rephrasing it in terms of moles make describing that process easier?

Or is the idea that, if you actually were describing a reaction between a mole of oxygen and two moles of hydrogen, the word “mole” is easier than saying “6.022… * 10^23 atoms of oxygen reacted with 6.022… * 10^23 pairs of atoms of hydrogen”? Since it is apparently a number that comes up a lot, it may as well have a convenient name, why not, but I don’t understand why that particular number name should be considered somehow different from other number names (like “four” or “dozen” or “thousand” or “googol”, which I presume are not considered units).

I suppose there is the problem that Avogadro’s number is only experimentally determined, with current conventions, so it’s not quite as precise a number name as those other examples. Seems it’d be easier to get rid of the prototype kilogram and define Avogadro’s number as exactly (some nice integer in the vicinity of the current bounds upon it), and then grams and kilograms in terms of that.

Not at all. A unit was needed to keep track of the quantities of things even before the atom was discovered. You couldn’t say you had so many atoms in a sample when you had no idea what an atom was. Since we were defining things in terms of moles long before an atom was even proven to exist, a unit was needed. The unit stuck.

That’s a technological limitation. It’s nice to have a standard for a unit that doesn’t depend on some physical artifact, but it’s more important that it be consistent from one measurement to another. With current technology, it’s possible to measure the mass of the lump of platinum-iridium in Paris (and the other lumps made in imitation of it) to a greater precision than it is to measure the mass of a carbon 12 atom, even once you allow for difficulties like keeping that platinum-iridium lump clean and not eroding it and so on. I understand some folks are trying to work on techniques to replace it (such as, for example, by counting out precisely some number of carbon 12 atoms and weighing them), but they’re not yet ready for prime time.

Ah, I see; I never thought about that, but that makes sense.

[nerd nitpick]
It’s in Sèvres`
[/nerd nitpick]

Got mole problems?

Call Avogadro: 602-1023

It may have been a technological limitation then, but by the time an atom was discovered, many constants and measurements defined by it. Changing these constants would be much worse than going metric since they are by now hidden everywhere. All of that effort is for zero benefit for anybody except those afraid of a little fifth grade math.