The Mole and Avogadro's Number
The name mole (German Mol) is attributed to Wilhelm Ostwald who introduced the concept in the year 1902. It is an abbreviation for molecule (German Molekül), which is in turn derived from Latin moles "mass, massive structure". (From the Wikipedia article on the mole unit.)
A good site that introduces the mole concept and includes sample calculations and practice problems can be found here, from John Park's excellent ChemTeam site, referred to many times in this site.
Don't miss the interview with Count Amedeo Avogadro and his wife the Countess Felicita, located here (thanks to Kory Tonouchi, Roosevelt High, Honolulu, HI).
Avogadro's 1811 publication, "Essay on a Manner of Determining the Relative Masses of the Elementary Molecules of Bodies, and the Proportions in Which They Enter into These Compounds", may be found here (thanks to Carmen Giunta, Le Moyne College, Syracuse, NY).
A fun "mole" page to visit is here. It is a collection of student projects from Carondelet High School. Back in '98 they had a mole mystery described here.
Did you know we have a mole day? 
Find out about it here. And for some corny mole jokes, don't miss the Dictionary of Mole Day Terms & Jokes here! 
Introduction (page 1 of 2)
A mole of objects contains Avogadro's number, 6.022 X 1023, objects. Just as a dozen apples is 12 apples, a mole of apples is 6.022 X 1023 apples. A mole of iron atoms is 6.022 X 1023 iron atoms. A mole of water molecules is 6.022 X 1023 water molecules.
The NIST 2007 value of Avogadro's number is 6.022 141 79 ± 0.000 000 30 X 1023 mol�?. For most calculations, a rounded value of 6.022 X 1023 (four significant figures) is satisfactory.
This is an incredibly large number. A mole of say, grapefruit, stacked together, would occupy the volume of the entire planet earth! And yet a mole of water molecules is in only about 18 milliliters of water. Since atoms and molecules are so small, there are gigantic numbers of them in ordinary gram quantities of substances such as what we weigh and use in the chemistry lab.
AMUs, Grams, and Moles
The value of Avogadro's number is actually chosen arbitrarily, based on the definition of the atomic mass unit, amu or u. By definition, a single carbon-12 atom weights 12 amu exactly. Therefore, one amu is one-twelfth the mass of a single carbon-12 atom.
Now, how many carbon-12 atoms would weigh exactly 12 grams?
From experiment, the actual mass of a single carbon-12 atom in grams has been determined. For example, using the method of mass spectrometry, the mass of a single carbon-12 atom has been measured to be about 1.993 X 10�?3 g. From this we can calculate the number of carbon-12 atoms in 12 grams of carbon-12:
12 g X 1 carbon-12 atom = 6.021 X 1023 carbon-12 atoms
1 1.993 X 10�?3 g
(from Clugston and Flemming,
Advanced Chemistry)
This is the basis of Avogadro's number. Better experimental methods have yielded the more accurate value of Avogadro's number we have today.
By definition, 12 grams of carbon-12 contain one mole, or Avogadro's number of, carbon-12 atoms.
We can also relate the two mass scales, grams and amu, as follows:
6.022 X 1023 atoms of carbon-12 X 12 amu = 6.022 X 1023 amu / g.
12 g 1 atom of carbon-12
That is, 1 g = 6.022 X 1023 amu.
The average weight of a carbon atom found in nature is a little more than 12 amu, actually 12.0107 amu, because there is a small amount of heavier carbon-13 atoms present.
We can calculate the average weight of one mole of carbon atoms as follows:
1 mole C X 6.022 X 1023 C atoms X 12.0107 amu X 1 g = 12.0107 grams
1 mole C C atom 6.022 X 1023 amu
As we have with carbon-12, the weight of a single carbon atom, on average, is 12.0107 amu, and one mole of carbon atoms weighs 12.0107 grams, the same number.
What about other elements, does the same relationship hold? Indeed yes, the proportions of the weight of a single atom of carbon compared to a single atom of, say, iron is the same, whether we are comparing the weights of single atoms, one dozen atoms, or one mole of atoms. For example, it is known from experiment that, on average, an iron atom is 4.6496 times more massive than a carbon atom, which is 55.845 amu per iron atom. By proportion, one dozen iron atoms will be 4.6496 times more massive than one dozen carbon atoms. Likewise, one mole of iron atoms will be 4.6496 times more massive than one mole of carbon atoms, which is 55.845 grams per mole of iron atoms.
We therefore have two types of units for atomic weights, molecular weights, and formula weights:
1) amu per single atom, molecule, or formula unit
A single iron atom weighs 55.845 amu / atom
A single water molecule weighs 18.0153 amu / molecule
A single "unit" or "formula unit" of NaCl weighs 58.443 amu / formula unit
2) grams per mole of atoms, molecules, or formula units
One mole of iron atoms weighs 55.845 g / mol
One mole of H2O molecules weighs 18.0153 g / mol
One mole of NaCl formula units weighs 58.443 g / mol
In general, we can refer to the weight of one mole of a pure substance as its molar mass. If the substance is an element such as iron, the molar mass is the atomic weight of the substance. If the substance is molecular, like H2O, we can call the molar mass the molecular weight of the substance. If the substance is ionic, rather than molecular, we can refer to the molar mass as the "formula weight" of the substance.
Since most chemical calculations involve converting between grams and moles, you should get into the habit of using g/mol units (and showing them in your work!). You will only occasionally need to use atomic mass units in calculations.
Important Formulas
The important calculation formulas to memorize are moles = grams / molar mass
and rearranging,
grams = moles X molar mass
We use these two formulas more than any others in chemistry, because so often we are required to convert from grams to moles and moles to grams in chemical calculations.
On the next page are some sample calculations using moles. –�?gt;
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