Lesson 9-1

The Mole

     The mole concept is one of the most feared and misunderstood concepts in all of chemistry.  Many adults who cringe at recalling Chemistry will tell you that they never understood that "mole stuff."  The funny thing about it is that the mole is really a very simple concept.  I think that many people never really understand it because it is not always presented clearly.  If explained correctly, I feel that the mole can be an easy concept to master.

     You know that there are 12 items in a dozen.   No matter what the item, a dozen is equal to 12.  A gross is another unit of grouping.  There are 144 items in a gross.  A score, another set group, is equal to 20 items.  You can have a score of years or a score of rocks, but it will always be 20 items.  Now, a mole is a unit of grouping, just like these examples.  The only difference is that a mole represents a large number of items, 602 000 000 000 000 000 000 000 (or 6.02 x 10²³) items to be more specific.  What the items are doesn't matter.  You can have a mole of molecules, a mole of ions, or a mole of stars.  The number of items in a mole will always be 6.02 x 10²³ .   This number is known as Avogadro's number.

     Why such a large number for the mole?   Well, why do they sell eggs in a dozen?  Maybe because no one wants to buy just one egg, and if you buy fifty, some will go bad before you eat them.  The reason we need so many items in a mole may be because we need to group molecules in very large groups in order to be able to get a measurable reading on our balances.  We can't find the mass of one atom, or even one gross of atoms, on our laboratory balances, the instruments are not sensitive enough.  We can, however, find the mass of one mole of atoms on our balance.

     The real reason for packing 6.02 x 10²³ items into a mole is because there are 6.02 x 10²³ u (or atomic mass units) in one gram.  This allows us to use the mass numbers on the periodic table for both the mass of an atom (atomic mass)  and the mass of a mole of atoms (molar  mass), we only need to change the units.  Table 5-2a will demonstrate what this idea:

     At this point, some students might say "Hey, I thought that a mole is always 6.02 x 10²³?  How can the molar mass of carbon and helium be different?"   That is like saying, "How can the weight of a dozen elephants be different than the weight of a dozen ants?  Shouldn't they both be 12?"  It is the number of items that is always the same, not the mass or weight or size of those items.

    Of course the periodic table can be used to determine the molar mass of molecules and formula units as well.  If the molecular mass is found in atomic mass units (u), the molar mass of that molecule will have the same value with the unit grams (g).  Table 5-2b has some examples.

     The mole allows us to do many of the important calculations that are required for quantitative analysis of samples in the lab.  Below we will go over examples of several types of these calculations.   When we use moles in calculations we will abbreviate the units to mol.  When we show the molar mass of a substance we will show the units in g/mole (read "grams per mole")

Please forward all questions, comments and criticisms to Gregory L. Curran.
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Last Modified February 07, 2008