The creation of formulas and the finding of their molecular masses
is only partially complete. What exactly does that number you have found for the molecular mass
mean? Is it the mass of one molecule, or a million, or ten billlion?
What unit is it measured in? |

The Carbon-12 Based Atomic Mass Scale |

In the years following Dalton's presentation of his atomic theory,
chemists worked very hard at determining a complete set of relative masses
for all the elements that were known. By knowing these relative masses
the chemists were able to select amounts of elements in grams needed for any
desired atom ratio. In establishing a table of atomic amsses,
it is necessary to have a reference point against which to compare the relative
masses. Currently, the agreed-upon reference is the most abundant isotope
of carbon, which is carbon-12. By definition, an atom of this isotope is
defined as having the mass of exactly 12.000 amu. (atomic mass units)
In other words, an amu is defined as 1/12th of the mass of one atom of carbon-12. |

The definition of the size of the atomic mass unit was quite arbitrary.
It could just as easily have been selected to be 1/24th of the mass of one
atom of a carbon atom, or 1/10th the mass of a calcium atom, or any other
value. Why 1/12th the mass of carbon-12? Carbon is a very common element,
available to any scientist and by choosing the amu to be of this size,
the atomic masses of nearly all the other elements are almost whole numbers,
with the lightest atom having a mass of approximately 1. (hydrogen-1 has
a mass of 1.007825 amu when carbon-12 is assigned a mass of exactly 12 amu.) |

The Mole |

Suppose we want to make molecules of carbon dioxide, CO_{2},
in such as way that there would be no extra carbon and oxygen atoms left.
If we took ten atoms of carbon and twenty atoms of oxygen we would make 10
molecules of carbon dioxide. Suppose we wanted to make more, lets say
40,000 molecules of carbon dioxide. Once again, we could count out
40,000 atoms of carbon and 80,000 atoms of oxygen, let them react and we'd
have 40,000 molecules of carbon dioxide. This sounds all very nice
and neat unless you realize the trap. Have you ever seen an atom?
Atoms are too tiny to count individually. Saying they are tiny
is even wrong because tiny can be seen. Even with the best scanning
tunneling electron microscope ever invented the largest atoms known, look
just like fuzzy cloud tops. We have never seen individual atoms.
There are no lenses with the resolving power or balances fine enough to measure
an individual atom. |

We get around this problem because each element has its own characteristic
atomic mass and each formula has its own unique molecular mass.
We know, that oxygen weighs in at 1.33 times that of carbon, because of this
we get a ratio of their masses:
12.0 amu (for one atom of C) 1 |

If we take a sample of oxygen and carbon in a ratio of 1.33 to 1,
we must obtain equal numbers of their atoms. That is, if we actually
had a balance that could measure amu directly we could mass out 32
amu of oxygen and 12 amu's of carbon - a mass ratio of 2.66 to 1 - we would
have exactly 2 atoms of oxygen for every atom of carbon and a 2 to 1 ratio
by atoms. |

When we mass out a sample of an element such that its mass in grams
is numerically equal to the element's atomic weight, we always obtain the
same number of atoms no matter what element we choose. Thus 12.0 g
of carbon has the same number of atoms as 16.0 grams of oxygen, or 32.1 g
of sulphur, or 55.8 g or iron. |

This relationship also extends to compounds. The formula mass of water,
H_{2}O, is 18.0 amu. If we take 18.0 grams of water then it should
have the same number of molecules as there were atoms in 12.0 grams of carbon.
The carbon-12 isotope, which makes up 98.89% of all naturally occuring carbon,
is the reference used by SI Metric in its definition of the base unit for
a chemical substance, the mole, abbreviated mol. |

This mole concept is the most important
in all of chemistry. Once this concept is grasped all the rest of chemistry
will appear easy. |

Avogadro's Number |

The mole is defined as 6.02 x 10 ^{23} units. It is called
Avogadro's number in honour of Italian scientist, Amadeo Avogadro (1776-1856).
It is a pure number with a special name, just like so many others.
For example:
2 = pair
12 = dozen
144 = gross
500 = ream
6.02 X 10 ^{23} = Avogadro's number |

This number is not an odd number at all. It became inevitable once
the amu was defined. The relationships needed are:
or
12 amu = 1 atom of C (Specific example)
1 mol of C = 6.02 X 10 ^{23} atoms of C
1 mol of C = 12 g of C |

Another useful relationship is that 1 amu = 1.66 X 10
^{-24}
g
1.66 X 10 ^{-24} grams |

The mass in grams of a substance, that equals one mole is often called
its molar mass, and the units are grams/mole or g/mol.
For example, aspirin has a molecular mass of 180 grams. Therefore if we massed
out exactly 180 grams of aspirin we would have Avogadro's number of aspirin
molecules. |