Empirical and Molecular Formulas

No, this post isn't overdue at all, no, no. It's just fashionably late :)
This lesson is after Percent Composition, if anyone is keeping track of order!

Off to the lesson now:

Empirical formulas are the SIMPLEST formula of a compound
For example, the molecular formula of Octane is C8H18
The empirical formula of this is C4H9 (because both of the subscripts above are divisible by 2)

Empirical formulas show only the simplest ratios and not the actual number of atoms of compounds, whereas Molecular Formulas give the actual number of atoms. Here are some lovely examples:

Molecular Formulas                               Empirical Formulas
P4O10                                                      P2O5

C6H18O3                                                C2H6O

C5H12O                                                  C5H12O    (cannot be simplified further)

To determine the empirical formula, we need to know the ratio of each element. This is simple when you get the molecular formula, like above, but sometimes you are just given the mass, moles, etc, and so being the incredible chemistry students that we are, we will determine the empirical formulas!:

Example: 

Atoms      Mass        Molar Mass    Moles       Smallest Mole       Ratio
C              50.5g        12 g/mol         4.2 mol            1.33                    4
H              5.26 g        1.0 g/mol       5.26 mol           1.66                   5

N              44.2 g        16 g/mol        3.16 mol            1.0                    3                                              

 Are you asking "Why, how did you do that, you incredible genius?" then here are some general steps:
first, depending on what you are given, you find the mass, molar mass, and moles of each element!
Then, you divide by the smallest amount of moles (in this case, the N because it is 3.16 mol), and that will give you a number, which you will place under the 'smallest mole' column. Now, often this number will be a BEAUTIFUL whole number with no decimals to worry about, in which case you rewrite those numbers under 'ratio' and BAM you've got your equation.
Well this example is different. The Smallest Mole numbers have come out all ugly and have to undergo changes! In this case, since the ending is either .33 or .66, then you multiply all of the smallest mole numbers by 3 to give you a whole number, which gives you the ratio. And so your compount is C4H5N3.
*note: if the ending were .5, you'd multiply by 2, if the ending were .25 or .75 you'd multiply by 4!

IMPORTANT MESSAGE: that rainbow chart took a long time to do. so appreciate the beautiful colors fully. please and thank you

Now what do you do if you're given the empirical formula and you are trying to change it to Molecular Formula?
First of all, you NEED to know the molar mass (in chem 11 it's usually given). Then you find the molar mass of the empirical, which will show you what you need to multiply the subscripts in the Empirical formula by. Here's an incredible example:

The empirical formula for a substance is CH2O and its molar mass is 60.0 g/mol. Determine the molecular formula:
Empirical Formula               Molecular Formula                      
CH2O                                     C2H4O2

30.0g/mol                                60.0 g/mol

In this case, the molar mass of the molecular formula was double the molar mass of the empirical formula, so we just have to multiply the subscripts of the Empirical formula by 2!

Here's a long, but good video on Molecular and Empirical Formulas:

http://www.youtube.com/watch?v=gfBcM3uvWfs

 This is what came up when I looked up "Empirical Formulas". I just had to put this up, cause it just reminds me of Mr Doktor so incredibly much.