# Find empirical formula when individual masses are known

I'd like to illustrate my question with an example:

Imagine nickel reacts with fluoride and you start with $0.766\ \mathrm{g}$ of nickel, an unknown amount of fluoride and you end with $1.261\ \mathrm{g}$ of nickelfluoride. The mass of fluoride in the end must then be $1.261\ \mathrm{g} - 0.766\ \mathrm{g} = 0.495\ \mathrm{g}$.

So, in the end there's $0.013$ moles of nickel and $0.026$ moles of fluoride. Now the question is, what is the empirical formula for nickel fluoride? ($\small\ce{NiF2}$). I see that for every mole of nickel you need $2$ moles of fluoride but then why is the empirical formula $\small\ce{Ni2F4}$ incorrect?

An empirical formula is, according to IUPAC’s Gold Book definition:

Formed by juxtaposition of the atomic symbols with their appropriate subscripts to give the simplest possible formula expressing the composition of a compound

Wikipedia goes even further:

the simplest positive integer ratio of atoms of each element present in a compound

and then

An empirical formula makes no reference to isomerism, structure, or absolute number of atoms.

So, NiF2 is an empirical formula, while Ni2F4 simply isn't. Another example is given by, again, Wikipedia:

For example, formaldehyde, acetic acid and glucose have the same empirical formula, CH2O

Their molecular formulas (which account for the absolute number of atoms in a molecule) differ, though: formaldehyde CH2O, acetic acid C2H4O2, glucose C6H12O6.

• And what about H2O2? – Jef Oct 5 '12 at 21:45
• H2O2 is a valid molecular formula, but not an empirical formula. The empirical formula is HO. – F'x Oct 5 '12 at 22:19