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](http://goldbook.iupac.org/E02063.html):



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


[Wikipedia](http://en.wikipedia.org/wiki/Empirical_formula) goes even further:



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


and then



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


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
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> 
> 


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


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