Find empirical formula when individual masses are known

Question

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 nickel fluoride. 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 are $0.013$ 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. Then why is the empirical formula $\small\ce{Ni2F4}$ incorrect?

Answer 1

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.