# Why not to consider hydrogen and oxygen moles to determine an empirical formula

A question says:

Find the empirical formula of an organic compound from the following composition:

34.62% C, 3.88% H, 61.50% O.

The answer is $\ce{C3H4O4}$. It was found by using the mass (percentages) divided by the molar mass of each element.

But they didn't consider the (2moles) of hydrogen and oxygen in calculations. Why? in this case it will be $\ce{CH3O3}$

On the other hand, in a similar question in the text book, they consider the moles of hydrogen and oxygen in calculations by multiplying by 2 after getting the moles.

Could anybody explain?

The answer is $\ce{C3H4O4}$. It was found by using the mass (percentages) divided by the molar mass of each element.
But they didn't consider the (2moles) of hydrogen and oxygen in calculations. Why? In this case it will be $\ce{CH3O3}$
It is completely irrelevant whether an element usually exists as a diatomic molecule in nature. If that would matter, how would you then treat sulfur-containing compounds? Note that natural sulfur typically consists of $\ce{S8}$ molecules! How would you treat elements that only appear in the form of compounds in nature but never as pure elements.