Can someone help me with this following chem problem?

A compound containing only xenon and fluorine can be obtained by shining light on a flask that contain .526 g of xenon and excess fluorine gas. When all the xenon was consumed, 678 of the new compound was obtained. Determine the empirical formula of the compound.

I first found the molar mass of both xenon and fluorine. Xe: 131.29 g/mol and F 18.99 g/mol. The # of moles is the mass in grams / molar mass.

I got:

.526/ 131.29 = .00401 moles for xenon.

.678/ 18.99 = .0357 moles for Fluorine.

Where do I go from here? Does the mole ratio need to be expressed as a whole number?

one mistake

if you had 0.526 g xenon, and all of it was consumed, then 0.526 g of the resulting compound weight is due to xenon. That means the weight due to whatever amount of fluorine was used must be the new compound weight minus the weight due to xenon, or 0.678g - 0.526g = 0.152g.

so the moles of fluorine come from this 0.152g, not from the overall mass of compound obtained. Use this mass to find moles, as you did before but with the wrong number.

then divide by the lowest moles #. The empirical formula is typically written as the lowest whole number ratio... so when you divide out you might not be done... simply multiply both numbers until you get a whole number ratio.