Xenon of mass 5.08 g reacts with fluorine to form 9.49 g of a xenon fluoride. What is the empirical formula of this
I am not familiar with these types of empirical formula problems. If I'm given the percentages of the elements in the compound, I know what to do, but I have no idea what to do here.
Chemistry - Olivia, Monday, March 23, 2009 at 10:54am
Oh, and the answer is A. I just don't understand how you get it.
Chemistry - Dr Russ, Monday, March 23, 2009 at 10:58am
This is very like the percentage problems. A useful way to do the percentage problems is to follow the steps
1. Percentages (make sure they add up to 100)
2. masses (assuming 100 g total so numerically same as percentages)
3. convert mass to moles (by dividing by relative atomic mass)
4. divide by smallest number of moles (which gives ratio of the atoms for empirical formula)
here we have the masses so start at step 3.
Chemistry - Olivia, Monday, March 23, 2009 at 11:04am
If I don't know the empirical formula, what am I supposed to use for the relative atomic mass here for xenon fluoride to get moles?
Chemistry - Olivia, Monday, March 23, 2009 at 11:07am
Never mind! :) I think I see what I need to do. I simply subtract the initial mass from the mass of the compound to get the mass of the Chlorine and go from there. Correct?
Chemistry - Dr Russ, Monday, March 23, 2009 at 11:49am
Yes, subtract the two so you have
4.41 g 5.08 g
Number of moles?
Divide by smallest.