This is the question that I'm trying to solve, can anyone help me out.

Elemental sulfur (1.256 g) is combined with fluorine, F2, to give a compound with the formula SFx, a very stable, colorless gas. If you have isolated 5.722 g of SFx, what is the value of x?

Molorov,

I may have a possible solution to your problem since it kind of resembles several, similar, problems I'm having problems with.

First thing I did was to state the chemical equation.  (It may or may not help.)
S₈(s) + F₂(g) --> SF_x(g)

Then I stated the masses of the given compounds:

mass SFx = 5.722g SFx
mass S = 1.256g S

Then I calculated the mass percentages of S and F in the compound:
%S in SFx = (1.256g / 5.722g) * 100% = 21.95%
%F in SFx = 100% - 21.95% = 78.05%

Next I converted those percentages to moles:
mol S = 21.95g S * (1 mol S / 32.07g S) = .6844 mol S
mol F = 78.05g F * (1 mol F / 19.00g F) = 4.108 mol F

Next I find the simplest whole number mole ratios between S and F by
by dividing both moles by smallest mole of the two:
S = .6844 / .6844 = 1
F = 4.108 / .6844 = 6

Finally I find the empirical formula for SFx: S₁F₆

Don't know if its right or not, but that's how I would do it.  

HDC