# Question a40ec

Sep 27, 2017

$\text{41.2 g}$

#### Explanation:

The balanced chemical equation that describes your reaction looks like this

${\text{S"_ (8(s)) + 24"F"_ (2(g)) -> 8"SF}}_{6 \left(g\right)}$

Notice that the reaction consumes $24$ moles of fluorine gas for every $1$ mole of octatomic sulfur that takes part in the reaction.

This means that in order for your sample of sulfur, which consists of

11.6 color(red)(cancel(color(black)("g"))) * "1 mole S"_8/(256.52color(red)(cancel(color(black)("g")))) = "0.04522 moles S"_8

to be completely consumed by the reaction, you need to provide

0.04522 color(red)(cancel(color(black)("moles S"_8))) * "24 moles F"_2/(1color(red)(cancel(color(black)("mole S"_8)))) = "1.085 moles F"_2#

To convert the mass of fluorine gas to moles, use the molar mass of diatomic fluorine

$1.085 \textcolor{red}{\cancel{\textcolor{b l a c k}{\text{moles F"_2))) * "37.997 g"/(1color(red)(cancel(color(black)("mole F"_2)))) = color(darkgreen)(ul(color(black)("41.2 g}}}}$

The answer is rounded to three sig figs, the number of sig figs you have for the mass of sulfur.