Question

Question #17cfc

Answer 1

Here's what I got.

Explanation:

The first thing you need to do here is to figure how many molecules of hydrogen sulfide you have in your sample.

To do that, use the molar mass of the compound to calculate how many moles it contains

`8 color(red)(cancel(color(black)("g"))) * ("1 mole H"_2"S")/(34.1color(red)(cancel(color(black)("g")))) = "0.2346 moles H"_2"S"`

Now, each mole of hydrogen sulfide will contain Avogadro's constant of hydrogen molecules.

`color(blue)(ul(color(black)("1 mole H"_ 2"S" = 6.022 * 10^(23)"molecules H"_2"S")))`

This means that your sample contains

`0.2346 color(red)(cancel(color(black)("moles H"_2"S"))) * (6.022 * 10^(23)"molecules H"_2"S")/(1color(red)(cancel(color(black)("mole H"_2"S"))))`

`= 1.413 * 10^(23)"molecules H"_2"S"`

Now focus on finding the number of electrons present in one molecule of hydrogen sulfide. Each hydrogen atom will contribute `1` electron to the total, so you have

`2 xx "1 e"^(-) = "2 e"^(-) ->` from hydrogen

A neutral sulfur atom has an atomic number equal to `16`, which means that it contains `18` protons inside its nucleus and `18` electrons surrounding its nucleus.

This means that you have

`1 xx "16 e"^(-) = "16 e"^(-) ->` from sulfur

The total number of electrons present in an `"H"_2"S"` molecule is

`"2 e"^(-) + "16 e"^(-) = "18 e"^(-)`

Your sample will thus contain

`1.413 * 10^(23) color(red)(cancel(color(black)("molecules H"_2"S"))) * "18 e"^(-)/(1color(red)(cancel(color(black)("molecule H"_2"S"))))`

`= color(darkgreen)(ul(color(black)(2.5 * 10^(24)"e"^(-))))`

I'll leave the answer rounded to two sig figs, but keep in mind that you only have one significant figure for the mass of the sample.