Question

Given reaction is $4Fe{S_2} + 11{O_2} \to 2F{e_2}{O_3} + 8S{O_2}$. Given $75.0$ grams of $Fe{S_2}$ , how many grams of sulfur dioxide are produced ( $MM$in $g/mol$: $Fe{S_2} = 119.99,{O_2} = 32,F{e_2}{O_3} = 159.7,S{O_2} = 64.07$ ) ?
A) $40.0g$
B) $70.0g$
C) $150g$
D) $80.0g$

Answer 1

Here we know the mass of $Fe{S_2}$ is given and the point is we have to convert it to mole. Next for finding the number of moles of sulfur dioxide we have to take a particular mole ratio between $Fe{S_2}$ and $S{O_2}$. Then we will be getting an exact answer which is the molar mass in a state of number of moles and we have to convert that molar mass from mole to mass as the question is asking about mass of $S{O_2}$.

Complete step-by-step answer:
First of all before proceeding to solution, we have to present down the given equation which is as follows,
$4Fe{S_2} + 11{O_2} \to 2F{e_2}{O_3} + 8S{O_2}$
Next we have to write down the mass of $Fe{S_2}$ which we shall convert to mole. For this particular conversion, actually we cannot convert it just simply as we need to find the molar mass of the compound which we can obtain from periodic table where there is an equivalence statement of $1molFe{S_2} = 119.98Fe{S_2}$ can be derived because we should need this to convert to moles.
The given mass of $Fe{S_2} = 75.0g$
Next we shall convert this to moles as now we are provided with all necessary quantities as in following way,
Number of moles of $Fe{S_2} = 75.0g \times \dfrac{{1mol}}{{119.99g}}$
Number of moles of $Fe{S_2} = 0.625mol$
Next we need to have a look on to the balanced equation . There we can see a mole ratio between $Fe{S_2}$ and $S{O_2}$ which is as in the ratio $4:8$ as there are four moles of $Fe{S_2}$ and eight moles of $S{O_2}$ in the given balanced reaction. Thus we can see a relation which converts number of moles of $Fe{S_2}$ to number of moles of $S{O_2}$ which will be as follows,
Number of moles of $S{O_2}$ $= 0.625molFe{S_2} \times \dfrac{{8molS{O_2}}}{{4molFe{S_2}}}$
Number of moles of $S{O_2} = 1.250mol$
Next we are proceeding to the final solution. We have to know that the question is asking about the mass of $S{O_2}$ which is produced in the reaction. So , in this case what we have to do is we need to again convert the value from number of mole to a particular mass as shown in below,
Mass of $S{O_2} = 1.250molS{O_2} \times \dfrac{{64.07gS{O_2}}}{{1molS{O_2}}}$
Mass of $S{O_2} = 80.09g$

Hence, option D is correct which is $80.09g$

Note: First of all we need to get familiar with the given balanced equation well so as to find out the relation of mole ratio between $Fe{S_2}$ and $S{O_2}$ . Then only we can proceed further to get the final answer. Next thing is that it is not easy to remember the molar mass of every compound. So, it will be easy as it can be obtainable from periodic tables where certain equivalence statements can be derived to move forward in calculation. An important point we have to know is about conversion of number of moles to particular mass as whenever in the question it will be asking to find out the mass of a given substance.