Answer :
To determine the empirical formula of a compound given the percentages of its elements, follow these steps:
1. Determine the mass percentages of the elements in grams:
Given that the compound contains 36.76% iron (Fe) and 21.11% sulfur (S), we can assume a 100-gram sample for simplicity. This implies we have:
- 36.76 grams of iron (Fe)
- 21.11 grams of sulfur (S)
2. Calculate the moles of each element:
We use the molar masses of iron and sulfur to convert the mass into moles. The molar masses are:
- Iron (Fe): 55.845 g/mol
- Sulfur (S): 32.06 g/mol
Using the formula [tex]\( \text{moles} = \frac{\text{mass (g)}}{\text{molar mass (g/mol)}} \)[/tex]:
For iron (Fe):
[tex]\[ \text{moles of Fe} = \frac{36.76 \text{ g}}{55.845 \text{ g/mol}} = 0.6582 \text{ mol} \][/tex]
For sulfur (S):
[tex]\[ \text{moles of S} = \frac{21.11 \text{ g}}{32.06 \text{ g/mol}} = 0.6585 \text{ mol} \][/tex]
3. Calculate the mole ratio:
To find the mole ratio, divide each element's moles by the smallest number of moles calculated:
- For Fe:
[tex]\[ \text{ratio of Fe} = \frac{0.6582}{0.6582} = 1.0 \][/tex]
- For S:
[tex]\[ \text{ratio of S} = \frac{0.6585}{0.6582} = 1.0 \][/tex]
4. Derive the empirical formula:
The mole ratio for both elements is approximately 1:1. Therefore, the empirical formula of the compound is:
[tex]\[ \text{Fe}_1\text{S}_1 \text{ or simply } \text{FeS} \][/tex]
So, the correct answer is:
- [tex]$Fe$[/tex] [tex]\( \square\)[/tex] = 1
- [tex]$S$[/tex] [tex]\( \square\)[/tex] = 1
The empirical formula is [tex]\( \text{FeS} \)[/tex].
1. Determine the mass percentages of the elements in grams:
Given that the compound contains 36.76% iron (Fe) and 21.11% sulfur (S), we can assume a 100-gram sample for simplicity. This implies we have:
- 36.76 grams of iron (Fe)
- 21.11 grams of sulfur (S)
2. Calculate the moles of each element:
We use the molar masses of iron and sulfur to convert the mass into moles. The molar masses are:
- Iron (Fe): 55.845 g/mol
- Sulfur (S): 32.06 g/mol
Using the formula [tex]\( \text{moles} = \frac{\text{mass (g)}}{\text{molar mass (g/mol)}} \)[/tex]:
For iron (Fe):
[tex]\[ \text{moles of Fe} = \frac{36.76 \text{ g}}{55.845 \text{ g/mol}} = 0.6582 \text{ mol} \][/tex]
For sulfur (S):
[tex]\[ \text{moles of S} = \frac{21.11 \text{ g}}{32.06 \text{ g/mol}} = 0.6585 \text{ mol} \][/tex]
3. Calculate the mole ratio:
To find the mole ratio, divide each element's moles by the smallest number of moles calculated:
- For Fe:
[tex]\[ \text{ratio of Fe} = \frac{0.6582}{0.6582} = 1.0 \][/tex]
- For S:
[tex]\[ \text{ratio of S} = \frac{0.6585}{0.6582} = 1.0 \][/tex]
4. Derive the empirical formula:
The mole ratio for both elements is approximately 1:1. Therefore, the empirical formula of the compound is:
[tex]\[ \text{Fe}_1\text{S}_1 \text{ or simply } \text{FeS} \][/tex]
So, the correct answer is:
- [tex]$Fe$[/tex] [tex]\( \square\)[/tex] = 1
- [tex]$S$[/tex] [tex]\( \square\)[/tex] = 1
The empirical formula is [tex]\( \text{FeS} \)[/tex].
Answer:
FeS
Explanation:
1. Determine the mass percentages of the elements in grams:
Given that the compound contains 36.76% iron (Fe) and 21.11% sulfur (S), we can assume a 100-gram sample for simplicity. This implies we have:
- 36.76 grams of iron (Fe)
- 21.11 grams of sulfur (S)
2. Calculate the moles of each element:
We use the molar masses of iron and sulfur to convert the mass into moles. The molar masses are:
- Iron (Fe): 55.845 g/mol
- Sulfur (S): 32.06 g/mol
Using the formula \( \text{moles} = \frac{\text{mass (g)}}{\text{molar mass (g/mol)}} \):
For iron (Fe):
\[
\text{moles of Fe} = \frac{36.76 \text{ g}}{55.845 \text{ g/mol}} = 0.6582 \text{ mol}
\]
For sulfur (S):
\[
\text{moles of S} = \frac{21.11 \text{ g}}{32.06 \text{ g/mol}} = 0.6585 \text{ mol}
\]
3. Calculate the mole ratio:
To find the mole ratio, divide each element's moles by the smallest number of moles calculated:
- For Fe:
\[
\text{ratio of Fe} = (0.6582)/(0.6582) = 1.0
\]
- For S:
\[
\text{ratio of S} = (0.6585)/(0.6582) = 1.0
\]
4. Derive the empirical formula:
The mole ratio for both elements is approximately 1:1. Therefore, the empirical formula of the compound is:
\[
\text{Fe}_1\text{S}_1 \text{ or simply } \text{FeS}
\]
So, the correct answer is:
- $Fe$ \( \square\) = 1
- $S$ \( \square\) = 1
The empirical formula is \( \text{FeS} \).
