Question

of a chemical reaction.

1. list several reasons why the actual yield from a chemical reaction is not usually equal to the theoretical yield.
2. explain how percent yield is calculated.
3. apply in an experiment, you combine 83.77 g of iron with an excess of sulfur and then heat the mixture to obtain iron (iii) sulfide.

$2\text{fe}(\text{s}) + 3\text{s}(\text{s}) \
ightarrow \text{fe}_2\text{s}_3(\text{s})$
what is the theoretical yield, in grams, of iron (iii) sulfide?

Explanation:

Step1: Calculate moles of Fe

Molar mass of Fe = 55.85 g/mol.
$n_{\text{Fe}} = \frac{83.77\ \text{g}}{55.85\ \text{g/mol}} = 1.500\ \text{mol}$

Step2: Relate moles of Fe to $\text{Fe}_2\text{S}_3$

From reaction: 2 mol Fe → 1 mol $\text{Fe}_2\text{S}_3$.
$n_{\text{Fe}_2\text{S}_3} = 1.500\ \text{mol} \times \frac{1\ \text{mol}\ \text{Fe}_2\text{S}_3}{2\ \text{mol}\ \text{Fe}} = 0.7500\ \text{mol}$

Step3: Calculate mass of $\text{Fe}_2\text{S}_3$

Molar mass of $\text{Fe}_2\text{S}_3$ = $2\times55.85 + 3\times32.07 = 207.91\ \text{g/mol}$.
$m_{\text{Fe}_2\text{S}_3} = 0.7500\ \text{mol} \times 207.91\ \text{g/mol} = 155.9\ \text{g}$

Step4: List reasons for yield difference

1. Side reactions consume reactants.
2. Incomplete reaction of reactants.
3. Product loss during separation/purification.
4. Impurities in starting materials.

Step5: Explain percent yield calculation

Percent yield compares actual to theoretical yield, expressed as a percentage.
Formula: $\text{Percent Yield} = \frac{\text{Actual Yield}}{\text{Theoretical Yield}} \times 100\%$

Answer:

32.

• Side reactions form unwanted products
• Reactants may not fully react
• Product is lost during collection/purification
• Starting materials contain impurities

1. Percent yield is calculated by dividing the actual amount of product obtained (actual yield) by the maximum possible amount of product (theoretical yield), then multiplying the result by 100%: $\text{Percent Yield} = \frac{\text{Actual Yield}}{\text{Theoretical Yield}} \times 100\%$
2. 155.9 g