`C_(7)H_(6)O_(3) + C_(4)H_(6)O_(3) rightarrow C_(9)H_(8)O_(4) + C_(2)H_(4)O_(2)`
What is the percent yield if 0.85 g of aspirin is formed in the reaction of 1.00 g of salicylic acid with excess acetic anydride ?

`{:("Substance", "Molar Mass"), (C_(7)H_(6)O_(3), 138.12g.mol^(-1)), (C_(4)H_(6)O_(3), 102.09g.mol^(-1)), (C_(9)H_(8)O_(4), 180.15g. mol^(-1)), (C_(2)H_(4)O_(2), 60.05g.mol^(-1)):}`

Options

A. `65%`
B. `77%`
C. `85%`
D. `91%`

Explanation

To calculate the percent yield of aspirin formed in the reaction between salicylic acid and acetic anhydride, we will follow these steps:

### Step 1: Calculate the number of moles of salicylic acid (C₇H₆O₃)
Given:
- Mass of salicylic acid = 1.00 g
- Molar mass of salicylic acid (C₇H₆O₃) = 138.12 g/mol

Using the formula for moles:
\[ \text{Number of moles} = \frac{\text{mass}}{\text{molar mass}} \]
\[ \text{Number of moles of C₇H₆O₃} = \frac{1.00 \, \text{g}}{138.12 \, \text{g/mol}} \approx 0.00724 \, \text{mol} \]

### Step 2: Determine the theoretical yield of aspirin (C₉H₈O₄)
From the balanced chemical equation, we see that 1 mole of salicylic acid produces 1 mole of aspirin. Therefore, the number of moles of aspirin produced will also be 0.00724 mol.

Now, we calculate the theoretical mass of aspirin:
- Molar mass of aspirin (C₉H₈O₄) = 180.15 g/mol

Using the moles calculated:
\[ \text{Mass of C₉H₈O₄} = \text{Number of moles} \times \text{Molar mass} \]
\[ \text{Mass of C₉H₈O₄} = 0.00724 \, \text{mol} \times 180.15 \, \text{g/mol} \approx 1.30 \, \text{g} \]

### Step 3: Calculate the percent yield
Given that the actual yield of aspirin is 0.85 g, we can calculate the percent yield using the formula:
\[ \text{Percent Yield} = \left( \frac{\text{Actual Yield}}{\text{Theoretical Yield}} \right) \times 100 \]
Substituting the values:
\[ \text{Percent Yield} = \left( \frac{0.85 \, \text{g}}{1.30 \, \text{g}} \right) \times 100 \approx 65.38\% \]

### Final Answer
The percent yield of aspirin is approximately **65%**.

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