# Question 383e6

Jun 6, 2015

Your tablet contains 0.33 g of aspirin, and has a percent composition of aspirin of 66%.

The key to this problem is the balanced chemical equation for this neutralization reaction

${C}_{9} {H}_{8} {O}_{4 \left(a q\right)} + N a O {H}_{\left(a q\right)} \to {C}_{9} {H}_{7} {O}_{4} N {a}_{\left(a q\right)} + {H}_{2} {O}_{\left(l\right)}$

Notice the $1 : 1$ mole ratio that exists between aspirin and sodium hydroxide. This tells you that, in order for a complete neutralization to take place, you need equal numbers of moles of each compound.

Since you know the volume and molarity of the sodium hydroxide solution used in the titration, you can determine the number of moles of base used

$C = \frac{n}{V} \implies n = C \cdot V$

${n}_{N a O H} = \text{0.10 M" * 18.30 * 10^(-3)"L" = "0.00183 moles}$ $N a O H$

This means that the tablet contained

0.00183cancel("moles"NaOH) * "1 mole aspirin"/(1cancel("mole"NaOH)) = "0.00183 moles" ${C}_{9} {H}_{8} {O}_{4}$

Use aspirin's molar mass to determine how many grams would contain that many moles

$0.00183 \cancel{\text{moles") * "180.16 g"/(1cancel("mole")) = color(green)("0.33 g}}$ ${C}_{9} {H}_{8} {O}_{4}$

The percent composition of aspirin in your tablet will be

(0.33cancel("g"))/(0.50cancel("g")) * 100 = color(green)("66%")#

SIDE NOTE Both answers ar rounded to two sig figs.