Acetylsalicylic acid (ASA) is the pain reliever in Aspirin. ASA is a weak acid with [tex]K_a = 3.2 \times 10^{-4}[/tex]. One Aspirin tablet was ground up and dissolved in 25.0 mL of water. It was titrated using 0.15 mol/L sodium hydroxide. To reach the endpoint, 98.3 mL of base was required.

Determine the concentration of the ASA solution.

To determine the concentration of the acetylsalicylic acid (ASA) solution, we first recognize that the titration involves a neutralization reaction between ASA, a weak acid, and sodium hydroxide (NaOH), a strong base. Given that the endpoint was reached with 98.3 mL of 0.15 mol/L NaOH, we can calculate the moles of NaOH used in the titration.

Step 1: Calculate the moles of NaOH used.

Moles of NaOH = Volume (L) × Concentration (mol/L)Step 2: Determine the moles of ASA.

Since NaOH and ASA react in a 1:1 molar ratio (as implied by the neutralization reaction), the moles of ASA at the endpoint will equal the moles of NaOH used.

Moles of ASA = 0.014745 mol

Step 3: Calculate the concentration of the ASA solution.

Concentration (mol/L) = Moles of solute / Volume of solution (L)

The original volume of the ASA solution was 25.0 mL (or 0.025 L). Therefore, the concentration of ASA in the solution is:

Concentration of ASA = 0.014745 mol / 0.025 L = 0.5898 mol/L

This calculation reveals that the concentration of the ASA solution was 0.5898 M.