The pH of a solution containing HIO3 (Ka = 1.7 × 10⁻¹) and KIO3 is 1.00. If the molarity of HIO3 is 0.025 M, what is the molarity of KIO3?

• A. 0.050 M
• B. 0.075 M
• C. 0.100 M
• D. 0.025 M

Answer: (C)

Verified step by step guidance

1. Start by understanding that HIO3 is a weak acid and KIO3 is its conjugate base. The solution is a buffer solution, and we can use the Henderson-Hasselbalch equation to find the molarity of KIO3.
2. The Henderson-Hasselbalch equation is: ${\mathrm{pH}}_{}={\mathrm{pK}}_a+\log \left(\frac{\mathrm{[A^-]}}{\mathrm{[HA]}}\right)$, where [A^-] is the concentration of the conjugate base (KIO3) and [HA] is the concentration of the acid (HIO3).
3. Calculate the pKa using the given Ka value: ${\mathrm{pK}}_a=-\log \left(K_a\right)$. Substitute Ka = 1.7 × 10⁻¹ into the equation to find pKa.
4. Substitute the given pH value (1.00), the calculated pKa, and the molarity of HIO3 (0.025 M) into the Henderson-Hasselbalch equation to solve for the molarity of KIO3.
5. Rearrange the equation to solve for [A^-] (KIO3): $\log \left(\frac{\mathrm{[A^-]}}{\mathrm{[HA]}}\right)$ = pH - pKa. Then, calculate [A^-] using the antilog function.