The problem asks to calculate the pH of a buffer system containing 1.5 M $\text{HClO}$ and 1.7 M $\text{NaClO}$, using the dissociation constant ($K_a$) of the acid $\text{HClO}$.

This is a typical buffer problem, and the pH of a buffer solution can be calculated using the Henderson-Hasselbalch equation:

[ \text{pH} = \text{p}K_a + \log \left( \frac{[\text{Base}]}{[\text{Acid}]} \right) ]

Where:

• $[\text{Base}]$ is the concentration of $\text{ClO}^-$ from $\text{NaClO}$,
• $[\text{Acid}]$ is the concentration of $\text{HClO}$,
• $\text{p}K_a = -\log(K_a)$.

From the problem, we know:

• $[\text{HClO}] = 1.5 \, \text{M}$,
• $[\text{NaClO}] = 1.7 \, \text{M}$.

We still need the $K_a$ value for hypochlorous acid $(\text{HClO})$ at 25°C. The typical value for $K_a$ of $\text{HClO}$ is around $2.9 \times 10^{-8}$. Now, let's proceed with the calculation.

1. Calculate pKa: $\text{p}K_a = -\log(2.9 \times 10^{-8}) \approx 7.54$

2. Apply the Henderson-Hasselbalch equation: $\text{pH} = 7.54 + \log \left( \frac{1.7}{1.5} \right)$ $\text{pH} = 7.54 + \log(1.133) \approx 7.54 + 0.054 = 7.59$

Therefore, the pH of the buffer system is approximately 7.59.

Would you like further details on this solution or any clarifications?

Related Questions:

1. How does the Henderson-Hasselbalch equation apply to buffer systems?
2. What factors influence the pH of a buffer solution?
3. Why is it important to use significant figures when calculating pH?
4. How does temperature affect the $K_a$ value of acids like $\text{HClO}$?
5. Can you calculate the pH if only $\text{HClO}$ or $\text{NaClO}$ was present?

Tip:

Always check the units and make sure concentrations are in molarity (M) when applying the Henderson-Hasselbalch equation.