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To calculate the pH of a solution with a hydronium ion concentration \([H_3O^+]\), the formula used is:

\[
\text{pH} = -\log[H_3O^+]
\]

Given in the image, \([H_3O^+] = 8.8 \times 10^{-5}\) M.

Let's calculate the pH:

\[
\text{pH} = -\log(8.8 \times 10^{-5})
\]

\[
\text{pH} \approx -(\log(8.8) + \log(10^{-5}))
\]
\[
\log(8.8) \approx 0.944
\]
\[
\log(10^{-5}) = -5
\]
\[
\text{pH} = -(0.944 - 5) = 4.056
\]

So, the pH of the solution is approximately **4.06**, considering two significant figures (based on the two significant figures in \(8.8\)).

Would you like a detailed breakdown or have any questions?

Here are five related questions for further exploration:
1. How does the concentration of hydronium ions affect pH levels?
2. What is the relationship between pH and pOH in a solution?
3. How would the pH change if the hydronium concentration were \(8.8 \times 10^{-4}\) M instead?
4. What does a pH of 4.06 suggest about the acidity of the solution?
5. How does temperature affect the pH of a solution?

**Tip**: Remember that the significant figures in a logarithmic calculation are determined by the number of significant figures in the original number used in the logarithm.

Calculate the pH of a solution that has [H₃O⁺] = 8.8 × 10⁻⁵ M. Be sure your answer has the correct number of significant figures.

Learn how to calculate the pH of a solution with a hydronium ion concentration of 8.8 × 10⁻⁵ M. This solution uses logarithmic properties and considers significant figures, ideal for high school chemistry students.

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