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Chapter 14: Problem 56

An antacid purchased at a local drug store has a pOH of \(2.3 .\) Calculate the \(\mathrm{pH},\left[\mathrm{H}^{+}\right],\) and \(\left[\mathrm{OH}^{-}\right]\) of this solution. Is the antacid acidic or basic?

Short Answer

Expert verified
The antacid has a pH of 11.7, making it basic. The concentration of H⁺ ions is approximately \(2.0 \times 10^{-12} \, M\), and the concentration of OH⁻ ions is approximately \(5.0 \times 10^{-3} \, M\).

Step by step solution

01

Calculate pH

To calculate the pH of the solution, we can use the formula: pH = 14 - pOH Given pOH = 2.3, we plug it into the formula: pH = 14 - 2.3 = 11.7
02

Calculate the concentration of H⁺ ions

To calculate the concentration of H⁺ ions, we use the equation: \([H^+]=10^{-\mathrm{pH}}\) Now, plug in the pH value calculated in step 1: \([H^+] = 10^{-11.7}\) \([H^+] \approx 2.0 \times 10^{-12} \, M\)
03

Calculate the concentration of OH⁻ ions

To calculate the concentration of OH⁻ ions, we will use the equation: \([OH^-] = 10^{-\text{pOH}}\) Now, plug in the given pOH value: \([OH^-] = 10^{-2.3}\) \([OH^-] \approx 5.0 \times 10^{-3} \, M\)
04

Determine if the antacid is acidic or basic

Now that we know that the pH of the antacid solution is 11.7, we can use the pH scale to determine if the solution is acidic, neutral, or basic: - A solution with pH less than 7 is acidic - A solution with pH equal to 7 is neutral - A solution with pH greater than 7 is basic Since the pH of the antacid solution is greater than 7, the antacid is considered to be basic.

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Most popular questions from this chapter

When someone hyperventilates, a condition known as respiratory alkalosis can occur. Explain the cause and effect of respiratory alkalosis. Hint: Reference Exercises 146 and 147 .

A typical sample of vinegar has a pH of \(3.0 .\) Assuming that vinegar is only an aqueous solution of acetic acid \(\left(K_{\mathrm{a}}=1.8 \times\right.\) \(10^{-5}\) ), calculate the concentration of acetic acid in vinegar.

A \(0.100-\mathrm{g}\) sample of the weak acid \(\mathrm{HA}\) (molar mass \(=\) 100.0 \(\mathrm{g} / \mathrm{mol} )\) is dissolved in 500.0 \(\mathrm{g}\) water. The freezing point of the resulting solution is \(-0.0056^{\circ} \mathrm{C}\) . Calculate the value of \(K_{\mathrm{a}}\) for this acid. Assume molality equals molarity in this solution.

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A solution is made by adding 50.0 \(\mathrm{mL}\) of 0.200\(M\) acetic acid \(\left(K_{\mathrm{a}}=1.8 \times 10^{-5}\right)\) to 50.0 \(\mathrm{mL}\) of \(1.00 \times 10^{-3} \mathrm{M} \mathrm{HCl}\) a. Calculate the pH of the solution. b. Calculate the acetate ion concentration.
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