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Chapter 14: Problem 57

Calculate the concentration of an aqueous HI solution that has \(\mathrm{pH}=2.50 .\) HI is a strong acid.

Short Answer

Expert verified
The concentration of the aqueous HI solution can be calculated using the given pH value and the formula [H₃O⁺] = 10^(-pH). With a pH of 2.50, [H₃O⁺] ≈ 3.16 × 10⁻³ M. Since HI is a strong acid and dissociates completely in water, the concentration of the HI solution is approximately 3.16 × 10⁻³ M.

Step by step solution

01

Understand the definition of pH

pH is defined as the negative logarithm of the H₃O⁺ ion concentration: pH = -log₁₀([H₃O⁺]) We are given the pH value (2.50), and we need to find the concentration of the H₃O⁺ ions, represented as [H₃O⁺].
02

Calculate the concentration of H₃O⁺ ions

To find [H₃O⁺], we can use the given pH value by rearranging the pH formula: [H₃O⁺] = 10^(-pH) Substitute the given pH value (2.50) into the formula: [H₃O⁺] = 10^(-2.50) Calculate the concentration of H₃O⁺ ions: [H₃O⁺] ≈ 3.16 × 10⁻³ M
03

Determine the concentration of the HI solution

Since HI is a strong acid, it dissociates completely in water. That means that the concentration of H₃O⁺ ions is equal to the concentration of the HI solution. Therefore, the concentration of the aqueous HI solution is approximately 3.16 × 10⁻³ M.

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Most popular questions from this chapter

What are the major species present in \(0.015 M\) solutions of each of the following bases? a. \(\mathrm{KOH}\) b. \(\mathrm{Ba}(\mathrm{OH})_{2}\) What is \(\left[\mathrm{OH}^{-}\right]\) and the \(\mathrm{pH}\) of each of these solutions?

Determine \(\left[\mathrm{OH}^{-}\right],\left[\mathrm{H}^{+}\right]\), and the \(\mathrm{pH}\) of each of the following solutions. a. \(1.0 \mathrm{M} \mathrm{KCl}\) b. \(1.0 \mathrm{M} \mathrm{KF}\)

A solution is prepared by adding \(50.0 \mathrm{~mL}\) of \(0.050 M \mathrm{HBr}\) to \(150.0 \mathrm{~mL}\) of \(0.10 \mathrm{M}\) HI. Calculate the concentrations of all species in this solution. HBr and HI are both considered strong acids.

A sample containing \(0.0500 \mathrm{~mol} \mathrm{Fe}_{2}\left(\mathrm{SO}_{4}\right)_{3}\) is dissolved in enough water to make \(1.00 \mathrm{~L}\) of solution. This solution contains hydrated \(\mathrm{SO}_{4}{ }^{2-}\) and \(\mathrm{Fe}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}{ }^{3+}\) ions. The latter behaves as an acid: $$\mathrm{Fe}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}^{3+}(a q) \rightleftharpoons \mathrm{Fe}\left(\mathrm{H}_{2} \mathrm{O}\right)_{5} \mathrm{OH}^{2+}(a q)+\mathrm{H}^{+}(a q)$$ a. Calculate the expected osmotic pressure of this solution at \(25^{\circ} \mathrm{C}\) if the above dissociation is negligible. b. The actual osmotic pressure of the solution is \(6.73\) atm at \(25^{\circ} \mathrm{C}\). Calculate \(K_{\mathrm{a}}\) for the dissociation reaction of \(\mathrm{Fe}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}^{3+}\). (To do this calculation, you must assume that none of the ions goes through the semipermeable membrane. Actually, this is not a great assumption for the tiny \(\mathrm{H}^{+}\) ion.)

At \(25^{\circ} \mathrm{C}\), a saturated solution of benzoic acid \(\left(K_{\mathrm{a}}=6.4 \times 10^{-5}\right)\) has a pH of \(2.80\). Calculate the water solubility of benzoic acid in moles per liter.
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