Calculate the [H⁺] in an aqueous solution that is 0.0755 M HF and 0.100 M NaF .The value of K_a for HF is 3.5 * 10^-⁴.

Question

Quizplus · Accepted Answer

We need to use the equilibrium expression for the reaction: HF + H2O ⇌ H3O+ + F- Since HF is a weak acid, it will not dissociate completely, so we can use the Henderson-Hasselbalch equation: pH = pKa + log([A-]/[HA]) where pKa is the acid dissociation constant, [A-] is the concentration of the conjugate base (F-), and [HA] is the concentration of the acid (HF). pKa = -log(Ka) = -log(3.5 x 10^-4) = 3.455 plugging in the values: pH = 3.455 + log(0.100/0.0755) = 3.655 [H3O+] = 10^-pH = 2.511 x 10^-4 M [F-] = 0.100 M [HF] = 0.0755 M Now we can use the expressions for [H^{+] and [}-]: [H^{+] = [HF] - [H3O+] = 0.0755 - 2.511 x 10^-4 = 0.0753 M
[}-] = [F-] = 0.100 M [H^+][-] = (0.0753)(0.100) = 0.00753 [H⁺] = [H^{+] + [}-] = 0.0753 + 0.100 = 0.1753 M However, we need to take into account the fact that F- can react with H3O+ to produce HF: F- + H3O+ ⇌ HF + H2O Kb = Kw/Ka = 1.00 x 10^-14/3.5 x 10^-4 = 2.857 x 10^-11 Let x be the concentration of H3O+ that reacts with F-. Then: x^2/[HF] = Kb x^2 = (2.857 x 10^-11)(0.0755) x = 1.256 x 10^-6 M This concentration is small compared to [H3O+], so we can neglect it in calculating the pH and in calculating [H⁺]. Thus, the final answer is: [H⁺] = 0.1753 M, or approximately 0.176 M. The closest answer choice is C.

Calculate the [H+] in an Aqueous Solution That Is 0

Calculate the [H⁺] in an Aqueous Solution That Is 0