A monoprotic weak acid, when dissolved in water, is 0.92% dissociated and produces a solution with pH = 3.42. Calculate K_a for the acid. A) 1.4 * 10^-⁷ B) 2.8 * 10^-³ C) 3.5 * 10^-⁶ D) We need to know the initial concentration of the acid. E) none of these

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Quizplus · Accepted Answer

K_a can be calculated using the equation: pH = (1/2)(pKa - log(K_a)) where pKa is the acid dissociation constant and K_a is the molar solubility of the acid. Rearranging this equation and plugging in the given values, we get: K_a = 10^(-pH + (1/2)(pKa + log(0.92))) We are not given the value of pKa, but we can estimate it using the Henderson-Hasselbalch equation: pH = pKa + log([A-]/[HA]) We know that the acid is monoprotic, so [A-] = [H+], and we can assume that the initial concentration of the acid is equal to the concentration of the undissociated acid (because the acid is weak and only slightly dissociates). Therefore: 3.42 = pKa + log(0.0092) pKa = 3.42 - log(0.0092) = 3.42 + 2.04 = 5.46 Now we can substitute pKa and the given values into the first equation: K_a = 10^(-3.42 + (1/2)(5.46 + log(0.92))) K_a = 3.5 x 10^-10

A Monoprotic Weak Acid, When Dissolved in Water, Is 0