Question

In aqueous solution, pyridine (C₅H₅N) is a weak base that accepts a proton from water to form the hydroxide ion (OH⁻) and the pyridinium ion (C₅H₅NH), according to the following equation: C₅H₅N(aq) + H₂O(l) ⇌ OH⁻(aq) + C₅H₅NH(aq). The base-dissociation constant (Kb) for this base is 1.7 × 10⁻⁹. If the pH of the solution is 9.25, what is the initial concentration of pyridine? (Assume that the temperature is 25°C.)

Answer 1

The initial concentration of pyridine in a solution with a pH of 9.25 can be determined by using the base-dissociation constant (Kb), the water autoionization constant (Kw), and pH to calculate hydroxide ion concentration and finally the pyridine concentration at equilibrium.

Step-by-step explanation:

To find the initial concentration of pyridine in an aqueous solution with a pH of 9.25, we can utilize the relationship between the base-dissociation constant (Kb) and the ionization of the base. Given that Kb for pyridine is 1.7 × 10⁻⁹, and the water's autoionization constant (Kw) is 1.0 × 10⁻¹⁴ at 25°C, we can calculate the pKa of the conjugate acid (pyridinium ion) using the equation pKa+pKb=pKw. The pH given allows us to find the concentration of hydroxide ions, [OH⁻], and subsequently the equilibrium concentration of pyridine.

First, we convert pH to [H³O⁺] using the equation pH = -log[H³O⁺]. Next, we find [OH⁻] using Kw = [H³O⁺][OH⁻]. With the equilibrium concentration for [OH⁻] and the given Kb, the initial concentration of pyridine can be determined by setting up an ICE table and solving for the initial concentration.