Multiple Choice
Calculate the pH of a 0.800 M NaCH3CO2 solution. Ka for acetic acid, CH3CO2H, is 1.8 × 10^-5.
17. Acid and Base Equilibrium
pH of Weak Acids
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Consider the titration of 25.00 mL of 0.150 M hydrazoic acid (Ka = 1.90 x 10^-5) with 0.200 M NaOH. (Hydrazoic acid is HN3, not to be confused with ammonia, NH3) What is the pH of the hydrazoic acid solution prior to beginning the titration?
A
pH = 2.67
B
pH = 4.12
C
pH = 5.00
D
pH = 3.45
Verified step by step guidance1
Identify the type of problem: This is a weak acid equilibrium problem, as we are dealing with hydrazoic acid (HN3), a weak acid, before any titration begins.
Write the equilibrium expression for the dissociation of hydrazoic acid: HN3 ⇌ H+ + N3-. The equilibrium constant expression (Ka) is given by Ka = [H+][N3-]/[HN3].
Set up an ICE (Initial, Change, Equilibrium) table to determine the concentrations of the species at equilibrium. Initially, [HN3] = 0.150 M, and [H+] and [N3-] are both 0 M.
Express the changes in concentration in terms of x, where x is the concentration of H+ ions at equilibrium. The equilibrium concentrations will be [HN3] = 0.150 - x, [H+] = x, and [N3-] = x.
Substitute the equilibrium concentrations into the Ka expression: 1.90 x 10^-5 = (x)(x)/(0.150 - x). Solve this equation for x, which represents the [H+], and then calculate the pH using the formula pH = -log[H+].
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