Weak Acids and Common-Ion Effect
This page delves deeper into weak acid calculations and introduces the common-ion effect.
The page starts with an example problem: calculating the [H+] and pH of a 0.05 F acetic acid solution. It demonstrates the step-by-step process for solving weak acid equilibrium problems.
Vocabulary: Formality (F) is equivalent to molarity (M) for these calculations.
The solution involves setting up an ICE (Initial, Change, Equilibrium) table and using the Ka expression to solve for the equilibrium concentrations.
Highlight: For weak acids where Ka ≤ 10^-4, the x in the denominator can be neglected to simplify calculations.
The page also includes an example of calculating the initial concentration of nitrous acid (HNO2) given a specific pH value.
Example: To find [HNO2] for a solution with pH 2.6, first calculate [H3O+] = antilog(-2.6) = 2.5119 x 10^-3 M, then use the Ka expression to solve for the initial concentration.
Finally, the page introduces the common-ion effect:
Definition: The common-ion effect is a phenomenon where the dissociation of a weak electrolyte is decreased by adding a strong electrolyte that has an ion in common with the weak electrolyte.
This effect shifts the equilibrium of the weak electrolyte towards the undissociated form, reducing its ionization.
Example: Adding sodium acetate (NaC2H3O2) to an acetic acid (HC2H3O2) solution will shift the equilibrium to the left, reducing the dissociation of acetic acid.
