Final answer:
To determine the pH of a 0.100 M acetic acid solution, an ICE table is used to calculate [H+] as 1.32 × 10⁻³ M. The pH is then found by taking the negative logarithm of [H+], resulting in a pH of 2.879.
Step-by-step explanation:
The subject of the question is to calculate the pH of a 0.100 M solution of acetic acid with a known Ka value. To solve it, we use an ICE table for the ionization of acetic acid.
First, express the ionization:
CH₃COOH(aq) → H⁺(aq) + CH₃COO⁻(aq)
Then set up an ICE table:
- Initial concentrations are [CH₃COOH]=0.100
- The change in concentration for [H⁺] and [CH₃COO⁻] is +x and -x for [CH₃COOH]
- At equilibrium, [CH₃COOH]=0.100-x, [H⁺]=x, [CH₃COO⁻]=x
Next, plug the values into the acid dissociation constant equation, Ka = 1.8 x 10⁻⁵ = x²/(0.100-x). Assuming x << 0.100, Ka ≈ x²/0.100. Solve for x to get [H⁺] = 1.32 × 10⁻³ M.
Finally, calculate the pH:
pH = -log(1.32 × 10⁻³) = 2.879