Final answer:
The correct result for the equilibrium state in a solution with 0.1M CH₃COOH and 0.1M HCl, considering that HCl is a strong acid providing a much higher concentration of [H+], is A. [H+] = 0.1M from HCl.
Step-by-step explanation:
To determine which of the following results is correct for the equilibrium state in a solution originally having 0.1M CH₃COOH (acetic acid) and 0.1M HCl, we must consider the strong acid's effect on the acetic acid dissociation. Since HCl is a strong acid, it will completely dissociate in water, providing a [H+] concentration of 0.1M due to the HCl alone. This high concentration of [H+] will suppress the dissociation of acetic acid, meaning that the acetic acid will not contribute significantly to the [H+] in solution.
The given Ka of CH₃COOH is 1.8 × 10⁵, which should be 1.8 × 10⁻⁵ if we assume it is written incorrectly. Normally, we would calculate [H+] at equilibrium and the degree of dissociation of acetic acid using the ice table method and the Ka value, but since the presence of HCl dominates, these calculations are not necessary for this scenario. Hence, the correct option is A [H+] = 0.1M, because the contribution of H+ from HCl will far exceed that from the weak acid CH₃COOH.