Below is a representation of an aqueous solution of a weak acid HA at equilibrium. (Each circle represents 1.0 ×10^-3 mol of atoms, and the volume of the box is 1.0 L. Solvent water molecules are not shown for clarity.) What is the pH of the solution? A) 2.4 B) 3.0 C) 2.1 D) 11.0 E) 11.6

Question

Quizplus · Accepted Answer

The equilibrium representation shows that some of the weak acid HA has dissociated into its conjugate base A- and H+. Since HA is a weak acid, it will not fully dissociate and a significant amount of HA will remain in solution. The pH of the solution can be calculated using the equation: pH = -log[H+]. Since the concentration of H+ in the solution is not given, we need to use the equilibrium constants to calculate it. The equilibrium constant expression for the dissociation of a weak acid is Ka = [A-][H+]/[HA]. We can set up an ICE table to solve for [H+]:
HA + H2O ⇌ H3O+ + A-
I       0.002            0            0   
C   -x              +x       +x 
E  0.002-x        x           x

Plugging these values into the Ka expression and solving for x, we get:
Ka = x^2/(0.002-x) = 1.6 x 10^-5
x = [H+] = 0.0040 M
Therefore, pH = -log(0.0040) = 3.0 approximately. Choice B is the best answer.

Below Is a Representation of an Aqueous Solution of a Weak