Now calculate Kb the base dissociation constant for [C2H3O2-], acetate anion, for each of your trials from the concentrations of species at the equivalence point.
Report 3 significant figures, e.g. 5.97E-10.
I don't have access to your data but I expect you need to utilize the OH^-/HAc values you found.
........Ac^- + HOH ==> HAc + OH^-
I.......C..............0.......0
C......-x.............x.........x
E......C-x.............x.........x
C under acetate is concn of the salt. I suspect you measured the pH of the solution, converted that to OH^- and plugged into Kb = (x)^2/C-x. That gives you an experimental value for Kb. You can compare that with the actual value by Kb = (Kw/Ka) = (1E-14/Ka for acetic acid) = 5.56E-10
To calculate the base dissociation constant, Kb, for the acetate anion ([C2H3O2-]), we need to know the concentrations of species at the equivalence point. However, you have not provided the concentrations for the trials. Could you please provide the concentrations of species ([C2H3O2-], [OH-], [CH3COOH]) at the equivalence point for each trial?
To calculate the base dissociation constant, Kb, for [C2H3O2-] (acetate anion) at the equivalence point, you would need the concentrations of the species involved.
Assuming you have the concentrations of [C2H3O2-] and the weak acid, let's call it HA, at the equivalence point, you can use the following steps to calculate Kb:
1. Write down the balanced chemical equation for the dissociation of the weak acid, HA, into its conjugate base, C2H3O2-, and the hydronium ion, H3O+.
HA ⇌ C2H3O2- + H3O+
2. Set up an ICE table (Initial, Change, and Equilibrium) to track the changes in the concentrations of the species. Since this is the equivalence point, the initial concentrations of the acid and its conjugate base are equal, and the initial concentration of the hydronium ion is zero.
Initial:
HA : [HA]
C2H3O2- : [C2H3O2-]
H3O+ : 0
Change:
HA : -x
C2H3O2- : +x
H3O+ : +x
Equilibrium:
HA : [HA] - x
C2H3O2- : [C2H3O2-] + x
H3O+ : x
3. Write the expression for the base dissociation constant, Kb, which is the ratio of the concentration of the conjugate base to the concentration of the weak acid at equilibrium.
Kb = ([C2H3O2-] + x) / ([HA] - x)
4. At the equivalence point, the concentrations of the acid and the conjugate base are equal, so [HA] - x = [C2H3O2-] + x.
5. Simplify the expression by substituting this equality into the Kb equation:
Kb = (2[C2H3O2-]) / [C2H3O2-]
Kb = 2
Therefore, for any weak acid at the equivalence point, the base dissociation constant, Kb, will always be 2.
In conclusion, the value of Kb for [C2H3O2-], acetate anion, at the equivalence point is 2, with 3 significant figures.