What is the pH of an aqueous solution of 0.364 M ethylamine (a weak base with the formula #C_2H_5NH_2#)?
1 Answer
Explanation:
In order to be able to solve this problem, you need to know the value of the base dissociation constant,
#K_b = 5.6 * 10^(-4)#
http://www.bpc.edu/mathscience/chemistry/table_of_weak_bases.html
So, ethylamine is a weak base, which means that it does not ionize completely in aqueous solution to form ethylammonium cations,
An equilibrium will be established between the unionized ethylamine molecules and the two ions that result from its ionization.
Use this equilibrium reaction as a base for an ICE table to find the equilibrium concentration of hydroxide anions
#" ""C"_ 2"H"_ 5"NH"_ (2(aq)) + "H"_ 2"O"_ ((l)) rightleftharpoons "C"_ 2"H"_ 5"NH"_ (3(aq))^(+) + "OH"_((aq))^(-)#
By definition, the base dissociation constant for this equilibrium will be
#K_b = ( ["C"_2"H"_5"NH"_3^(+)] * ["OH"^(-)])/(["C"_2"H"_5"NH"_2])#
In your case, this will be equivalent to
#5.6 * 10^(-4) = (x * x)/(0.364 - x) = x^2/(0.364 - x)#
Rearrange to get a quadratic equation
#x^2 + 5.6 * 10^(-4) * x - 2.0384 * 10^(-4) = 0#
This quadratic will produce two solutions, one positive and one negative
#x_1 = color(red)(cancel(color(black)(-0.01456)))" "# and#" "x_2 = 0.014 color(white)(a)color(green)(sqrt())#
Since
#x = 0.014#
This means that the equilibrium concentration of hydroxide anions will be
#["OH"^(-)] = "0.014 M"#
At this point, you can calculate the pOH of the solution by using
#color(purple)(|bar(ul(color(white)(a/a)color(black)("pOH" = - log(["OH"^(-)]))color(white)(a/a)|)))#
You will have
#"pOH" = - log(0.014) = 1.85#
In aqueous solution at room temperature, you have the following relationship between pOH and pH
#color(purple)(|bar(ul(color(white)(a/a)color(black)("pOH " + " pH" = 14)color(white)(a/a)|)))#
This means that the pH of the solution will be
#"pH" = 14 - 1.85 = color(green)(|bar(ul(color(white)(a/a)12.15color(white)(a/a)|)))#