Question #62d6f
1 Answer
Explanation:
The idea here is to use the ammonium ion's acid dissociation constant,
The relationship between the base dissocaition constant, the acid dissociation constant, and water's self-ionization constant,
#K_W = K_a xx K_b#
In your case, the base dissociation constant of ammonia will be
#K_b = K_W/K_a = 10^(-14)/(5.62 * 10^(-10)) = 1.78 * 10^(-5)#
To determine the pH of the solution, you first need to know the
Ammonia will react with water to form ammonium and hydroxide ions. Use an ICE table to determine the equilibrium concentration of hydroxide ions
#"NH"_text(3(aq]) " "+" " "H"_2"O"_text((l]) " "rightleftharpoons" " "NH"_text(4(aq])^(+) " "+" " "OH"_text((aq])^(-)#
By definition, the base dissociation constant will be
#K_b = (["NH"_4^(+)] * ["OH"^(-)])/(["NH"_3]) = (x * x)/(0.1 - x)#
Because the value of
#(0.1 - x) ~~ 0.1#
This means that you have
#K_b = x^2/0.1 = 1.78 * 10^(-5)#
#x = sqrt(0.1 * 1.78 * 10^(-5)) = 0.001334#
This means that you have
#["OH"^(-)] = x = "0.001334 M"#
The
#p"OH" = -log(["OH"^(-)])#
#p"OH" = -log(0.001334) = 2.87#
The pH of the solution will thus be
#p"H" = 14 - p"OH" = 14 - 2.87 = color(green)(11.13)#