For an arbitrary nonelectrolyte, calculate the freezing point when #"0.0192 mols"# of it is dissolved in #"5.00 g"# of water? #K_f = 1.86^@ "C/m"# for water.
1 Answer
Jan 26, 2017
The idea is that dissolving anything into a solvent will decrease its freezing point. It is expected that you already know
Recall:
#DeltaT_f = T_f - T_f^"*" = -iK_fm# where:
#i = 1# is the van't Hoff factor for a non-electrolyte. It simply says that only one nonelectrolyte particle exists for every nonelectrolyte particle that is placed into solution.#K_f = 1.86^@ "C/m"# is the freezing point depression constant for water.#m# is the molality of the solution, i.e.#"mol solute"/"kg solvent"# .#T_f# is the freezing point of the solution.#T_f^"*"# is the freezing point of the pure solvent.
First, calculate what
#DeltaT_f = -(1)(1.86^@ "C/m")(("0.0192 mols solute")/(5.00xx10^(-3) "kg H"_2"O"))#
#= -7.14^@ "C"#
Therefore, the final freezing point is:
#color(blue)(T_f) = DeltaT_f + T_f^"*" = -7.14^@ "C" + 0^@ "C" = color(blue)(-7.14^@ "C")#
