Why do increasing concentrations of zinc chloride not linearly decrease the freezing point?

Question

I did an experiment on how different concentrations of $\ce{ZnCl2}$ affect the freezing point of demi water. My results were that till a concentration of about $10\ \mathrm{mol\,dm^{-3}}$ the freezing point decreased (more or less in a linear manner). After that at concentrations of $15~\mathrm{mol\,dm^{-3}}$ the freezing point started to increase again (it again approached $0~\mathrm{^\circ C}$).

The graph below shows the freezing point depression of different $\ce{ZnCl2}$ solutions. The $x$-axis is the concentration of $\ce{ZnCl2}$ in $\mathrm{mol\,dm^{-3}}$, the $y$-axis is the freezing point depression in kelvin.

Why does the freezing point depression initially decrease, and then increase again after a certain concentration has been reached?

Answer 1

You've reached the other side of the eutectic.

See, that graph of lowering freezing temperature is always a part of a bigger picture.

Once you are past that minimum, it is no longer freezing point that you are measuring. Instead, it is the point where the solution becomes saturated and $\ce{ZnCl2*6H2O}$ (or whatever hydrate is stable at those conditions) starts crystallizing.

You'd see a similar pattern with pretty much any soluble salt. Zinc chloride, if anything, is good in that it dissolves immensely well, so the graph is wide, and its features can be seen more clearly.

Answer 2

For the freezing point depression to keep increasing with increasing solute the $\ce{Zn^{2+}}$ and $\ce{Cl^-}$ ions need to be free of interactions with each other. In concentrated solutions this isn't true for two reasons.

First there aren't enough water molecules. Each ions actually has severance layers of water molecules in a hydration cluster.

Second the cations and anions attract in concentrated solutions. So the activity of the ions doesn't increase linearly with concentration.