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In my lab, I mixed 8:2 (mass:mass) of Hexane:Ethanol

(Hexane: 0 ppm of water content, analytical grade, & ethanol: 4% water content, pharmacy grade)

At room temperature, and at -20ºC, there was a clear phase separation.

I repeated the experiment, but using anhydrous ethanol instead (1300 ppm of water content).

Now, at room temperature, and at -20ºC, there was no phase separation.

I added enough water to the mixture to make it equal to the first mixture, and the phase separation came clearly into view.

My question is, is this kind of very specific data of miscibilities available somewhere? Ethanol 96% is what is commonly used, yet pure ethanol behaves differently in this case.

Any link or insight into the importance of such little water is appreciated.

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    $\begingroup$ With anhydrous ethanol and hexane you have a binary system (hexane:ethanol), with 96% Ethanol and hexane you have a ternary system (hexane:ethanol:water). Yes they have different phase diagrams. Why would you expect anything different? $\endgroup$ – MaxW Apr 1 at 17:53
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    $\begingroup$ I'm just surprised that a small amount of water makes such a difference. Plus, any polar solvent is going to have a decent amount of water in it. So next time I see one used in a text, I'll be doubting just how anhidrous it is. $\endgroup$ – Juan Perez Apr 1 at 18:15
  • $\begingroup$ Welcome to StackExchange Chemistry. Water definitely does not mix with hexane, so there will be two phases for sure. Now, ethanol has a choice between the more polar phase (water) and the less polar (hexane). Most of it will go to the polar phase. $\endgroup$ – Karsten Theis Apr 1 at 18:17
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    $\begingroup$ This is kind of an extreme case. Anhydrous ethanol and hexane are "just barely miscible"; the next step up in polarity (anhydrous methanol) is not miscible with hexane in all proportions. Furthermore, 4% water represents a significant amount of "contamination" with the least miscible common solvent. Even non-anhydrous solvents are often purchased/used with <1000 ppm water, and 1% is 10000 ppm. For most purposes solvent miscibility tables work fine even with slightly impure solvents, but for these aqueous mixtures, you may just have to test things yourself on a small scale. $\endgroup$ – Nicolau Saker Neto Apr 1 at 21:39
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Theory and Application of solubility Parameters are beyond our scope. However, an interesting viewer should have at least read through collection of research by Charles M. Hansen, published during 60s (Ref.1). One of the parameters discussed in this book is Hildebrand solubility parameter ($\delta$), which is a good indication of solubility, according to Wikipedia. The principal utility of Hildebrand solubility parameter is that it provides simple predictions of phase equilibrium, based on a single parameter. These predictions are often useful, particularly for nonpolar and slightly polar systems without hydrogen bonding. According to the tabulated $\delta$ values, ethanol is polar, and has relatively higher value ($\pu {12.92 cal^{\frac{1}{2}}cm^{\frac{3}{2}}}$) than that of hexane ($\pu {7.24 cal^{\frac{1}{2}}cm^{\frac{3}{2}}}$) (Ref.1). Yet, it is interesting to note that hexane is soluble in absolute ethanol, but is insoluble in methanol ($\delta_{\ce{MeOH}}=\pu {14.28 cal^{\frac{1}{2}}cm^{\frac{3}{2}}}$), albeit the difference of relevant Hildebrand solubility parameters (${\Delta\delta}$ is just $\pu {1.36 cal^{\frac{1}{2}}cm^{\frac{3}{2}}}$ (from Ref.1 and 2).

It is an interesting aspect of the Hildebrand solvent parameters that the $\delta$ value of a solvent mixture can be determined by averaging the $\delta$ values of the individual solvents by volume. For example, if you mix $V_\mathrm{A}$ of solvent A with $V_\mathrm{B}$ of solvent B, the mixture will have Hildebrand value of $\delta_\mathrm{Mix}$ which iswould be given by: $$\delta_\mathrm{Mix} = \delta_\mathrm{A}\left( \frac{V_\mathrm{A}}{ V_\mathrm{A}+ V_\mathrm{B}}\right) + \delta_\mathrm{B}\left( \frac{V_\mathrm{B}}{ V_\mathrm{A}+ V_\mathrm{B}}\right)$$

Let's see what is the $\delta$ of 96% ethanol ($\delta_{\ce{H2O}}=\pu {23.5 cal^{\frac{1}{2}}cm^{\frac{3}{2}}}$): $$ 23.5\left(\frac{4}{100}\right) + 12.92\left(\frac{96}{100}\right) = 13.34 = \delta_{96\%\ce{EtOH}} \lt \delta_{\ce{MeOH}}$$

Let's see now, what is the value of $\delta$ of 90% ethanol: $$ 23.5\left(\frac{10}{100}\right) + 12.92\left(\frac{90}{100}\right) = 14.0 = \delta_{90\%\ce{EtOH}} \lt \delta_{\ce{MeOH}}$$

Even though still $\delta_{90\%\ce{EtOH}} \lt \delta_{\ce{MeOH}}$, the concept is clear. The value is more than $\delta_{\ce{EtOH}}$ and getting closer by each addition of water droplet and to been not miscible with hexane. Therefore, it is safe to say that these data would gives you some insight to what you are looking for.

Note: For a review of solution systems: See ref.3


References:

  1. C. M. Hansen, In The Three Dimensional Solubility Parameter and Solvent Diffusion Coefficient: Their Importance in Surface Coating Formulation; Danish Technical Press: Copenhagen, Denmark, 1967.
  2. C. M. Hansen, "The Three Dimensional Solubility Parameter Key to Paint Component Affinities: 1. Solvents Plasticizers, Polymers, and Resins," Journal of Paint Technology 1967, 39(505), 104-117.
  3. R. Wiśniewski, E. Śmieszek, E. Kamińska, " Three-dimensional solubility parameters: simple and effective determination of compatibility regions," Progress in Organic Coatings 1995, 26(2-4), 265-274 (https://doi.org/10.1016/0300-9440(96)81583-9).
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