We know that like dissolves like. And dichloromethane is a polar solvent and water is also a polar solvent. Also there ought to be a strong hydrogen bonding between the chlorine and hydrogen atoms. So why is it immiscible?

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    $\begingroup$ The quick way to answer is that DCM isn't polar enough, as well as the strength of the H-bond it might form in water. A detailed answer involves the enthalpy end entropy terms analysis of the dissolution but I am afraid it will remain rather qualitative. And see also chemistry.stackexchange.com/questions/38137/… where interestingly an azeotrop is mentioned. It is quite possible that ideally mixing results in a ordered phase, so that the heat of solubilisation doesn't compensate for the reduced entropy. $\endgroup$
    – Alchimista
    Mar 30 '19 at 9:07
  • $\begingroup$ @Alchimista What do you mean by ordered phase? $\endgroup$ Mar 30 '19 at 9:12
  • $\begingroup$ @Alchimista. Dipole moment of DCM is 1.47 and that of water is 1.85. $\endgroup$ Mar 30 '19 at 9:14
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    $\begingroup$ It can be that entropy is lower as for the water might have to arrange more strictly to accommodate the dcm molecules. This would give a positive delta G as for the intermolecular forces are weaker for sure. $\endgroup$
    – Alchimista
    Mar 30 '19 at 9:16
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    $\begingroup$ For an immiscible compound, DCM is quite soluble in water (~20g will dissolve in a litre of water). As will chloroform (~10g/L) which is enough to make a recreational drink (or was in Victorian times before we worked out how toxic it was). $\endgroup$
    – matt_black
    Mar 30 '19 at 20:37

In the following figure the free energy computed from solubilities available in the wikipedia are plotted against T. The positive values of $\Delta G$ imply that formation of a 1 molal solution from the pure components is an endergonic process (requires energy input), consistent with the limited solubility of DCM in water.

enter image description here

From the slope and intercept we see that the enthalpy of solubilization of DCM in water is negative (exothermic) while the entropy change is also negative. The negative entropy change, as explained in the comments, opposes mixing. The exothermic character means that increasing T actually discourages solubilization (the negative entropic effect is emphasized at higher T).

This does not address the mechanism of solubilization, of course. However, it supports the argument in the comments that formation of water cages or conformationally restrictive hydrogen bonding to the solute imposes an entropic penalty upon solubilization.

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    $\begingroup$ Nice formulation for what I tried to convey. I am just surprised by the fact that the mixing has an exothermic character, as for I would have expected water to water interaction to be stronger as well, so that the entropy term could impose at the reltively small concentrations also mentioned in comments. $\endgroup$
    – Alchimista
    Mar 31 '19 at 18:26
  • $\begingroup$ Thanks @Alchimista This is admittedly just an extension of the argument you raised. I did not come across a more systematic description of the solvation process in this particular system, the details are missing. $\endgroup$
    – Buck Thorn
    Mar 31 '19 at 19:50

The physical causes of solubility of substances in each other can semiquantitatively be expressed by the Hansen solubility parameters.

$\begin{array}{} \text{Water}:&&\delta_d=15.6&\delta_p=16.0&\delta_h=42.3&R_0=47.8\\ \text{Dichloromethane}:&&\delta_d=18.2&\delta_p=6.3 & \delta_h=6.1&R_0=20.2 \end{array}$

$$R_0=20.2,\ \ \ R_a=37.8$$


RED < 1 would show solubility of both substances in each other. But RED > 1, so both substances are little soluble in each other.

The solubility parameters above show, the energies from polar and from hydrogen bridging bond intermolecular attractions are higher in water.

Generally, chlorinated hydrocarbons have lower energies from polar and hydrogen bridging bond intermolecular attractions.

The strongest hydrogen bridging bonds are formed between the atoms of the strongly electronegative elements $\ce{F}$, $\ce{O}$ and $\ce{N}$. But the atoms of other electronegative elements, e.g. $\ce{Cl}$-, $\ce{S}$- and $\ce{C}$-atoms, are capable of weak hydrogen bridging bonds at suitable molecular structure.

The atomic radii of the third period elements ($\ce{P}$, $\ce{S}$, $\ce{Cl}$) are much larger than those of the second period elements ($\ce{N}$, $\ce{O}$, $\ce{F}$). $\ce{Cl}$-atoms have lower charge density therefore. They are weaker proton acceptors therefore.

Hydrogen bridging bonds can be proved e.g. by molecular spectroscopy.

Li Bian: Proton Donor Is More Important Than Proton Acceptor in Hydrogen Bond Formation: A Universal Equation for Calculation of Hydrogen Bond Strength. J. Phys. Chem. A 107 (2003) (51) 11517–11524

Weinhold, F.; Klein, R. A.: What is a hydrogen bond? Mutually consistent theoretical and experimental criteria for characterizing H-bonding interactions. Mol. Phys. 110 (2012) (9-10) 565-579

Gilli, G.; Gilli, P.: The Nature of the Hydrogen Bond: Outline of a Comprehensive Hydrogen Bond Theory. Oxford University Press, Oxford, 2013


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