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I understand that the solubility (in terms of moles/volume) of group 2 halides increase with increase in anion size, i.e.

$$\ce{MF2 < MCl2 < MBr2 < MI2},$$

where $\ce{M = Mg, Ca, Sr, Ba},\dots$ due to large decreases in lattice enthalpy.

But what is the explanation for the following discrepancies?

  • $\ce{BeF2 > MgF2 = CaF2 < SrF2 < BaF2}$
  • $\ce{BeCl2 < MgCl2 < CaCl2 > SrCl2 > BaCl2}$

In other words, how can I determine the solubility order of $\ce{BX2, MgX2, CaX2, SrX2, BaX2}$ (where $\ce{X = F, Cl, Br, I}$)? The given orders of solubility have been verified with data.

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closed as too broad by Nilay Ghosh, M.A.R., Jan, Tyberius, Todd Minehardt Nov 9 '17 at 1:17

Please edit the question to limit it to a specific problem with enough detail to identify an adequate answer. Avoid asking multiple distinct questions at once. See the How to Ask page for help clarifying this question. If this question can be reworded to fit the rules in the help center, please edit the question.

  • $\begingroup$ Think about what influences lattice enthalpy. $\endgroup$ – gsurfer04 Oct 21 '17 at 11:13
  • $\begingroup$ What do you mean by solubility? Mass per volume/mass, moles per volume/moles? What about hydrates? $\endgroup$ – aventurin Oct 21 '17 at 13:25
  • $\begingroup$ My data was based on mass per volume. But your right, when comparing solubility, the solubility should be in terms of moles per volume. I am seeing if some change is required for the solubility orders. No i'm not talking about hydrates. $\endgroup$ – Pranoy De Oct 21 '17 at 17:19
  • $\begingroup$ No, the order remains the same. $\endgroup$ – Pranoy De Oct 21 '17 at 17:54
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To me it does not make too much sense to try to compare and explain the solubility (mass concentration) of anhydrous group 2 chlorides.

First of all, mass concentration does not seem to be the adequate physical quantity for comparison. Some of the chlorides have such high solubility that their saturated solution contains significant less water per volume than others. Molar fractions of solute and water would be a better choice.

Second, all of the chlorides precipitate as the hydrated salt from water. At normal temperature these are

  • $\ce{BeCl2\cdot 4 H2O}$
  • $\ce{MgCl2\cdot 6 H2O}$
  • $\ce{CaCl2\cdot 6 H2O}$
  • $\ce{SrCl2\cdot 6 H2O}$
  • $\ce{BaCl2\cdot 2 H2O}$
  • $\ce{RaCl2\cdot 2 H2O}$

Besides differences in structure their different stoichiometry makes a comparison difficult. And, most important, the formation of hydrates shows that we probably should not discuss the solubility of anhydrous chlorides at all.

Third, some anhydrous halides are more covalent than ionic compounds ($\ce{BeF2, BeCl2}$). This makes it difficult to argue with lattice energies and ionic radii.

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  • $\begingroup$ So, because of their significant differences in structure and properties (ionic or covalent), lattice enthalpy and hydration enthalpy are not adequate/appropriate to explain the solubility of anhydrous group 2 halides. Am I right? $\endgroup$ – Pranoy De Oct 26 '17 at 5:49

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