Why, out of the main halogens, is bromine the most soluble in water?

Why, out of the main halogens, is bromine $(\ce{Br2})$ the most soluble in water?


Why is there such a drastic decrease for iodine $(\ce{I2})?$


Here is a table from my inorganic chemistry textbook with the numbers that prompted this question.


$$
\begin{array}{lr}
\hline
\text{Halogen} & s/\pu{g L^-1}~\ce{H2O} \\
\hline
\ce{F2} & \text{explos. reac.} \\
\ce{Cl2} & 6.5 \\
\ce{Br2} & 33.6 \\
\ce{I2} & 0.33 \\
\hline
\end{array}
$$




In the case of iodine versus bromine, the sharp drop in solubility may be attributed in part to iodine showing incipient intermolecular bonding -- as a metal would have between its atoms. Such a shift towards metallic character would be expected to lead to reduced solubility in nonmetallic liquids such as water or even carbon tetrachloride (see figures quoted [here](https://chemistry.stackexchange.com/questions/113138/solubility-of-iodine-in-some-organic-solvents)).


When bromine is solidified, the bond length within a diatomic molecule is nearly the same as in the gas phase. [Wikipedia](https://en.wikipedia.org/wiki/Bromine), citing Reference [1], gives the solid phase bond length as 227 pm versus the gas phase bond length of 228 pm. With iodine, however, condensation leads to a lengthening of the intramolecular bond, indicating a delocalization of the bonding into intermolecular regions, From [Wikipedia](https://en.wikipedia.org/wiki/Iodine), also citing Reference [1]:



> 
> Gaseous iodine is composed of I2 molecules with an I–I bond length of 266.6 pm. ... It is even longer (271.5 pm) in solid orthorhombic crystalline iodine, which has the same crystal structure as chlorine and bromine. As such, within the iodine molecule, significant electronic interactions occur with the two next-nearest neighbours of each atom, and these interactions give rise, in bulk iodine, to a shiny appearance and semiconducting properties. Iodine is a two-dimensional semiconductor with a band gap of 1.3 eV (125 kJ/mol): it is a semiconductor in the plane of its crystalline layers and an insulator in the perpendicular direction.
> 
> 
> 


The tendency of iodine towards adopting a metallic structure is also evident from studies of the transformation of the solid element under high pressure. [Duan et al. [2]](https://doi.org/10.1103/PhysRevB.76.104113) report that bromine is calculated to undergo an insulator-metal transition at about 55 GPa pressure. With iodine, [Pasternak et al. [3]](https://doi.org/10.1103/PhysRevLett.58.575) identify formation of zigzag chains and resultant netallization with 16 GPa pressure, less than one-third the calculated requirement for bromine metallization.


A final consideration: Given the role of incipient metallic character and intermolecular bond delocalization just discussed, the solubility of iodine in water might be improved by partially recovering these characteristics in the solute. This is where the well-known impact of an iodide salt comes in: iodine is much more soluble in water that contains iodide because the triiodide ion, $\ce{I3^-}$, with its delocalized three-center four-electron bond, is formed in the solution. Such a behavior is not significantly carried over to any lighter halogens.


**References**


1.
Greenwood, Norman N.; Earnshaw, Alan (1997). *Chemistry of the Elements* (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8. pp 800-804.


2.
Defang Duan, Yanhui Liu, Yanming Ma, Zhiming Liu, Tian Cui, Bingbing Liu, and Guangtian Zou (2002). "Ab initio studies of solid bromine under high pressure".
*Phys. Rev. B* **76**, 104-113. <https://doi.org/10.1103/PhysRevB.76.104113>


3.
M. Pasternak, J. N. Farrell, and R. D. Taylor (1987). "Metallization and structural transformation of iodine under pressure: A microscopic view".
*Phys. Rev. Lett.* **58**, 575. <https://doi.org/10.1103/PhysRevLett.58.575>


 7

Solubility is the confluence of many factors and can be looked at through many chemical frameworks; it might first be a good idea to lay out our framework for understanding solubility before tackling the question:


One relatively simple-but-useful framework for solubility can be to see it as the thermodynamic stability of the substance *when solvated* compared to the thermodynamic stability of the substance when interacting *only with itself*. We'll start there.


Knowing that pressure and temperature affect the phases a substance or mixture—i.e. the way and degree to which it self-interacts—the conditions we are looking at are going to be an important consideration whenever we are asking a question like this. At any given pressure and temperature, if the substance energetically prefers self-intermolecular interactions (a.k.a. cohesion,) then it will not be favorable to solvate. If the molecule is energetically ambivalent to interactions with itself vs. interactions with the solvent, then this would be an example of *miscibility.* This ambivalence is what defines an "ideal solution" when generalizing the ideas of ideal gases to other phases. Following this trend, energetic *favorability* of a solvated state over a non-solvated state at a certain pressure and temperature will thus increase solubility.


So the question we are asking under this framework specifically is "Why, at standard temperature and pressure, are interactions between elemental bromine and water significantly more stabilizing than interactions between water and elemental chlorine or elemental iodine?"


This is where the molecular complexity begins to have a big impact on the answer: unlike more fundamental qualities like electronegativity or size, there is no singular trend across the halogens that can be used as a smoking gun for the solubility of their elemental forms. I'm not fully convinced that the phase of the elemental form aids in this solubility much ("bromine is a liquid, so it mixes easier" e.g.,) specifically because these figures define the **concentration of a saturated solution** and not the ease/rate of solvation which would be more of a kinetic effect than a thermodynamic effect.


Whenever I see an optimum in the middle of any trend (high solubility in the middle of the halogens, e.g.,) I usually look for two "forces" of similar magnitude that work in opposite directions to each other. The two in this question I'm most drawn to (there could be more, and probably is) are molecular size and the σ-hole size/halogen-bonding (XB) ability (we'll get to this part if you haven't heard of it, hang tight.)


Let's start with size. The larger the surface-area of a non-polar molecule is, the more susceptible it is to cohesion via van der Waals interactions. By this logic, chlorine should be the most soluble, followed by bromine, and then iodine. Because we see bromine breaking this trend, there must be a trend that works in the opposite direction.


A viable trend working in the opposite direction regards the (usually ignored) non-homogeneity/anisotropy of electron-density in singly bound halogens: it turns out that, instead of electrons smearing equally across the sphere of the halogen atom, there is a concentration of positive charge collinear with the axis of the σ-bond and the electron-density moves to form a belt orthogonal to the σ-bond.


This effect is more pronounced for larger halogens than smaller halogens (pictured below) for many reasons, but mostly it's a combination of: larger polarizability of the electron distribution; low *s*-orbital character in the σ-bond; and a lower stability of the σ-bond decreasing the energy of the σ\* orbital, leading to some destructive interference in the molecular orbital. The effect this has on intermolecular interactions has led to the phenomenon being called **"halogen-bonding (XB)"**, as to relate it to hydrogen-bonding.


Here are two examples of trifluoromethylated halogens that emphasize how the prominence in positive-charge (colored red) can differ between them, with the chlorine motif on the left and the bromine motif on the right:


[![CClF_3 molecule calculated with DFT](https://i.stack.imgur.com/VvobIm.png)](https://i.stack.imgur.com/VvobI.png) [![CBrF_3 molecule calculated with DFT](https://i.stack.imgur.com/FU7Jdm.png)](https://i.stack.imgur.com/FU7Jd.png)


The lone pair on a water molecule can interact with this area of positive charge (called a "σ-hole") similar to interacting with the partially positive charge of a hydrogen—hence the name halogen-bonding—and this can stabilize solvation. The belt of electron-density carries partial negative charge, and thus can also interact with the σ-hole; this results in water coordination to the σ-hole being in competition with the self-intermolecular interaction. The crystal structure of orthorhombic iodine suggests that this interaction is strong in the solid-phase, so it is likely at least partially responsibility for the lower solubility of iodine.


If you'd like to learn more about XB interactions, I'll leave the DOI for a quick 2007 overview I recommend here: 10.1007/s00894-006-0154-7. I pulled those images directly from it.


My current master's thesis work is on exploring XB interactions, so my response comes primed with the fact that I think about these interactions a lot. There are countless other factors that perturb solubility, which is why we compress it into one simple figure and measure it empirically—I do believe with my best intuition, however, that these are the factors of greatest influence.


 6