Solubility and Solvent Systems
‘If it dissolves, your problem is solved’ – a simple adage, and also one that chemists take quite literally. Finding a solvent that can dissolve your substrates is a key step for any synthesis. In many reactions, solvents also play a big role in forming the products. An understanding of solubility can help to better plan reactions, as well as improve on existing conditions and yields.
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What’s in a Solution?
A solvent is a medium in which a reaction takes place, and takes up the bulk of the system (a container such as a beaker, or even a cell). In biological processes, for example, water is usually the solvent – that’s why we need to drink so much of it! A solute is something that dissolves in the solvent, usually your reagents – the stuff that you want to react to produce something new. Gases can also be solutes, as seen in carbonated drinks where carbon dioxide is dissolved in water.
If the solute fully dissolves in the solvent they together form a solution. Partial dissolution of a solid results in a dispersion, in which small particles may be seen floating about. These particles can be so tiny that they might seem dissolved at first glance, requiring techniques like dynamic light scattering to accurately determine this.
Importance of Solubility
Solubility is defined as ‘the analytical composition of a saturated solution, expressed in terms of the proportion of a designated solute in a designated solvent’1. But why is it so important for the substrates to dissolve, anyway?
By definition, solids are energetically stable because of the bonds holding the molecules together. Remember that bonding between two molecules usually lowers their overall energy, resulting in a more stable system. Hence, bonding lowers the reactivity of all the molecules in a solid!
Furthermore, solids aren’t ‘free’ to move about in solution; only the surface molecules are able to react. By dissolving, individual particles can spread out to increase the overall surface area. This, in turn, leads to better reactivity.
Dissolving Different Compounds
It can help to explain solubility if we classify compounds into ionic species and non-ionic species. By working through the fundamentals of chemical bonds, perhaps we can shed some light on the fascinating chemistry of dissolution.
Ions with opposing charges – due to differences in electronegativity – attract, forming electronically neutral ionic, or polar, compounds. The individual ions that make up these compounds are able to dissociate in solution, releasing ‘free’ ions. However, this process of dissolution depends on the strength of the solvent to pull the ions apart.
We can use Coulomb’s law to explain the interaction between ions in solution:
where the force of attraction F between two ions depends on k (Coulomb’s constant, 8.9875 x 109 Nm2C-2), q1 and q2 (the electronic charges of the ions), r (the distance between ions). Another factor that has to be taken into account is ε (the dielectric constant), which depends on the solvent that the ions are dissolved in.
To obtain ‘free’ ions in solution, we first need to overcome the force holding them together (F). But what can help with this? From the equation, k is a constant, while r, q1 and q2 are unalterable. The only thing that we can vary is ε, which depends on our choice of solvent. From Coulomb’s law, increasing ε should lower the overall force that holds these ions together, making them easier to dissolve!
The dielectric constant ε is actually a measure of the bond strength between the solvent molecules themselves. This value correlates with its dipole moment. Greater separation of electric charge in a molecule (due to electronegativities) results in a larger dipole moment.
Before we get carried away with dipoles, let’s first take a look at non-ionic (or non-polar) compounds. Unlike ionic compounds, non-ionic ones do not have a formal charge. This makes Coulomb’s law unsuitable for describing their attractive forces.
They can, however, possess temporary dipoles – electron density is never stationary. This ‘charge’ tends to propagate through nearby molecules, such that they develop their own dipoles. We denote these informal, weak ‘charges’ with a lowercase delta (δ). The weak attraction that occurs between δ+ and δ- of neighboring molecules are known as London dispersion forces.
Even though we classify London dispersion forces as ‘weak’, breaking these compounds up is not as simple as we might think. We can see this in hydrocarbons, being notoriously insoluble in water and forming visibly distinct layers.
The real driving force for the dissolution of non-ionic compounds is entropy. We can think of this as the spread of particles over time, and the 2nd law of thermodynamics tells us that entropy always increases. What this means is that when a substance is put into a liquid, it wants to spread out to increase its entropy (i.e dissolve).
What prevents this is the attractive forces between the solvent molecules themselves. The interactions between a non-ionic compound and a polar solvent are weaker than the solvent-solvent bonds themselves! This means the solvent is happier to maintain its bonds with another solvent molecule, rather than introduce a non-ionic compound into its bonding.
What about non-ionic/non-polar solvents? In this case, the bonds between solvent molecules aren’t as strong. This means that it is happier to bond with a non-ionic compound – if entropy allows it, of course. Hence, non-ionic compounds prefer non-ionic solvents, i.e. solvents with a small dielectric constant.
Choosing the Correct Solvent
A basic understanding of solubility provides us with a platform from which we can plan and predict reactions, especially when this knowledge is used in conjunction with existing literature. In drug discovery, for example, a certain extent of water solubility is required for a molecule to be biologically active. In ion exchange and some chromatography techniques, choosing a suitable solvent is arguably the most important aspect of the process.
Understanding the concepts of dissolution and the dielectric constant, we can now predict the ability of a solvent to dissolve compounds. Ionic compounds will dissolve better in solvents with a high ε value, while non-ionic compounds tend to prefer solvents with a lower ε.
Unfortunately, achieving solvation isn’t as simple as choosing a solvent with a suitable dielectric constant. There are other weak forces of attraction that, along with London dispersion forces, can change the behavior of the system. These include pi stacking, magnetic dipoles and hydrogen bonding. The consequence of this is that in many chemical reactions, solvent choice remains very much a process of trial and error.
Choosing a solvent for your reaction is oftentimes not a straightforward process; in an ideal world you’d have a solvent system that’s cheap, not too toxic, and most importantly doesn’t screw with the reagents to produce unwanted side reactions. Efficient and environmentally friendly solvent use is of increasing importance, in an area known as green chemistry.
Cover graphic: artwork of DMSO solvent in a substitution reaction by Melanie (@nanoclustering)
- Wilkinson, A. D. M. A., & McNaught, A. (1997). IUPAC Compendium of Chemical Terminology, (the” Gold Book”). International Union of Pure and Applied Chemistry.