IFAP - Solubility (Lesson)

Solubility

Dissolving is a spontaneous process for soluble substances. By spontaneous, we mean a process that occurs under specified conditions without the requirement of energy from some external source. Sometimes things can be done to speed up the dissolving process, but even without mixing or heat for example the solution would still form on its own. As an example, when containers of helium and argon are connected, the gases spontaneously mix due to diffusion and form a solution. The formation of this solution clearly involves an increase in matter dispersal, since the helium and argon atoms occupy a volume twice as large as that which each occupied before mixing. These two of the factors that drive the spontaneous formation of a mixture: 

1. A decrease in the internal energy of the system (enthalpy)
2. An increased dispersal of matter in the system (entropy)

For a solution to form, the particles must intermingle. In order for this to occur, the solute and solvent particles must have attractions for each other that are stronger than (or at least as strong as) their attraction for particles of their own kind. Look at the example above and assume the purple solute particles are polar molecules, perhaps methanol molecules. Both molecules exhibit hydrogen bonding intermolecular forces. If water were more attracted to another water molecule than it was to methanol then it would move the methanol molecule out of the way to experience the attraction to another water molecule. But, water is attracted to methanol almost as equally as it is to other water molecules. Since their intermolecular attractions are similar, they mix freely together and are said to be miscible, or soluble in one another.  When the strengths of the intermolecular forces of attraction between solute and solvent species in a solution are similar to those present in the components, the solution is called an ideal solution.

Ionic compounds will tend to dissolve in polar solvents because the cations interact with the negative portion of water's dipoles while anions interact with the positive portion.  

What about nonpolar molecules like CCl₄ and benzene (a symmetrical molecule to the right)?

   

The only IMF present are the dispersion forces, (temporary induced dipoles).  The forces of attraction between CCl₄ molecules are about the same as those between benzene molecules. So since neither benzene nor carbon tetrachloride is MORE attracted to itself than the other molecule, they freely mix together.

• Molecular compounds that do not have dipoles and have predominantly dispersion forces will tend to dissolve in nonpolar solvents. The larger and more polarizable the electron cloud is, the more interactions will occur with a polar solvent.  
• Polar molecules like sucrose dissolve in polar water due to dispersion or possibly hydrogen bonding IMFs.

You Try It! Solubility

The amount of a chemical you can dissolve in a specific solvent is limited. At some point the solution becomes saturated. This means that if you add more of the compound, it will not dissolve anymore and will remain solid instead.  

For gases, the solubilities of nonpolar gases in water generally increase as the molecular mass of the gas increases.  As the gas molecules become larger, the strength of the solvent–solute interactions due to London dispersion forces increases, approaching the strength of the solvent–solvent interactions. As the temperature of the solvent increases, less gas will dissolve.  The added energy of the solvent creates strong collisions with the gas molecules, which then escape the solution.  This is why carbonated drinks go flat when warm, but stay bubbly longer when cold.

In a solution where the solute is a solid or liquid, changing the temperature changes the solubility of the solute. When the temperature of a saturated solution in dynamic equilibrium is increased, the equilibrium shifts so that the heat is used up. For most solids and liquid solutes, this equates to more solute dissolving in the solution. For a few solutes, when the temperature increases, the solute becomes less soluble. Ce₂(SO₄)₃  is an example of this. Solubility curves show the changes in solubility of several solutes in 100 g of water as the temperature is increased.

Look at NaCl at room temperature, approximately 25^oC, as an example. If we take 30.0 g NaCl and add it to 100 mL of water, we will have an unsaturated solution.

     

If we take 40.0 g NaCl and add it to 100 mL of water, we will have a saturated solution with 4.0 g of NaCl undissolved at the bottom of the solution.

   

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