So how does adding hydroxyl groups increase the solubility of a hydrocarbon in water? To understand this we return to the two components of the free energy equation, enthalpy and entropy. For a solute to dissolve in a liquid the solute molecules must be dispersed in that liquid. Solubility will depend upon how many solute molecules can be present within a volume of solution before they begin to preferentially associate with themselves (rather than the solvent molecules). When the solute molecules are dispersed, whatever bonds or attractions holding the particles together in the solute will be replaced by interactions between solvent and solute molecules. You might now reasonably suspect that one reason that diamonds are not soluble in water is that the C-C bonds holding a carbon atom within a diamond are much stronger (that is, take more energy to break) than the possible interactions that could be formed between carbon atoms and water molecules.

6.1 Solutions
6.2 Solubility
6.3 H-bonds
6.4 Free Energy
6.5 Polarity
6.6 Temperature

For diamond to dissolve in water a chemical reaction would have to take place in which multiple covalent bonds were broken. Based on this idea, we can begin to generalize – the stronger the interaction between the solute particles, the less favored it will be for the solute to dissolve in water. At the same time, the stronger the interactions between solute and solvent molecules, the more solubility should increase.

So do intermolecular interactions explain everything about solubility? Does it explain the differences between the solubilities in water of hexane, hexanol, and hexanediol? Hexanediol  (HO(CH₂)₆OH) is readily soluble, and if we consider its structure we would predict that interactions between hexanediol molecules would include hydrogen bonding (involving the two hydroxyl groups) and van der Waals interactions. We can approach this from a more abstract perspective. If we indicate the non-hydroxyl (–OH) part of a molecule as “R”, then an alcohol molecule can be represented as R–O-H, and a diol could be represented as HO-R-OH. Note that In water, R represents H! All alcohols have the ability to hydrogen bond with each other - and with water. So when an alcohol dissolves in water, the interactions between the alcohol molecules are replaced by interactions between alcohol and water molecules - an interaction similar to that between water molecules. Like water molecules, alcohols have a dipole (unequal charge distribution), with a small negative charge on the oxygen(s) and small positive charges on the hydrogen (bonded to those oxygen atoms). It makes sense that molecules that have similar structures will interact in similar ways - and so small molecular weight alcohols can dissolve in water. But if you look again at the table, notice that hexanol (a 6 carbon chain with one O–H group, is much less soluble than hexanediol (a 6 carbon chain with two O-H groups - one at each end). As the non-polar carbon chain lengthens, the solubility typically decreases. However if there are more O–H groups present, there are more possible interactions with the water.

Entropy and solubility or why don’t oil and water mix?

The fact that oil and water don’t mix is well known, it has even become a common metaphor for other things that don’t mix (people, faiths, etc). What is not quite so well known is, Why? Oil is a generic name for a group of compounds, many of which are hydrocarbons or contain hydrocarbon-like regions. Oils, for want of a better word are oily - they are slippery (and to risk being tedious) do not mix with water. The molecules in olive oil or corn oil typically have a long hydrocarbon chain of about 16-18 carbons. These molecules often have polar groups called esters (that is: groups of atoms that contain C–O bonds) at one end, but once you get more than 6 carbons in the chain, these groups do not greatly influence solubility in water - just as the O–H groups in alcohols do not influence solubility. So, oily molecules are primarily non-polar and interact with one another, and with other molecules (including water molecules), primarily through London dispersion forces. When oil molecules are dispersed in water, the interactions between them and the water molecules will include both LDF’s and water dipole - induced oil molecule dipole interactions. Such dipole - induced dipole interactions are common and can be significant, but if we were to estimate the enthalpy change associated with dispersing oily molecules in water, we would discover ΔH is approximately zero for many systems. That is the energy required to separate the molecules in the solvent and solute is about equal to the energy released when the new solvent-solute interactions are formed.

Remember that up to now, we have seen that the entropy change associated with mixing molecules is positive - entropy increases. So, if the enthalpy change associated with mixing oils and water is approximately 0, and the entropy of mixing is usually positive why then do oil and water not mix? It appears that the only possibility left is that the change in entropy associated with dissolving oil molecules in water must be negative (thus making ΔG positive). More than that, if we were able to disperse oil molecules throughout an aqueous solution, we would find that the mixed system would (spontaneously) separate (unmix.) That seems to be a process that involves work - what force drives it?

Rest assured, there is a non-mystical explanation, but one that requires us to think at both the molecular and the system level. When hydrocarbon molecules are dispersed in water, the water molecules rearrange to maximize the number of H-bonds they make with one another. They form a cage-like structure around each hydrocarbon molecule. This cage of water molecules around each hydrocarbon molecule is a more ordered arrangement than that found in pure water, particularly when we count up and add together all of the individual cages! It is rather like the arrangement of water molecules in ice (although restricted to regions around the hydrocarbon molecule). This more ordered arrangement results in a decrease in entropy. The more oil molecules dispersed in the water, the larger the decrease in entropy. On the other hand, when the oil molecules clump together, the area of “ordered water” is reduced, fewer water molecules are affected. There is therefore an increase in entropy associated with the clumping of oil molecules, a totally counterintuitive idea! This increase in entropy leads to a negative value for –TΔS (because of the negative sign). The interactions between oil and water molecules are minimized, which leads to the formation of separate oil and water phases. Depending on the relative densities of the substances, the oily phase can be either above or below the water phase. This entropically driven separation of oil and water molecules is commonly (and unfortunately) referred to as the hydrophobic effect. Why unfortunate? Oil molecules are not afraid (phobic) of water and they do not repel water molecules; remember all molecules (unless they have a permanent and similar electrical charge) attract each other via London Dispersion Forces.

Because the insolubility of oil into water is controlled primarily by changes in entropy, it is directly influenced by the temperature of the system. At low temperatures it is possible to stabilize mixtures of water and hydrocarbons. In such mixtures, which are known as clathrates, the hydrocarbon molecules are surrounded by stable cages of water molecules (ice). Recall that ice has relatively large open spaces within its crystal structure - it is within these holes that the hydrocarbon molecules fit (which means you might well be able to predict the maximum size of the hydrocarbon molecules that can form clathrates). For example some oceanic bacteria generate CH₄ (methane) which is then “dissolved” in the (cold) water to form methane clathrates. It is estimated that between two and ten times the current estimates of conventional natural gas resources are present as methane clathrates.

Solubility of ionic compounds: salts

Polar compounds tend to dissolve in water, and we can extend that generality to the most polar compounds of all - ionic compounds. Sodium chloride, the most common ionic compound, is soluble in water (360 g/L). Recall that NaCl is an ionic compound – a salt crystal is not composed of discrete molecules of NaCl, but rather is best described as an extended array of Na+ and Cl- ions bound together through electrostatic interactions. Therefore, when an ionic compound dissolves in water the electrostatic interactions within the crystal must be broken. In contrast, when molecular compounds dissolve in water, it is the intermolecular forces that are disrupted.

One might imagine that the breaking of ionic bonds would require a very high energy input (after all we have already seen that diamonds will not dissolve in water because actual covalent bonds would have to be broken). That would be true if all we considered was the energy required to break the ionic bonds - as indicated by the fact that NaCl melts at 801ºC and boils at 1413ºC. That said, we know that substances like table salt (NaCl) dissolve readily in water, so clearly there is something else going on. The trick is, of course, that when NaCl dissolves we have to consider the whole system, just like we did for molecular species. We need to consider the interactions that are broken and those that are formed. These changes in interactions are reflected in the ΔH term (of ΔG = ΔH – TΔS).

Now let us turn to changes in entropy. When a crystal of NaCl comes into contact with water, the water molecules interact with the Na+ and Cl- ions on the crystal’s surface. The positive ends of water molecules (the Hs) interact with the chloride ions, while the negative end of the of the water molecules (the O) interacts with the sodium ions. That is: an ion on the surface of the solid interacts with water molecules from the solution; these water molecules form a dynamic cluster around the ion. Thermal motion then moves the ion and its water shell into solution.

Question to answer:

• Draw a molecular level picture of a solution of NaCl, be sure to show all the kinds of particles and interactions present in the solution.
• When we calculate and measure thermodynamic quantities (such as ΔH, ΔS and ΔG) why is it important to specify the system and the surroundings.
• When a substance dissolves in water, what is the system and what is the surroundings? Why?
• What criteria would you use to specify the system and surroundings?
• For a solution made from NaCl and water, what interactions must be overcome as the NaCl goes into solution? What new interactions are formed in the solution?
• If the temperature goes up when the solution is formed what can we conclude about the relative strengths of the interactions that are broken and those that are formed?
• What about if the temperature goes down?
• When you measure the temperature are you measuring the system or the surroundings?
  

Questions for later:

• Why is the water shell around an ion not “stable”?
• What are the boundaries of a biological system?