So, since this is my last semester of undergraduate studies here at Virginia Tech, I had a relatively easy schedule set up for myself. In this schedule, I had no classes before 10:00 AM, indeed, on Mondays and Wednesdays I didn't have class until 2:30, and I didnt' even have class at all on Friday.
So then I got to thinking, which is generally a bad sign, and I decided I should take a biochemistry course, since everything I am interested in regarding microbes is on the molecular level. Well, the only biochemistry course that I could take, not being either a biochemistry major nor a graduate student was offered in a single section at 8:00 AM MWF. . . UGH. . .
However, after only one class period, my view has changed, and I now think that it is all worth it. What, you may be asking yourself, could have caused such a drastic change? Well I'll tell you: Hydrophobic Interactions.
Yes you heard me right. One thing that is really pounded into bio-major's heads is the whole 'like dissolves like' creed, which says that polar solvents can dissolve polar solutes, but not nonpolar ones, and vice versa. The reasoning behind polar solvents (such as water) being able to dissolve their polar brethren is always given and is pretty understandable:
The oxygen atom in a water molecule carries a partial negative charge due to increased electron density and thus is attracted to the positive end of a polar solute, say NaCl. thus the O side of the water molecule attacks the Na side of NaCl, and the Hydrogen side, bearing a partial positive attracts the negative Cl side, thus pulling them apart and dissolving the NaCl.
But what about solutes such as glucose, that aren't ionic and separable? Same idea, only water molecules surround the entire molecule binding via hydrogen bonds to the partial negative and positive charges located throughout the molecule, thus dissolving it whole, rather than splitting it apart.
So that's all well and good, but what makes nonpolar molecules 'hydrophobic?' I undersand why they aren't dissolved like polar compounds, but what makes them 'afraid' of the water molecules and want to 'get away' from them? And what makes them congregate? If they don'thave significant partial charges, there is nothing to attract them to each other. Why don't nonpolar molecules simply float around randomly, simply ignoring each other and the surrounding water molecules?
These are questions that had been at the back of my mind for quite some time now, but thanks to an assigned reading in my biochemistry textbook (Chapter 2 in Essential Biochemistry by Pratt and Cornely if you're wondering) they have been answered.
The answer lies in the Second Law of Thermodynamics, which states that, in a closed system entropy must always increase, or in this case free energy, which is related to entropy, must always decrease. The free energy G = H -TS (where H is enthalpy, T is absolute temperature, and S is entropy) must decrease when a nonpolar molecule is hydrated, or surrounded my water molecules. In this case the change in enthalpy during hydration contributes much less to the overal change in free energy than the change in entropy does, so we'll focus on the latter. Now in order to decrease free energy, we'll need to increase entropy (because it's negative in the relation above)
So when a nonpolar molecule is hydrated, the layer of water molecules sourrounding it cannot participate in normal hydrogen bonding, so they must 'lay flat' against the molecule such that neither the partially positive nor the partially negative ends stick out. This conformation, in restricting the possible orientations of the water molecules decreases their entropy, which can't happen. So what happens, is the nonpolar molecolues aggregate into one giant blob so that a minimum amount of water molecules have decreased entropy, because, as size increases, the surface area-to volume ratio decreases.
So there you have it. One major relevation after only the first day of class. I can't wait until the next class!