I'm confused about why hydrophobic molecules, which do not have high polarity, would have a tendency to attract and cluster with themselves. It is easier to understand the hydrophilic as long as one believes that positive charges attract negative charges. Anyone have any insights?
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asked Sep 14, 2015 at 3:04
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$\begingroup$ related: chemistry.stackexchange.com/questions/26544/… $\endgroup$– MithoronCommented Sep 14, 2015 at 11:12
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2 Answers
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The driving force behind the hydrophobic effect is entropy. Suppose I put several hydrophobic molecules in water. These molecules have two choices. (1) The molecules can stay separate from each other and randomly move around in solution. (2) They can group together and form a cluster. It turns out that it is entropically favorable to cluster. The reason is that it takes more water to individually solvate the separate molecules. It takes less water to solvate a cluster of the molecules. When water solvates these molecules, it is ordering itself. In fact, it is believed that water forms cages around hydrophobic molecules in solution. Therefore, if the molecules cluster, less water will have to order itself. Less ordering means greater entropy (randomness).
It also true that London dispersion forces will occur in these clusters. However, this is not the driving force for the hydrophobic effect. The driving force is entropy.
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There are other intermolecular forces besides dipole-dipole forces (which is what polar molecules primarily use) that allow molecules to come together. For nonpolar molecules, they use dispersion (London) forces. Dispersion forces come from random movements of electrons which cause temporary dipoles. Due to the large amount of electrons in the molecules, lots of these temporary dipoles exist at all times, keeping the molecules together.
