Saturday, August 25, 2012

42. The Secondary Chemical Interactions

The covalent, ionic, and metallic interactions I described in Part 41 are called the primary interactions. They result in strong chemical bonds. A number of secondary interactions or bonds among atoms also exist, which are substantially weaker than the primary interactions. Particularly ubiquitous and important among these is the hydrogen bond. 

Take the example of water, H2O or H-O-H. The oxygen atom forms covalent bonds with the two hydrogen atoms. Each such covalent bond (O-H) has two electrons associated with it, one coming from hydrogen and one from oxygen. The electron distribution around the hydrogen nucleus in such a bond is not like that in a symmetrical bond like C-C in the structure of diamond. The oxygen nucleus has a charge number (Z) equal to 8, which is much more than the charge number 1 of H, so it hogs (attracts towards itself) a larger share of the electron charge cloud associated with the covalent bond; we say the oxygen atom is very electronegative. This makes the nucleus of the hydrogen atom somewhat less shielded by its electron than when there was no bonding of any kind. For similar reasons, the oxygen nucleus and its charge cloud of electrons are together a little more negative than they would be in an isolated atom of O. The end result is that the O-H bond in a water molecule is like a little dipole. There are two such bonds in the H-O-H molecule, so there are two positive ends and a negative end. The upshot is that, because of this charge separation, the entire water molecule has a net 'dipole moment'.

The water molecules, being dipoles, tend to orient themselves such that a positive end (a hydrogen end) of one molecule points towards the negative end (the oxygen end) of another molecule. So we speak of hydrogen bonds, denoted in this example by O-H…O.

The most crucial aspect of the hydrogen bond in the evolution of chemical and biological complexity is that it is of intermediate strength, not as strong as the covalent bond, and yet not as weak as the so-called van der Waals interaction (or the London dispersive interaction):

The van der Waals interaction is very weak, and it is always present between any two atoms. Quantum-mechanical fluctuations in the electronic charge cloud around an atom can result in a transient charge separation or dipole or multipole moment, and the electric field of this multipole induces a multipole moment on any neighbouring atom. This results in a small attraction between the two atoms, the so-called van der Waals attraction, interaction, or bond.

The energy required to break a chemical bond is a measure of its strength. The melting point of a solid is an indicator of the strength of the weakest bonding in it. The covalent bond is the strongest, with a typical bonding energy of ~400 kilocalories (kcal). The ionic bond is typically half as strong as the covalent bond. The metallic bond shows a wide range of strengths, two extreme examples being the bonding in mercury on one extreme, and the bonding in tungsten on the other. The strength of a hydrogen bond is typically 14 kcal. And van der Waals bonding involves energies below 1 kcal.

The most relevant fact for our purpose here is that the energy involved in hydrogen bonding is typically only ~10 times larger than the energy of thermal fluctuations, but is still much lower than the energy of a typical covalent bond.

At typical temperatures at which biological systems exist, it is difficult for thermal fluctuations to break covalent bonds, but there is a fairly good chance that they can break hydrogen bonds.

We have seen above that water is an aggregate of tiny dipoles. We say that it is a polar material. By contrast, there are a large number of ‘hydrocarbons’ which are nonpolar materials. [A hydrocarbon is a compound made predominantly of hydrogen and carbon atoms.] In contrast to the O-H bond in water, which is a bond with a dipole moment, the C-H bond in a hydrocarbon is largely nonpolar: The two electrons forming the C-H covalent bond are shared almost equally between C and H (there are quantum-mechanical reasons for this). Thus, a C-H bond hardly results in the creation of a dipole, and therefore it does not readily form a hydrogen bond with a water molecule.

Now suppose we mix a nonpolar fluid with a polar fluid like water. Segregation will occur. The nonpolar molecules will tend to huddle together because they cannot take part in the hydrogen bonding of water. They have a kind of ‘phobia’ for water molecules, and so we speak of the hydrophobic interaction. Since the hydrogen bond is of intermediate strength, the hydrophobic interaction is also of intermediate strength.

There are many types of organic compounds that are predominately of hydrocarbon (i.e. nonpolar) structure, but have polar functional groups attached to them. Examples of this type are cholesterol, fatty acids and phospholipids. Such molecules have a nonpolar or hydrophobic end, and a polar or hydrophilic end. When put in water, they self-aggregate such that the hydrophilic ends point towards water, and the hydrophobic ends get tucked away, avoiding interfacing with water. This is why oil does not mix with water.

By contrast, alcohol and water mix so readily that no stirring is needed; both are polar liquids. As the king said: ‘I do not care where the water flows, so long as it does not enter my wine!’

Beautiful high-symmetry self-assemblies like micelles, liposomes, and bilayer sheets may ensue because of the hydrophobic interaction. Art without artist!

The second law of thermodynamics for open systems is the only self-organization principle there is; subject to the constraints of the first law of thermodynamics, of course. Much of the symmetry we see in Nature is a consequence of this law (Wadhawan 2011).

Here is an interesting video on the polarity of water molecules: