Organisms, cells, organelles, proteins, and even individual amino acids and more are held togeather through a variety of means. Included in these means are of course covalent bonds, but there are also many other different forces at work holding togeather the organic in question. What are these forces and what to they mean? Why do we need them?

Well, the non-covalent Water-Solute and Solute-Solute interactions are very important in biochemistry. They are necessary to maintain the three dimensional structures necessary for protein, membrane, and nucleic acid function. These mysterious forces also drive the binding processes that regulate cell physiology. Examples include drug or substrate binding to enzymes, the binding of an antagonist to the cell receptor in intercellular communication, and so on. These forces are vital in the development, day-to-day life, and even the death of cells, organs, and the like. So by now I'm sure you're asking, what are these Fabulous Forces?

Before we can answer that, though, we have to look at one other major facet of cellular life: water. Water is a very large component of cells; most are mostly water, spend their life in some environment composed largely of water, and most need it this way or they could not survive. Why is this? What makes water unique?

The keys to the water question lie in water's high polarity and it's high dielectric constant. From this stems the wonderful fact that water is able to form Hydrogen bonds with itself and with others. Hydrogen bonds are an attractive electrostatic interaction between a hydrogen atom bonded covalently to an electronegative atom such as oxygen, nitrogen, or sulfur (which, togeather with the hydrogen is called the donor group) and another atom with unshared electrons (such as oxygent) called the acceptor group. These bonds are relatively weak (only up to about 20 kilojoules per mole, compared to 40 for ionic bonds and from 200 - 500 for covalent single bonds). In addition, hydrogen bond lengths can be up to 0.3 nanometers in length, compared to approximately 0.1 nm for the comparable covalent bond. Knowing now what we do about water, let's investigate the effect of water on the electrostatic interactions between ions. In other words, what happens to ionic bonds in water? Well, think first of salt. Salt is sodium and chloride held ionically togeather. It is fairly stable in air, but once placed in water, it dissolves into sodium and chloride ions. Why? Well, water engages in hydrogen bonding with the atoms in the ionic substance, which there's nothing wrong with; it's just a natural extension of electrostatic attraction. However, remember that hydrogen bonds can be quite large. And there can be a lot more water than there are sodium or chloride ions. Because of this water is able to "screen" the charge on ions, thus reducing ionic interactions. With it's high dielectric constant, water can also reduce the attractive forces of ions. So by coming between ions, water effectively increases the distance between oppositely charged species, reducing the force (defined as F = (charge Na ion * charge Cl ion) / (r^2 * the dielectric constant of water, approx. 80). What does this mean? It means there is less of an attractive force between the sodium and chloride ions in water than in air, so they dissociate, and entropy effects help along the way.

Now finally then we are capable of intellegent discussion regarding the different non-covalent bond types: Hydrogen bonds, Ionic bonds, van der Waals interactions, and hydrophobic interactions. For those of you reading this enrolled in BMB401 at the Pennsylvania State University, an excellent table comparing these figures exists as table 4-4, while a good one comparing these to covalent bonds lies in table 3-5.

Hydrogen bonds have already been discussed, and ionic skimmed over (they are not overly important in most organic systems. However, we have not touched on van der Waals yet. What are van der Waals forces? Well, to start with, they're weak. There are three types. From strongest to weakest, they are dipole-dipole, dipole-induced dipole, and induced dipole-induced dipole. It should be noted that the induced-induced force is also called a London dispersal force. For van der Waal's forces, the maximum force is at the radii of the two interacting atoms. At this location the force of attraction is greatest. The farther apart the atoms, the weaker the force. If the atoms get nearer than the radii, there is a force of repulsion stronger than the van der Waals force.

Hydrophobic interactions are NOT bonds, hence the name hydrophobic interactions as opposed to hydrophobic bonds. These interactions are based on the fact that the molecules in question do not like water (hence the name hydrophobic). Because of this dislike for water, they stay close to one another, forming perhaps micelles (like a soap bubble) or lipid bilayers (like in cells). This can be seen in figures seven a and b in the text, and is a driving force for many 3 - D structures in the cell.

Before wrapping up this section, we should consider the fact that there are polar, non-polar, and amphpathic substances in the body. Examples of polar molecules include glucose, glycine, aspartic acid, lactic acid, and glycerol. Good examples for any non-polar compounds are any waxes. And amphipathic compounds includ phenylalanine and phosphatidylcholine.

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