Water and Hydrogen Ions in Biology

The importance of the ionization of water and the concentration of hydrogen ions to biology cannot be overstated. Most cells and unicellular organisms maintain the [H+] of their cytoplasm within a very narrow range necessary for survival. Organelles are maintained at different [H+] necessary to optimize specialized functions. And blood also has a very narrow pH range or approximately 7.4. Derangement of this pH either way is deleterious to the organism. All of this is true because the hydrogen ion concentration of the solvent effects the rates and equilibrium constants of most biochemical reactions. In other words, most enzymes can function effectively only in a narrow range of hydrogen ion concentration.

From all of this we can see that hydrogen ions are very important to cells and biochemistry. But where to all of these hydrogen ions come from? Cells don't exactly split dihydrogen molecules, do they? Well, they don't, but there are other sources of hydrogen ions, which we'll investigate below.

Water, as everyone knows, can be ionized. It does this naturally also, although not to a large extent. Kw is a constant defined as the ion product of water. For pure water this is 1.0 * 10^-14, with both the hydrogen ion concentration and the hydroxide ion concentration equaling 1.0 * 10^-7. The product of these two ions, it should be noted, always yields Kw, regardless of the conditions.

Weak acids, being acids, can also be ionized to give hydrogen ions. Organic acids are only partially dissociated in water (hence weak, as dissociation is not near 100%, unlike mineral acids such as HCl). Ka, a dissociation constant, can be crafted for all organic acids much like Kw was derived for water. It is the product of the concentrations of the products (no pun intended) divided by the product of the reactants, or (Ka = [H+]*[A-]) / [HA] for the reaction equation HA <> H+ + A-. (Again here I would like to apologize for the current format. Thank you for being patient).

By calculating the pKa (or the -log (Ka)) and knowing the concentrations, then the solution pH can be calculated. This can be done by setting the concentration of free acid the undissociated acid) equal to the concentration of the total acid (the amount of acid that would be present had none dissociated). This can be done as long as the concentration of the total acid is much greater than the Ka, which is usually the case (as Ka's are generally no higher than 10^-3 and the concentration is generally at least 0.X Molar). Similar treatment can be used for any Bronsted acid.

This, of course, begs the next question: what is a Bronsted acid? Simply put, a Bronsted Acid is any compound which will dissociate to lose a proton (called a proton donor). A Bronsted Base then is any compound which will accept a proton (called a proton acceptor, duh).

From the study of weak acid dissociation, two gentlemen, one named Henderson and one named Hasselbalch determined a mathematical formula, surprisingly named the Henderson-Hasselbalch Equation. Given the chemical equation HA <> H+ + A-, they derived the following IMPORTANT equation:

pH = pKa + log(base/undis. acid)

see figure 4 - 10. This equation can be used to determine, given the right information, the pH, the pKa, the Ka, [H+], [A-], or [HA] and so on. It also allows for the development of equivalence, a sideshow topic not discussed here. However, do know that after adding 0.5 equivalents of strong base to any weak acid the concentration of the undissociated acid is equal to that of the conjugate base, or pH = pKa.

The HH equation paved the way for the next recognition in weak acids, that of the buffer. A buffer is a mixture of a weak acid and its conjugate base. The purpose of a buffer is to hold the pH of a solution nearly constant despite additions of hydrogen or hydroxide ions. Most buffers are most effective when the pH = pKa, and are generally useful when the pH is only 0.5 units from this situation. The capacity (or strength) of the buffer is determined by buffer concentration, which is the sum of the concentrations of the acid and the conjugate base forms of the buffer. This concludes, for now, our discussion on hydrogen ions and water.

So far this lecture we have discussed hydrogen ions and water as well as a brief introduction to buffers. What we didn't talk about so far, and where we pick up now, are the major physiological buffers. These are the buffers found in most organisms in large amounts doing important tasks. In other words, they are essential to the system in which they function. We shall take a short look at some of these now, and then begin looking at amino acids next time.

Our first example for a major physiological buffer is phosphoric acid, H2PO4. This, with it's conjugate base of phosphorous acid, HPO4, forms an effective buffer around the pH of seven. The pKa for this solution is, depending on the source, approximately 7.2 at infinite dilution or 6.86. This buffer occurs in most organisms, and can be better referenced in the book (which shall be done here in the coming weeks).

The second example involves carbon dioxide and water forming H2CO3, which further is broken down to HCO3-. The pKa for the first reaction is around 6.1 while that for the second is nearer 10.2. How exactly does this system work? Well, it's interesting to note that this system functions in the blood and is one of the big reasons we don't keel over and die when we chug a coke. An enzyme called carbonic anhydrase is found in red blood cells. This catalyzes hydration in carbon dioxide and the equivalent reverse reaction. Why is this important? Well, the amount of carbon dioxide in a system can vary greatly depending on the physical exertion being exercised. So too can the amount of acid being introduced into the blood. Therefore this enzyme is necessary to keep an excellent equilibrium (which it does). Another interesting facet of this system is that the H2CO3 (which is kept at a concentration of 25 milliMolar)- HCO3 (at about 1.25 mM) solution is not an ideal buffer. We had earlier stated that a buffer works at X + or - 0.5 or, at best, + or - 1. Blood plasma has a pH of about 7.4, and the buffer we're using works best at 6.1. But it still works. Why? It works because we have a virtually unlimited supply of carbon dioxide to constantly replenish the system when needed. And because the enzyme, carbonic anhydrase, works so well.

Also a popular organic buffer are Imidazole groups. These are derived from the side chain of the amino acid Histidine (HIS), which is present in most proteins. The pKa of the reaction (which this format does not allow me to display) is 6.0. Again, more information will be put here on this subject as time and information permit.

Finally, miscellaneous organic acids and nucleotides are used, as well as pH regulation through kidney excretion of NH4+. And thus ends this topic until the final information is fixed.