Water, pH, and Non-Covalent Bonding
Ionization of Water
Sometimes the hydrogen of one water molecule will "jump" to another water molecule:
H2O + H2O H3O+ + OH-
This proton hopping is called the ionization of water (an ion is a positively or negatively charged atom or molecule). This ionization creates a H3O+ and a OH- molecule. The H3O+ is often written as simply H+. This is because a H3O+ is just a H+ that jumps from one water molecule to another:
H2O H+ + OH-
So remember, H3O+ = H+
Looking at either of the two chemical equations above, it is important to note that the reverse reaction is also occurring.
How much H+ and OH- exist in water? Very, very little! The ratio of either H+ or OH- to H2O in neutral water is 1:1,000,000,000! Since this is such a small amount of either H+ or OH-, they rarely meet and neutralize each other.
The equilibrium constant, Keq describes the ionization equilibrium of water:
Keq = [H+][OH-]
Because of this relationship it is important to note that if the [H+] goes up then the [OH-] must go down, and vice-versa, for the value for the Keq of water must remain constant. For neutral water, the Keq is 1 x 10-14 M and the concentrations of [H+] and [OH-] are each 1 x 10-7 M . Let's look at that last number without the exponent:
This is obviously a very small number. A more manageable way to discuss small numbers such as this is to take the negative logarithm. For the concentration of [H+], this is called the pH. In this case:
-log(0.0000001 M) = 7
The pH of a solution is simply the negative logarithm of [H+]. The pH of a solution describes the acidity of a solution. Acidic solutions are those with a pH of less than 7 and basic solutions have a pH greater than 7. A solution, like H2O, with a pH = 7 is neutral. Similarly, the pOH could be used to describe a solution in terms of its OH- concentration. pOH is the negative logarithm of the OH- concentration. One useful thing to remember is:
pH + pOH = 14.
In the body, the pH of blood is 7.4. This corresponds to a [H+] of about 40 nM. This value can only vary from 37 nM to 43 nM without serious metabolic consequences.
In living systems, much of the chemistry involves interactions between acids and bases. Acids are H+ donors and bases are H+ acceptors.
HA H+ + A-
Acid and base reactions are made up of conjugate acid-base pairs. Strong acids are those that readily give up a H+. Strong bases readily accept a H+. Weak acids do not readily give up a H+, but will under the right conditions. A weak base is one that does not easily take up a H+. The conjugate base for a strong acid is a weak base. In contrast, a strong base has a weak acid as its conjugate. Think of a strong acid as a person with an ugly wig. This person can't wait to get rid of this! So now we have a wig (H+) and a person without the wig (the weak base). You can understand that the person without the wig doesn't really want to take it back!
However, what about strong base-weak acid conjugate pairs?
A- + H+ HA
Think of it in similar terms: a person in dire need of a wig will grab just about anything. This is our strong base. Now on the right side of the equation we have someone with a wig who really doesn't want to give it up, just as a weak acid does not want to part with its H+.
* We can think of a strong acid or base as having a "flip side", which is its weak base or weak acid conjugate, respectively.
How readily an acid gives up its H+ is expressed by the acid dissociation constant, or Ka:
The pKa is the negative logarithm of the Ka. Strong acids have small pKas. Looking at the above equation, it can be seen that when [A-] = [HA], then Ka = [H+]. Then the pKa = pH. This is the basis for the Henderson-Hasselbalch equation:
When acids or base are added to a solution, the pH changes. The controlled addition of an acid or base to a solution is called a titration. A titration curve can be made by plotting the pH changes against the volume of acid or base added to a solution. Many titration curves are made by plotting pH on the y axis and the volume of base added to the solution on the x axis. For example, pH changes to acetic acid, CH3COOH, and hydrochloric acid, HCl, with 0.1M sodium hydroxide, NaOH, is shown to the left.
Notice the shape of the curves. The CH3COOH curve starts out at a higher pH than the HCl curve. This is because CH3COOH is a weaker acid than HCl. As NaOH is added to the solution of CH3COOH, the pH rises. However, the addition of NaOH to the HCl solution does not significantly change the pH. It takes less NaOH to change the pH of a weak acid solution than it does to change the pH of a strong acid solution. Eventually all of the acid in both solutions will have reacted with base, such that adding more base quickly increases the pH. This is called the equivalence point. This is the point where all the acid has reacted with the base, such that further additions of base quickly raise the pH. The equivalence point for a strong acid-strong base titration is pH = 7. For weak acids-strong base titrations, the equivalence point is pH >7.
As shown above, the pH of a solution is dependent on the concentration of H+ ions. Addition or removal of H+ ions, then, can greatly affect the pH of a solution. In the body, the pH of cells and extracellular fluids can vary from pH 8 in pancreatic fluid to pH 1 in stomach acids. The average pH of blood is 7.4, and of cells is 7 - 7.3. Although there is great variation in pH between the fluids in the body, there is little variation within each system. For example, blood pH only varies between 7.35 - 7.45 in a healthy individual. Large changes in pH can be life threatening. How does the body maintain a constant blood pH? The body uses a buffer system to withstand changes in pH. Buffers are made up of a mixture of a weak acid with its conjugate base or a weak base with its conjugate acid. Remember that an acid donates a H+. A weak acid does not donate its H+ as easily. Similarly, a weak base will not accept a H+ as well as a strong base.
Buffers maintain pH by binding H+ or OH- ions. This stabilizes changes in pH. The bicarbonate buffer system maintains blood pH near pH 7.4. The carbonic acid, H2CO3, in the blood is in equilibrium with the carbon dioxide (CO2), in the air.
Buffers are most effective in a pH range near its pKa. This is where the titration curve is most shallow, and where the pH is least affected by added acid or base. For example, in the titration curve of phosphoric acid (another blood buffering system):
The region in the rectangle is the buffering region of biological importance. Note that the slope of the curve is shallow within one pH unit of the pKa value. This is where the buffering range is most effective, meaning that the buffer is able to resist changes in pH.
Copyright 2002, John Wiley & Sons Publishers, Inc.