Below are several problems using the Hendersen-Hasselbach equation with the bicarbonate buffer system, the ammonia buffer system, and aspirin. For the most part, these problems were put on the board in class.

Example 1: Be able to draw the equations that show how the bicarbonate buffer system works in blood. What is the respiratory compensation when the blood pH drops to 7.3? What is the respiratory compensation when the blood pH rises to 7.5?
CO2 + H2O = H2CO3 = H+ + HCO3-

pH = pKa + Log [HCO3-]/[H2CO3]

If metabolism produces excess acid (H+), the H+ combines with and lowers the concentration of bicarbonate ([HCO3-]).

If excess base is added, carbonic acid dissociates to provide protons to neutralize the base. As a result, the concentration of bicarbonate increases.

If enough acid or base is added to change the pH, not only will the bicarbonate ion concentration change but respiratory compensation will occur. Breathing faster and deeper will remove carbonic acid and more shallow breathing will increase the concentration of carbonic acid. This is compensatory respiratory alkalosis and acidosis respectively. Note that in both cases, the pH moves back toward 7.4.

Also, note that the body does not wait for a large change in bicarbonate or pH before beginning compensation.  That is, respiratory compensation begins as soon as the pH changes so that, by the time you see the patient, changes in both bicarbonate and carbonic acid have occurred.

Note that PaCO2 is the partial pressure of CO2 in arterial blood. It is reported as mmHg and can be converted to mM by multiplying by 0.03. In this class, the units for the concentrations in the Henderson-Hasselbalch equation will always be in mM (mmol/L) (meq/ml).

Remember that the [H2CO3] is regulated by the brain and lungs.

You should memorize the following normal values for the bicarbonate buffer system in a normal person and, in this course, always use a pKa of 6.1:

pH = pKa + Log [HCO3-]/[H2CO3]

7.4 = 6.1 + Log [24mM]/[1.2mM]
Notice that if you are given any pH and pKa, you can determine the ratio of [HCO3-]/[H2CO3]

Example 2: Draw the formula for the dissociation of the ammonium ion.
If the pKa =9.25, what form is found at pH = 7.4?:
NH4+ = H+ + NH3

pH = pKa + Log [NH3]/[NH4+]

7.4 = 9.25 + Log [NH3]/[NH4+]

-1.85 = Log [NH3]/[NH4+]

1.85 = Log [NH4+]/[NH3]
Note, put 1.85 into your calculator and depress the 10X button. 71 will appear.

71 = [NH4+]/[NH3]

Ammonium ion predominates. The concentration of ammonium ion is 71 times grater that the concentration of ammonia.

Example 3: Concerning Dennis Veere: If the pKa for acetylsalicylic acid to acetylsalicylate is 3.5, is aspirin a weak or strong acid? Which form is prevalent in the stomach at a pH of 1? Which form is prevalent in blood at pH of 7.4? Prove it using the Henderson-Hasselbalch equation.
Acetylsalicylic acid = Acetylsalicylate + H+, pKa =3.5

pH = pKa + Log [Acetylsalicylate]/[Acetylsalicylic acid]

In stomach, pH = 1; pKa = 3.5.
1 = 3.5 + Log [Acetylsalicylate]/[Acetylsalicylic acid]

-2.5 = Log [Acetylsalicylate]/[Acetylsalicylic acid]

2.5 = Log [Acetylsalicylic acid]/ [Acetylsalicylate]

316 = [Acetylsalicylic acid]/ [Acetylsalicylate]

In blood, pH = 7.4;

7.4 = 3.5 + Log [Acetylsalicylate]/[Acetylsalicylic acid]

3.9 = Log [Acetylsalicylate]/[Acetylsalicylic acid]

7943 = [Acetylsalicylate]/[Acetylsalicylic acid]

Is aspirin a weak or strong acid? Any acid with a pKa is a weak acid.