How do buffer solutions allow our bodies to maintain homeostasis
consistently in the blood (keep the blood's pH constant)?
- Question by Michael Newton
Michael produced a question that is significance to what we have learned over the course this semester. It comes from Chapter 16, Aqueous Ionic Equilibrium, where we learned the concept of buffers and solubility equilibria.
A buffer resists pH change by neutralizing added acid or added base. Our blood contains several buffering systems, with the most important one being of carbonic acid and the carbonate ion. Normal blood has pH of 7.4, this can be found by using the Henderson-Hasselbalch equation:
[HCO3-] = 0.024 M and [H2CO3] = 0.0012 M
pKa for carbonic acid at body temperature = 6.1
pH=pKa + log ([base]/[acid])
= 6.1 + log ([HCO3-]/[H2CO3])
= 6.1 + log (0.024M/0.0012M)
pH= 7.4
The concentration of the bicarbonate ion is 20 times higher than the concentration of carbonic acid and the pH of the buffer is more than one pH unit away from pKa... this is due to the higher bicarbonate ion concentration in blood makes the buffer capacity of blood greater for acid than for base, which is necessary because the products of metabolism that enter the blood are mostly acidic. A great example of this is when we exercise. Our bodies produce lactic acid (HC3H5O3), and when this enters our bloodstream it must be neutralized. The bicarbonate ion must neutralizes the lactic acid, and then an enzyme, carbonic anhydrase, then catalyzes the conversion of carbonic acid into carbon dioxide and water. Our bodies eliminate the carbon dioxide from our blood when we breathe... the larger the amount of lactic acid, the heavier we breathe.
All of this is in thanks to the buffers in our blood. Without them we would not be able to maintain a constant pH, which could result in life-threating issues.
Source:
Tro, Nivaldo J. "Aqueous Ionic Equilibrium." Chemistry: A Molecular Approach. Upper Saddle Ranch, NJ: Pearson Prentice hall, 2011. 712+. Print.
