12.2: Physiological Buffers

Although the lung's ability to expel CO₂ and the kidney’s ability to excrete or absorb hydrogen ions allow close regulation of pH, their responses alone are not sufficient to prevent immediate local changes in pH at the tissue. This is the role of the buffering systems.

Buffering systems are chemicals within tissue and the blood that have the ability to absorb either hydrogen ions and/or hydroxyl ions. Once these ions are removed from solution (albeit temporarily) then their effect on pH is diminished. We will deal with buffers in the context of acids, as this is the most common physiological situation.

If you need an analogy for the function of buffers, imagine them as a chemical mop—they soak up the hydrogen ions and stop them from making a cellular mess, but the hydrogen ions, although contained, remain in the system. It is the role of the lungs and kidneys to "rinse the mop" and get rid of the hydrogen ions from the system.

There are three major chemical buffering groups in the body:

1. the bicarbonate system,
2. the phosphate system, and
3. intra- and extracellular proteins.

We will deal with the bicarbonate system as it involves the respiratory system and is also the major extracellular buffer.

Bicarbonate buffering: A buffering system consists of a weak base capable of absorbing a strong acid and a weak acid capable of absorbing a strong base. As such, the bicarbonate system involves two components: sodium bicarbonate (a weak base) and carbonic acid (a weak acid). Let us look at how it works and put it in the context of the lungs.

First let us see how a weak acid (carbonic acid) deals with a strong base, in this example, sodium hydroxide (equation 12.2.1).

Buffering a strong base using a weak acid:

\[NaOH = \textcolor{red}{H_2CO_3} \nonumber \]

Sodium hydroxide is a strong base as it rapidly dissociates into a hydroxyl ion and a sodium ion.

\[\substack{Na^{+}\\\textcolor{red}{\textbf{OH}^{-}}} + \textcolor{red}{H_2CO_3} \nonumber \]

The hydroxyl ion is the potential threat to physiological function so must be buffered. This is achieved by the carbonic acid dissociating into a hydrogen ion and bicarbonate (a process you are familiar with).

These dissociated ions now bind to form new partnerships as water and sodium hydroxide (a weak base) (equation 12.2.3).

\[\substack{Na^{+}\\\textcolor{red}{\textbf{OH}^{-}}} + \substack{HCO_{3-}\\{H}^{+}} \Rightarrow H_2O + \textcolor{blue}{NaHCO_3} \nonumber \]

So there are a couple of things to notice here beyond watching the ions move and form new components. First, the buffering process has taken a situation with the threat from a strong base (NaOH) and toned it down to a situation with a weak base (NaHCO₃); the problem has not gone away, it has just been reduced (or buffered). Second, you will see that both of the components of the bicarbonate system, carbonic acid and sodium bicarbonate, appear in the equation—we have just shifted from one to the other.

Let us look at the opposite situation to see what happens when the buffering system is faced with a strong acid. This time a strong acid (hydrochloric acid) is faced with our weak base (sodium bicarbonate) (equation 12.2.4).

Buffering a strong acid using a weak base:

\[HCl + \textcolor{blue}{NaHCO_3} \nonumber \]

The hydrochloric acid rapidly dissociates into a hydrogen ion and a chloride ion. The hydrogen ion now threatens physiological function and must be buffered.

Our weak base dissociates into sodium and bicarbonate ions. Again our ions recombine, this time to produce harmless sodium chloride and carbonic acid (equation 12.2.5).

\[\substack{\textcolor{red}{\textbf{H}^{+}}\\{Cl}^{-}} + \substack{HCO_{3-}\\{Na}^{+}} \Rightarrow NaCl + \textcolor{red}{H_2CO_3} \nonumber \]

Notice again we have reduced but not removed the threat as we have gone from the presence of a strong acid to a weak one. Also notice that our two components in the bicarbonate system appear in the equation, and we have switched from one to the other. This should now make you realize that these two components are part of a reversible equation, and this reversible equation, even after the addition of sodium to one end, should look rather familiar (equation 12.2.6).

\[CO_2 + H_2O \leftrightarrow \textcolor{red}{H_2CO_3} \leftrightarrow H^+ + HCO_{3-} + Na^+ \leftrightarrow \textcolor{blue}{NaHCO_3} \nonumber \]

As CO₂ is at one end of the equation you should appreciate how alveolar ventilation can influence the bicarbonate buffering system.

Because of their critical role in maintaining blood pH, bicarbonate ions are routinely measured along with arterial blood gases. Knowing what the blood pH, arterial CO₂, and bicarbonate levels are provides a very powerful and commonly used diagnostic measure allowing us not only to determine the pH status of the patient, but also the source of the problem and whether the renal or pulmonary systems are achieving compensation. Because of its power and common use, we are going to go through some fundamentals, and I am afraid that means looking at the bane of many a medical student: the Henderson–Hasselbalch equation. For those with a background in chemistry you might skip the next section, but for the rest of us, we are going to go through this step-by-step.