Carbonic acid/bicarbonate buffer

Palaghias, G. The Role of Phosphate and Carbonic Acid-bicarbonate Buffers in the Corrosion Processes of the Oral Cavity , Dental Materials, 1, 139-144 (1985)... 

Buffering refers to the ability of a solution to resist change in pH after the addition of a strong acid or base. The body s principal buffer system is the carbonic acid/bicarbonate (H2C03/HC03 ) system. 

C. The carbonic acid-bicarbonate system is the most important buffer of the... 

The carbonic acid-bicarbonate buffer system has a of 6.1, yet is still a very effective buffer at pH 7-4 because it is an open buffer system, in which one component, CO2, can equilibrate between blood and air. 

The excessive amount of bicarbonate in the blood means that blood has a much greater capacity to neutralize acids. Many acids accumulate in the blood during strenuous activity, for example lactic acid. Excretion of bicarbonate through the kidneys and the removal of carbon dioxide through respiration also regulate the carbonic acid/ bicarbonate blood buffer. 

Daily body activities are quite sensitive to large pH changes, and must be kept within a small range of H30 and OH concentrations. Human blood, for example, has a pH of approximately 7.4 maintained by a buffer system. If our blood pH drops below 7.35, it can cause symptoms such as drowsiness, disorientation and numbness. If the pH level drops below 6.8, a person can die. To maintain pH stability, there is a carbonic acid - bicarbonate buffer system in the blood. 

Buffers stabilize a solution at a certain pH. This depends on the nature of the buffer and its concentration. For example, the carbonic acid-bicarbonate system has a pH of 6.37 when the two ingredients are at equimolar concentration. A change in the concentration of the carbonic acid relative to its conjugate base can shift the pH of the buffer. The Henderson-Hasselbalch equation below gives the relationship between pH and concentration. 

HA] is the concentration of the acid and [A-] is the concentration of the conjugate base. The pKa of the carbonic acid-bicarbonate system is 6.37. When equimolar conditions exist, then [HA] = [A ]. In this case, the second term in the Henderson-Hasselbalch equation is zero. This is so because [A ]/[HA] = 1, and the log 1 = 0. Thus at equimolar concentration of the acid-conjugate base, the pH of the buffer equals the pKa in the carbonic acid-bicarbonate system this is 6.37. If, however, we have ten times more bicarbonate than carbonic acid, [A ]/[HA] = 10, then log 10 = 1 and the pH of the buffer will be... 

Carbonic Acid/Bicarbonate Buffer H2CO3/HCO3... 

Again, buffers are essential in keeping pH within defined limits so that biochemical reactions can occur with maximum efficiency (i.e. maintaining so-called homeostasis). If, for example, the pH gets too low, the acidic environment can damage (or denature) and thus disable enzymes. An important example is the buffering of blood by the carbonic acid-bicarbonate buffer system illustrated below ... 

The pH of the plasma may be considered to be a function of two independent variables (1) the PCO2, which is regulated by the lungs and represents the acid component of the carbonic acid/bicarbonate buffer system, and (2) the concentration of titratable base (base excess or deficit, which is defined later), which is regulated by the kidneys. The plasma bicarbonate concentration is generally taken as a measure of the base excess or deficit in plasma and ECF, although it is recognized tliat conditions exist in which bicarbonate concentration may not accurately reflect the true base excess or deficit. 

The most important extracellular buffer is the carbonic acid-bicarbonate system ... 

Thus, the physiologic regulation of both PCO2 and [HCO/] permit the carbonic acid/bicarbonate system to provide more effective buffering of the extracellular fluids than could be achieved on the basis of chemical buffering alone. 

The distribution of the carbon-containing species (carbonic acid, bicarbonate and carbonate) depends upon pH, as shown in Fig. 3.28 (see Box 3.4). Seawater is slightly alkaline, and its pH rarely falls outside the range 7.5-8.5 owing to the buffering effect of the equilibrium between bicarbonate and carbonate ions. Consequently, seawater contains negligible carbonic acid, and OH ions are more abundant than H+, so the main equilibria can be represented by Eqns 3.9a and 3.9b, which are combined to give Eqn 3.9c ... 

A pH of 7.4 is maintained in blood partly by a carbonic acid-bicarbonate buffer system based on the following equilibrium ... 

The CO2 in the blood also participates in the carbonic acid-bicarbonate buffer equilibrium. Carbon dioxide reacts with water in the blood to form carbonic acid ... 

A carbonic acid/bicarbonate (H2C03/HC03 ) buffer is used to control the pH of blood. An important feature of this buffer is that carbonic acid decomposes to CO2 and H2O ... 

Carbonic acid has a pTf of 6.1 at physiological temperature. Is the carbonic acid/bicarbonate buffer system that maintains the pH of the blood at 7.3 better at neutralizing excess acid or excess base ... 

This sequence of reactions illustrates the buffering of the blood by the carbonic acid-bicarbonate conjugate acid-base system. The protons generated at the tissue level by the ionization of carbonic acid are removed at the lungs when bicarbonate ion reenters the red blood cell and reacts with them to produce carbonic acid. H2CO3 then rapidly dissociates, under the influence of carbonic anhydrase, and the resultant CO2 diffuses into the alveolar space because its concentration in the plasma is higher than in the alveoli. 

There are also several mechanisms by which our body maintains the pH around 7.4. Some of these mechanisms use simple standard chemistry, some are more complex. These mechanisms are (i) the carbonic acid-bicarbonate buffer system, (ii) the protein buffer system, and (iii) the phosphate buffer system. Apart from these buffers, the pH of our body is also maintained by exhalation of carbon dioxide, elimination of hydrogen ions via the kidneys, etc. 

The major buffer system used to control blood pH is the carbonic acid—bicarbonate buffer system. Carbonic acid (H2CO3) and bicarbonate ion (HCO3 ) are a conjugate acid ase pair. In addition, carbonic acid decomposes into carbon dioxide gas and vrater The important equilibria in this buffer system are... 

In the body, conjugate acid-base pairs are more common. In the blood, for example, the carbonic acid/bicarbonate pair helps to control the pH. This buffer can be overcome, though, and some potentially dangerous situations can arise. If a person exercises strenuously, lactic acid from the muscles is released into the bloodstream. If there s not enough bicarbonate ion to neutralize the lactic acid, the blood pH drops, and the person is said to be in acidosis. Diabetes may also cause acidosis. On the other hand, if a person hyperventilates (breathes too fast), she breathes out too much carbon dioxide. The carbonic acid level in the blood is reduced, causing the blood to become too basic. This condition, called alkalosis, can be very serious. 

This system is referred to as the carbonic acid-bicarbonate buffer system, and it regulates/buffers the blood pH by addressing high acid (H+) levels in the blood by... 

The renal system is another major regulator of pH balance. The kidneys can control pH by secreting H+ from the body or retaining it to reverse an acidosis or alkalosis. The renal mechanism can correct an acidosis by reabsorbing CO, which then combines with water to form carbonic acid and bicarbonate, which is released into the bloodstream, and H+, as noted earlier in the carbonic acid-bicarbonate buffer system. The renal system can can correct alkalosis by excreting the CO, resulting in less bicarbonate formation. 

---