Carbonic acid-bicarbonate system

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

The answer is C. Ingestion of an acid or excess production by the body, such as in diabetic ketoacidosis, may induce metabolic acidosis, a condition in which both pH and HCOj become depressed. In response to this condition, the carbonic acid-bicarbonate system is capable of disposing of the excess acid in the form of CO2. The equilibrium between bicarbonate and carbonic acid shifts toward formation of carbonic acid, which is converted to COj and HjO in the RBC catalyzed by carbonic anhydrase, an enzyme found mainly in the RBC. The excess CO2 is then expired by the lungs as a result of respiratory compensation for the acidosis (Figure 1-2). The main role of the kidneys in managing acidosis is through excretion of H" rather than CO2. 

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... 

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. 

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. 

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. 

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. 

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 bicarbonate-carbonic acid buffer system plays a major role in regulating the pH of fluids in tissue spaces outside blood vessels. This fluid, commonly referred to as interstitial fluid and separated from plasma by the membrane barrier known as the capillary endothelium, primarily... 

The three Salem samples show increased leaching rates as a function of increased acidity, but not the rate values predicted by simple chemical stoichiometry. A pH decrease from 5.6 to 4.0 is a 39.8x increase in acidity, while a change from 4.0 to 3.0 is a 10x acidity increase. The observed changes were factors of 2.88x and 1.58x, respectively. These discrepancies can be attributed to the complex equilibrium interactions involved in the solubilities of the metal carbonates. These two solubility equilibria are further complicated by the two acid equilibria for the carbonic acid/bicarbon-ate/carbonate system in addition to the equilibrium solubility of congas in water. The solution of these simultaneous equilibria processes to determine the relationship between carbonate solubility and acid concentration is a non-trivial one (sixth degree in concentration). This solubility problem has been approached from several different viewpoints (31-35), the most convenient being a graphical solution of the solubility as a function of initial solution and final solution pH. From this method, it can be theoreti-... 

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

The solubihty of proteins in blood requires a pH in the range of 7.35 to 7.45. The bicarbonate-carbonic acid buffer system of blood (HCOj +... 

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 ... 

An average rate of metabolic activity produces roughly 22,000 mEq acid per day. If all of this acid were dissolved at one time in unbuffered body fluids, their pH would be less than 1. However, the pH of the blood is normally maintained between 7.36 and 7.44, and intracellular pH at approximately 7.1 (between 6.9 and 7.4). The widest range of extracellular pH over which the metabolic functions of the liver, the beating of the heart, and conduction of neural impulses can be maintained is 6.8 to 7.8. Thus, until the acid produced from metabolism can be excreted as CO2 in expired air and as ions in the urine, it needs to be buffered in the body fluids. The major buffer systems in the body are the bicarbonate-carbonic acid buffer system, which operates principally in extracellular fluid the hemoglobin buffer system in red blood cells the phosphate buffer system in all types of cells and the protein buffer system of cells and plasma. 

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. 

One very important buffer solution is human blood An equilibrium between carbonic acid (H2CO3) and its conjugate base bicarbonate (HCOsi helps blood to maintain a relatively constant pH of around 7.4. The carbonic acid buffer system is created by carbon dioxide (CO2) dissolved in blood carbon dioxide reacts with water (H2O) to form carbonic acid. Since the amount of carbon dioxide in the blood depends on the rate at which you breathe, your blood pH is influenced by your breathing rate. Your body can... 

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... 

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. 

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