Buffer systems of the blood

You might be wondering why the bicarbonate buffer can buffer effectively at pH 7.4 when its pKa is 6.1. The answer is that it doesn t buffer all that well. What makes it unique and the major buffer system of the blood is that C02, being a gas, can be exhaled by the lungs. Exhaling C02 is equivalent to exhaling protons. 

Blood has several buffer systems that work together to maintain a narrow pH range between 7.35 and 7.45. A pH value above or below these levels can be lethal, primarily because cellular proteins become denatured, which is what happens to milk when vinegar is added to it. The primary buffer system of the blood is a combination of carbonic acid and its salt, sodium bicarbonate, shown in Figure 10.21. Any acid that builds up in the bloodstream is neutralized by the basic action of sodium bicarbonate, and any base that builds up is neutralized by the carbonic acid. 

The buffer systems of the blood (mainly the bicarbonate/ carbonic acid buffer) minimize changes in pH. In acidoses, the bicarbonate concentration decreases to give a ratio of cHC03/cdC02 of <20 1. The respiratory compensatory mechanism responds to correct the ratio with increased rate and depth of respiration to eliminate CO2. Table 46-3 depicts expected compensation in both acidoses and alkaloses and corresponding laboratory values. 

The Henderson-Hasselbalch equation was developed independently by the Ameriean biological chemist L. J. Henderson and the Swedish physiologist K. A. Hasselbaleh, for relating the pH to the bicarbonate buffer system of the blood (see below). In its general form, the Henderson-Hasselbalch equation is a useful expression for buffer caleulations. It can be derived from the equilibrium constant expression for a dissociation reaction of the general weak acid (HA) in Equation (1.3) ... 

A buffer system prevents marked changes in the pH of a solution when an acid or base is added. Three major buffer systems of the blood are the bicarbonate buffer, the phosphate buffer, and the plasma proteins. The most important of these is the bicarbonate buffer system, consisting of a mixture of bicarbonate ions (HCO3 ) and carbonic acid (H2CO3). 

List three major buffer systems of the blood. 

Another formulation variable that must be considered is that of the solution pH and bulfer capacity. Since the anterior chamber fluid (aqueous humor) contains essentially the same buffering systems as the blood, products with a pH outside the physiological range of 7.0-7.4 are converted to this range by the buffering capacity of the aqueous humor if a relatively small volume of the solution is introduced. Often,... 

Buffers are extremely important in biological systems. The pH of arterial blood is about 7.4. The pH of the blood in your veins is just slightly less. If the pH of hlood drops to 7.0, or rises above 7.5, life-threatening problems develop. To maintain its pH within a narrow range, blood contains a number of buffer systems. The most important buffer system in the blood depends on an equilibrium between hydrogen carbonate ions and carbonate ions. Dissolved carbon dioxide reacts with water to form hydrogen carbonate ions. 

The second dissociation step in phosphate (H2P04/HP04 ) also contributes to the buffering capacity of the blood plasma. Although the pKa value of this system is nearly optimal, its contribution remains small due to the low total concentration of phosphate in the blood (around 1 mM). 

One of the most important buffer systems in the human body is that which keeps the pH of blood around 7.4. If the pH of blood fall below 6.8 or above 7.8, critical problems and even death can occur. There are three primary buffer systems at work in controlling the pH of blood carbonate, phosphate, and proteins. The primary buffer system in the blood involves carbonic acid, H COj and its conjugate base bicarbonate, HCO3. Carbonic acid is a weak acid that dissociates according to the following reaction ... 

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. 

H2P04, and in Equation 7.61, P04 is the salt of the acid HP04 . So each of these pairs constitutes a buffer system, and orthophosphate buffers can be prepared over a wide pH range. The optimum buffering capacity of each pair occurs at a pH corresponding to its Ka. The HP04 /H2P04 couple is an effective buffer system in the blood (see below). 

Note that these equilibria and Equation 7.88 hold although there are other buffer systems in the blood. The pH is the result of all the buffers and the [HCO3"]/ H2CO3] ratio is set by this pH. 

Fig. 4.9. Buffering systems of the body. COj produced from cellular metabolism is converted to bicarbonate and H in the red blood cells. Within the red blood cell, the is buffered by hemoglobin (Hb) and phosphate (HP04 ). The bicarbonate is transported into the blood to buffer Regenerated by the production of other metabolic acids, such as the ketone body acetoacetic acid. Other proteins (Pr) also serve as intracellular buffers.

The level of total ketone bodies in Lofata Burne s blood greatly exceeds normal fasting levels and the mild ketosis produced during exercise. In a person on a normal mealtime schedule, total blood ketone bodies rarely exceed 0.2 mM. During prolonged fasting, they may rise to 4 to 5 mM. Levels above 7 mM are considered evidence of ketoacidosis, because the acid produced must reach this level to exceed the bicarbonate buffer system in the blood and compensatory respiration (Kussmaul s respiration) (see Chapter 4). 

The control of pH within narrow limits is critically important in many chemical applications and vitally important in many biological systems. For example, human blood must be maintained between pH 7.35 and 7.45 for the efficient transport of oxygen from the lungs to the cells. This narrow pH range is maintained by buffer systems in the blood. 

In blood, the main buffer system is bicarbonate at a concentration of [HCOj"] = 0.02-0.03 M (20-30 mEq/ I). Hemoglobin provides a further 10 mEq/l buffer capacity, and phosphate makes a small contribution of 1.5 mEq/l. The 5 liters of blood in an average adult human are thus able to absorb about 0.15 mole before the pH becomes dangerously low. The major buffers of the body are, however, present in other tissues. The total musculature of the body, for example, can neutralize about 5 times as much acid as the blood, and the blood HCO37CO2 system represents only about a tenth of the total buffer capacity of the body. Since all the buffer systems of the body are able to interact and buffer each other, all changes in the acid/ base balance of the body are reflected in the blood. This mutual buffering by the shift of H from one body system to another is known as the isohydric principle. 

The body has three lines of defence buffering by the buffer systems of the body (blood, extracellular fluid, intracellular fluid and bone), respiratory compensa-... 

The first line of defence against a change in [H ] is provided by the buffer systems of the body. The buffers of the blood and extracellular fluid are immediately available whereas those of intracellular fluid take a matter of minutes to become effective. Because of its relatively low blood flow, bone, with an immense buffering power, requires hours or days to become available. 

The production of H+ ions, if left unchecked, would lower the pH of the blood and cause acidosis . This may disrupt some body functions and eventually lead to coma. The equilibrium between carbon dioxide and hydrogencarbonate is the most important buffering system in the blood. 

The system HCOa" H2CO3 is the most important buffer system of the entire organism. At the pH of blood (7.4), HCOs" and H2CO3 are at a ratio of 20 1. 

Some of the sulphonamides can be used as diuretics. The mechanism of their action relates to carbonic add excretion. Carbon dioxide generated from catabolic processes is carried to the lung and then removed by exhalation. However, part of the carbon dioxide is still dissolved in the blood. Hie dissolved carbon dioxide produces carbonic acid and its conjugated base (i.e. bicarbonate). This mixture of the weak acid and its conjugate base is one of the important buffer systems in the blood. The dissolved carbon dioxide is excreted in the urine. The processes of the conversion of carbon dioxide to carbonic acid and then... 

The important buffer system of blood plasma is the bicarbonate/carbonic acid couple ... 

The pH of the blood is maintained by a finely tuned buffering system, consisting primarily of hydrogen carbonate ion (HC03 ) and H30+ in equilibrium with water and C02. 

Moreover, several buffer systems exist in the body, such as proteins, phosphates, and bicarbonates. Proteins are the most important buffers in the body. Protein molecules contain multiple acidic and basic groups that make protein solution a buffer that covers a wide pH range. Phosphate buffers (HPO T /H2P07) are mainly intracellular. The pK of this system is 6.8 so that it is moderately efficient at a physiological pH of 7.4. The concentration of phosphate is low in the extracellular fluid but the phosphate buffer system is an important urinary buffer. Bicarbonate (H2C03/HC0 3) is also involved in pH control but it is not an important buffer system because normal blood pH 7.4 is too far from its pK 6.1 [144],... 

The buffering capacity of a buffer system depends on its concentration and its pKg value. The strongest effect is achieved if the pH value corresponds to the buffer system s pKa value (see p. 30). For this reason, weak acids with pKa values of around 7 are best suited for buffering purposes in the blood. 

The ketone bodies are carboxylic acids, which ionize, releasing protons. In uncontrolled diabetes this acid production can overwhelm the capacity of the blood s bicarbonate buffering system and produce a lowering of blood pH called acidosis or, in combination with ketosis, ketoacidosis, a potentially life-threatening condition. 

No studies on body burden reduction methods were located. The state of definitive knowledge of white phosphorus metabolism is too limited to permit extensive speculation on methods for reducing body burden. However, it is possible that increasing selective excretion of phosphate may increase the rate of inorganic conversion of white phosphorus to phosphate (this conversion is described in detail in Section 2.3). Since phosphate is a naturally occurring component of the blood s buffering system, this would effectively deactivate the phosphorus. No methods for selectively increasing phosphate excretion were located. 

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