Bicarbonate buffer equation

Henderson-Hasselbalch Equation Titration Curves p/—Isoelectric Point The Bicarbonate Buffer Imbalance in Blood pH Acidosis and Alkalosis... 

There are two ways of dealing with the bicarbonate buffer system. The first uses the Henderson-Hasselbalch equation and an effective pKa of 6.1. If there is more base (HCO 3) than acid (C02), the pH will always be bigger than the pKa. This is usually the case physiologically (pH = 7.4 pKa = 6.1) so that on a molar basis there is always more than 10-fold more HCO 3 than C02. 

The buffer value (p) is defined as the amount of base required to cause a change in pH of one unit. The buffer value of the bicarbonate buffer in plasma is 55.6 mmol/L. Derivation of this value is obtained by taking partial differentials of the Henderson-Hasselbalch equation, which is presented in detail in the second edition of this textbook. ... 

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) ... 

R and Ki as Functions of 7r, pH, and Buffer Composition. R (Equation 1) and Ki (Equation 2) for palmitic acid varied directly with tt and pH (Figure 6). Desorption occurred, yielding measurable R and values even at pH 9 and tt at 17 dynes/cm where the film was condensed (Figure 4). Furthermore, both R and K were lower for the fatty acid spread on tris buffers than for the fatty acid spread on bicarbonate buffers (Figure 6), showing that apparent film expansion on tris buffers (Figure 3) was, indeed, a desorption rate effect. 

H CO, can therefore be replaced in the mass action equation by PCO,. At this point, an understanding of the role of the bicarbonate buffer system in assessing clinical acid-base disorders can be achieved simply by reference to the relation.ship ... 

Equation 4.6 The Henderson-Hasseibaich equation for the bicarbonate buffer system... 

Determination of Alkaline Phosphatase Activity. Vomeronasal and olfactory epithelia were removed and transferred to 0.9% NaCl solution where they were maintained at 4°C. All soluble forms of enzymes were washed out from the surface of the receptor tissue (see Chukhrai et al., 1992 for details). Using procedures that are described in detail in Chukhrai et al. (1992), alkaline phosphatase activity in olfactory and vomeronasal epithelia was determined at pH=8.3 in 0.9% bicarbonate buffer. Disodium p-ni-trophenol phosphate (Sigma) was used as a substrate its initial concentration was 8.1 x 10" M in the buffer. The increase of p-nitrophenol (the product of p-nitrophenylphosphate hydrolysis) was measured with a double-beam spectrophotometer at 400 nm every 30 sec. The velocity of the reaction was determined by estimating the angle of the slope (optical units [OU]/min). The obtained values were divided into the extinction coefficient (C ) under corresponding pH values to obtain the reaction velocity values in pM/min. Effective parameters of the Michaelis-Menten equation were compared using standard procedures (Chukhrai et al., 1992). 

This bicarbonate buffer system is quantitatively the most important of the three because of the high bicarbonate concentration in blood which results from the large amounts of CO2 produced by the metabolic activity of cells. Loss of CO2 in the lungs displaces the equation to the left once more and prevents build up of hydrogen ions. The p/sT value of this system is considerably below pH 7-4, the normal physiological value for blood, and can be evaluated by applying the Henderson-Hasselbalch equation. The normal plasma bicarbonate value is 27 m-equiv per litre, approximately twenty times the total solubility of carbon dioxide of 1-35 mmol per litre (at a partial pressure of 40mmHg) From Equation (3-11)... 

Equation (3-13) shows that the equilibrium pH of the bicarbonate buffer system of plasma can be to some extent controlled by varying the partial pressure of carbon dioxide in the air to which the blood is exposed (i.e. in the lungs). Reduction in partial pressure results in carbon dioxide leaving the blood with a rise in the last term of Equation (3-13) provided that the bicarbonate concentration remains constant. The equilibrium pH of the buffer system and hence the pH of the blood consequently rise. Conversely an increase in partial pressure of carbon dioxide in the alveolar air will result in a fall in the pH of blood plasma. In practice the partial pressure of carbon dioxide in the alveolar air is controlled by the rate of pulmonary ventilation in relation to the rate of production of carbon dioxide by metabolic oxidation within the body. Increased ventilation (i.e. hyperventilation) will lower the partial pressure of carbon dioxide and raise the blood pH, while decreased ventilation raises the partial pressure, making the blood more acid (metabolic acidosis). Normally the respiratory centre controls the rate of ventilation to keep the partial pressure of carbon dioxide close to the normal value of 40 mmHg. 

Equation (6.5.2-3) does not contain ki, so that, when NaAv and Pa are measured and k, Da, and H are known. Ay can be calculated. To satisfy the above conditions, use is made of the reaction between CO2 and a carbonate-bicarbonate buffer containing arsenite [Sharma and Danckwerts, 1963 Roberts and Danckwerts, 1962] or between CO2 and aqueous amines [Sharma, 1965). 

From this equation, it can be seen that when the reaction is allowed to go to completion there will result a disappearance of 50% of the acid-labile phosphate. Because of the fact that in the above reaction an acid equivalent is liberated, owing to the strongly acid character of the new OH group of ADP, Colowick and Kalckar have developed a manometric method for following the course of this reaction. If the reaction is conducted in a bicarbonate buffer, the evolution of CO2 can be used to estimate the transfer of phosphate and thus the original ATP concentration. 

Henderson-Hasselbalch Equation n A formula relating the pH value of a solution to the pK value of the acid in the solution and the ratio of the acid and the conjugate base concentrations pH = pIQ + log( [A—] / [HA]), where [A—] is the concentration of the conjugate base and [HA] is the concentration of the protonated acid. For the bicarbonate buffer system in blood,... 

It is important to realize that the serum HCO, concentration may be affected by the presence of unmeasured endogenous acids (lactic acidosis or ketoacidosis). Bicarbonate will attempt to buffer these acids, resulting in a 1 mEq loss of serum HCO, for each 1 mEq of acid titrated. Because the cation side of the equation is not affected by this transaction, the loss of serum HC03 results in an increase in the calculated anion gap. Identification of an increased anion gap is very important for identifying the etiology of the acid-base disorder. The concept of the increased anion gap will be applied later in the case studies section. 

Dissolved carbon dioxide produces carbonic acid, which ionizes to bicarbonate and carbonate ions, the reactions for which are shown in Figure 5.2 (equations 1-3). This reaction sequence is extremely important because bicarbonate is a counterion to many cations, is active in buffering the soil solution, and is involved either directly or indirectly in many soil chemical reactions. Bicarbonates are generally more soluble than carbonates, which are generally insoluble. Adding acid to carbonates or bicarbonates results in the release of carbon dioxide and the formation of the salt of the acid cation. The acid is thus neutralized. 

The Henderson-Hasselbach equation allows the ratio of ionized un-ionized compound to be found if the pH and pKa are known. Consider carbonic acid (H2CO3) bicarbonate (HC03 ) buffer system... 

Limited pH changes may occur if water electrolysis reactions (Equations 3 and 4) occur at the same rate and efficiency. In a completely mixed reactor, the proton produced at the anode should neutralize the hydroxyl ion produced at the cathode. However, the results indicated that the pH decreased to less than 5.5 even under completely mixed conditions in fed-batch reactors. The pH drop indicate less hydroxyl production at the cathode, either because different electrolysis reactions occurred (other than Equation 4) or because of biochemical reactions in the reactor. The type and concentrations of ions in the solution will impact the pH changes and require further investigation. Sodium bicarbonate was used and was effective in buffering the system for the range of electric field strengths studied. 

All liming materials, whether oxide, hydroxide, or carbonate, react with soil water and carbon dioxide to yield the bicarbonate form when applied to acid soil. The partial pressure of carbon dioxide in the soil usually is several hundred times greater than that in atmospheric air, and drives the reaction that produces Ca(HC03)2, which is very important in buffering the soil solution (see Equation 3). 

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... 

Effect of Holding One s Breath on Blood pH The pH of the extracellular fluid is buffered by the bicarbonate/carbonic acid system. Holding your breath can increase the concentration of C02(g) in the blood. What effect might this have on the pH of the extracellular fluid Explain by showing the relevant equilibrium equation(s) for this buffer system. 

Brpnsted theory, 23 Definition of Kb, 38 Lewis theory, 24 HSAB theory, 12 Base saturation (%), 163 Basic organic compounds, 356 Bicarbonate, 30-33 Biotite, 104, 108 Boltzmann equation, 143 Bonding, 6-12 Covalent, 7 Ionic, 7 Boron, 127 Buffer capacity, 86... 

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