Acid-Base 1, 2, and 3 Linda Costanzo, Ph.D.

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1 Acid-Base 1, 2, and 3 Linda Costanzo, Ph.D. OBJECTIVES: After studying this lecture, the student should understand: 1. The relationship between hydrogen ion concentration and ph. 2. Production of acid in the body: volatile and fixed. 3. The derivation of the Henderson-Hasselbalch equation. 4. Functions of extracellular and intracellular buffers. 5. Graphical presentations of the Henderson-Hasselbalch equation. 6. How the bicarbonate buffer works when fixed acid is added to the body. 7. The relationship between hydrogen and calcium ions in binding to plasma albumin, and the effect of acid-base changes on ionized calcium concentration. 8. How hemoglobin operates as a buffer for hydrogen ions. 9. How hydrogen ions enter and leave cells. 10. The principles of the four simple acid-base disorders with respect to: effects on blood ph, P CO2, and bicarbonate concentration, buffering, and expected compensatory responses. 11. How compensations for acute and chronic respiratory disorders differ. Optional Reading: Physiology, Saunders, Costanzo; W.B. Saunders, 2006; Chapter 7. Physiology, Berne and Levy; Mosby, 2004; Chapter 38 I. INTRODUCTION In normal human beings, the [H + ] of the body fluids is about 40 neq/liter (40x 10-9 Eq/liter), corresponding to a ph of 7.4. Recall that: ph = -log 10 [H + ] Use of ph, instead of concentration has the advantage of allowing us to deal with manageable numbers from 0-14, representing a range of [H + ] from Eq/liter. The figure shows the relationship between ph and [H + ] over the physiological range.

2 Figure 1. Because the ph scale is logarithimic, equal changes in ph reflect unequal changes in [H + ]. For example, a change in ph from 7.0 to 7.1 represents a change in [H + ] from 100 to 79 neq/liter; however, a ph change from 7.3 to 7.4 represents a smaller change in [H + ]. The ph of the blood is normally maintained within a very narrow range, This tight control of ph is essential for virtually all normal cellular function. The extreme range of blood ph which is compatible with life is II. ACID PRODUCTION IN THE BODY The body must maintain a normally alkaline ph in spite of the daily production of large amounts of two kinds of acid. A. Volatile acid. Volatile acid is CO 2. Between 13,000 and 20,000 mmoles of CO 2 is produced per day from oxidative metabolism. This CO 2 yields H + ions as a result of one or both of the following reactions: CO 2 + H 2 O W H 2 CO 3 W H + + HCO - 3 or C.A. CO 2 + OH - W HCO - 3 C.A. 8 HOH 6 H + C.A. = Carbonic anhydrase

3 The final products of both reactions are H + and HCO - 3. There is still debate over which form is correct. For the sake of simplicity, we shall use only the first form as this is given in most textbooks. We will consider, also for simplicity, that carbonic anhydrase catalyzes the reaction of CO 2 and H 2 O to form H 2 CO 3 and the reverse reaction. Valtin's book uses the convention shown in Equation lb. The CO 2 produced from oxidative metabolism is of course carried to the lungs via the blood. About 90% of this metabolic CO 2 transiently generates H + (Equation la), threatening H + balance. How this metabolically produced H + is buffered in transit to the lungs is important and will be covered in these lectures. B. Non-volatile or fixed acid usually comes from amino acid and phospholipid metabolism. When meat constitutes a large part of the diet, there is daily production of mmoles of inorganic and organic acid. 1. Sulfuric acid is produced from the catabolism of proteins having the sulfur-containing amino acids cysteine, cystine and methionine. The example of methionine is shown in then next equation. 2 C 5 H 11 NO 2 S + 15 O 2 W 4H SO 4-2 (methionine) + CO 2 (NH 2 ) H 2 O + 9 CO 2 (urea) 2. Phosphoric acid is formed from the catabolism of phospholipids. 3. During normal exercise or in shock, the production of fixed (nonvolatile) acids may increase due to excess lactic acid production. 4. Certain pathologic conditions are associated with increased fixed acid production. For example, aceto-acetic acid and ß - hydroxybutyric acid are produced in excess in uncontrolled diabetes mellitus; lack of insulin increases lypolysis in fat cells. The free fatty acids enter the liver where they are converted to ketoacids, weak acids which exist mainly as H + and the ketoanion at ph 7.4. III. EQUILIBRIUM CONCEPT AND LAW OF MASS ACTION The behavior of acids and bases in biological fluids conforms to the kinetics of reversible reactions, i.e. K 1 HA W H + + A - K 2

4 The rate of dissociation of HA into H + and A - is characterized by a rate constant, K 1 ; the rate of the reverse reaction, association of H + and A - to form HA, is characterized by a second rate constant, K 2. When the rates of these two opposing reactions are exactly equal, then a state of chemical equilibrium exists and there is no further net change in the number of either species, HA or A -. The Law of Mass Action states that at equilibrium: Rearranging: K 1 [HA] = K 2 [H + ] [A - ] K 1 = [H + ] [A - ] K 2 [HA] The ratio of constants can be combined into one constant K', so now: K' = [H + ] [A - ] [HA] K' is the equilibrium constant of this reaction. The "prime" sign signifies that concentrations rather than activities are being used. This equation can be rearranged to solve for [H + ]: [H + ] = K' [HA] [A - ] To express [H + ] as ph, we take the negative log 10 of both sides: Remember: -log[h + ] = -log K' - log [HA] [A - ] -log [H + ] = ph -log K' = pk' -log [HA] = log [A - ] [A - ] [HA]

5 Making these substitutions: ph = pk' + log [A - ] [HA] This is the familiar Henderson - Hasselbalch equation. We will use the Henderson- Hasselbalch equation repeatedly in acid-base physiology to calculate the ph of solutions of weak acids and their conjugate bases. Note that strong acids have very high equilibrium constants, low values for pk' and dissociate almost completely; most of the acid is in the form of A - + H +. Conversely weak acids have low equilibrium constants, high values for pk', and most of the acid is not dissociated, but in the form of HA. IV. BUFFERING A. Definition When weak acids and bases are dissolved in aqueous solutions, only partial dissociation occurs. The resulting solution will contain both the acid and base forms of the parent molecule. Such solutions can resist a change in ph following addition of acid or base. This property is called buffering. The acid-base pairs are buffers. When a strong acid is added to a buffered solution, some of the added H + combine with the base (A - ) form of the buffer to yield more of the acid form (HA), rather than remaining free in solution. Thus, the increase in free [H + ] is much less than following the addition of an equivalent amount of acid to an unbuffered solution. B. Buffering in vivo. As already noted, normal acid-base balance in human beings requires defense of an alkaline blood ph against the daily production of volatile acid from CO 2 and fixed acid from protein and phospholipid catabolism. Buffering occurs in two locations: 1. Extracellular. Bicarbonate is the most important extracellular buffer. Inorganic phosphate and plasma proteins also contribute a small buffering capacity. 2. Intracellular. Buffering occurs by proteins (such as hemoglobin in the erythrocyte), organic and inorganic phosphates. In bone, CO 3-2 represents a large buffer store, whereby H + are exchanged for Na + and Ca +2 in bone matrix. Finally, most cells can exchange K + for H + to help buffer extracellular H +. Each of these forms of buffers will be discussed in detail.

6 The ability of a particular buffer to protect against ph changes depends on its concentration and its pk' in relation to the body fluids ph. The greater the buffer concentration and the closer the pk' to the ph of the body fluids, the more effective the buffering. Figure 2. The figure illustrates the buffering property of plasma. In the experiment, 156 ml of 1N HCl was infused intravenously into a dog. The dog's plasma ph fell from 7.44 to 7.14, a severe, but not fatal, acidosis. When the same amount of HCl was slowly added to 11.4 liters of distilled water (same volume as dog's total body water), the ph of the water fell precipitously after just a few meq of H + were added. The final ph was 1.84, one that would have been clearly fatal in the dog. Obviously, the dog's plasma had buffering capabilities that water did not. This example illustrates one type of buffering that occurs in humans, that of added fixed acid. The fast, immediate buffering of added fixed acid was via the H 2 CO 3 / HCO - 3 buffer pair in extracellular fluid (see next lecture). V. EXTRACELLULAR BUFFERS A. Buffering by H 2 CO 3 /HCO -3 (extracellular fluid) The first line of defense in buffering fixed acid is fast, physicochemical buffering in extracellular fluid. The H 2 CO 3 - HCO - 3 is by far the most important extracellular buffer because (a) its concentration is high, (b) its pk' is fairly close to the ph of extracellular fluid and (c) the acid form (H 2 CO 3 ) is in equilibrium with CO 2 which can be rapidly excreted by the lungs.

7 1. Applying the Henderson-Hasselbalch Equation to the HCO - 3/CO 2 system If HCl is buffered by the bicarbonate buffer system then: In this reaction, the amount of a strong acid (HCl) is reduced, the amount of a weak acid (H 2 CO 3 ) is increased; the buffer salt (HCO - 3) is depleted. The decrease in ph is not totally prevented, but minimized. ph = pk' + log [HCO - 3] [H 2 CO 3 ] H 2 CO 3 is in equilibrium with CO 2. Remember: CO 2 + H 2 O W H 2 CO 3 W H + + HCO - 3 C.A. In the presence of carbonic anhydrase, most H 2 CO 3 is in the form of dissolved CO 2 (400 dissolved CO 2 for every 1 H 2 CO 3 ). Thus a more realistic form of equation 2 is: ph = pk' + log [HCO - 3] [dissolved CO 2 + H 2 CO 3 ] The concentration of dissolved CO 2 in plasma is proportional to the partial pressure of CO 2 (PCO 2 ). The proportionality constant which converts PCO 2 in mm Hg to CO 2 concentration in mm is 0.03 (0.03 x mmhg = mm). H 2 CO 3 concentration is negligible and can be ignored. The final equation most useful to us is: ph = log [HCO - 3] 0.03 x PCO 2

8 Normal values for arterial plasma: [HCO - 3] = 24 mm PCO 2 = 40 mm Hg pk' = 6.1 at 37 C Substituting: ph = log 24mM 0.03 x 40 mm Hg ph = log 24mM 1.2mM ph = log 20 ph = 7.40 As you have undoubtedly noted, the Henderson-Hasselbalch equation correctly predicts the normal ph of arterial blood when the normal values for the HCO - 3 buffer pair are substituted. Likewise any of the three variables can be calculated when values for the other two are known. 2. Graphs of Henderson-Hasselbalch Equation. Graphical presentations of the relationship between PCO 2, bicarbonate concentration and ph (or [H + ]) are used in some texts. One such graphical representation is shown in the figure below.

9 Figure 3. The lines radiating from the origin show the relationship between arterial PCO 2 and [HCO - 3] at various values of ph. These are isohydric lines. The normal range of values for arterial blood are given by the ellipse in the center of the figure. The terms acidemia and alkalemia denote decreases and increases in plasma [H + ]. Notice that normal values of ph may obtain from certain abnormal combinations of PCO 2 and [HCO - 3].

10 3. Adding fixed acid; buffering by HCO - 3. If 12 mmoles of HCl were added to each liter of extracellular fluid, then buffering by HCO - 3 would occur: Let us assume for a moment that this newly generated CO 2 cannot be expired by the lungs. The ph of arterial blood would drop to a fatal level, 6.06: ph = log 12 mm 1.2mM + 12mM ph = log 12mM 13.2 mm ph = 6.06 However, the ventilatory response to the excess CO 2 produced from buffering occurs within seconds or minutes after the administration of HCl. If all the excess CO 2 is excreted by the lungs, then the ph will fall, but to a level compatible with life: ph = log 12 mm 1.2 mm ph = log 10 ph = 7.1

11 Respiratory compensation usually goes even further as alveolar ventilation increases in response to the CO 2 stimulus (compensatory hyperventilation). As a result, arterial PCO 2 might fall to 23mm Hg (instead of just back to 40mm Hg): ph = log 12 mm 0.03 x 23mm Hg ph = log 12 mm 0.69 mm ph = 7.34 Consequently, the ph has been restored almost to normal as a result of buffering by extracellular HCO - 3, excretion of the resulting CO 2, and additional compensatory hyperventilation by the lungs. The ph is not perfect, however, because the [HCO - 3] in arterial plasma is only ½ of its normal value. The kidneys are needed to replenish the depleted HCO - 3 stores by excretion of H + and reabsorption of newly synthesized HCO - 3. This is much slower and requires several days. 4. Real life buffering of fixed acid by extracellular HCO - 3 The above example is an extreme one, occurring only under experimental conditions or in disease. It is valuable, though, because it illustrates the buffering and compensation that occurs normally in response to the daily loads of fixed acid added from metabolism (e.g. H 2 SO 4 from sulfur-containing amino acids). Realistically, each meal might result in the release of 16 mmoles of H 2 SO 4. The extracellular fluid volume of a 70 kg man is 14 liters. Thus the [H + ] of extracellular fluid would increase by 32 mmoles/ 14 liters or about 2 mmole/liter. Sulfuric acid is the major fixed acid so quantitatively the reaction would be: 2H + + SO Na HCO - 3 W 2Na + + SO Na HCO H 2 CO 3 ^ 2CO 2 + 2H 2 0 After buffering, but before hyperventilation the ph would be: ph = log 22 mm 1.2 mm + 2 mm ph = 6.9

12 After the excess CO 2 is eliminated by the lungs: ph = log 22 mm 1.2 mm ph = 7.36 This might cause a small, imperceptible increase in alveolar ventilation. However, in the steady state the kidneys are continuously excreting H + and reabsorbing filtered and newly synthesized HCO 3 - so that the plasma HCO 3 - stays at 24 mm. Hence, the arterial ph remains at B. Buffering by HPO 4-2 /H 2 PO 4 - (extracellular fluid and urine) The phosphate buffer pair plays a small role in extracellular buffering capacity. Theoretically, phosphate should be an excellent buffer in the physiological ph range, as its pk' is 6.8. Compare buffering by CO 2 /HCO - 3 to buffering by HPO -2 4 /H 2 PO 4- shown in the figure below. Figure 4. The change in ph per meq of H + added or subtracted is smallest in the linear portion of each titration curve, i.e. this is the region of most effective buffering. The linear portion of the CO 2 /HCO - 3 curve extends from ph 5.1 to 7.1. The linear portion of the phosphate curve extends from ph 5.8 to 7.8, coinciding with the range compatible with life. In spite of this, phosphate only plays a small buffering role in extracellular fluid, because its concentration is low compared to HCO - 3. Also bicarbonate has the special property of being in equilibrium with a volatile weak acid,

13 CO 2, which can be rapidly excreted or retained by the lungs. Phosphate is a more important buffer in tubular fluid as its concentration can get quite high in the distal tubule. C. BUFFERING BY NH 3 /NH 4 + (urine) NH 4 + is the very weak acid of the base NH 3 : NH 4 + W NH 3 + H + The pk' of this reaction is about 9.2 so the NH 4 + /NH 3 pair does not function as a buffer for extracellular fluid in the conventional sense; even at very alkaline ph, most will be in the NH 4 + form. However, the NH 4 + /NH 3 buffer pair plays a special role in the renal excretion of H + as you will learn shortly. VI. INTRACELLULAR BUFFERS A. Buffering by organic phosphate These are a heterogeneous group, comprised of 2,3-DPG, glucose-1- phosphate, AMP, ADP, and ATP. The phosphate moiety functions much as the HPO 4-2 /H 2 PO 4-2 pair. Their pk's are in the very effective range of 6.0 to 7.5 The organophosphates contribute to the buffer capacity of the intracellular compartment as their concentration is high. B. Buffering by proteins Proteins can serve as buffers because they contain a large number of acidic or basic groups, such as -COOH, -NH 2 and -NH 3. Table 1 gives the pk' values for various dissociable groups found in proteins. Table 1. pk' values for dissociable groups on proteins Group (amino acid) pk' "-carboxyl 3.7 $-carboxyl (aspartic acid) 4.0 '-carboxyl (glutamic acid) 4.0 *imidazole (histidine) "-amino sulfhydryl (cysteine) 9.0 '-amino (lysine)

14 Of course, for a given protein, buffer capacity reflects the composite effect of all its dissociable groups. In mammals, most of the buffer capacity of proteins can be attributed to the imidazole group of histidine. The pk' of the imidazole group is 6.0, but is listed as in various proteins because of varying electrostatic forces. Figure 5. N-terminal a-amino groups feature a pk' in the optimal range, also. Other dissociable groups in proteins play a smaller role because their pk's are outside the useful physiological range. Largely because of the multiplicity of pk' values on a given protein, the titration curve may approach a straight line, rather than the typical sigmoid shape of the bicarbonate or phosphate curves shown in the previous lecture. 1. Plasma proteins (not intracellular, of course). Albumin accounts for much of the buffer capacity exhibited by plasma proteins because it contains 16 histidine residues/molecule. Globulins have considerably less buffer capacity.

15 Figure 6. H + and Ca 2+ compete for negatively charged binding sites on plasma proteins. In acidosis, there is increased H + in blood, H + binding to negative sites on albumin increases and displaces Ca 2+ ; thus, the ionized or free Ca 2+ increases. Conversely, in alkalosis, there is decreased H + in blood, H + binding to negative sites on albumin decreases, and the ionized or free Ca 2+ decreases, often causing symptoms of hypocalcemia. 2. Hemoglobin. Hemoglobin is by far the most important nonbicarbonate buffer of whole blood because of its high concentration inside the erythrocyte and its high buffer capacity. Most of the buffering capacity of hemoglobin in the physiological ph range results from its 36 histidine residues (9 on each of its 4 polypeptide chains). There are also α-amino groups which are sufficiently exposed to buffer H +. Hemoglobin also has some special attributes as a buffer. The pk' for deoxygenated hemoglobin is about 7.9. The pk' of the oxygenated form is lower than of the deoxygenated form. Hb n- n- HbO 2 1. pk' K' lower higher 3. acidity weaker stronger 4. electrostatic bonds more less Thus, HbO 2 n- is a stronger acid and gives up H + more readily than does Hb n-. (This property also forms the basis for the Bohr effect.)

16 The mechanism is as follows. Upon oxygenation, the Hb molecule undergoes a marked conformational change. The groups responsible for the change in acidity (or pk') are C-terminal histidines in the chain. These histidine residues participate in electrostatic bonds when the subunits are deoxygenated. Binding of O 2 causes a conformational change in Hb, such that the distances are now too great for electrostatic bonding to occur. The electrostatic binding doesn't occur and these H + are more readily dissociated; thus the stronger acid and lower pk'. The implications of this change in pk' on acid-base balance are illustrated in the figure below. Figure 7. First notice that oxygenated Hb is an excellent buffer for added H +, as expected (lower line). Deoxygenated Hb is an even better buffer because for every H + added a smaller reduction in ph occurs (upper line). Thus as Hb becomes deoxygenated at the arteriolar end of the capillary it becomes a more effective buffer. Conveniently, this is the precise moment when CO 2 (and thus H + ) are being added from the tissue cells to the capillary blood. The reduction of HbO 2 n- to Hb n- would have caused a tremendous increase in ph were it not for the H + being added to the blood at this time. The increase in pk' of Hb is such that 1.3 mmoles of CO 2 can be added to each liter of capillary blood

17 without a change in ph. 95% of the CO 2 added to each liter of blood, about 1.6 mmoles, is rapidly converted to H +. Since 1.3 mmoles can be buffered by Hbn- without a change in ph, the drop in venous blood ph due to addition of CO 2 is minimized. H + buffering by hemoglobin will be discussed again in the context of CO 2 transport in venous blood from the tissue to the lungs. C. Other intracellular proteins. Various cell proteins contribute to intracellular H + buffering, largely due to the presence of the imidazole group of histidine. D. How H + enters cells to be buffered How do H + enter and leave the ICF in order to take advantage of the vast supply of intracellular buffers? In respiratory disturbances, where there is an excess or deficit of CO 2, CO 2 readily crosses cell membranes; thus, in respiratory acidosis CO 2 enters the cells and in respiratory alkalosis CO 2 leaves the cells. In metabolic disturbances, where there is an excess or deficit of fixed H +, H + can cross cell membranes in exchange for another cation, K +, or it can cross with an organic anion such as lactate, β-ohbutyrate etc. Either way, electroneutrality must be preserved. If H + exchanges for K + (H + -K + shift), then a disturbance of K + metabolism occurs acidemia producing hyperkalemia and alkalemia producing hypokalemia. More commonly, there is an accompanying organic anion to move with H + (e.g., lactic acidosis) and an H + -K + shift is not required. Figure 8.

18 VII. CO 2 TRANSPORT IN BLOOD As stated before, some 13,000 to 20,000 mmoles of CO 2 are produced daily in an adult as a result of metabolism. This CO 2 can generate H + and potentially disturb H + balance. Elimination of CO 2 by the lungs prevents acidosis, but the transport of CO 2 in the venous blood to the lungs presents a major buffering problem. We know that buffering of CO 2 -generated H + in blood is very effective because the venous blood ph is seldom less than 7.37, only 0.03 ph units more acid than arterial blood. The buffering events in its transport to the lungs are as follows: Figure 9. Transport of carbon dioxide (CO 2 ) in the blood. CO 2 and H 2 O are converted to H - and HCO 3 inside red blood cells. H + is buffered by hemoglobin (Hb - H) inside the red blood cells. HCO 3 exchanges for Cl and is transported in plasma. The circled numbers correspond to the lettered steps discussed in the text below. A. CO 2 diffuses from tissues cells into the capillary blood because pco 2 is much higher in the metabolizing cells than in the blood. B. The red blood cell membranes are highly permeable to CO 2 and hence CO 2 will also enter the erythrocytes. Both in the plasma and in the red cells, CO 2 reacts with H 2 O to form H 2 CO 3 and thereafter H + and HCO - 3. In plasma the reaction is slow due to the lack of carbonic anhydrase; in red cells it is fast because the enzyme is present. What little H + is produced in plasma is buffered by plasma proteins and phosphate. In the red cell, the H+ produced is buffered by hemoglobin which has a large buffering capacity. You will recall that hemoglobin is an even more effective buffer in its deoxygenated form.

19 C. The HCO - 3 formed from the reaction of CO 2 and H 2 O in the red cells diffuses back into the plasma down a concentration gradient and Cl - moves into the red cells to maintain electrical neutrality. This is Cl-HCO - 3 exchange. Most of the CO 2 generated in tissues is carried to the lungs as HCO - 3 in plasma. D. Some of the CO 2 which enters the red cells is not converted to H + and HCO - 3. Instead, it combines with hemoglobin to form carbaminohemoglobin and is carried to the lungs in that form. A small amount of CO 2 is carried to the lungs dissolved in plasma or in the red blood cell water. E. In the lungs, the entire process operates in reverse as PCO 2 is lower in the alveoli than in the capillary blood perfusing the lung tissue. Thus all reaction are driven toward the reformation of CO 2 so that it can be eliminated. VIII. SIMPLE ACID-BASE DISORDERS There are four "simple" acid-base disturbances (where "simple" means one acidbase disorder at a time): metabolic acidosis, metabolic alkalosis, respiratory acidosis, and respiratory alkalosis. The metabolic disturbances involve gain or loss of fixed H +. The respiratory disturbances involve gain or loss of CO 2. In this section, we will compare the relative contributions of ECF and ICF buffers, respiratory compensation (if it occurs), and renal compensation for each of the four simple disorders. A. Metabolic acidosis Metabolic acidosis is caused by excess fixed acid, either due to ingestion of acid, overproduction of acid (e.g., diabetic ketoacidosis), or loss of HCO 3 - (diarrhea). The excess fixed H + is buffered by ECF HCO 3 - and, as a result, the blood HCO 3 - concentration decreases. This causes the arterial ph to decrease (H-H equation) and the acidemia stimulates central chemoreceptors. There is ensuing hyperventilation, which is the respiratory compensation for metabolic acidosis. This hyperventilation results in decreased PCO 2, which tends to normalize the ph (i.e., is a "compensation"). Buffering also takes place in ICF, utilizing protein and organic phosphate buffers and there is a decrease in intracellular ph. Ultimately, the kidneys excrete the excess fixed H + and synthesize new HCO 3 - to replace the HCO 3 - that was consumed in buffering the fixed H +.

20 B. Metabolic alkalosis Metabolic alkalosis is caused by loss of fixed H + (e.g., vomiting). Loss of H + causes the blood HCO 3 - concentration to increase and the arterial ph to increase (H-H equation). The alkalemia inhibits central chemoreceptors and there is hypoventilation, which is the respiratory compensation for metabolic alkalosis. The hypoventilation causes increased PCO 2, which tends to normalize the ph. Buffering also occurs in ICF and intracellular ph increases. Ultimately, the kidneys excrete the excess HCO 3 - and there is restoration of acid-base balance. C. Respiratory acidosis Respiratory acidosis is caused by hypoventilation leading to retention of CO 2, increased PCO 2, and decreased ph (H-H equation). There is no respiratory compensation for respiratory acidosis (that's impossible); thus, in the acute phase, ph may be markedly decreased. Buffering takes place almost exclusively in ICF. Within 2-3 days, the kidneys increase HCO 3 - reabsorption, which increases the blood HCO 3 - concentration. This increase in blood HCO 3 - is the renal compensation for respiratory acidosis and tends to normalize the ph. For the same value of PCO 2, chronic respiratory acidosis will have a higher ph (more normal) than acute respiratory acidosis. D. Respiratory alkalosis Respiratory alkalosis is caused by hyperventilation leading to loss of excess CO 2, decreased PCO 2, and increased ph (H-H equation). There is no respiratory compensation for respiratory alkalosis, and in the acute phase ph may be markedly increased. Buffering is exclusively in the ICF. Within 2-3 days, the kidneys decrease HCO 3 - reabsorption (the renal compensation), which decreases the HCO 3 - concentration and tends to normalize ph. For the same value of PCO 2, chronic respiratory alkalosis will have a lower ph (more normal) than acute respiratory alkalosis.

21 E. Summary of contributions of extracellular and intracellular buffering in the acid-base disorders. Extracellular Intracellular Metabolic acidosis 45% 55% Metabolic alkalosis 70% 30% Respiratory acidosis 3% 97% Respiratory alkalosis 1% 99%