Acid-Base Physiology PDF

PAGE 1 https://www.umfst.ro Lecture no 4 https://edu.umch.de Acid base balance 2024 May Lecturer Florina Gliga Introduction PAGE 2...

PAGE 1 https://www.umfst.ro Lecture no 4 https://edu.umch.de Acid base balance 2024 May Lecturer Florina Gliga Introduction PAGE 2 Although present in exceedingly low concentrations in most body fluids, protons nevertheless have a major impact on biochemical reactions → on a variety of physiological processes → critical for the homeostasis of the entire body and individual cells. Introduction PAGE 3 Acid-base balance is concerned with maintaining a normal hydrogen ion concentration in the body fluids. There must be a balance between the intake or production of H + and net removal of H + from the body to achieve homeostasis. This balance is achieved by – utilization of buffers in extracellular fluid and intracellular fluid, – respiratory mechanisms that excrete carbon dioxide, – renal mechanisms that reabsorb bicarbonate and secrete hydrogen ions. Introduction PAGE 4 Acids and Bases—Their Definitions and Meanings A hydrogen ion (H + ) is a single free proton released from a hydrogen atom. Molecules containing hydrogen atoms that can release hydrogen ions in solutions are referred to as acids. – A strong acid is one that rapidly dissociates and releases especially large amounts of H + in solution. An example is HCl. – Weak acids are less likely to dissociate their ions → release H + less vigorous. An example is H 2 CO 3. Introduction PAGE 5 Acids and Bases—Their Definitions and Meanings A base or alkali is an ion or a molecule that can accept an H +. – A strong base is one that reacts rapidly and strongly with H + → quickly removes H + from a solution. A typical example is OH − , which reacts with H + to form water (H 2 O). – A weak base binds with H + much more weakly than does OH −. A typical weak base is HCO 3 − The proteins in the body also function as bases because some of the amino acids that make up proteins have net negative charges that readily accept H +. – hemoglobin in the red blood cells - important basic activity Introduction PAGE 6 Strong and Weak Acids and Bases. Most acids and bases in the extracellular fluid that are involved in normal acid-base regulation are weak acids and bases. The most important ones are carbonic acid (H 2 CO 3) and HCO 3 − base. pH of Body Fluids PAGE 7 The hydrogen ion (H + ) concentration of the body fluids is extremely low. In arterial blood the H + concentration is 40 × 10 −9 equivalents per liter (or 40 nEq/L), which is more than six orders of magnitude lower than the sodium (Na + ) concentration. Because it is inconvenient to work with such small numbers, H + concentration is routinely expressed as a logarithmic function called pH: pH=−log10[H+] Calculated and converted to pH as follows: pH =−log10[40×10−9Eq/L]=7.4 pH of Body Fluids PAGE 8 When using pH instead of H + concentration, there are two points of caution. 1. Because of the minus sign in the logarithmic expression, a mental reversal is necessary: – As H + concentration increases, pH decreases – As H + concentration decreases, pH increases. 2. The relationship between H + concentration and pH is logarithmic, not linear - equal changes in pH do not reflect equal changes in H + concentration. – In other words, a given change in pH in the acidic range (pH < 7.4) reflects a larger change in H+ concentration than the same change in pH in the alkaline range (pH > 7.4). pH of Body Fluids PAGE 9 Image – the lack of linearity the relationship between H + concentration and pH over the physiologic range in body fluids. an increase in pH from 7.4 to 7.6 (0.2 pH units) reflects a decrease in H + concentration of 15 nEq/L; a decrease in pH from 7.4 to 7.2 (also 0.2 pH units) reflects a larger increase in H + concentration of 23 nEq/L. Costanzo, Linda Physiology, Chapter 7 Copyright © 2018 by Elsevier, Inc. All rights reserved.  pH of Body Fluids PAGE 10 The normal range of arterial pH is 7.37–7.42/ 7.35-7.45 When arterial pH is less than 7.35, it is called acidemia. When arterial pH is greater than 7.45, it is called alkalemia. – The pH range compatible with life is 6.8–8.0. Every organ system of the human body relies on pH balance A pH at this level is ideal for many biological processes, one of the most important being the oxygenation of blood. Also, many of the intermediates of biochemical reactions in the body become ionized at a neutral pH, which causes the utilization of these intermediates to be more difficult.  pH of Body Fluids PAGE 11 Intracellular pH is approximately 7.2 and slightly lower than extracellular pH. – Transporters in cell membranes regulate intracellular pH. Na + -H + exchangers extrude H+ from cells - alkalinized intracellular fluid (ICF). Cl − -HCO 3 − exchangers extrude HCO 3 − tends to acidify ICF. – Depending on the type of cells, the pH range of intracellular fluid is between 6.0 and 7.4. – Hypoxia and poor blood flow to the tissues can cause acid accumulation and decreased intracellular pH. pH of Body Fluids PAGE 12 H + Concentration (mEq/L) pH Extracellular fluid Arterial blood 4.0 × 10 −5 7.40 Venous blood 4.5 × 10 −5 7.35 Interstitial fluid 4.5 × 10 −5 7.35 Intracellular fluid 1 × 10 −3 to 4 × 10 −5 6.0-7.4 Urine 3 × 10 −2 to 1 × 10 −5 4.5-8.0 Gastric HCl 160 0.8 pH and H + Concentration of Body Fluids Guyton and Hall Textbook of Medical Physiology pH of Body Fluids PAGE 13 The mechanisms that contribute to maintaining pH in the normal range include: – buffering of H + in both extracellular fluid (ECF) and ICF, – respiratory compensation, – renal compensation. The mechanisms for buffering and respiratory compensation occur rapidly, within minutes to hours. The mechanisms for renal compensation are slower, requiring hours to days. pH of Body Fluids PAGE 14 Acid Production in the Body Arterial pH is slightly alkaline (7.4) despite the production of large amounts of acid on a daily basis. This acid production has two forms: – volatile acid (carbon dioxide, CO 2) – nonvolatile, or fixed, acid. Both volatile and fixed acids are produced in large quantities and present a challenge to the normally alkaline pH. pH of Body Fluids PAGE 15 Acid Production in the Body CO 2 CO 2 , or volatile acid, is the end product of aerobic metabolism in the cells and is generated at a rate of 13,000–20,000 mmol/day. CO 2 itself is not an acid. However, when it reacts with water (H 2 O), it is converted to the weak acid carbonic acid, H 2 CO 3 : CO2 +H2O ↔ H2 CO3 ↔ H + + HCO 3 − carbonic anhydrase H 2 CO 3 dissociates into H + and HCO 3 − , and the H + generated by this reaction must be buffered. pH of Body Fluids PAGE 16 Acid Production in the Body CO 2 CO 2 is produced by the cells is added to venous blood, converted to H + and HCO 3 − within the red blood cells, and carried to the lungs. In the lungs, the reactions occur in reverse and CO 2 is regenerated and expired. (CO 2 is therefore called a volatile acid.) Thus buffering of the H + that comes from CO 2 is only a temporary problem for venous blood. pH of Body Fluids PAGE 17 Acid Production in the Body CO 2 Effect of blood pH on the alveolar ventilation rate. Guyton and Hall Textbook of Medical Physiology, Chapter 31, Copyright © 2016 by Elsevier, Inc. All rights reserved. pH of Body Fluids PAGE 18 Acid Production in the Body Fixed Acid Catabolism of proteins and phospholipids results in the production of approximately 50- 80 mmol/day of fixed acid. Proteins – with the sulfur-containing amino acids generate sulfuric acid, – phospholipids generate phosphoric acid when they are metabolized. – both, are non volatile acids (in contrast with CO 2 - expired by the lungs). Therefore, fixed acids first must be buffered in the body fluids until they can be excreted by the kidneys. Sulfuric and phosphoric acids are produced from normal catabolic processes, in physiologic states. pH of Body Fluids PAGE 19 Acid Production in the Body Fixed Acid Other fixed acids can be produced in excessive quantities in certain pathophysiologic states, as a result of abnormal catabolic processes. – β-hydroxybutyric acid and acetoacetic acid, both ketoacids that are generated in untreated diabetes mellitus, – lactic acid, which may be generated during strenuous exercise or when the tissues are hypoxic. other fixed acids may be ingested, such as: – salicylic acid (from aspirin overdose), – formic acid (from methanol ingestion), – glycolic and oxalic acids (from ethylene glycol ingestion). Overproduction or ingestion of fixed acids causes metabolic acidosis. Regulation of Acid-Base Balance PAGE 20 Three primary systems regulate the H + concentration in the body fluids to prevent acidosis or alkalosis: (1) The buffer systems of the body fluids – react within seconds to minimize these changes. – they combine with an acid or a base to prevent excessive changes in H + concentration. – buffer systems do not eliminate H + from or add H + to the body but only keep them tied up until balance can be re-established. (2) The respiratory system, – acts within a few minutes to eliminate CO 2 and, therefore, H 2 CO 3 from the body. – the respiratory center, regulates the removal of CO 2 (and, therefore, H 2 CO 3 ) from the extracellular fluid Regulation of Acid-Base Balance PAGE 21 These first two lines of defense keep the H + concentration from changing too much until the more slowly responding third line of defense: (3) The kidneys, – can eliminate the excess acid or base from the body – are relatively slow to respond compared with the other defenses, over a period of hours to several days, – they are by far the most powerful of the acid-base regulatory systems. Regulation of Acid-Base Balance PAGE 22 Overview of acid-base balance. The lungs and kidneys work together to maintain acid-base balance. The lungs excrete CO 2 (volatile acid), and the kidneys excrete acid (renal net acid excretion [RNAE]) equal to net endogenous acid production (NEAP), which reflects dietary intake, cellular metabolism, and loss of acid and alkali (e.g., HCO 3 − loss in feces) from the body. Berne and Levy Physiology, 37, Copyright © 2018 by Elsevier, Inc. All rights reserved Fluid Buffers PAGE 23 Principles of Buffering A buffer is any substance that can reversibly bind H+ (a mixture of a weak acid and its conjugate base or a weak base and its conjugate acid). The general form of the buffering reaction is Buffer+ H+ ⇄ HBuffer The two forms of the buffer are called the buffer pair. In Brønsted-Lowry nomenclature, – for a weak acid, the acid form is called HA and is defined as the H + donor. the base form is called A − and is defined as the H + acceptor. Fluid Buffers PAGE 24 Principles of Buffering A buffered solution resists a change in pH. An aqueous solution consisting of a mixture of a weak acid and its conjugate base in the form of a salt, therefore resist changes in [H+]. Thus H + can be added to or removed from a buffered solution, but the pH of that solution will change only minimally. – For example, when H + is added to a buffered solution containing a weak acid, it combines with the A − form of the buffer and converts it to the HA (the acid form ). – Conversely, when H + is removed from a buffered solution (or OH − is added), H + is released from the HA form of the buffer, converting it to the A − (the base form). Fluid Buffers PAGE 25 Principles of Buffering The body fluids contain a large variety of buffers, which constitute an important first defense against changes in pH. This buffering capacity was demonstrated by two experiments: by injecting 150 mEq of H + (as hydrochloric acid, HCl) into a dog whose total body water was 11.4 L → caused the blood pH to decrease from 7.44 to 7.14—the dog was acidemic but alive. by addition of 150 mEq of H + to 11.4 L of distilled water → caused the pH to drop abruptly to 1.84, a value that would have been instantly fatal to the dog. Conclusion: – the dog's body fluids contained buffers protected its pH from the addition of large amounts of H +. The added H + combined with the A − form of these buffers, and a strong acid was converted to a weak acid. – the change in the dog's body fluid pH was minimized, although not totally prevented. – the distilled water contained no buffers and had no such protective mechanisms.  Fluid Buffers PAGE 26 Henderson-Hasselbalch Equation It is a formula used to calculate the pH of a buffered solution. K1 HA ⇄ H + + A K2 H2CO 3 ⇄ H + + HCO 3 − The forward reaction, the dissociation of HA into H + and A − , is characterized by a rate constant, K 1 , The reverse reaction is characterized by a rate constant, K 2. − When the rates of the forward and reverse reactions are exactly equal, there is a state of chemical equilibrium, in which there is no further net change in the concentration of HA or A− The ratio of rate constants can be combined into a single constant, K, called the equilibrium constant, as follows: K=[H+][A−]/[HA] Fluid Buffers PAGE 27 Henderson-Hasselbalch Equation Thus the final form of the Henderson-Hasselbalch equation is as follows: pH=pK + log 10 [A−]/[HA] where – pH=−log10[H+](pH units) – pK=−log10K(pH units) – [A−]=Concentration of base form of buffer(mEq/L) – [HA]=Concentration of acid form of buffer(mEq/L) Therefore the pH of a buffered solution can be calculated with the following information: the pK of the buffer, the concentration of the base form of the buffer [A − ], the concentration of the acid form of the buffer [HA]. Fluid Buffers PAGE 28 Principles of Buffering pK is a characteristic value for a buffer pair. Factors that may determine this value: – K= K1/K2 pK=−log10K Therefore, – strong acids such as HCl are more dissociated into H + and A − , and they have high equilibrium constants (K) and low pKs (because pK is minus log 10 of the equilibrium constant). – weak acids such as H 2 CO 3 are less dissociated and have low equilibrium constants and high pKs. Fluid Buffers PAGE 29 Principles of Buffering Titration curves are graphic representations of the Henderson-Hasselbalch equation - ex: the titration curve of a hypothetical weak acid (HA) and its conjugate base (A−) in solution. As H + is added or removed, the pH of the solution is measured. The pK of this hypothetical buffer is 6.5. At low (acidic) pH, the buffer exists primarily in the HA form. At high (alkaline) pH, the buffer exists primarily in the A − form. When the pH equals the pK, there are equal concentrations of HA and A − - half of the buffer is in the HA form and half in the A − form. Costanzo, Linda Physiology, Chapter 7 Copyright © 2018 by Elsevier, Inc. All rights reserved. Fluid Buffers PAGE 30 Principles of Buffering A striking feature of the titration curve is its sigmoidal shape. In the linear portion of the curve, only small changes in pH occur when H + is added or removed; the most effective buffering occurs in this range. Outside the effective buffering range, pH changes drastically when small amounts of H + are added or removed. – For this buffer, when the pH is lower than 5.5, the addition of H + causes a large decrease in pH; when the pH is higher than 7.5, the removal of H + causes a large increase in pH.  Fluid Buffers PAGE 31 Extracellular Fluid Buffers The major buffers of the ECF are bicarbonate and phosphate. For bicarbonate, the A − form is HCO 3 − and the HA form is CO 2 (in equilibrium with H2CO 3). When acid is added, it is buffered by HCO 3 − , which is then converted into dissolved CO 2 , decreasing the ratio of HCO 3 −/CO 2 and decreasing the pH. When base is added to the system, part of the dissolved CO 2 is converted into HCO 3 − , causing an increase in the ratio of HCO 3 − /CO 2 and increasing the pH. Costanzo, Linda Physiology, Chapter 7 Copyright © 2018 by Elsevier, Inc. All rights reserved. Fluid Buffers PAGE 32 Extracellular Fluid Buffers Extracellular Fluid Buffers HCO 3 − /CO 2 Buffer The most important extracellular buffer is HCO 3 − /CO 2. It is utilized as the first line of defense when H + is gained or lost from the body. The following characteristics account for the preeminence of HCO 3 − /CO 2 as an ECF buffer: (1) The concentration of the A − form, HCO 3 − , is high at 24 mEq/L. (2) The pK of the HCO 3 − /CO 2 buffer is 6.1, which is fairly close to the pH of ECF. (3) CO 2 , the acid form of the buffer, is volatile and can be expired by the lungs. Fluid Buffers PAGE 33 Extracellular Fluid Buffers HCO 3 − /CO 2 Expressing the anterior example (infusion of HCl to the dog) try to assume that ECF is a simple solution of NaHCO 3 , ignoring its other constituents. When HCl is added to ECF, H + combines with some of the HCO 3 − to form H 2 CO 3. Thus a strong acid (HCl) is converted to a weak acid (H 2 CO 3 ). H 2CO 3 then dissociates into CO 2 and H 2O, both of which are expired by the lungs. The pH of the dog's blood decreases, but not as dramatically as if no buffer were available: Fluid Buffers PAGE 34 Extracellular Fluid Buffers HCO 3 − /CO 2 This reaction is slow, and exceedingly small amounts of H 2CO 3 are formed unless the enzyme carbonic anhydrase is present. This enzyme is especially abundant: – in the walls of the lung alveoli, where CO 2 is released; – in the epithelial cells of the renal tubules, where CO 2 reacts with H 2 O to form H 2 CO 3. The HCO 3 − concentration is regulated mainly by the kidneys By increasing the rate of respiration, the lungs remove CO 2 from the plasma, and by decreasing respiration, the lungs elevate P co 2.  pH of Body Fluids PAGE 35 Extracellular Fluid Buffers HCO 3 − /CO 2 An acid-base map shows the relationships between Pco2 , HCO 3 − concentration, and pH The lines radiating from the origin on the map - mean same H + concentration or same pH; each line gives all of the combinations of Pco 2 and HCO 3 − that yield the same value of pH. The ellipse in the center shows the normal values for arterial blood. Any point on the graph can be calculated by substituting the appropriate values into the Henderson-Hasselbalch equation. Costanzo, Linda Physiology, Chapter 7 Copyright © 2018 by Elsevier, Inc. All rights reserved. Fluid Buffers PAGE 36 Extracellular Fluid Buffers HPO 4 −2 /H 2 PO 4 − Inorganic phosphate also serves as a buffer. Its titration curve can be compared with that for HCO 3 −. Recall that the pK for HCO 3 − /CO 2 is 6.1, with the linear portion of the titration curve extending from pH 5.1 to 7.1; technically, the linear portion is outside the buffering range for a pH of 7.4. On the other hand, the pK of the HPO 4 −2 /H 2 PO 4 − buffer is 6.8, with the linear portion of its curve extending from pH 5.8 to 7.8. It seems that inorganic phosphate would be a more important physiologic buffer than HCO 3 , because its effective buffering range is closer to 7.4, the pH of blood. Costanzo, Linda Physiology, Chapter 7 Copyright © 2018 by Elsevier, Inc. All rights reserved. Fluid Buffers PAGE 37 Extracellular Fluid Buffers However, two features of the HCO 3 − /CO 2 buffer make it the more effective buffer in ECF: (1) HCO 3 − is in much higher concentration (24 mmol/L) than phosphate (1–2 mmol/L). (2) The acid form of the HCO 3 − /CO 2 buffer is CO 2 , which is volatile and can be expired by the lungs. Fluid Buffers PAGE 38 Extracellular Fluid Buffers HPO 4 −2 /H 2 PO 4 − In contrast to its minor role as an extracellular buffer, the phosphate buffer is especially important in the tubular fluids of the kidneys for two reasons: (1) phosphate usually becomes greatly concentrated in the tubules, thereby increasing the buffering power of the phosphate system, (2) the tubular fluid usually has a considerably lower pH than the extracellular fluid does. Fluid Buffers PAGE 39 Extracellular Fluid Buffers – plasma proteins Plasma proteins also buffer H + In an acid-base disturbance negatively charged groups on plasma proteins (e.g., albumin) can bind either H + or Ca 2+. (40% of total Ca 2+ is protein-bounded) of Ca 2+. A relationship exists between plasma proteins, H + , and calcium (Ca 2+ ), which results in changes in ionized Ca 2+ concentration when there is an acid-base disturbance. In acidemia, there is an excess of H + in blood. Because more H + is bound to plasma proteins, less Ca 2+ is bound, producing an increase in free Ca 2+ concentration. In alkalemia, there is a deficit of H + in blood. Because less H + is bound to plasma proteins, more Ca 2+ is bound, producing a decrease in free Ca 2+ concentration (hypocalcemia). Fluid Buffers PAGE 40 Intracellular Fluid Buffers There are vast quantities of intracellular buffers, which include organic phosphates and proteins. A rapid equilibrium occurs in the red blood cells; in this case CO 2, can rapidly diffuse through all the cell membranes; hemoglobin (Hb) buffers, as follows: H + + Hb ⇄ HHb However, except for the red blood cells, the slowness with which H + and HCO 3 − move through the cell membranes often delays for several hours the maximum ability of the intracellular proteins to buffer extracellular acid-base abnormalities.  Fluid Buffers PAGE 41 Intracellular Fluid Buffers Organic Phosphates The phosphate buffer system is also important in buffering intracellular fluid: – because the concentration of phosphate in this fluid is many times that in the extracellular fluid – the pH of ICF is lower than that of ECF. Organic phosphates in ICF include (ATP), (ADP), (AMP), glucose-1-phosphate, 2,3-diphosphoglycerate (2,3-DPG) - H + is buffered by the phosphate moiety of these organic molecules. The pKs for these organic phosphates range from 6.0 to 7.5, ideal for effective physiologic buffering. Renal Mechanisms in Acid-Base Balance PAGE 42 The kidneys control acid-base balance by excreting either: – acidic urine – basic urine They play two major roles in the maintenance of normal acid-base balance: – reabsorption of HCO 3 − - so that this important extracellular buffer is not excreted in urine. – excretion of fixed H + that is produced from protein and phospholipid catabolism: (1) excretion of H + as titratable acid (i.e., buffered by urinary phosphate) (2) excretion of H + as NH 4 +. Excretion of H + by either mechanism is accompanied by reabsorption and synthesis of new HCO − Renal Mechanisms in Acid-Base Balance PAGE 43 The overall mechanism by which the kidneys excrete acidic or basic urine : Large numbers of HCO 3 − are filtered continuously into the tubules, and if they are excreted into the urine, this removes base from the blood→ decrease pH Large numbers of H + are also secreted into the tubular lumen by the tubular epithelial cells, thus removing acid from the blood→ increase pH. If more H + is secreted than HCO 3 − is filtered → a net loss of acid from the extracellular fluid. Conversely, if more HCO 3 − is filtered than H + is secreted → a net loss of base. Renal Mechanisms in Acid-Base Balance PAGE 44 Secretion of H + and Reabsorption of HCO 3 − by the Renal Tubules Each day the body produces about 80 mEq of nonvolatile acids, mainly from the metabolism of proteins - nonvolatile because they are not H 2 CO 3 - cannot be excreted by the lungs. – The primary mechanism for removal is renal excretion. Each day the kidneys filter about increased amounts HCO 3 − (180 L/day × 24 mEq/L) - normal conditions, almost all this is reabsorbed from the tubules, – prevent the loss of bicarbonate in the urine - more important than the excretion of nonvolatile acids Renal Mechanisms in Acid-Base Balance PAGE 45 The HCO 3 − balance is realised by three distinct renal mechanisms: 1. Reabsorption of filtered HCO 3 − (the Na/H antiporter) 2. Generation of new HCO 3 − by titratable acid (TA) excretion 3. Formation of HCO 3 − − from generation of NH 4 + Renal Mechanisms in Acid-Base Balance PAGE 46 Renal Mechanisms in Acid-Base Balance Reabsorption of bicarbonate in different segments of the renal tubule. The percentages of the filtered load of HCO 3 − absorbed by the various tubular segments are shown, as well as the number of milliequivalents reabsorbed per day under normal conditions. Guyton and Hall Textbook of Medical Physiology, Chapter 31, Copyright © 2016 by Elsevier, Inc. All rights reserved. Regulation of Acid-Base Balance PAGE 47 Renal Mechanisms in Acid-Base Balance Reabsorption of HCO 3 − Bicarbonate ions do not readily permeate the luminal membranes of the renal tubular cells; – HCO 3 − is reabsorbed by a special process in which it first combines with H + to form H 2CO 3 , which eventually becomes CO 2 and H 2O. – The CO 2 can move easily across the tubular membrane; therefore, it instantly diffuses into the tubular cell, – Inside the cell it recombines with H 2O, under the influence of carbonic anhydrase, to generate a new H 2CO 3 molecule. – This H 2 CO 3 in turn dissociates to form HCO 3 − and H + ; the HCO 3 − then diffuses through the basolateral membrane into the interstitial fluid and is taken up into the peritubular capillary blood. Regulation of Acid-Base Balance PAGE 48 Renal Mechanisms in Acid-Base Balance H + is Secreted by Secondary Active Transport in the Early Tubular Segments Cellular mechanisms for (1) active secretion of H + into the renal tubule; (2) tubular reabsorption of HCO 3 − by combination with H + to form carbonic acid, which dissociates to form carbon dioxide and water; (3) sodium ion reabsorption in exchange for H + secreted. Guyton and Hall Textbook of Medical Physiology, Chapter 31, Copyright © 2016 by Elsevier, Inc. All rights reserved. Regulation of Acid-Base Balance PAGE 49 Renal Mechanisms in Acid-Base Balance Generation of New HCO 3 − by the Ammonia Buffer System Glutamine (metabolism of amino acids in the liver→ transported into the epithelial cells→ each molecule of glutamine is metabolized → form two NH 4 + and two HCO 3 −. NH 4 + is secreted into the tubular lumen in exchange for sodium. The HCO 3 − is transported across the basolateral membrane, along with the reabsorbed Na +. Guyton and Hall Textbook of Medical Physiology, Chapter 31, Copyright © 2016 by Elsevier, Inc. All rights reserved. Regulation of Acid-Base Balance PAGE 50 Renal Mechanisms in Acid-Base Balance H + is Secreted by Secondary Active Transport in Kidney Tubules The epithelial cells of the proximal tubule, the thick segment of the ascending loop of Henle, and the early distal tubule all secrete H + into the tubular fluid by sodium-hydrogen counter-transport. About 95 percent of the bicarbonate is reabsorbed in this manner, requiring about 4000 mEq of H + to be secreted each day by the tubules Regulation of Acid-Base Balance PAGE 51 Renal Mechanisms in Acid-Base Balance HCO 3 − Is “Titrated” Against H + in the Tubules H + is excreted in combination with phosphate buffer and the mechanism by which new HCO3 − is added to the blood. Buffering of secreted H + by filtered phosphate (NaHPO 4 ). Note that a new HCO 3 − is returned to the blood for each NaHPO 4 that reacts with a secreted H +. Guyton and Hall Textbook of Medical Physiology, Chapter 31, Copyright © 2016 by Elsevier, Inc. All rights reserved.  PAGE 52 Boron and Boulapaep Concise Medical Physiology, Chapter 39, Copyright © 2021 by Elsevier, Inc. All rights reserved. Acid-Base Disorders PAGE 53 Disturbances of acid-base balance are among the most common conditions in all of clinical medicine. Acid-base disorders are characterized by an abnormal concentration of H + in blood, reflected as abnormal pH. Acidemia is an increase in H + concentration in blood (decrease in pH) and is caused by a pathophysiologic process called acidosis. Alkalemia, on the other hand, is a decrease in H + concentration in blood (increase in pH) and is caused by a pathophysiologic process called alkalosis. Acid-Base Disorders PAGE 54 Disturbances of acid-base balance are described as either metabolic or respiratory, depending on whether the primary disturbance is in HCO 3 − or CO 2. Acid-Base Disorders PAGE 55 Metabolic acid-base disturbances are primary disorders involving HCO 3 −. Metabolic acidosis is caused by a decrease in HCO 3 − concentration. – Low pH. – This disorder is caused by gain of fixed H + in the body (through overproduction of fixed H + , ingestion of fixed H + , or decreased excretion of fixed H + ) or loss of HCO 3 −. Metabolic alkalosis is caused by an increase in HCO 3 − concentration. – High pH. – This disorder is caused by loss of fixed H + from the body or gain of HCO 3 −. Acid-Base Disorders PAGE 56 Respiratory acid-base disturbances are primary disorders of CO2 (i.e., disorders of respiration). – Respiratory acidosis is caused by hypoventilation, which results in CO 2 retention, increased Pco 2 , decreased pH. – Respiratory alkalosis is caused by hyperventilation, which results in CO 2 loss, decreased Pco2, increased pH. Acid-Base Disorders PAGE 57 When there is an acid-base disturbance, several mechanisms are utilized in an attempt to keep the blood pH in the normal range. The first line of defense is buffering in ECF and ICF. In addition to buffering, two types of compensatory responses attempt to normalize the pH: respiratory compensation and renal compensation Acid-Base Disorders PAGE 58 If the acid-base disturbance is metabolic (i.e., disturbance of HCO 3 − ), then the compensatory response is respiratory to adjust the P co 2 ; If the acid-base disturbance is respiratory (i.e., disturbance of CO 2 ), then the compensatory response is renal (or metabolic) to adjust the HCO 3 − concentration. Acid-Base Disorders PAGE 59 Guyton and Hall Textbook of Medical Physiology, Chapter 31, Copyright © 2016 by Elsevier, Inc. All rights reserved. References PAGE 60 1. Linda S. Costanzo, Physiology Sixth Edition 2. Guyton and Hall, Textbook ot Medical Physiology thirteen edition, John E. Hall 3. Boron, Walter F., MD, PhD; Boulpaep, Emile L., MD, Medical Physiology, Third Edition, 4. Berne and Levy Physiology, Koeppen, Bruce M., MD, PhD; Stanton, Bruce A., PhD. Published January 1, 2018. 5. Netter's Essential Physiology, Mulroney, Susan E., PhD; Myers, Adam K., PhD. Published January 1, 2016 6. https://www.youtube.com/watch?v=oJ5jc0nH3h0