REGULATION OF ACID BASE BALANCE| Tutorial B-1 RS

Hydrogen Ion Concentration Is Precisely RegulatedPrecise H+ regulation is essential because the activities of almost all enzyme systems in the body are influenced by H+ concentration. Therefore, changes in hydrogen concentration alter virtually all cell and body functions. Normal concentration of H+ is 0.00004 mEq/L (3.5 million times smaller than concentration of sodium in extracellular fluid 142 mEq/L)

Acids and BasesTheir Definitions and Meanings Hydrogen ion: single free proton released from a hydrogen atom Acid: Molecules containing hydrogen atoms that can release hydrogen ions in solutions. Example: hydrochloric acid (HCl), which ionizes in water to form hydrogen ions (H+) and chloride ions (Cl) Base: Ion or a molecule that can accept an H+. Examples: HCO3 is a base because it can combine with H+ to form H2CO3. 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+ The terms base and alkali are often used synonymously. Alkali: molecule formed by the combination of one or more of the alkaline (sodium, potassium, lithium, and so etc) with a highly basic ion such as a hydroxyl ion (OH). The base portion of these molecules reacts quickly with H+ to remove it from solution Strong and Weak Acids and BasesStrong acid: one that rapidly dissociates and releases especially large amounts of H+ in solution. Example: HClWeak acids: have less tendency to dissociate their ions and release H+ with less vigor. Example: H2CO3Strong base: one that reacts rapidly and strongly with H+ and reacts with H+ to form water. Example: OH-Weak base: binds with H+ much more weakly than does OH. Example: HCO3-

Normal Value of Hydrogen Ion Concentration and pH of Body Fluids Blood H+ concentration: 0.00004 mEq/L (40 nEq/L). Normal variations: 3 to 5 nEq/L. Extreme conditions (without causing death): 10 nEq/L to as high as 160 nEq/LBecause concentration ion hydrogen are in very small numbers, it is customary to express H+ concentration on a logarithm scale, using pH units:

Therefore, the normal pH = log [0.00000004] = 7.4 pH normal of arterial blood: 7.4 pH normal venous blood and interstitial fluid: 7.35 because of the extra amounts of carbon dioxide (CO2) acidosis < 7.4 (tolerated 0.05) < alkalosis 6.8 < person can live more than a few hours < 8.0 pH in every tissue and fluid is different

Defenses Against Changes in Hydrogen Ion Concentration: Buffers, Lungs, and KidneysThere are 3 primary systems that regulate the H+ concentration in body fluids to prevent acidosis/alkalosis: (1) Chemical acid-base buffer systems of the body fluids, which immediately combine with acid or base to prevent excessive changes in H+ concentration react within a fraction of a second to minimize these changes. Buffer systems do not eliminate H+ from or add them to the body but only keep them tied up until balance can be reestablished.(2) Respiratory center, which regulates the removal of CO2 (indirectly H2CO3) from the extracellular fluid. Second line of defense, acts within a few minutes(3) Kidneys, which can excrete either acid or alkaline urine. Relatively slow to respond compared with others, (over a period of hours to several days) but the most powerful of the acid-base regulatory systemsOBJECTIVE: keep pH to balance either added by strong acid or baseBuffering of Hydrogen Ions in the Body Fluids Buffer: any substance that can reversibly bind H+ General form: When the H+ concentration increases, the reaction is forced to the right and more H+ binds to the buffer. . Conversely, when the H+ concentration decreases, the reaction shifts toward the left, and H+ is released from the buffer IMPORTANCE: low concentration of H+ in the body fluids and the relatively large amounts of acids produced by the body each day

Bicarbonate Buffer System Consists of a water solution that contains two ingredients: (1) a weak acid, H2CO3 and (2) a bicarbonate salt, such as NaHCO3. H2CO3 is formed in the body by the reaction: Exceedingly small amounts of H2CO3 are formed unless the enzyme carbonic anhydrase is present Carbonic anhydrase abundant in the walls of the lung alveoli (CO2 is released) and in the epithelial cells of the renal tubules (CO2 reacts with H2O to form H2CO3) H2CO3 ionizes weakly: Bicarbonate salt (commonly NaHCO3) in the extracellular fluid NaHCO3 ionizes almost completely:

System formed: Because of the weak dissociation of H2CO3, the H+concentration is extremely small When a strong acid (ex: HCl) is added to the bicarbonate buffer solution, H+ released from the acid is increased (buffered by HCO3-): result more H2CO3 formed CO2 and H2O. Excess CO2 greatly stimulates respiration, which eliminates the CO2 from the extracellular fluid When a strong base (ex: NaOH) is added to the bicarbonate buffer solution: OH (from the NaOH) + H2CO3 form additional HCO3- & concentration of H2CO3 (because reacts with NaOH), causing CO2 to combine with H2O to replace the H2CO3

CO2 levels in the blood inhibits respiration and the rate of CO2 expiration HCO3 that occurs is compensated for by renal excretion of HCO3-

Acid base disturbancesMetabolic acid-base disorders: a primary change in extracellular fluid bicarbonate concentration a primary decrease in bicarbonate concentration: metabolic acidosis a primary increase in bicarbonate concentration: metabolic alkalosis increase in PCO2: respiratory acidosis decrease in PCO2: respiratory alkalosis

Phosphate Buffer System Not important as an extracellular fluid buffer, but plays a major role in buffering renal tubular fluid and intracellular fluids. The main elements of the phosphate buffer system are H2PO4 and HPO42- When a strong acid (ex: HCl) is added: (hydrogen accepted by the base HPO42- and converted to H2PO4) Strong acid (HCl) is replaced by an additional amount of a weak acid, NaH2PO4 & pH minimized When a strong base (ex: NaOH) is added to the buffer system: (the OH is buffered by the H2PO4 to form additional amounts of HPO42- + H2O a strong base (NaOH) is traded for a weak base, NaH2PO4, causing only a slight increase in pH

The phosphate buffer system has a pK of 6.8, which is not far from the normal pH of 7.4 in the body fluids allows the system to operate near its maximum buffering power. However, its concentration in the extracellular fluid is low, only about 8 per cent of the concentration of the bicarbonate buffer. The phosphate buffer is especially important in the tubular fluids of the kidneys, for: (1) Phosphate usually becomes greatly concentrated in tubules, thereby buffering power of phosphate system(2) The tubular fluid usually has a considerably lower pH than the extracellular fluid does, bringing the operating range of the buffer closer to the pK (6.8) of the system.

Proteins: Important Intracellular Buffers Approximately 60 to 70 per cent of the total chemical buffering of the body fluids is inside the cells, and most of this results from the intracellular proteins This diffusion of the elements of the bicarbonate buffer system causes the pH in intracellular fluid to change when there are changes in extracellular pH. For this reason, the buffer systems within the cells help prevent changes in the pH of extracellular fluid but may take several hours to become maximally effective. The slowness with which H+ and HCO3 move through the cell membranes often delays for several hours the maximum ability of the intracellular proteins to buffer extracellular acid-base abnormalities Example: Hemoglobin, the reaction:

Respiratory Regulation of Acid-Base Balance CO2 is formed continually in the body by intracellular metabolic processes diffuses from the cells into the interstitial fluids and blood flowing blood transports it to the lungs diffuses into the alveoli pulmonary ventilation transferred to the atmosphere Metabolic formation of CO2 pCO2 extracellular balancing: pulmonary ventilation The higher the alveolar ventilation, the lower pCO2, and conversely. Increasing alveolar ventilation to about 2x normal raises the pH of extracellular fluid by about 0.23 Conversely a decrease in alveolar ventilation to reduces the pH by 0.45 Feedback control of H+ concentration by the respiratory system: Respiratory control cannot return the H+ concentration all the way back to normal when a disturbance outside the respiratory system has altered pH. Ordinarily, the respiratory mechanism for controlling H+ concentration has an effectiveness between 50 and 75%, corresponding to a feedback gain of 1 to 3. That is, if the H+ concentration is suddenly by adding acid to the extracellular fluid and pH falls from 7.4 to 7.0, the respiratory system can return the pH to a value of about 7.2 to 7.3. This response occurs within 3 to 12 minutes. Respiratory regulation of acid-base balance is a physiologic type of buffer system because it acts rapidly and keeps the H+ concentration from changing too much until the slowly responding kidneys can eliminate the imbalance. In general, the overall buffering power of the respiratory system is one to two times as great as the buffering power of all other chemical buffers in the extracellular fluid combined. That is, one to two times as much acid or base can normally be buffered by this mechanism as by the chemical buffers. Impairment of lung function can cause respiratory acidosis. Ex: emphysema ability of lungs to eliminate CO2

Renal Control of Acid-Base Balance The kidneys control acid-base balance by excreting either an acidic or a basic urineMechanism: Large numbers of HCO3 are filtered continuously into the tubules, and if they are excreted into the urine, this removes base from the blood Large numbers of H+ are also secreted into the tubular lumen by the tubular epithelial cells, thus removing acid from the blood See net loss. If there is lot base gone pH decrease, if more acid gone pH increaseThus, the kidneys regulate extracellular fluid H+ concentration through three fundamental mechanisms: (1)secretion of H+, (2) reabsorption of filtered HCO3-, and (3) production of new HCO3-

130110110177|Gabriella Chafrina| 13/12/13