Acids and Bases their definitions and meanings

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1 Acids and Bases their definitions and meanings Molecules containing hydrogen atoms that can release hydrogen ions in solutions are referred to as acids. (HCl H + Cl ) (H 2 CO 3 H + HCO 3 ) A base is an ion or a molecule that can accept a hydrogen ion. (HPO 4 2 is base because it can accept hydrogen ion to form H 2 PO 4 ) The proteins in the body also as bases because some of the amino acids that make up proteins have negative charges that readily accept hydrogen ions. Alkalosis refers to excess removal of hydrogen ions from the body fluids. Acidosis refers to the excess addition of hydrogen ions in the body fluids. A strong acid is one that rapidly dissociates and releases large amounts of H + in solution (HCl) A week acid have less tendency to dissociate its ions and, therefore release H + (H 2 CO 3 )

2 Control of AcidBase Balance 1. There must be a balance between the production of H + and the net removal of H + from the body. 2. Precise H + regulation is essential because the activities of almost all enzyme systems in the body are influenced by H + concentration. 3. Na + = 142 mmol/l, H + = mmol/l (40 nmol/l) 4. ph = log [H + ] = log[ ] = 7.4 (The lower limit of ph at which a person can live more than a few hours is about 6.8 and the upper limit is about 8.0) 5. There are three primary systems that regulate the H + concentration in body fluids to prevent acidosis: A/ Chemical acidbase buffer systems of the body fluids (seconds) B/ Lungs (few minutes) C/ Kidneys (hours to days)

3 Metabolic Sources of Acids and Bases A. Reactions producing CO 2 (Merely a Potential Acid) 1. Complete oxidation of neutral carbohydrated and fat CO 2 + H 2 O 2. Oxidation of most neutral amino acids Urea + CO 2 + H 2 O B. Reactions producing nonvolatile acids 1. Oxidation of sulfurcontaining amino acids Urea + CO 2 + H 2 O + H 2 SO 4 2H + + SO 42 (examples: methionine, cysteine) 2. Metabolism of phosphorouscontaining compounds H 3 PO 4 H + + H 2 PO Oxidation of cationic amino acids Urea + CO 2 + H 2 O + H + (examples: lysine+, arginin+) 4. Production of nonmetabolizable organic acids HA H + + A (examples: uric acid, oxalic acid) 5. Incomplete oxidation of carohydrate and fat HA H + + A (examples: lactic acid, ketoacidosis) C. Reactions producing nonvaletile bases 1. Oxidation of anionic amino acids Urea + CO 2 + H 2 O + HCO 3 (examples: glutamate, aspartate ) 2. Oxidation of organic anions CO 2 + H 2 O + HCO 3 (examples: lactate, acetate )

4 Buffering of Hydrogen Ions in the Body Fluids Daily production of H + = 80 mmol, Body fluid concentration = mmol/l Buffer + H + H Buffer In this example, a free H + combines with the buffer to form a weak acid (H Buffer) Bicarbonate Buffer System H + + HCO 3 H 2 CO 3 CO 2 + H 2 O From these reactions, one can see that the hydrogen ions from the strong acid react with HCO 3 to form the very weak acid (H 2 CO 3 ), which in turn forms CO 2 and H 2 O. The excess of CO 2 stimulates respiration

5 CO 2 + H 2 O H 2 CO 3 HCO 3 + H NaOH Na The weak base NaHCO 3 replaces the strong base NaOH. At the same time the concentration of H 2 CO 3 decreases (because it reacts with NaOH), causing more CO 2 to combine with H 2 O, in order to replace the H 2 CO 3. The net result is a tendency for the CO 2 levels in the blood to decrease, but it is prevented by the decreased ventilation. The rise in blood HCO 3 is compensated by increased renal excretion of HCO3.

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7 Phosphate Buffer System It plays a major role in buffering renal tubular fluid and intracellular fluid HPO H + H 2 PO 4 Proteins: Important Intracellular Buffer H + + Hemoblobin HHemoglobin Approximately 60 to 70 percent of the total chemical buffering of the body fluids is inside the cells, and most of this results from the intracellular proteins.

8 H 2 CO 3 H + + HCO 3 For any acid, the concentration of the acid relative to its dissociated ions is defined by the dissociation constant K. K = H + x HCO 3 H 2 CO 3 This equation indicates that in an H 2 CO 3 solution, the amount of free H + is equal to: H + = K x H 2 CO 3 HCO 3 The concentration of undissociated H 2 CO 3 cannot be measured in solution because is rapidly dissociated. However, the amount of CO 2 dissolved in the blood is directly proportional to the amount of undissociated H 2 CO 3. H + = K x CO 2 HCO 3 The dissociation constant (K) is only about 1/400 of K because the proportionality ratio between H 2 CO 3 and CO 2 is 1:400. In addition, we don t measure the total amount of CO 2 in blood. Fortunately, the amount of CO 2 in the blood is linear function of pco 2 multiplied by the solubility coefficient.

9 The solubility coefficient for CO 2 is 0.03 mmol/mmhg. This means that 0.03 millimole of H 2 CO 3 is present in the blood for each mmhg pco 2 measured (0.03 x pco 2 ) H + = K x HCO 3 It is customary to express H + concentration in ph units rather than in actual concentration. ph = log H +. The dissociation constant can be expressed in a simila manner, pk = log K. log H + = log pk log (0.03 x pco 2 ) HCO 3 ph = pk log (0.03 x pco 2 ) HCO 3 Rather than work with a negative logarithm, we can change the sign of the logarithm and invert the numerator and denominator. ph = pk + log HCO x pco 2

10 Isohydric Principle: all buffers in a common solution are in equilibrium with same H + concentration HA 1 HA 2 HA 3 H + = K 1 x = K 2 x = K 3 x A 1 A 2 A 3 All buffer systems work together, because H + is common to the reactions of all these systems. Any condition that changes the balance of one of the buffer system also changes the balance of al the others.

11 HendersonHasselbalch Equation: ph = log HCO x pco 2 ph = Constant + Kidney Lung

12 HendersonHasselbalch Equation: ph = log HCO x pco 2 1. Increase in bicarbonate ion concentration causes the ph to rise. 2. Increase in pco 2 causes the ph to decrease. 1. Bicarbonate concentration is regulated mainly by the kidneys. 2. pco 2 concentration is regulated by the rate of respiration. 1. When disturbances of acidbase balance results from a primary changes in extracellular fluid bicarbonate concentrations are referred to as metabolic acidbase disorders. 2. When disturbances of acid base balance results from a primary changes in pco 2 are referred as respiratory acidbase disorders.

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14 Respiratory regulation of acidbase balance Respiratory regulation of acidbase balance is a physiological type of buffer system, because it acts rapidly and keeps the hydrogen ion concentration from changing too much until the much slowly responding kidneys can eliminate the imbalance. H + Alveolar ventilation pco 2 If hydrogen ion concentration is suddenly increased by adding acid to extracellular fluid and ph falls from 7.4 to 7.0, the respiratory systém can retrun the ph to a value of 7.2 to 7.3 within 3 to 12 minutes. Respiratory system has an efectivness between 50 to 75 per cent.

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17 The kidneys regulate extracellular fluid H + concentrations thought three fundamental mechanisms: 1. Reabsorption of filtered HCO 3 2. Secretion of H + 3. Production of new HCO 3 Ad L/day x 24 mmol/l = 4320 mmol of HCO 3

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19 Proximal tubule, thick ascending loop of Henle, early distal tubule

20 Thus, each time a hydrogen ion is formed in the tubular epithelial cells, a bicarbonate ion is also formed and released back into the blood. The net effect of these reactions is a reabsorption of bicarbonate, although the bicarbonate ions that actually enter the extracellular fluid are not the same. The transport of HCO 3 accros the basolateral membrane is facilitated by: 1. Na + HCO 3 cotransporter 2. Cl HCO 3 exchange

21 Although the secretion of hydrogen ions in the late distal tubule and collecting duct accounts for only percent of the total hydrogen secreted, this mechanism is important in forming a maximally acidic urine. In the proximal tubules, hydrogen ion concentration can increase only about threefold (compared to the filtered load), in the collecting tubule the hydrogen concentration can be increased as 900fold.

22 Late distal tubule and collecting tubules (intercalated cells)

23 Phosphate and Ammonia Buffers Minimal urine ph is 4.5, corresponding to an H + concentration 0.03 mmol/l. In order, to excrete the 80 mmol of nonvolatile acid formed each day, about 2667 liters of urine would have to be excreted if the H + remained free in solution. 500 mmol/day of H + must be sometimes excreted.

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25 Therefore, whenever an H + secreted into the tubular lumen combines with a buffer other than, HCO 3 the net effect is addition of a new HCO 3 to the blood. A second buffer system in the tubular fluid that is even more important quantitatively than the phosphate buffer system is composed of ammonia (NH 3 ) and the ammonium ion (NH 4+ ).

26 Collecting duct

27 Proximal tubule, thick ascending limb of the loop of Henle, distal tubule

28 Glutamine Glutarate αketoglutarate 2 2NH HCO 3 2NH 4 + LIVER (Metabolism) KIDNEY (Excretion) 2NH CO 2 Urea + H H + Save 2HCO 3 by excretion of 2NH 4 + Loss of 2HCO 3 by buffering of 2H +

29 Under conditions of chronic acidosis, the rate of NH 4 + excretion can increase to as much 500 mmol/day. Therefore, with chronic acidosis, the dominant mechanism by which acid is eliminated from the body is excretion of NH 4+.

30 Quantifying Renal AcidBase Excretion Net acid excretion = NH + 4 excretion + Urinary titratable acid bicarbonate excretion Titratable acid represents the nonbicarbonate, nonnh 4 + buffer excreted in the urine (phosphate and other organic buffers) The most important stimuli for increasing H + secretion by the tubules are: 1. An increase in pco 2 of extracellular fluid. 2. An in H + concentration in extracellular fluid.

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34 HCO 3 from 24 to 14 mm ph from 7.4 to 7.17 ΔH + = ( ) = mm = mm Even though we have added 10 mm, H + increased by only 28 nanomols. Therefore, open system has neutralized mmol of the added 10 mmol H +. In this system only depletion of HCO 3 limits the neutralization of H +. The buildup of CO 2 is not a limiting factor because the atmosphere is an infinite sink for newly produced CO 2. HendersonHasselbalch Equation: ph = log HCO x pco 2

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37 [HCO 3 ] (mm) ph P CO2 = 20 mmhg P CO2 = 40 mmhg P CO2 = 80 mm Hg mm 12 mm 24 mm

38 Characteristics of Primary AcidBase Disturbances ph H + pco 2 HCO 3 Normal nmol/l 40 mmhg 24 mmol/l

39 In respiratory acidosis is the initial cause an increase in pco 2. The compensatory response is an increase in plasma HCO 3, caused by addition of new bicarbonate to the extracellular fluid by the kidney. The rise in HCO 3 helps to offset the increase in pco 2, thereby returning the plasma ph toward normal.

40 Isohydric hypercapnia (i.e. the same ph at a higher P CO2 )

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42 Characteristics of Primary AcidBase Disturbances ph H + pco 2 HCO 3 Normal nmol/l 40 mmhg 24 mmol/l Respiratory Acidosis

43 In respiratory alkalosis the initial cause is a decrease in plasma pco 2 caused by hyperventilation. The compensatory response is a reduction in plasma HCO 3 caused by increased renal excretion of bicarbonate.

44 Characteristics of Primary AcidBase Disturbances ph H + pco 2 HCO 3 Normal nmol/l 40 mmhg 24 mmol/l Respiratory Acidosis Respiratory Alkalosis

45 In metabolic acidosis, is also a decrease in ph (as in respiratory acidosis) and rise in extracellular fluid hydrogen ion concentration, however in this case the primary abnormality is a decrease in plasma HCO 3. The primary compensations include increased ventilation, which reduces pco 2 and renal compensation, which by adding new bicarbonate to the extracellular fluid helps to minimize the fall in extracellular fluid HCO 3 concentration. 1. Diabetes mellitus 2. Diarrhea (loss of large amounts of sodium bicarbonate into the feces) 3. Vomiting of intestinal contents 4. Renal tubular acidosis (a/ impairment bicarbonate reabsorption, b/ inability to secrete H + ) 5. Chronic renal failure

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47 Characteristics of Primary AcidBase Disturbances ph H + pco 2 HCO 3 Normal nmol/l 40 mmhg 24 mmol/l Respiratory Acidosis Respiratory Alkalosis Metabolic Acidosis

48 In metabolic alkalosis the primary cause is a rise in HCO 3 in extracellular fluid. The primary compensation is decreased ventilation, which raises pco 2 and the secondary compensation is increased renal bicarbonate excretion. 1. Vomiting of gastric content 2. Administration of diuretics (except the carbonic anhydrase inhibitor) (diuretics cause increases in tubular fluid volume and delivery of sodium to distal parts of nephron. This in turn leads in increased reabsorption in these parts of the nephron. Reabsorption is coupled with H + secretion)

49 Characteristics of Primary AcidBase Disturbances ph H + pco 2 HCO 3 Normal nmol/l 40 mmhg 24 mmol/l Respiratory Acidosis Respiratory Alkalosis Metabolic Acidosis Metabolic Alkalosis

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53 Buffer Base (BB) equal the sum of all the conjugate bases in 1 L of arterial whole blood: 1. HCO 3 = 24 mmol/l 2. Protein = 15 mmol/l 3. Hemoglobin = 9 mmol/l BB = 48 mmol/l Base Excess (BE) = the observed BB minus the normal BB (BE = 2 to + 2 mmol/l) Anion gap (AG) = (Na + + K + ) (HCO 3 + Cl ) = (142 mmol/l + 4 mmol/l) (25 mmol/l+ 102 mmol/l) = 16 mmol/l Organic acids increase AG, because the lost of HCO 3 is not replaced by routinely unmeasured anions. (Metabolic acidosis) Respiratory acidosis does not increase AG, because excess H + is derived from H 2 CO 3 pool, not the noncarbonic acid pool. Osmolar gap refers to the disparity between the measured and calculated serum osmolarity. Increased osmolar gap provides a reasonable good screening procedure for toxins. Other causes of metabolic acidosis do not affect osmolar gap, since the metabolic acid simply replaces HCO 3 P oms = 2 [Na + ] + [BUN] + [glucose]