ACID-BASE PHYSIOLOGY 1 Dr. Ana-Maria Zagrean
Acid-base physiology consists in all the processes inside the body which keep the H + concentration within normal values, thus maintaining the proper balance between acids and bases Hydrogen ions (H + ) / protons have a major impact on biochemical reactions and on a variety of physiological processes that are critical for the homeostasis of individual cells and for the entire body. There are sophisticated and powerful systems to maintain [H + ] values within narrow and precise ranges in the blood plasma, intracellular fluid, and other body compartments. Outside these ranges, proteins are denatured, enzymes activity is compromised, nerve and cardiac function is altered. Acid base balance depends on: -water and ion balance (renal excretion of acid, reabsorption of bicarbonate) -blood gas homeostasis (pulmonary excretion of CO2) -buffers availability (HCO3-, HPO4, protein anions, carbonate, etc)
Maintenance of Acid-Base (AB) balance by 2 major mechanisms: 1. Buffer systems - composed of a weak acid and it s salt with a powerful base, which have two origins: plasmatic/extracellular fluid (ECF) and intracellular fluid (ICF) (mosly erythrocyte); they both fight against sudden shifts in AB balance (act in seconds) 2. Biological mechanisms - in which lungs (regulate AB in minutes) and kidneys play a major role (regulate AB balance in days)
ph and buffers An acid is any chemical substance (e.g., CH3COOH, NH4+) that can donate an H+. Hydrogen ions do not truly exist as free protons in aqueous solutions, but as protons surrounded by a shell of water molecules, forming H3O+ (hydronium ion). For practical purposes, we refer to protons as if they are free. A base/alkali is any chemical substance (e.g., CH3COO, NH3) that can accept an H+. ph values vary enormously among different intracellular and extracellular compartments: [H+] varies over a large range in biological solutions, from >100 mm in gastric secretions to <10 nm in pancreatic secretions. Blood acidity may be expressed by H+ concentration ([H+]) or ph: For [H+] = 10 7 M, the ph is 7.0. The higher the [H+], the lower the ph Normal ph = 7.4 (7.35 7.45)
Relationship between [H+] and ph Values A 10-fold change in [H+] corresponds to a ph shift of 1, whereas a 2-fold change in [H+] corresponds to a ph shift of ~0.3. Note that log 2 = 0.3
Because the ph of neutral water at 37 C is 6.81, most major body compartments are alkaline.
Acids and bases/alkali Acids are H+ donors. Bases are H+ acceptors. Acids and bases can be: Strong dissociate completely in aqueous solution (HCl, NaOH) Weak dissociate only partially in aqueous solution (lactic acid, carbonic acid)
Acid and alkali load ACID - Acid containing foods + production from metabolism ALKALI - Alkali containing foods + production from metabolism their buffering leads to extra acid load that must be buffered and then excreted
Acid load - fixed versus volatile acid FIXED = NON VOLATILE Dietary acids Daily production of acids= 50-100 meq of H+- under physiological conditions - from cell metabolism Fixed acids from catabolism of proteins (amioacids, uric acid, sulphuric acid, phosphoric acid); carbohydrates (pyruvic acid, succinic acid, lactic acid (anaerobiosis)); fats (fatty acids, ketoacids (diabetes/starvation) - acetoacetic acid, beta-hydroxybutyric acid ) VOLATILE ACID - CARBONIC ACID (H2CO3), in equilibrium with its dissolved gaseous component CO2 that can be excreted through ventilation
Carbonic acid Metabolism of fats and carbohydrates result in the production of 15-20 mol of CO 2 per day Before elimination by the lungs, most of the CO 2 is taken up by red blood cells, reacting with H 2 O to form carbonic acid as shown below: CO 2 + H 2 O H 2 CO 3 H + + HCO 3 - CA (carbonic anhidrase intracellular) HAMBURGER RBC bicarbonate synthesis REVERSED HAMBURGER RBC carbonic acid synthesis
Small changes in ph can have substantial physiological consequences Most biologically important molecules contain chemical groups that can either donate an H+ (e.g., R COOH R COO + H+) and thereby act as a weak acid, or accept an H+ (e.g., R NH2 + H+ R NH+ 3 ) and thus behave as a weak base. Thus, a ph shift (change in the [H+]) causes a change in net electrical charge (valence) of biologically molecules, that can alter biological activity: - directly (e.g., by altering the affinity for a charged ligand) or - indirectly (e.g., by altering molecular conformation).
Biological importance of ph ph-sensitive molecules -enzymes -receptors and their ligands, -ion channels, -membrane transporters (Na-K pump activity falls by about half when the ph shifts by ~1 ph unit from the optimum ph) -structural proteins ph-sensitive cellular processes: -enzymes activity (e.g. the activity of phosphofructokinase, a key glycolytic enzyme, falls by ~90% when ph falls by only 0.1) -membrane permeability -control of respiration -heart activity -oxygen Hb dissociation curve -nerve excitability & action potential of myelinated nerve -cell proliferation in response to mitogenic activation (fall as much as 85% when intracellular ph falls by 0.4).
Enzymes are affected by changes in ph The most favorable ph value - the point where the enzyme is most active, known as the optimum ph
ph and synaptic transmission Alkalosis increases synaptic transmission- alkalosis > 7.8 seizures Acidosis decreases synaptic transmission - acidosis < 7 coma (uremic / diabetic- ketone bodies ) ph and heart activity High [H+] in blood H+ diffuses in the cells electroneutrality law K+ diffuses out of the cells: Hyperpolarisation of heart muscle Low excitability Hyperkalemia
ph effect on hemoglobin dissociation curve - Bohr effect
ph and ventilation Chemoreceptors: Peripheral (carotid/ aortic body) Central (medulla oblongata)
ph and ventilation Low ph hyperventilation (ventilation increases 4-5 x when ph is 7) High ph hypoventilation CO2 formed by tissue metabolism is eliminated through respiration CO2 passes the blood brain barrier, it hydrates (CA) and forms H2CO3 H+ + HCO3- H+ influences central chemoreceptors CO2 regulates ventilation rate and depth indirectly by H+ increase
Physiological variations of ph Circadian rhythm: in the morning and at night - more acidic ph, as CO2 accumulates during sleep Age: newborn babies and children have a ph closer to the superior limit of the interval (7.42)- this favors anabolic processes (growth), while older people have more acidic ph due to the catabolic processes Physical effort (lactic acid production) Digestion phases: gastric digestion, there is a more basic ph -alkalosis (elimination of H+ in the gastric secretions) intestinal digestion, the ph is more acid - acidosis (HCO3- in the intestinal secretions) Altitude: more basic ph - physiologic alkalosis, as low po2 causes hyperventilation and thus low CO2. Temperature variations (lower ph with higher temp.)
Buffers reduce the free [H+] to minimize the size of the ph changes produced by adding acid or alkali to a solution, but cannot remove H+ ions from the body A buffer is any substance that reversibly consumes or releases H+ buffers help to stabilize ph Buffers do not prevent ph changes, they only help to minimize them. A Buffer B has a deprotonated form B(n), with valence n, that is in equilibrium with its protonated form HB(n+1), with a valence of n + 1: HB(n+1) is a weak acid because it does not fully dissociate; B(n) is its conjugate weak base Conversely, B(n) is a weak base and HB(n+1) is its conjugate weak acid. The total buffer concentration, [TB], is the sum of the concentrations of the protonated and unprotonated forms: Whenever a buffering reaction occurs, the concentration of one member of the pair increases while the other decreases.
Examples of buffer pairs The valence of the acidic (i.e., protonated) buffer form can be: - positive - zero - negative NH4+ (ammonium), H2CO3 (carbonic acid), and H2PO 4 ( monobasic inorganic phosphate) are all weak acids, whereas NH3 (ammonia), HCO3 (bicarbonate), and HPO 4 2 ( dibasic inorganic phosphate) are the respective conjugate weak bases.
Each buffer reaction is governed by a dissociation constant - K If we add to a physiological solution a small amount of HCl (a strong acid, that fully dissociates), the buffers in the solution consume almost all added H+: For each H + buffered, one B(n) is consumed. H + that is not buffered remains free in solution and is responsible for the ph decrease. If we instead titrate this same solution with a strong base such as NaOH, H + derived from HB(n+1) neutralizes almost all the added OH : For each OH buffered, one B(n) is formed. The tiny amount of added OH that is not neutralized by the buffer equilibrates with H + and H 2 O and is responsible for an increase in ph.
The buffering power (β) is a useful measure of the strength of a buffer β represents the number of moles of strong base (e.g., NaOH) added to 1 liter of solution to increase ph by 1 ph unit. This value is equivalent to the amount of strong acid (e.g., HCl) added to decrease the ph by 1 ph unit. In the absence of CO2 /HCO3 : -the buffering power of whole blood (plasma + RBCs, leukocytes, and platelets) is ~25 mm/ph unit, known as the non-hco3 - buffering power (β non-hco3- ). 25 mmol of NaOH would need to be added to 1 L of whole blood to increase the ph by 1 unit, assuming that β is constant over this wide ph range and in the absence of CO2 /HCO3. -for blood plasma, which lacks the cellular elements of whole blood, β non-hco3- is only ~5 mm/ph unit, which means that only about 5 mmol NaOH/1 L of plasma would be needed to produce the same ph increase.
Henderson-Hasselbalch equation: ph depends on the ratio [CO2]/[HCO3-] The most important physiological buffer pair is CO2 and HCO3. The impressive strength of this buffer pair is due to the volatility of CO2, which allows the lungs to maintain stable CO2 concentrations in the blood plasma despite ongoing metabolic and buffer reactions that produce or consume CO2. Exemple: a beaker contains an aqueous solution of 145 mm NaCl (ph = 6.81), but no buffers. When this solution is exposed to an atmosphere containing CO2, then the concentration of dissolved CO 2 ([CO 2 ] Dis ) can be calculated using Henry s law: Note that the pk of a buffer system is highest when pk=ph (when buffer components ratio is 1:1) The solubility coefficient s is ~0.03 mm/mm Hg at 37 C in blood plasma. Because the alveolar air with which arterial blood equilibrates has a P CO2 of ~40 mm Hg, [CO 2 ] Dis in arterial blood is 1.2 mm:
Interaction of CO 2 with water CO 2 itself is neither an acid nor a base, but CO 2 reacts with H 2 O to form carbonic acid: fast with CA CO 2 hydration reaction is very slow, but the enzyme carbonic anhydrase (CA) catalyzes a reaction that effectively bypasses this slow hydration reaction. Carbonic acid is a weak acid that rapidly dissociates into H+ and HCO 3, thus ph decreases as the reaction forms H+ (1 : 1 stoichiometry)! Even though the dissociation of H 2 CO 3 leads to generation of a weak conjugated base, HCO 3, ph decreases because H+ forms along with the weak base. In the absence of CA, the slow CO 2 hydration reaction limits the speed at which increased [CO 2 ]Dis leads to the production of H+.
The CO2 hydration and H2CO3 dissociation reactions could be considered as one reaction: The dissociation constant for this equilibrium is In logarithmic form, this equation becomes Considering Henry s law: [CO2] = s PCO2, then pk = -log 10 K This is the Henderson-Hasselbalch equation, a logarithmic restatement of the CO2 /HCO3 equilibrium showing that ph depends not on [HCO3 ] or PCO2 per se, but on their ratio.
The Henderson-Hasselbalch equation correctly predicts the normal ph of arterial blood. Human arterial blood has a PCO 2 of ~40 mm Hg and an [HCO 3 ] of ~24 mm pk = 6.1 at 37 C for the CO 2 /HCO 3 equilibrium (K = 10-6.1 M)
CO 2 /HCO 3 - has a far higher buffering power in an open than in a closed system The buffering power of a buffer pair such as CO2 /HCO3 depends on three factors: 1. Total concentration of the buffer pair, [TB]. buffering power β is proportional to total buffer concentration [TB]. 2. The ph or [H+] of the solution. 3. Whether the system is open or closed. - Open system: one member of the buffer pair equilibrate between the system (solution in which the buffer is dissolved) and the environment (everything else) - Closed system: neither member of the buffer pair can enter or leave the system, then HB(n+1) can become B(n), and vice versa, but [TB] is fixed. Examples: -inorganic phosphate in a beaker of water, or -a titratable group on a protein in blood plasma.
In a closed system, the buffering power of a buffer pair is ᵝ At a given [TB], βclosed has a bell-shaped dependence on ph. βclosed is maximal when [H+] = K (i.e., when ph = pk). Most non-hco3 buffers in biological fluids behave as if they are in a closed system. The total βnon-hco3 in a mixture of non-hco3 buffers is the sum of their βclosed values.
βnon-hco 3 - for the whole blood Whole blood contains a mixture of many non-hco 3 buffers: -hemoglobin -other proteins -small molecules: inorganic phosphate. The buffering power of whole blood is nearly constant near the physiological ph.
A physiologically important condition under which a buffer can function is in an open system. CO2 dissolved in blood plasma equilibrates with gaseous CO2 in the alveoli [CO2]Dissolved is fixed (1.2 mm) during buffering reactions in open system. Total CO2 = [CO2] + [HCO3 ] can vary widely and serve as a powerful buffer Whether CO2 /HCO3- neutralizes an acid or a base, the open-system buffering power is βopen does not have a maximum and rises exponentially with ph when PCO2 is fixed. βnon-hco3 = buffering power of all non-hco3 buffers in whole blood is ~25 mm/ph unit. βopen is ~55 mm/ph unit in normal arterial blood with [HCO3 ] = 24 mm. βopen represents more than two thirds of the total buffering power. The relative contribution of βopen is far more striking in interstitial fluid, which lacks the cellular elements of blood and also has a lower protein concentration.
Buffering reactions for bicarbonate buffer alone in an open system One liter of a solution having a ph of 7.4, a PCO2 of 40 mm Hg (1.2 mm CO2), and a [HCO3 ] of 24 mm, but no other buffers, will neutralize an added 10 mmol of HCl by the available [HCO3 ], forming nearly 10 mmol H2CO3, and then nearly 10 mmol CO2 plus nearly 10 mmol H2O. The CO2 that forms does not accumulate, but evolves to the atmosphere so that [CO2]Dis is constant. [HCO3 ] decreases by 10 mm (i.e., the amount of added H+), from 24 to 14 mm. The final ph falls from 7.40 to 7.17, as predicted by the Henderson-Hasselbalch equation, corresponding to an increase in free H+ of 28 nm: = 28 nm Even though we have added 10 millimoles of HCl to 1 L, [H+] increased by only 28 nm. The open-system buffer has neutralized 9.999,972 mm of the added 10 mm H+. The buffering provided by CO2 /HCO3 in an open system (βopen) is so powerful because only depletion of HCO3 limits neutralization of H+. The buildup of CO2 is not a limiting factor because the atmosphere is an infinite sink for newly produced CO2.
Bicarbonate buffer in an open system For keeping ph = 7.4 Lungs eliminate CO2 (volatile acid) and determine the concentration of H2CO3 by regulating pco2 at 40 mmhg pco2 dissolved CO2 by hydration is in equillibrum with H2CO3 The kidneys maintain [HCO3-] at 24mEq/l
Buffering reactions for blood bicarbonate buffer special conditions For normal arterial blood in a closed-system (as in a capped syringe): buffering power βco2 /HCO3 =2.6 mm/ph unit (< 5% of the βopen=55 mm/ph unit) A physiological example in which the CO2 /HCO3 system is poorly open is ischemia, in which a lack of blood flow minimizes the equilibration of tissue CO2 with blood CO2. Also, lack of O2 determines high production of lactic acid. Thus, ischemic tissues are especially susceptible to large ph shifts.
ACID-BASE CHEMISTRY WHEN CO2/HCO3- IS THE ONLY BUFFER In the absence of other buffers, in an open system, doubling* PCO2 causes ph to fall by 0.3, but causes almost no change in [HCO3-] The resulting increase in [CO2]Dis causes the CO2 /HCO3 equilibrium to shift toward formation of H+ and HCO3 example of a CO2 titration, initiated by altering PCO2. it is a respiratory acidosis acidosis because ph falls, and respiratory because pulmonary problems are the most common causes of an increase in the PCO2 of arterial blood *log2=0.3
ACID-BASE CHEMISTRY WHEN CO2/HCO3- IS THE ONLY BUFFER In the absence of non-hco3 buffers, how far ph falls during respiratory acidosis depends on the initial ph and PCO2, as well as on the final PCO2. For example, doubling PCO2 from 40 mm Hg to 80 mm Hg causes [CO2]Dis to double to 2.4 mm. At this point, the 1-L system is far out of equilibrium and can return to equilibrium only if some CO2 (X mmol) combines with X H2O to form X H+ and X HCO3. X 40 nmol (0.000,040 mmol) X represents the flux of CO2 that passes through the reaction sequence CO2+ H2O H2CO3 H + + HCO3- to re-establish the equilibrium (no other buffers present) CO2 from the atmosphere replenishes the CO2 consumed in this reaction, so that [CO2]Dis remains at 2.4 mm after the new equilibrium is achieved.
The flux X is so small, as no other buffers are present and every H+ formed remains free in solution. Thus, only a minuscule amount of H+ need be formed before [H+] nearly doubles from 40 to nearly 80 nm, while [HCO3 ] undergoes only a tiny fractional increase, from 24 to 24.000,040 mm. The final ph in this example of respiratory acidosis is: Or: The opposite acid-base disturbance, in which PCO2 would fall, is respiratory alkalosis.
In the absence of non-hco3- buffers, doubling [HCO3-] causes ph to rise by 0.3 In 1 L of solution that has the ionic composition of arterial blood, adding 24 mmol of HCO3 (e.g., NaHCO3) drives the CO /HCO3 equilibrium toward CO2. The new equilibrium is achieved when X mmol of HCO3 combines with X H+ to produce X CO2 and X H2O. Because the system is open to CO2, the generation of X mmol CO2 causes no change in [CO2]Dis, as the newly formed CO2 simply evolves into the atmosphere. X 0.000,020 mm. in the absence of other buffers, an initial doubling of [HCO3 ] from 24 to 48 mm causes [H+] to fall by nearly half, from 40 to 20 nm. Example of metabolic alkalosis
The ph of a CO 2 /HCO 3 solution does not depend on [HCO 3 ] or PCO 2 per se, but on their ratio. Because it is the kidney that controls [HCO 3 ] in the blood plasma, and because it is the lung that controls PCO 2, the ph of blood plasma is under the dual control of both organ systems, a concept embodied by a simplified variant of the Henderson-Hasselbalch equation:
Acid excretion Lungs excrete volatile acid (CO2) Major source of rapid acid excretion 13000 meq/day of carbonic acid Kidneys- excrete fixed acids 40-80 meq/day Fixed acids may increase to 300 meq /day if necessary Base excretion Only kidney regulated Primary base in the organism HCO3- The kidney can retain or excrete bicarbonate as needed
Lungs, kidneys, and RBCs contribute to the acid-base balance: The lungs control the gas exchange with the atmospheric air. CO2 generated in tissues is transported in plasma as bicarbonate; the hemoglobin in RBCs contributes to CO2 transport. Hemoglobin also buffers the H+ derived from carbonic acid. The kidneys reabsorb filtered bicarbonate in the proximal tubules and generate new bicarbonate in the distal tubules, where there is a net secretion of H+
The Major Body Buffer Systems Site Buffer System Comment Interstitial fluid Bicarbonate For metabolic acids Phosphate Protein Not important because its concentration is too low Not important because its concentration is too low Blood Bicarbonate Important for metabolic acids Haemoglobin Plasma protein Phosphate Important for carbon dioxide Minor buffer Concentration too low Intracellular fluid Proteins Important buffer Phosphates Important buffer Urine Phosphate Responsible for most of 'Titratable Acidity' Ammonia Important - formation of NH 4 + Bone Ca carbonate Used in prolonged metabolic acidosis