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11/22/15 Dr. Chris Doumen Collin College BIOL 2402 Acid Base Homeostasis 1 Acid -Base Balance and Regulation Acid-Base Balance refers to the precise regulation of free hydrogen ion concentration in the body fluids Free hydrogen ions determine the acidity of the body fluids and pH is used as a specific H+ indicator pH = -log [H+] H2O H+ + OH- where [H+] is the molar concentration [H+] = 10-7 M [H2O] = 55 M pH = - log [10-7 ] = 7 2 1 11/22/15 Acid -Base Balance and Regulation In physiology, 7.4 is considered neutral because it reflects the average blood pH ( concentration of H+ = 40 nM; compare that with [Na+] = 140 mM … ) • Acidosis or acidemia: A blood pH below 7.35 • Alkalosis or alkalemia : A blood pH above 7.45 Death occurs within seconds when blood pH falls below 6.8 or above 8.0. Regulation of blood pH, more specifically H+, is thus a very important homeostatic factor in life. 3 Acid -Base Balance and Regulation Sources of H+ in the body VOLATILE ACIDS : Carbon dioxide and carbonic acid • CO2 and H2CO3 FIXED ACIDS : Inorganic acids (non carbonic acids ) from diet • phosphoric acid, sulfuric acid, ammonia • ~ 1 - 1.5 mmoles of H+ /kg/day ORGANIC ACIDS : resulting from metabolism • citric acid, lactic acid, pyruvic acid, ketone body acids 4 2 11/22/15 Acid -Base Balance and Regulation Quantities of H+ produced per day Diet : 70 000 000 nmoles/day Metabolism : 5 000 000 000 nmoles/day But if we calculate the amount of free protons in the blood stream ( = pH 7.4) is comes out to be 40 nmol /L ( = 10-7.4 M) Our body is thus constantly challenged by an enormous overload of protons. It requires different mechanism to keep the proton levels under control. Small uncontrolled changes in pH can have drastic effects on our metabolic sanity. 5 Lines of defense against changes in pH The body’s different mechanisms to deal with changes in protons are 3 fold : • Chemical Buffers • Respiratory mechanism • Renal mechanism Chemical buffers act immediately ; they bind excess protons ( or release protons) but do not eliminate H+ from the body. They can only soak up extra H+ depending on the concentration of the chemical buffers present. When capacity is full, the 6 additional H+ needs to be removed from the body 3 11/22/15 Lines of defense against changes in pH The lungs and kidneys aid in the removal of acids from the body. They act however more slowly, with the lungs being faster compared to the kidneys. 7 Lines of defense against changes in pH 1. Chemical Buffers Chemical Buffers are composed out of compounds that minimize pH changes when acids or bases are added. The compounds come in pairs : • a weak acid (HY below) : releases H+ • a weak base (Y- below): binds H+ HY Weak acid H+ + Y Weak base 8 4 11/22/15 Lines of defense against changes in pH 1. Chemical Buffers HY Weak acid H+ + Y Weak base These reactions have a certain equilibrium status and the direction of the reactions are influenced by the presence of their compounds on either side of the arrows : if there is a sudden increase in HY, the reaction will proceed to the right and more H+ will be created On the other hand, if more protons are added, the reaction will proceed to the left and the extra prtons will be removed from a solution by the weak base, Y9 Lines of defense against changes in pH A. The most important buffer in ICF (inside cells) are the proteins Proteins contain both acid and basic groups and can thus bind H+ and/or release protons quite easily. 10 5 11/22/15 Lines of defense against changes in pH • If pH climbs, the carboxyl group of amino acid acts as a weak acid • If the pH drops, the amino group acts as a weak base Hemoglobin in RBC is important in binding H+ produced by tissues and thus in buffering blood pH. It is the only intracellular buffer system with an immediate effect on ECF pH . 11 Lines of defense against changes in pH B. The most important buffer in ECF is the bicarbonate buffer CO2 + H2O H2CO3 H+ + HCO3- • Has the following limitations: • It cannot protect the ECF from pH changes due to increased or depressed CO2 levels • Only functions when respiratory system and control centers are working normally • It is limited by availability of bicarbonate ions (bicarbonate reserve) • Bicarbonate ion shortage is rare Due to large reserve of sodium bicarbonate 12 6 11/22/15 Lines of defense against changes in pH 13 Lines of defense against changes in pH c. Phosphate buffer is an important intracellular and urinary buffer system Na2HPO4 + H+ NaH2PO4 + Na+ Humans consume more phosphate than needed. The excess is filtered into the nephrons and is not re-absorbed by the kidney. The phosphate helps to buffer urine pH in the nephron. It binds the secreted protons and keeps the pH above 5. If it were not for this buffer, urine pH would be extremely acidic very fast ( below 4.5) and prevent the nephron from secreting H+ . Since phosphates are used extensively in side the cell, this buffer system also contributes to intracellular buffering of pH. 14 7 11/22/15 Lines of defense against changes in pH Buffer Systems occur in Intracellular fluid (ICF) Phosphate Buffer System The phosphate buffer system has an important role in buffering the pH of the ICF and of urine. Extracellular fluid (ECF) Protein Buffer Systems Protein buffer systems contribute to the regulation of pH in the ECF and ICF. These buffer systems interact extensively with the other two buffer systems. Hemoglobin buffer system (RBCs only) Amino acid buffers (All proteins) Carbonic Acid– Bicarbonate Buffer System The carbonic acid– bicarbonate buffer system is most important in the ECF. Plasma protein buffers 15 Lines of defense against changes in pH • Limitations of Chemical Buffer Systems • Provide only temporary solution to acid–base imbalance • They do not eliminate H+ ions • Supply of chemical buffer molecules is limited • Maintenance of Acid–Base Balance • For homeostasis to be preserved, captured H+ must: • Be permanently tied up in water molecules • Through CO2 removal at lungs • Be removed from body fluids • Through secretion at kidney 16 8 11/22/15 Lines of defense against changes in pH 2. Respiratory Mechanism of H+ regulation Regulation occurs by CO2 removal and involves the bicarbonate reaction • If not enough CO2 is expelled by the lungs, more CO2 stays behind in the blood • CO2 drives the bicarbonate reaction to the left and forms more Bicarbonate and protons and pH drops CO2 + H2O H2CO3 H+ + HCO3- • The opposite occurs when too much CO2 is expelled CO2 + H2O H2CO3 H+ + HCO317 Lines of defense against changes in pH This reaction is so important that it can provide direct information on the status of the body by analyzing the components of this reaction in the blood stream. Lets look at this reaction in a simplified version and think in acid-base terms. H+ + HCO3- CO2 + H2O Weak acid Weak base Every equilibrium reaction has an equilibrium set point know as the Keq value . The K value is the ratio of products over substrates at equilibrium ( the value of water is ignored since it is huge compared to the others). Keq = [H+] x [HCO3-] [H2O] x [CO2] = [H+] x [HCO3-] [CO2] 18 9 11/22/15 Lines of defense against changes in pH 1) Let’s now take the log of this equation log Keq = log [H+] x [HCO3-] [CO2] : = log [H+] + log [HCO3-]/[CO2] 2) Swap log Keq and log [H+] around and we get: -log [H+] = -log Keq + log [HCO3-]/[CO2] 3) Since -log [H+] = pH and define pKa as -log Keq we obtain : pH= pKa + log [HCO3-]/[CO2] Or in general when dealing with weak acids and bases pH = pKa + log{[Weak Base]/[ Weak Acid]} Lines of defense against changes in pH This is known as the Henderson-Hasselbalch equation for weak acids/bases and it can be used to determine pH levels pH = pKa + log {[Base]/[Acid]} pH = pKa + log {[HCO3-]/[CO2]} Since the CO2 levels in blood relate to the partial gas pressure for CO2 in blood, the concentration of CO2 in blood at body temp. is 0.03 times the partial pressure for CO2. pH = pKa + log {[HCO3-]/0.03 PCO2} 20 10 11/22/15 Lines of defense against changes in pH pH = pKa + log {[HCO3-]/0.03 PCO2} Healthy blood pH=7.4 and pKa for Bicarbonate reaction is a constant equal to 6.1 If we substitute now we get 7.4 = 6.1 + log {[HCO3-]/0.03 PCO2} Or 7.4 -6.1 = 1.3 = log {[HCO3-]/0.03 PCO2} You can get rid of a log expression by putting it to the power 10. 10(pH - pKa) = 10(1.3) =[HCO3-]/0.03 PCO2 = 20 21 Lines of defense against changes in pH So the ratio of [HCO3-] to {0.03 PCO2 } determines blood pH ! And since 101.3 = 20, the ratio of [HCO3-] to {0.03 PCO2 } in healthy blood should be in a ration of 20 : 1 ! The respiratory system uses this to adjust pH by regulating CO2. Changes can occur within minutes ; changing AVR by 2 ( or 1/2) will change pH of the blood by 0.2 units Anything that impairs respiratory system may thus affect the acid-base balance of the body. When a change in acid-base balance is due to a problem with the respiratory system, it is referred to as Respiratory Acidosis or 22 Respiratory Alkalosis. 11 11/22/15 Lines of defense against changes in pH 23 Figure 23-26a The Chemoreceptor Response to Changes in PCO2. Increased arterial PCO2 a An increase in arterial PCO2 stimulates chemoreceptors that accelerate breathing cycles at the inspiratory center. This change increases the respiratory rate, encourages CO2 loss at the lungs, and decreases arterial PCO2. Stimulation of arterial chemoreceptors Stimulation of respiratory muscles Increased PCO2, decreased pH in CSF Stimulation of CSF chemoreceptors at medulla oblongata Increased respiratory rate with increased elimination of CO2 at alveoli HOMEOSTASIS DISTURBED Increased arterial PCO2 (hypercapnia) HOMEOSTASIS RESTORED HOMEOSTASIS Normal arterial PCO2 Start Normal arterial PCO2 24 12