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Acid - base equilibrium pH concept pH = - log [H+] pH 1 [H+] 100 mmol/L D = 90 mmol/L 2 10 mmol/L D = 9 mmol/L 3 1 mmol/L 2 pH = - log [H+] 3 pH = - log [H+] pH pH of capillary blood norm: 7,35 – 7,45 Sorensen unit – 7,7 44,7 nmol/L – 35,5 nmol/L pathology: 6,9 126 nmol/L pH 7,4 – 7,1 7,4 – 7,7 (0,3 unit) (0,3 unit) – Sorensen unit 20 nmol/L [H +] 40 – 80 nmol/L (D = 40 nmol/L) 40 – 20 nmol/L (D = 20 nmol/L) 4 Buffer solutions Mixtures: weak acid and its salt with strong base CH3COOH + CH3COO- weak base and its salt with strong acid NH3 aqueous + NH4+ two salts of polyprotic acids -H2PO4+ -HPO42- 5 Henderson – Hasselbalch equation AH A- + H+ [A-] [H+] Ka = [AH] [A-] [H+] - log Ka = - log [AH] because - log Ka = pKa [A-] [H+] pKa = - log [AH] 6 Henderson - Haselbalch equation AH A- + H+ [A-] [H+] pKa = - log [AH] pKa = - log [A-] – log [H+] + log [AH] -log [H+] = pKa + log[A-] – log [AH] -log [H+] = pH [A-] pH = pKa + log [AH] 7 Henderson – Hasselbalch equation [A-] pH = pKa + log [AH] pH of the buffer mixture depends on: kind of the acid ratio of the concentrations of buffer components pH of the buffer mixtures does not change with dilution of the solution! 8 Buffer solutions – mechanism of action acetate buffer: CH3COOH i CH3COO- at equilibrium: CH3COOH + H2O CH3COOH CH3COO- + H3O+ CH3COO- + H+ acetic acid is a weak acid, acetate ions come entirely from salt (acetate) add some amount of strong acid: +H+ CH3COOH CH3COO- + H+ strong acid removes weak acid from its salt 9 Buffer capacity Dc b = DpH b– buffer capacity Dc – amount of strong acid or strong base added to the buffer solution mol/L DpH – observed change of pH Buffer capacity depends on concentration of components: increases with their increment decreases with dilution of the buffer 10 Buffer capacity Dc b = DpH [A-] pH = pKa + log [AH] is the highest when pH = pK Excess of acid in the buffer – better buffering of bases Excess of a salt – better buffering of acids. During addition of acid or base buffer capacity decreases is equal to zero when whole salt in the buffer is converted to weak acid or whole weak acid is converted to salt. 11 Intracellular pH cytoplasm 6,0 mitochondria endoplasmatic reticulum nucleus 7,0 – 7,4 average pH of intracellular fluid: 6,95 (112 nmol/L) differences in pH of extracellular fluid between cells from different organs: erytrocytes 7,20 renal tubular epithelial cells cells of skeletal muscles. 6,9 7,32 12 Mechanisms of pH regulation in organism Organ’s regulation regulation by kidney regulation by lungs regulation by bones Buffer regulation proteinate buffer H-proteins Proteinates phosphate buffer -H2PO4-HPO42- bicarbonate buffer HCO3H2CO3 Hemoglobin buffer 13 Hemoglobin buffer is the most important proteinate buffer 1. Hemoglobin makes about ¾ of all blood’s protein 2. Hemoglobin has an acidic character because of the presence of majority of acidic groups of hem over basic groups of globins. Therefore, hemoglobin (Hb) has a large capacity for base binding. 3. Acidity of hemoglobin changes significantly with degree of oxidation. Hemoglobin transports not only oxygen from lungs to cells but also carbon dioxide from cells to lungs. 14 Proteinate buffer-hemoglobin buffer Hemoglobine buffer is an intracellular buffer. The Hb and oxygen connection is reversible and the Hb molecule changes its conformation when it binds O2. HHb (Hb without oxygen) and HHbO2 (Hb connected with oxygen, oxyhemoglobin) differ in their ability to donate or accept H+ ions. HbO2 + H+ stronger acid HHb + O2 In tissues weaker acid HHb + O2 + HCO3- HbO2 + H2O + CO2 In Lungs Buffering properties depend on the equilibrium between oxy-hemoglobin and hemoglobin. Hemoglobin is less acidic than oxy-hemoglobin. With increasing CO2 pressure and H+ concentration, the amount of HHb increases and buffering capacity increases as well. 15 Proteinate buffer-hemoglobin buffer is based on binding of H + ions within capillary vessels of all cells and its release in capillary vessels in lungs. In acidic environment: carboxyl groups of amino acids do not dissociate basic groups (amine, imidasole) are proton acceptors In basic environment: carboxyl groups are proton donors and they neutralize hydroxide ions in weak basic environment - pH 7,4, proteins are in form of anions. concentration of proteins in the blood is 16 mEq/l Buffer capacity of proteinate buffer is: 5 mmols/ l/one pH unit 16 Proteinate Buffer – hemoglobin buffer I system HHbO2 KHbO2 II system HHb KHb hemoglobin is an important buffer in blood hemoglobin makes ¾ of whole proteins in the blood. hemoglobin is acidic because of excess of acidic groups of heme over basic groups from globin. acidity of hemoglobin can change depending on oxygenation It is indispensable for full effectiveness of carbonate buffer in an open system 17 Phosphate buffer – important urinary buffer phosphate buffer in blood : H2PO4K H2PO4- = 6,2 x 10-8, HPO42- + H+ pK2 = 7,21 HPO42- / H2PO4- = 4/1 phosphate buffer in urine: HPO42- / H2PO4- = 1/4 HPO42- + H+ H2PO4- main intracellular buffer optimal pK for this buffer is 6,8. in phosphate buffer in urine (pH about 6,0) ratio of hydrogen phosphate to dihydrogen phosphate is 1/4 Difference between ratios of phosphates in blood and urine is because hydrogen phosphate of urine binds protons secreted by distal tubules of kidneys and is converted to dihydrogen phosphate. 18 Bicarbonate buffer – an extracellular buffer; in equilibrium with atmospheric air. The most important buffer sytem in the blood is: HCO3-/H2CO3 organism removes by lungs a product of dehydration of carbonic acid - carbon dioxide. This buffer acts in an open system. H2CO3 i CO2 dissolved in water phase remain at equilibrium with CO2 which is in gasous phase. CO2 in blood circulating in lungs remains at equilibrium with CO 2 in lung vesicles. buffer in an open system has several times higher capacity in comparison to a closed system 19 Bicarbonate buffer is responsible for the physiological pH of blood, 7.35-7.45 CO2 + H2O 99% gas liquid H+ + HCO3- H2CO3 1% liquid liquid liquid % - expresses amounts in equilibrium state [A] pH = pKH CO + log 2 3 [ AH] [ HCO3- ] pH = pKH2CO3 + log [CO2 ] 20 Gasometric parameters of blood (physiological levels): pKH2CO3 = 6,11 [HCO3- ] = 24 mmol/L [CO2 ] = a x p a – coefficient of solubility of CO 2 in plasma a = 0,225 mmol/L p – partial pressure of CO2 in lung vessels pCO2 = 5,32 kPa 21 Bicarbonate buffer CO2 + H2O 99% H+ + HCO3- H2CO3 1% [ HCO3- ] pH = pKH2CO3 + log [CO2 ] pKH2 CO3 = 6,11 [HCO3 - ] = 24 mmol/L [CO2 ] = a x p a = 0,225 mmol/L pCO2 = 5,32 kPa pH of bicarbonate buffer in blood 24 = 7,4 pH = 6,11 + log 0,225 x 5,32 22 Bicarbonate buffer CO2 + H2O 99% H2CO3 1% H+ + HCO3- Bicarbonate buffer – mechanism of action +H+ CO2 + H2O H2CO3 H+ + HCO3- The kidneys regulate HCO-3 ion concentration in blood plasma and protect the organism against metabolic acidosis. The role of kidneys is: • reabsorption of HCO-3 ions in kidney tubules • excretion of HCO-3 ions when their level changes in extracellular fluid • regeneration of lost HCO-3 ions or in reactions with H+ derived from water (in proximal tubular cells and collecting tubule cells) 23 Bicarbonate buffer most impotant buffer in blod, in acid-base balance acts in open system in normal conditions HCO3/CO2 is 20:1 metabolic component [HCO3- ] pH = pKH2CO3 + log pCO2 x a respiratory component 24 Disturbances of acid-base balance pH = pKH2CO3 + log [HCO3- ] pCO2 x a metabolic component respiratory component Too much acid in the body resulting from accumulation of acid or depletion of alkaline reserves leads to acidosis – abnormally low pH. May be caused by : diabetic ketoacidosis, lung disease, kidney disease. The condition opposite to acidosis is – alkalosis – pH is to high due to excess base or insufficient acid in the body. 25 Disorder pH HCO3- pCO2 Metabolic acidosis Respiratory acidosis Metabolic alkalosis Respiratory alkalosis primary changes secondary changes pH = pKH2CO3 + log [HCO3- ] pCO2 x a 26 Bicarbonate buffer To 1 liter of plasma 10mmols of strong acid were added . Calculate pH change when the system is: • open • closed. pKH2CO3 = 6,11; [HCO3- ] = 24 mmol/L; a = 0,225 mmol/l /kPa pCO2 = 5,32 kPa; pH=7,4 Open system 24 HCO3- + 10 H+ -------> 10 CO2 + 10 H2O + 14 HCO3- pH = 6,11 + log 24 – 10 0,225 x 5,32 14 = 6,11 + log --------- = 7,17 1,2 Closed system 24 HCO3- + 10 H+ -------> X 10 CO2 + 10 H2O + 14 HCO324 – 10 pH = 6,11 + log 1,2 + 10 14 = 6,11 + log --------- = 6,2 11,2 27 Red blood cells’ role in acid-base regulation of blood . Reaction in capillary tissues O2 Erythrocyte HbO2- + H+ CO2 + H2O To tissues H-Hb + O2 H2CO3 H+ + HCO3- carbonic anhydrase Cl- HCO3Cl- CO2 From tissues Offfset chloride 28 Red blood cells role in blood’s acid-base regulation; reaction in lung capillaries O2 from lungs Erythrocyte H-Hb + O2 HCO3 + H+ Cl- HbO2- + H+ H2CO3 H2O + CO2 carbonic anhydrase HCO3- Cl- CO2 Exhaled 29 THE END 30