Acid-Base Balance Dr J.O Akande Introduction • The importance of acid–base homeostasis cannot be overstated because of its importance in keeping hydrogen ion (H+) balance under control. • Acid - Base balance is primarily concerned with two ions: – Hydrogen (H+) - – Bicarbonate (HCO3 ) • H+ ion must be maintained within narrow ranges in order to be compatible with living systems • Enzymes, hormones and ion distribution are all affected by Hydrogen ion concentrations • Excitability of nerve and muscle cells are also affect Definitions • Acid – Any substance that can yield a hydrogen ion (H+) or hydronium ion when dissolved in water – Release of proton or H+ • Base – Substance that can yield hydroxyl ions (OH-) – Accept protons or H+ • The relative strengths of acids and bases, their ability to dissociate in water, are described by their dissociation constant (also ionization constant K value) Cont’d • pK/ pKa – Negative log of the ionization constant of an acid – Strong acids would have a pK <3 – Strong base would have a pK >9 • pH – Negative log of the hydrogen ion concentration – pH= pK + log([base]/[acid]) – Represents the hydrogen concentration • For acids, raising the pH above the pK will cause the acid to dissociate and yield an H. • For bases, lowering the pH below the pK will cause the base to release OH. • Many species have more than one pK, meaning they can accept or donate more than one H. Buffer • Combination of a weak acid and /or a weak base and its salt • Resists changes in pH • Buffering is a process by which a strong acid (or base) is replaced by a weaker one, with a consequent reduction in the number of free H+ – • H+Cl + NaHCO3 →H2CO3 + NaCl (strong acid) (buffer) (weak acid) (neutral salt) • Therefore the change in pH, after addition of acid, is less than it would be in the absence of the buffer. • Effectiveness depends on • pK of buffering system • pH of environment in which it is placed Terms • Acidosis – pH less than 7.35 • Alkalosis – pH greater than 7.45 • Note: Normal pH is 7.35-7.45 • Acidosis is commoner than alkalosis because metabolism tends to produce H+ rather than OH– • The pH is a measure of H+ activity. • It is log 10 of the reciprocal of [H+] in mol/L. • log 100 = log 102 = 2, and log 107 = 7 Cont’d • If [H+] is 10–7 (0.0000001) mol/L, • then –log [H+] = 7. • A blood pH of 7.0 indicates a severe acidosis. A blood pH of 7.7 similarly indicates a severe alkalosis. • Urinary pH is much more variable than that of blood, and [H+] can vary 1000-fold. The Henderson–Hasselbalch equation • Expresses the relation between pH and a buffer pair – that is, a weak acid and its conjugate base. • The equation is valid for any buffer pair • The pH being dependent on the ratio of the concentration of base to acid. • Note that pKa is the negative logarithm of the acid dissociation constant (Ka), and, the larger the value of pKa, the smaller the extent of acid dissociation Cont’d Cont’d – • In this equation the base is bicarbonate (HCO3 ) and the acid is carbonic acid (H2CO3). • It is not possible to measure H2CO3 directly • However, it is in equilibrium with dissolved CO2, of which the partial pressure (PCO2) can be estimated. • The concentration of H2CO3 is derived by multiplying this measured value by the solubility constant (S) for CO2 • If the PCO2 is expressed in kPa, S = 0.23, or in mmHg, S = 0.03. The overall pKa of the bicarbonate system is 6.1. • Therefore if PCO2 is in kPa Law of Mass Action CA + - CO2 + H2O H2CO3 H + HCO3 AT EQUILIBRIUM, + - [H ] x [HCO3 ] K = [H2CO3] + - [H ] x [HCO3 ] = SOLUBILITY x PCO2 + - [H ] x [HCO3 ] K = 0.03 x PCO2 WHICH, AFTER TAKING LOGS, BECOMES: - [HCO3 ] pH = pK + Log( ) 0.03 x PCO2 - THE HENDERSON - HASSELBALCH EQUATION Derivation of normal pH Cont’d – • Plasma [HCO3 ] is controlled largely by the kidneys and PCO2 by the lungs. • In acid–base disturbances due to respiratory problems the kidneys are essential for compensation • Conversely, in metabolic (nonrespiratory) causes of acid–base imbalance the compensation is due mainly to changes in pulmonary function. • Despite considerable fluctuations in the rate of release of H+ into the ECF, • The [H+], and therefore pH, is relatively tightly controlled in blood ACID-BASE REGULATION • Maintenance of an acceptable pH range in the extracellular fluids is accomplished by three mechanisms: • Chemical Buffers • React very rapidly (less than a second) • Respiratory Regulation • Reacts rapidly (seconds to minutes) • Renal Regulation • Reacts slowly (minutes to hours) 15 16 Chemical Buffers • Immediate acting • Bicarbonate buffer - mainly responsible for • Combine with offending buffering blood and acid or base to interstitial fluid neutralize harmful • Phosphate buffer - effects until another effective in renal tubules system takes over • Protein buffers - most plentiful - hemoglobin Respiratory System • Lungs regulate blood levels of CO2 • CO2 + H2O = Carbonic acid • High CO2 = slower breathing (hold on to carbonic acid and lower pH) • Low CO2 = faster breathing (blow off carbonic acid and raise pH) • Twice as effective as chemical buffers, but effects are temporary Kidneys • Reabsorb or excrete • Adjustments by the excess acids or bases kidneys take hours to into urine days to accomplish • Produce bicarbonate • Bicarbonate levels and pH levels increase or decrease together Buffering System • Four Major Buffer Systems – Protein Buffer systems • Amino acids • Hemoglobin Buffer system – Phosphate Buffer system – Bicarbonate-carbonic acid Buffer system Protein Buffer System – Originates from amino acids • ALBUMIN- primary protein due to high concentration in plasma – Buffer both hydrogen ions and carbon dioxide Hemoglobin Buffer System – Haemoglobin is an important blood buffer. – It only works effectively in cooperation with the bicarbonate system – Roles • Binds CO2 • Binds and transports hydrogen and oxygen • Participates in the chloride shift • Maintains blood pH as hemoglobin changes from oxyhemoglobin to deoxyhemoglobin Erythrocytes • Erythrocytes produce little CO2 as they lack aerobic pathways. • Plasma CO2 diffuses along a concentration gradient into erythrocytes, where CD catalyses its reaction with water to form carbonic acid (H2CO3), which then dissociates. + – • Much of the H is buffered by Hb, and the HCO3 diffuses out into the extracellular fluid along a concentration gradient. • Electrochemical neutrality is maintained by diffusion of Cl– in the opposite direction into cells. • This movement of ions is known as the ‘chloride shift Oxygen Dissociation Curve Curve B: Normal curve Curve A: Increased affinity for hgb, so oxygen keep close Curve C: Decreased affinity for hgb, so oxygen released to tissues Phosphate Buffer System • Phosphate is normally the most important buffer in the urine because its pKa is relatively close to the pH of the glomerular filtrate • Assists in the exchange of sodium for hydrogen • It participates in the following reaction -2 + – • HPO 4 + H H2PO 4 • Essential within the erythrocytes • In a mild acidosis, more phosphate ions are released from bone than at normal pH; • Increased urinary H+ secretion is linked with increased buffering capacity in the glomerular filtrate owing to the increase of phosphate. Cont’d • At a urinary pH below 5.5, most of the filtered phosphate is converted to dihydrogen phosphate. • Therefore, at low pH, urinary phosphate cannot maintain the essential buffering of continued H+ secretion. Ammonia • Urinary ammonia probably allows H+ secretion, and – therefore HCO3 formation, to continue after other buffers have been depleted. • Ammonia, produced by hepatic deamination of amino acids, is rapidly incorporated into urea, with a net production of H+. • However, as the systemic [H+] increases, there is some shift from urea to glutamine (GluCONH2) synthesis, with a slight fall in hepatic H+ production. • Glutamine is taken up by renal tubular cells, where it is hydrolysed by glutaminase to glutamate (GluCOO–) and NH4+ Cont’d – + • H2O + GluCONH2 → GluCOO + NH4 • Ammonia and NH4+ form a buffer pair with a pKa of about 9.8 • PH= 9.8 log • Ammonia can diffuse out of the cell into the tubular lumen much more rapidly than NH4+. • If the luminal fluid is acidic, NH3 will be retained within the lumen by avid combination with H+ derived from the CD mechanism. • This allows H+, produced in the kidneys, to be excreted as + – ammonium chloride (NH4 Cl ); – • Thus, in severe acidosis, HCO3 formation can continue even when phosphate buffering power has been exhausted. – • There is a net gain of HCO3 . Bicarbonate/carbonic acid buffer system • Function almost instantaneously • Cells that are utilizing O2, produce CO2, which builds up. • Thus, more CO2 is found in the tissue cells than in nearby blood cells. • This results in a pressure (pCO2). • Diffusion occurs, the CO2 leaves the tissue through the interstitial fluid into the capillary blood – 5 - 8% transported in dissolved form – A small amount of the CO2 combines directly with the hemoglobin to form carbaminohemoglobin – 92 - 95% of CO2 will enter the RBC, and under the following reaction + - CO2 + H20 ↔ H + HCO3 – Once bicarbonate formed, exchanged for chloride Cont’d Excreted by lungs Conjugate Carbonic base acid Bicarbonate Excreted in urine Cont’d • A buffer pair is most effective at maintaining a pH near its pKa. • The optimum pH of the ECF is about 7.4, and the pKa of the bicarbonate system is 6.1. • This may seem to be disadvantageous, the bicarbonate system is the most important buffer in the body • it accounts for more than 60 per cent of the blood buffering capacity • H+ secretion by the kidney depends on it • It is necessary for efficient buffering by haemoglobin Lungs/respiratory • Quickest way to respond, takes minutes to hours to correct pH • Eliminate volatile respiratory acids such as CO2 • Doesn’t affect fixed acids like lactic acid • Body pH can be adjusted by changing rate and depth of breathing “blowing off” • Provide O2 to cells and remove CO2 • The rate of respiration, and therefore the rate of CO2 elimination, is controlled by chemoreceptors + • The receptors respond to changes in the [CO2] or [H ] of plasma or of the cerebrospinal fluid.