Introduction to Acid-Base Disorders PDF

Introduction to Acid-base Disorders DR ABDULSALAM I ENESI Dept. of Medical Biochemistry Fed. University Lokoja  Outline Introduction Acids and bases pH Buffers Acid-base balance in the body Respiratory regulation of pH Renal regulation of pH Respiratory acidosis...

Introduction to Acid-base Disorders DR ABDULSALAM I ENESI Dept. of Medical Biochemistry Fed. University Lokoja  Outline Introduction Acids and bases pH Buffers Acid-base balance in the body Respiratory regulation of pH Renal regulation of pH Respiratory acidosis Metabolic acidosis Respiratory alkalosis Metabolic alkalosis  Introduction The word acid base balance refers to maintenance of stable level of pH of body fluids. During metabolic processes both acids or bases are formed. Under normal conditions they are neutralized by specific systems involved in maintenance of pH level. Under pathological conditions excessive amounts of acids or bases may accumulate in body fluids and tissues leading to disturbances in acid base balance. In a normal healthy person, the blood pH ranges from 7.35-7.45. Maintenance of appropriate concentration of hydrogen ion (H+) is critical to normal cellular function. This Hydrogen (H+) Homeostasis or regulation involves the body buffers, the respiratory system (lungs) and the kidney. By the combined action of these systems constant H+ concentration is maintained in the body. Hydrogen ion The role of hydrogen ion is numerous in our body: It mediate the gradient of H+ concentration between inner and outer mitochondrial membrane thus acts as the driving force for oxidative phosphorylation. Secondly, the surface charge and physical configuration of proteins are affected by changes in hydrogen ion concentration. Also, Hydrogen ion concentration decides the ionization of weak acids and thus affects their physiological functions. The medical essence of ensuring Proper pH are 1. Ensure optimal action of enzymes and for the transport of molecules within the body and between cells and its surroundings. 2. Proper pH is required for the maintenance of structure of nucleic acids, proteins, coenzymes and various metabolites. 3. Acidosis and alkalosis are two important disorders of acid base balance  Acids and bases According to the definition proposed by Bronsted, Acids are substances that are capable of donating protons and bases are those that accept protons. i.e Acids are proton donors and bases are proton acceptors Acid is therefore, a compound that dissociates in aqueous solution to produce proton (H+) and a conjugate base (A-). HA = H+ + A- Acid may dissociate partially (weak acid) or completely (strong acid) in solution. In solution, weak acid establishes equilibrium between the proton and its conjugate base. Base is a compound that accepts proton in aqueous environment, eg ammonia reacts with a proton to produce an ammonium ion Acids and bases  Weak and Strong Acids The extent of dissociation decides whether they are strong acids or weak acids. Strong acids (HCl) dissociate completely in solution, while weak acids (acetic acid), ionize incompletely Since the dissociation of an acid is a freely reversible reaction, at equilibrium the ratio between dissociated and undissociated particle is a constant. The dissociation constant (Ka) of an acid is given by the formula where [H+] is the concentration of hydrogen ions, [A–] = the concentration of anions or conjugate base, and [HA] is the concentration of undissociated molecules. The pH at which the acid is half ionized is called pKa of an acid which is constant at a particular temperature and pressure. Strong acids will have a low pKa and weak acids have a higher pKa. pH, pOH and pKa According to Sorensen, The pH of a solution is simply defined as the negative logarithm of H ions conc, pH = - log [H+] = log [H+] Thus the pH value is inversely proportional to the acidity or H+ conc. Each pH unit represents a factor of 10 differences in [H+]. The pOH of a solution can though be defined as the negative logarithm of the OH ion conc. pOH = -log[OH-] The relationship between pH, pKa, concentration of acid and conjugate base (or salt) is expressed by the Henderson-Hasselbalch equation, Therefore, when the concentrations of base and acid are the same, then pH is equal to pKa. Thus, when the acid is half ionized, pH and pKa have the same values. The pH scale Common Terms used Buffers A buffer is a solution that can resist changes in pH upon addition of acid or basic components. Buffers are able to neutralize small amount of added acid or base to maintain the pH of the solution relatively stable. They are responsible for the maintenance of pH of plasma, ICF, ECF and tissues of the body. Most buffers are solutions composed of approximately equal amounts of a weak acid and salt of its conjugate base. Body fluids such as blood, cerebrospinal fluid, saliva etc have constant pH under normal physiological conditions. This is possible due to the presence of buffer in these fluids.  Buffers The pH of a buffer system is related to concentration of its weak acid as well as salt or conjugate base of weak acid and pK of weak acid. In logarithmic form, the relationship is expressed as Henderson- Hasselbalch equation When the concentrations of weak acid and its conjugate base are equal the equation becomes pH = pK. A buffer is most effective when the concentrations of salt and acid are equal or when pH = pKa. The pH of an ideal Buffer is affected by The value of pK (The lower the value of pK, the lower is the pH of the solution) and The ratio of salt to acid concentrations  Buffer The effective range of a buffer is 1 pH unit higher or lower than pKa. Since the pKa values of most of the acids produced in the body are well below the physiological pH, they immediately ionize and add H+ to the medium. This would necessitate effective buffering. Phosphate buffer is effective at a wide range, because it has 3 pKa values. Thus pK value can be defined as pH at which the concentration of acid and its conjugate base (salt) are equal. In simple words pK is the pH at which acid is half dissociated or neutralized. Based on titration of weak acid against base it was found that each buffer has maximum buffering action at its pK value. Features of buffer Composition of a Buffer are of two types: a. Mixtures of weak acids with their salt with a strong base or b. Mixtures of weak bases with their salt with a strong acid. Examples are i. H2CO3/NaHCO3 (Bicarbonate buffer) (carbonic acid and sodium bicarbonate) ii. CH3COOH/CH3COO Na (Acetate buffer) (acetic acid and sodium acetate) iii. Na2HPO4/NaH2PO4 (Phosphate buffer) Features of buffer Characteristics of a buffer are i. It has a definite pH value ii. Its pH value doesn’t change even with the addition of a small amount of strong acid or base. iii. Its pH value doesn’t change on keeping for a long time. iv. Its pH value doesn’t change on dilution.  Buffers of the body fluids Buffers are the first line of defense against acid load. Other lines of defense are respiratory and renal regulations in that order. The buffers are effective as long as the acid load is not excessive, and the alkali reserve is not exhausted. Once the base is utilized in this reaction, it is to be replenished to meet further challenge Example of buffer system are 1. Bicarbonate Buffer System 2. Phosphate Buffer System 3. Protein Buffer System Buffer systems of the body Buffers Buffer Bicarbonate Phosphate Protein System Mixture NaHCO3–/H2CO3 H2PO4/HPO4-2 haemoglobin (Hb–/HHb) Constituent acid part, carbonic acid acid part, (H2PO4) Histidine Imidazole (H2CO3) base part, (HPO4–), base part, bicarbonate (HCO3–), The pKa 6.1. 6.8 6.1 Normal plasma 24mmol/l 20-21mmol/l 14gm/dl level Buffering Highest (highly High (wide range of Low (close to body capacity distributed) ph) ph) Location Plasma Intracellular Plasma Regulators of Renal< metabolic Renal Respiratory the components Respiratory< respiratory  Lungs: respiratory regulation of ph The respiratory system responds to any change in pH immediately, but it cannot proceed to completion. This is the second line of defense achieved by changing the pCO2. When there is a fall in pH of plasma (acidosis), the respiratory rate is stimulated resulting in hyperventilation. This would eliminate more CO2, thus lowering the H2CO3 level. The rate of respiration (rate of elimination of CO2) is controlled by the chemoreceptors in the respiratory center which are sensitive to changes in the pH of blood. The activity of the carbonic anhydrase increases in acidosis and decreases with alkalosis through which hemoglobin serves to transport the CO2 formed in the tissues. The reverse occurs in the lungs during oxygenation and elimination of CO2. When the blood reaches the lungs, the Bicarbonate re enters the erythrocytes by reversal of chloride shift. It combines with H+ liberated on oxygenation of hemoglobin to form carbonic acid which dissociates into CO2 and H2O. CO2 is thus eliminated by the lungs.  Renal regulation of ph An important function of the kidney is to regulate the pH of the extracellular fluid. This is the third line of defense Normal urine has a pH around 6; this pH is lower than that of extracellular fluid (pH = 7.4). This is called acidification of urine. The pH of the urine may vary from as low as 4.5 to as high as 9.8, depending on the amount of acid excreted. The major renal mechanisms for regulation of pH are: A. Excretion of H+; Generation of Bicarbonate B. Reabsorption of bicarbonate (recovery of bicarbonate) C. Excretion of titratable acid (net acid excretion) D. Excretion of NH4+ (ammonium ions) Excretion of hydrogen ions in the proximal Reabsorption of tubules; CA = Carbonic anhydrase bicarbonate from the tubular fluid; CA = Carbonic anhydrase Phosphate mechanism Ammonia mechanism in tubules  Acid-Base Balance In normal life, the variation of plasma pH is very small. The normal pH of plasma is 7.4 (average hydrogen ion concentration of 40 nanomoles/liter). The pH of plasma is maintained within a narrow range of 7.38 to 7.42. The pH of the interstitial fluid is generally 0.5 units below that of the plasma. If the pH is below 7.38, it is called acidosis. Acidosis is the clinical state, where acids accumulate or bases are lost When the pH is more than 7.42, it is alkalosis. A loss of acid or accumulation of base leads to alkalosis Life is threatened when the pH is lowered below 7.25 and it is very dangerous if pH is increased above 7.55. Acidosis leads to CNS depression and coma. Death occurs when pH is below 7.0. Alkalosis induces neuromuscular hyperexcitability and tetany. Death occurs when the pH is above 7.6. Classification of Acid-Base Disturbances Disturbances in Acid base balance are grouped into acidosis and alkalosis. 1. Acidosis (fall in pH) It is due to accumulation of acids and blood pH is below 7.4 a. Respiratory acidosis: Primary excess of carbonic acid. b. Metabolic acidosis: Primary deficit of bicarbonate 2. Alkalosis (Rise in pH) It is due to accumulation of alkali and blood pH is above 7.4. c. Respiratory alkalosis: Primary deficit of carbonic acid.  Acid-base Disturbances The Types of acid-base disturbances are  Disturbances in Acid base balance Acidosis or alkalosis due to less or more of bicarbonate are called as metabolic acidosis or metabolic alkalosis respectively. Like wise acidosis or alkalosis due to more or less of carbonic acid are called as respiratory acidosis and respiratory alkalosis respectively. These disturbances can be acute or chronic and always body attempts to restore normal acid base balance or H CO–3 / H2 CO3 ratio by changing the removal of CO2 by lungs or by altering the reabsorption of H CO–3 and H+ removal in kidney. If the normal ratio of H CO–3, H2CO3 is restored then the acidosis or alkalosis is compensated. If body fails in its  Compensatory responses Each of the acid-base disturbances will be followed by a secondary compensatory change in the counter acting variable, e.g. a primary change in bicarbonate involves an alteration in pCO2. The compensatory (adaptive) responses are: a. A primary change in bicarbonate involves an alteration in pCO2. The direction of the change is the same as the primary change and there is an attempt at restoring the pH to 7.4. b. Adaptive response is always in the same direction as the primary disturbance. Primary decrease in arterial bicarbonate involves a reduction in arterial blood pCO2 by alveolar hyperventilation. c. Similarly, a primary increase in arterial pCO2 involves an increase in arterial bicarbonate by an increase in bicarbonate reabsorption by the kidney. The compensatory change will try to restore the pH to normal. However, the compensatory change cannot fully correct a disturbance. Clinically, acid-base disturbance states may be divided into: i. Uncompensated ii. Partially compensated iii. Fully compensated  Metabolic acidosis It is the most common acid base disturbance. In this condition the plasma bicarbonate level is low. Metabolic acidosis may result from 1) Excess production of acids which occurs in diabetes mellitus, starvation, phenyl ketonuria and maple syrup urine disease. 2) Intense muscular exercise may lead to accumulation of lactic acid ie lactic acidosis. 3) Ingestion of mineral acids eg certain drugs. 4) Loss of HCO3 eg in vomiting, diarrhea, loss of pancreatic fluids or upper intestinal contents due to intestinal obstruction. 5) Decreased H+ secretion in kidney eg nephritis. Increased elimination of HCO3 by kidney eg renal failure Metabolic acidosis and Anion Gap There is always a difference between the measured cations and the anions. The unmeasured anions constitute the anion gap. This is due to the presence of protein anions, sulphate, phosphate and organic acids The sum of cations and anions in ECF is always equal, so as to maintain the electrical neutrality. Sodium and potassium together account for 95% of the cations whereas chloride and bicarbonate account for only 86% of the anions. Only these electrolytes are commonly measured. The anion gap is calculated as the difference between (Na+ + K+) and (HCO3– + Cl–). Normally this is about 12 mmol/liter. The anion gap could be High Anion Gap (HAG), Normal Anion Gap (NAG) or Decreased Anion Gap (DAG) Metabolic acidosis Depending on the cause, the anion gap is altered. The types of this are 1. High Anion Gap Metabolic Acidosis (HAGMA) , 2. Normal Anion Gap (NAG), 3. Decreased Anion Gap (DAG)  Metabolic acidosis Metabolic acidosis is compensated by lungs and kidney. Increased respiration eliminates CO2 faster and carbonic acid content diminishes. The renal compensatory mechanism involves excretion of more ammonia and acid phosphates. These compensatory mechanisms may restore pH of blood. If acidosis is not compensated the pH falls and patient may go into coma. Chronic metabolic acidosis cases are treated by administration of sodium lactate or citrate.  Metabolic alkalosis. Metabolic alkalosis is rare. It is due to more bicarbonate in plasma. Causes for metabolic alkalosis are 1. Excessive loss of HCl eg prolonged vomiting eg pyloric obstruction, 2. Ingestion of salts of acids like sodium lactate or citrate and sodium bicarbonate, 3. Excessive production and excretion of ammonia. In metabolic alkalosis, the condition is compensated by pulmonary and renal mechanisms. Pulmonary compensatory mechanism is hypoventilation. Respiratory rate is decreased, CO2 accumulates in plasma and carbonic acid formation increases. At the same time kidney compensates alkalosis by increasing elimination of H CO–3 and decreasing H+ secretion. By the combined action of these organs, the blood pH come back to normal. If metabolic alkalosis is not compensated, tetany develops and convulsive seizures may occur in children.  Respiratory acidosis Respiratory acidosis is due to more plasma PCO2 level. Causes for respiratory acidosis are 1. Depression of respiration (Hypoventilation). Hypoventilation occurs due to excessive dosage of respiratory depressants eg morphine, barbiturates and other. 2. Obstruction to air passage eg pneumonia, emphysema, asthma and tracheal obstruction. Mainly renal mechanism compensate this condition by absorbing more HCO–3 and eliminating more H+ and ammonia in urine.  Respiratory alkalosis this acid base imbalance in which plasma PCO2 level is low in. Respiratory alkalosis may result from Hyperventilation. In hyperventilation, there is Stimulation of respiratory centre in the brain. Typically, It occurs in fever, head injury, anxiety, hysteria, salicylate poisoning, hot climate and high altitude. Kidney compensates this imbalance by elimination of more HCO–3 and decreasing H+ secretion. Causes of Acid-Base Disturbances Stages of compensation Acid-base abnormalities  Laboratory diagnosis of acid base disturbances Determination of the type of acidosis or alkalosis can be made by measuring plasma pH, PCO2 and HCO–3. Various blood parameters in acid base disturbances are given below Clinical presentation of Acid-base Disturbances Thanks EASSAY QUESTION 1. Describe the role of body buffers in maintenance of body fluids pH or acid base balance 2. What is H+ homeostasis, name the systems involved and highlight its clinical relevance 3. Highlight 5 major similarities and differences between the three level of acid base balance 4. How does the body compensated for acid base disorders and what are the limitations 5. Name the major buffer systems and give 10 differences between them