Functions of magnesium Regulator or cofactor for a number of enzymes, including phosphatases and 26g phosphokinases involved in energy storage and utilisation NatKt-ATPase: the sodium pump Protein synthesis denyl cyclase: involved in cyclic AMP production lycolytic enzymes Regulator of slow calcium channels. Mediates hypocalcaemia-induced parathyroid hormone (PTH)release and is involved in maintenance of end-organ sensitivity to PTH and vitamin d Acts as an endogenous calcium antagonist, competing with calcium for intra- cellular entry via cytoplasmic channels and by activating cell membrane pumps that actively pump out calcium. Muscle contraction Membrane stabilisation: regulator of membrane excitability Structural component of cell membranes and the cytoskeleton. Cell division acntesaenla This is associated with Reduced intake Renal losses: acute tubular necrosis; post-obstructive diuresis; diuretics; renal tubular acidosis Gastrointestinal losses: malabsorption; secretory diarrhoea. Redistribution into cells BIBLIOGRAPHY King, L S, Agre, P Pathophysiology of the aquaporin water channels. Annu. Rev. Physiol. 58(1996),619-48
Functions of magnesium Regulator or cofactor for a number of enzymes, including phosphatases and phosphokinases involved in energy storage and utilisation: NaþKþ -ATPase: the sodium pump Ca2 þ-ATPase Proton pumps Protein synthesis Adenyl cyclase: involved in cyclic AMP production Glycolytic enzymes Regulator of slow calcium channels. Mediates hypocalcaemia-induced parathyroid hormone (PTH) release and is involved in maintenance of end-organ sensitivity to PTH and vitamin D. Acts as an endogenous calcium antagonist, competing with calcium for intracellular entry via cytoplasmic channels and by activating cell membrane pumps that actively pump out calcium. Neuromuscular transmission. Muscle contraction. Membrane stabilisation: regulator of membrane excitability. Structural component of cell membranes and the cytoskeleton. Cell division. Hypomagnesaemia This is associated with Reduced intake. Renal losses: acute tubular necrosis; post-obstructive diuresis; diuretics; renal tubular acidosis. Gastrointestinal losses: malabsorption; secretory diarrhoea. Redistribution into cells. BIBLIOGRAPHY King, L. S., Agre, P. Pathophysiology of the aquaporin water channels. Annu. Rev. Physiol., 58 (1996), 619–48. Water and electrolyte balance 30
Acid-base balance 3 ■ Introduction Acid-base homeostasis is essential to allow normal tissue and organ system function. Specifically, intracellular enzyme systems require an appropriate pH for maintenance of activity. Extracellular fluid pH is normally maintained around a level of 7.35-7.45. This is achieved through reversible chemical buffer systems and through homeostatic responses mediated by the lungs and the kidneys The pH scale This scale for measurement was introduced by Sorenson in 1909 to facilitate deal g with hydrogen ion concentrations. This avoids the need to deal with negative indices and to accommodate the very wide range of H* and OH solutions that untered in acid-base reactions when measured in The pH is the negative logarithm to base 10 of H ion activity. Acidic solutions rays have a pH less than 7.0. The pH scale is an exponential scale, where a 0.3 unit fall in pH reflects a doubling of hydrogen ion concentration I Bronsted-Lowry definitions An acid is a proton donor to a base a base is a proton acceptor from an acid cid-base reactions are proton transfer reactions. An acid losing a proton to a base forms a base itself. a base accepting a proton forms an acid. Every acid has its conjugate base, and every base its conjugate acid An acid-base reaction is a dynamic equilibrium between two conjugate acid-base pairs
Acid–base balance & Introduction Acid–base homeostasis is essential to allow normal tissue and organ system function. Specifically, intracellular enzyme systems require an appropriate pH for maintenance of activity. Extracellular fluid pH is normally maintained around a level of 7.35–7.45. This is achieved through reversible chemical buffer systems and through homeostatic responses mediated by the lungs and the kidneys. & The pH scale This scale for measurement was introduced by Sorenson in 1909 to facilitate dealing with hydrogen ion concentrations. This avoids the need to deal with negative indices and to accommodate the very wide range of Hþ and OH solutions that are encountered in acid–base reactions, when measured in nanomoles per litre. The pH is the negative logarithm to base 10 of H ion activity. Acidic solutions always have a pH less than 7.0. The pH scale is an exponential scale, where a 0.3 unit fall in pH reflects a doubling of hydrogen ion concentration. & Bronsted–Lowry definitions An acid is a proton donor to a base. A base is a proton acceptor from an acid. Acid–base reactions are proton transfer reactions. An acid losing a proton to a base forms a base itself. A base accepting a proton forms an acid. Every acid has its conjugate base, and every base its conjugate acid. An acid–base reaction is a dynamic equilibrium between two conjugate acid–base pairs. 3 31
Table 3.1. Relation of ph to hydrogen ion activity H ion activity(nmol/D) 6777777777 20 096002506 70 A substance that can act as both an acid and as a base is called an amphoteric substance. Water has amphoteric properties. The strength of an acid is specified by its dissociation constant. A strong acid has a dissociation constant greater than that of H3o* and is almost completely ionised in aqueous solutions. A weak acid has a dissociation constant less than that of H3o*(I in aqueous solutions) and is only partially ised in aqueous solutions. Factors affecting body acid-base balance Body acid production creates a hydrogen ion load, mainly from oxidation of proteins, carbohydrates and other organic molecules. This amounts to 50-100 mmol of hydrogen ions per day in the adult, and includes volatile acids(such as carbonic acid) and fixed acids(such as phosphoric and sulphuric acids The body is a net producer of acid Acids produced by the body consist of carbonic acid(respiratory acid), and various metabolic acids, including sul- phuric acid, phosphoric acid, lactic acid, citric acid, ammonium ions and ketone bodies(acetoacetic acid and beta-hydroxybutyric acid Buffering of the hydrogen ion load by extracellular and intracellular buffer Excretion of the hydrogen ion load due to loss of bicarbonate Respiratory regulation of volatile acid excretion by variations in alveolar ventilation affecting the excretion of carbon dioxide. Respiratory compensation is achieved by control of alveolar ventilation through the intervention of Peripheral chemoreceptors in the carotid and aortic bodies; Central chemoreceptors in the ventro-lateral medulla
A substance that can act as both an acid and as a base is called an amphoteric substance. Water has amphoteric properties. The strength of an acid is specified by its dissociation constant. A strong acid has a dissociation constant greater than that of H3Oþ and is almost completely ionised in aqueous solutions. A weak acid has a dissociation constant less than that of H3Oþ(1 in aqueous solutions) and is only partially ionised in aqueous solutions. & Factors affecting body acid–base balance * Body acid production creates a hydrogen ion load, mainly from oxidation of proteins, carbohydrates and other organic molecules. This amounts to 50–100 mmol of hydrogen ions per day in the adult, and includes volatile acids (such as carbonic acid) and fixed acids (such as phosphoric and sulphuric acids). The body is a net producer of acid. Acids produced by the body consist of carbonic acid (respiratory acid), and various metabolic acids, including sulphuric acid, phosphoric acid, lactic acid, citric acid, ammonium ions and ketone bodies (acetoacetic acid and beta-hydroxybutyric acid). * Buffering of the hydrogen ion load by extracellular and intracellular buffer systems. * Excretion of the hydrogen ion load due to loss of bicarbonate. Respiratory regulation of volatile acid excretion by variations in alveolar ventilation affecting the excretion of carbon dioxide. Respiratory compensation is achieved by control of alveolar ventilation through the intervention of: Peripheral chemoreceptors in the carotid and aortic bodies; Central chemoreceptors in the ventro-lateral medulla. Table 3.1. Relation of pH to hydrogen ion activity pH H ion activity(nmol/l) 6.90 126 7.0 100 7.10 79 7.20 63 7.30 50 7.40 40 7.50 32 7.60 25 7.70 20 7.80 16 Acid–base balance 32
Acid production by the body Volatile: carbonic acid: 15-20 moles per da Non-volatile or fixed: 80 moles per day Phosphoric acid: oxidation of phosphoproteins, phospholipids hosphoglycerides and nucleic acids Sulphuric acid: oxidation of methionine and cysteine Buffers a buffer is a conjugate acid-base pair. It is represented by a mixture of sub stances in aqueous solution, usually a weak acid and its corresponding sodium salt(a conjugate base), which acts as a proton acceptor for the corresponding weak acid. A buffer system can resist changes in H ion concentration when small amounts of strong(completely ionised) acids or bases are added. It can thus be considered as a pair of substances that can donate or accept(reversibly bind) H ions in order to minimise changes in H ion concentration. The bodys buffer systems are primarily involved with the neutralisation of acids. Activity of the buffer system depends on a simple equilibrium constant deter- mined by the unique dissociation constant (pk of the buffer. The system is most effective at a concentration identical to that of the concentration of its conjugate base or acid In a mixture of buffer pairs, each buffer pair is in equilibrium with all others at given pH. By controlling one buffer pair, all pairs in equilibrium with this pair can be controlled, minimising changes in pH. The larger the concentration of acid and conjugate base (or base and conjugate acid) the greater the buffering capacity of the system. Henderson-Hasselbach equation The pK of an acid is numerically equal to the ph of the solution when the molar concentrations of the acid and its conjugate base are equal. CO2+H2O=H2 CO3=HT+HCO3 H2 CO3)=K(H*)(HCO3) I/H*=K(HCO3 )(H2 CO3 Taking logarithms on both sides pH=pK logl0(HCO3 )(H2CO3) By the equation, pH remains constant provided the ratio ofHCO3 /CO2 remains constant. The equation can be used to calculate the ph of a buffer solution, the amount of acid or salt required to make a buffer solution of desired pH, or the effect on the pH of a buffer solution when a small amount of acid or base is added
Buffers A buffer is a conjugate acid–base pair. It is represented by a mixture of substances in aqueous solution, usually a weak acid and its corresponding sodium salt (a conjugate base), which acts as a proton acceptor for the corresponding weak acid. A buffer system can resist changes in H ion concentration when small amounts of strong (completely ionised) acids or bases are added. It can thus be considered as a pair of substances that can donate or accept (reversibly bind) H ions in order to minimise changes in H ion concentration. The body’s buffer systems are primarily involved with the neutralisation of acids. Activity of the buffer system depends on a simple equilibrium constant determined by the unique dissociation constant (pK) of the buffer. The system is most effective at a concentration identical to that of the concentration of its conjugate base or acid. In a mixture of buffer pairs, each buffer pair is in equilibrium with all others at a given pH. By controlling one buffer pair, all pairs in equilibrium with this pair can be controlled, minimising changes in pH. The larger the concentration of acid and conjugate base (or base and conjugate acid) the greater the buffering capacity of the system. Henderson–Hasselbach equation The pK of an acid is numerically equal to the pH of the solution when the molar concentrations of the acid and its conjugate base are equal. CO2 þ H2O ¼ H2CO3 ¼ Hþ þ HCO3 (H2CO3) ¼ K(Hþ)(HCO3 ) I/Hþ ¼ Kþ(HCO3 )(H2CO3) Taking logarithms on both sides, pH ¼ pK + log10(HCO3 )(H2CO3) By the equation, pH remains constant provided the ratio of HCO3 /CO2 remains constant. The equation can be used to calculate the pH of a buffer solution, the amount of acid or salt required to make a buffer solution of desired pH, or the effect on the pH of a buffer solution when a small amount of acid or base is added. Acid production by the body Volatile: carbonic acid: 15–20 moles per day Non-volatile or fixed: 80 moles per day Phosphoric acid: oxidation of phosphoproteins, phospholipids, phosphoglycerides and nucleic acids Sulphuric acid: oxidation of methionine and cysteine Factors affecting body acid–base balance 33
Physiological buffer systems Buffer systems comprise the first line of the bodys response to changes in pH They include: NaHCO3/H,CO3 Haemoglobin and its potassium salt Plasma proteins: plasma protein and sodium proteinate Phosphate(H2 POA/HPO 2 Extracellular fluid and cerebrospinal fluid HCO3/CO Proteins Phosphate (H2POA /HPO4) Intracellular fluid Proteins Organic phosphates HCO3 /CO2 Buffering capacity The buffering capacity of a buffer pair depends on their total concentration rather than on their ratio. The pH of a solution containing a buffer pair depends they on their ratio, and concentrating or diluting the system does not alter ntracellular proteins account for 75% of the bodys buffering capacity. The protein and phosphate systems buffer changes in carbon dioxide. The bicarbo- nate system can only buffer metabolic acids. Bicarbonate and non-bicarbonate buffer systems are in equilibrium with each other The total buffering capacity of blood at a pH of 7. 4 and constant paCO2 is about 5 mmol/. The total buffer base is the sum of the concentrations of all buffer anions that can take up Ht ions, and is about 45 mmol/l A scheme for the assessment and management of acid-base abnormalities e Look at the pao Correct rapidly if abnormal, with supplemental oxygen and ventilatory support where appropriate. The paO2 needs to be interpreted in relation to the inspired oxygen concentration, and calculation of the alveolar-arterial gradient may be necessary to unmask hypoxia under these circumstances
Physiological buffer systems Buffer systems comprise the first line of the body’s response to changes in pH. They include: Blood NaHCO3/H2CO3 Haemoglobin and its potassium salt Plasma proteins: plasma protein and sodium proteinate Phosphate(H2PO4 /HPO4 2) Extracellular fluid and cerebrospinal fluid HCO3 /CO2 Proteins Phosphate(H2PO4 /HPO4 2) Intracellular fluid Proteins Phosphate Organic phosphates HCO3 /CO2 Buffering capacity The buffering capacity of a buffer pair depends on their total concentration rather than on their ratio. The pH of a solution containing a buffer pair depends wholly on their ratio, and concentrating or diluting the system does not alter the pH. Intracellular proteins account for 75% of the body’s buffering capacity. The protein and phosphate systems buffer changes in carbon dioxide. The bicarbonate system can only buffer metabolic acids. Bicarbonate and non-bicarbonate buffer systems are in equilibrium with each other. The total buffering capacity of blood at a pH of 7.4 and constant paCO2 is about 75 mmol/l. The total buffer base is the sum of the concentrations of all buffer anions that can take up Hþ ions, and is about 45 mmol/l. A scheme for the assessment and management of acid–base abnormalities * Look at the paO2. Correct rapidly if abnormal, with supplemental oxygen and ventilatory support where appropriate. The paO2needs to beinterpretedin relation to the inspired oxygen concentration, and calculation of the alveolar–arterial gradient may be necessary to unmask hypoxia under these circumstances. Acid–base balance 34