Disturbances of water, pH and electrolyte balance


      Water and electrolytes are subject to a rapid turnover due mainly to dietary intake and renal excretion. Other causes of water (and electrolytes) elimination are faeces, respiration and sweating. Under pathological conditions water and electrolyte loss may occur because of several reasons that include: diarrhoea (e.g. cholera, bacterial enteritis or cholitis); vomiting, inability of the kidney to concentrate the urine (e.g. diabetes mellitus, diabetes insipidus).
      As a general rule, water is absorbed or excreted passively, thus it follows the main extracellular electrolytes, sodium and chloride, and water deficits usually occur together with sodium and chloride deficits. It is important to remark, however, that the intracellular water compartment is usually most affected and that symptoms may occur also under conditions in which the water and electrolyte deficits one may estimate from blood composition are relatively mild. The central nervous system is particularly sensitive to water and electrolyte unbalance and symptoms (e.g. irritability, convulsions, coma) may be severe.

Loss from GI tract: vomiting, diarrhoea, medical practices (e.g.: colostomy)
Loss form the skin: excess sweating
Loss due to hemodialysis or peritoneal dialysis
Acute or chronic renal failure: "salt wasting renal disease
Excess diuretic therapy
Diabetes mellitus
Adrenal disease: Addison disease (glucocorticoid deficiency); hypoaldosteronism

      Measurement of electrolyte concentration in the blood is usually effected by potentiometric methods, using the appropriate electrodes. Other methods are less commonly employed, e.g. flame atomic absorption.
      Hyponatremia is a condition in which sodium losses exceed water losses; decrease of the electrolyte concentrations and hypo-osmolarity follow. This condition may occur in several types of renal failure, in heart failure because of the sodium retention in peripheral oedema, in several endocrine disease and in neoplastic diseases, in the course of diuretic therapy. An uncommon but important cause is the (usually paraneoplastic) syndrome of inappropriate secretion of antidiuretic hormone (ADH). Water intoxication occurs when the plasma osmolarity falls below 240 mOsm/Kg (normal value approx 300 mOsm / Kg).
      Hypernatremia (i.e. increased concentration of serum electrolytes, especially sodium) is typically due to water deprivation or increased and unreplaced water losses (e.g. profuse sweating). It requires prompt rehydration, in severe cases by intravenous administration of glucose isotonic solution.
      Hypokalemia is due to excessive loss of potassium in the urine, faeces or sweat. Renal waste of potassium occcurs in Bartter syndrome (a disease of unknown origin) and in the presence of excess secretion of mineralocorticoid hormones (e.g. because of a benign tumour of the adrenal glands) and in primary disturbances of the kidney involving the proximal and distal tubuli. It may be iatrogenic, due to excess diuretic therapy (e.g. using thiazides or furosemide) and it is always advisable in prolonged diuretic therapy to associate a K-saving diuretic (e.g. spironolactone). Normal potassiemia is 5 mEq/L and in severe depletion (3 mEq/L or less) muscular weakness, and severe cardiac arrythmias, both due to the fact that excitable tissues function is compromised if the intra- extra-cellular potassium gradient is altered. Severe hypokalemia requires (careful !) intravenus administration of potassium.
      Hyperkalemia may occur in type I diabetic patients or in the course of severe acute kidney disease (e.g. glomerulonephritis, acute renal failure), given that the kidney excretes potassium very efficiently and, if functional prevents or corrects this condition. Pseudohyperkalemia is the transient increase of potassium concentration due to release of intracellular potassium by the red or white blood cells or by platelets. True hyperkalemia is a dangerous condition that requires prompt treatment, given that at potassiemia > 6.5 mEq/L severe arrythmias occur and ventricular fibrillation is possible.

      Disturbances of Calcium metabolism. The concentration of calcium in the blood serum in the healthy adult is approx. 2.5 mMoles/L (or 5 mEq/L or 10 mg/dL) and is regulated by two hormones and one vitamin: the Parathyroid hormone (PTH) is a protein secreted by the parathyroid glands; it causes reabsorption of calcium from the skeletal deposits, from the urine and from the gut and causes the calcemia to increase. Thyrocalcitonin is a protein hormone produced by th eparafollicular cells of the thyroid and causes calcium phophate deposition in the bone matrix, opposing the effect of PTH. Vitamin D promotes absorption of calcium from the diet. The serum concentration of calcium is so carefully regulated because this ion is essential for the excitability of nervous and muscular tissues and changes in its concentration cause severe symptoms. Hypocalcemia is an infrequent finding; it may depend on several causes, among which: (i) hypoparathyroidism (often associated to surgical removal of the parathyroid in the course of a thyroidectomy); (ii) vitamin D deficiency (e.g. rickets); (iii) renal tubular disease or renal failure; (iv) acute pancreatitis (calcium chelation by lipolytic products). Clinical symptoms include reduced cerebral function (pseudo-dementia), possibly with psychotic symptoms, and muscular tetany. Hypercalcemia is usually caused by hyperparathyroidism (often of neoplastic origin), and is associated to excessive bone matrix reabsorption (osteoporosis). Other neoplastic diseases, unrelated to the parathyroids, can cause osteolysis with hypercalcemia and osteoporosis, because of secretion of osteoclast activating factors (so called "humoral hypercalcemia of malignancy"). Hypervitaminosis D is another possible cause of hypercalcemia.
      Hypophosphatemia. Calcium phosphate is the main mineral component of the bone tissue, and mobilization of phosphate usually follows that of calcium: e.g. PTH causes reabsorption of both calcium and phosphate. However, in the intracellular milieu and in the diet calcium and phosphate have different distributions and thus changes in the serum concentrations of phosphate does not necessarily follow those of calcium. Hypophosphatemia is not uncommon but rarely severe or even symptomatic. The main cause is reduced renal reabsorption.
      Disturbances of Magnesium metabolism. In the healthy adult the serum magnesium concentration is approx. 2 mEq/L and is regulated mainly at the level of urinary and fecal excretion. Hypomagnesemia may result from prolonged poor dietary intake or reduced intestinal absorption (diarrhoea, malnutrition, etc.) it may be aggravated by some physiological conditions of excess consumption (e.g. lactation). Hypermagnesemia may occur, together with other electrolyte disturbances, in chronic or acute renal failure.


      The alterations of the blood pH are called acidoses (if arterial pH < 7.35) and alkaloses (if arterial pH > 7.45). They are due to abnormal production or excretion of acidic or basic solutes in the serum, and of course are counteracted by the blood buffers. An interesting and well-written tutorial may be found at this link.

BufferConcentration (ionizable group equivalents)pKProduction/excretion
CO2 / bicarbonate 26-28 mM (CO2: 1.2 mM - 40 mmHg) 6.1 Krebs cycle; excreted by the lung (CO2) and the kidney (bicarbonate)
Hemoglobin 5 mM 7.0 (HbO2) - 7.8 (Hb) produced during the red blood cell differentiation; degraded by macrophages
Phosphate 1.2 mM 7.0 in equilibrium with calcium phosphate in the bone matrix; excreted in the urine

      The blood buffers behave in quite a complicated manner. The most concentrated buffer is that composed by carbon dioxide and bicarbonate, whose reactions are as follows:
1) CO2 gas <==> CO2 aq
2) CO2 aq + H2O <==> H2CO3
3) H2CO3 + H2O <==> HCO3- + H3O+
Reaction 1 is the equilibration of carbon dioxide between the aqueous and gaseous phases; in the organism it only occurs at the level of lung capillaries and is governed by the Henry's law, with the absorption coefficient of 0.031 mM/mmHg. Reaction 2 is the hydration of carbon dioxide to carbonic acid; it occurs spontaneously, but is accelerated by the enzyme carbonic anhydrase. The equilibrium strongly favors CO2 over H2CO3 by approximately 700:1. Reaction 3 is the spontaneous dissociation of carbonic acid to hydrogen ion plus bicarbonate, and is governed by the pKa of carbonic acid (close to 4). The further dissociation of bicarbonate to carbonate can be safely neglected because of its high pKa (over 11). The chemical system thus described is complex but can be simplified by combining equations 2 and 3 and neglecting the concentration of carbonic acid:
4) CO2 aq + 2 H2O [ <==> H2CO3 + H2O ] <==> HCO3- + H3O+
The equilibrium constant of eq. 4 equals the product of the equilibrium constants for reactions 2 and 3, and corresponds to a pKa of 6.1 at 37 C (the pKa being - log Ka). If one adopts this simplification, the physiologically relevant equilibria can be described with only two equations (eq.1 and eq.4), that can be combined into just one Henderson and Hasselbalch equation:
pH = 6.1 + log ([HCO3-] / 0.031 P CO2)
where [HCO3-] is expressed in mMoles/L and P CO2 in mmHg. Given that pH is approx. 7.4, the ratio [HCO3-] / [CO2] equals 20; in numbers [HCO3-] is approx. 26 mM and [CO2] approx. 1.2 mM (corresponding to a P CO2 of 40 mmHg). The above values hold for the arterial blood of healthy adults.

      At a ratio of 20:1 for the [salt]/[acid] fraction, the bicarbonate buffer is not working in the blood under condition of optimal buffer capacity. However, the tissues produce CO2, whereas the lungs excrete the gas and the kidneys excrete bicarbonate. As a consequence the [salt]/[acid] ratio can be regulated and maintained constant, giving very high buffering capacity to the system. The lung alone eliminates some 20 moles CO2 / day, a massive amount.

      The second most concentrated buffer in human blood is given by hemoglobin (Hb). Hb's concentration is in the order of 14 mg/dL (over 30 mg/dL in the erythrocytes), which corresponds to 9 mM (calculated on a heme basis). Since the stoichiometric ratio between bound O2 and hydrogen ions exchanged is 0.5 (2 hydrogen ions per tetramer) the effective concentration of the Hb buffer is 4.5 mM. OxyHb behaves as a weak acid with pKa=7.0; deoxyHb behaves as a weak acid with pKa=7.8. Since the pKa of oxyHb is lower than the blood pH, oxyHb is predominantly de-protonated in the blood (with a ratio Cs/Ca=3:1); the opposite occurs for deoxyHb. As a consequence in the lung Hb combines with O2 and releases hydrogen ions, whereas in the tissues it releases O2 and binds hydrogen ions. The pH changes induced in the blood by Hb are of opposite sign with respect to those due to bicarbonate (which prevail due to the higher concentration of the latter): this contributes to minimize the pH changes due to respiration, as shown in the following scheme:

      Disturbances of the blood pH and buffer concentrations are called respiratory if caused by altered functioning of the lungs, metabolic otherwise. It is important to stress that disfunction of the lung can be to some extent compensated by the kidney and vice versa; thus each organ tends to oppose the disfunction of the other. The key features of the different forms of acidosis and alkalosis are as follows:

AbnormalityBlood pH (normal values: venous 7.36; arterial 7.4)Pressure of blood CO2 (normal values: venous 43 mmHg; arterial 40 mmHg) total CO2 (i.e. CO2 + bicarbonate; normal values 21-28 mM)
Metabolic alkalosis increased increased increased
Metabolic acidosis decreased decreased (compensatory) decreased
Respiratory alkalosis increased decreased (primary) decreased (compensatory)
Respiratory acidosis decreased increased (primary) increased (compensatory)

      The following remarks will help explaining the above table: (i) Pressure of blood CO2 indicates the concentration of pure CO2, i.e. it does not include bicarbonate; total CO2 is dominated essentially by bicarbonate ion concentration (the ratio [HCO3-] / [CO2] being approximately 20). (ii) The metabolism produces mainly acids (CO2 and lactic acid, acetoacetic acid, etc.): thus acidosis is more frequent, more severe and more varied; alkalosis is often caused by excess loss of acidic fluids (e.g. severe vomiting). (iii) Acids can be volatile (CO2), excreted by the lung at a very fast rate, and non-volatile (lactic acid, acetoacetic acid), excreted mainly by the kidney at a slower rate. (iv) Metabolic alkalosis is most often due to loss of acids (e.g. vomiting), excess intake of alkaline substances (e.g. gastric antiacids, bicarbonate), and diuretics; it is corrected mainly by the kidney that excretes the excess bicarbonate. (v) Metabolic acidosis, caused by overproduction of non-volatile acids (e.g. diabetic ketoacidosis) or by their impaired renal excretion, stimulates respiration that excretes the volatile acid (CO2): hence compensatory hyperpnea and hypocapnia (reduced P CO2). It is interesting to remark that pulmonary correction of metabolic acidosis is more effective than of metabolic alkalosis because respiration frequency can be increased to a more significant extent than it can be decreased. (vi) Respiratory alkalosis is a consequence of hyperventilation (loss of CO2), but this only occurs in some types of CNS disturbances or under unusual environmental conditions (e.g. muscular effort at high altitude, where atmospheric P O2 is decreased - air hunger). (vii) Respiratory acidosis is a common consequence of impaired gas exchanges (e.g. depression of respiratory centers in the CNS, insufficient mechanical ventilation in polyomyelitis or tuberculosis, or ventilation perfusion imbalance in chronic obstructive pulmonary disease, emphysema, etc.).

      The hemogas analysis is the measurement of the pH and the concentrations and partial pressures of O2 and CO2 in a sample of the patient's blood drawn in a gas tight syringe. The measure is usually effected by means of potentiometric methods, used gas-specifci electrodes.
      As a general rule, a hemogas analysis will indicate if an abnormality is present and will give some indication of its possible cause; the fundamental indications are as follows:
Normal to diminished pH and diminished total CO2 = metabolic acidosis with respiratory compensation (e.g. diabetic ketoacidosis).
Normal to diminished pH and increased total CO2 = respiratory acidosis; metabolic component unassessed (severely compromised gas exchange in the lung or neurological depression of ventilation).
Normal to increased pH and increased total CO2 = metabolic alkalosis; respiratory component unassessed (e.g. vomiting).
Normal to increased pH and decreased total CO2 = respiratory alkalosis; metabolic component unassessed (uncommon; neurological hyperpnea?).
      Once the presence of an acid-base imbalance has been detected, whose diagnosis is not obvious, more refined measurements are indicated in order to separate the respiratory and metabolic cotributions. Special attention is required in the elderly given that metabolic and respiratory conditions of equal or opposite sign may coexist (e.g. pulmonary emphysema, causing chronic respiratory acidosis, may be present together with vomiting, causing acute metabolic alkalosis, or with diabetes, causing chronic metabolic acidosis). Several clinical concepts (and measurements) have been developed to discriminate the metabolic and respiratory components of blood buffers inbalance, as listed below:

Standard pH, historically the first concept introduced to rationalize complex deviations from the healthy conditions of blood buffers balance was introduced into the clinical practice by Hasselbalch in 1916. It is the pH of the patient's arterial blood measured under standard conditions (P CO2=40 mmHg, hemoglobin fully saturated with O2, T = 37 C). Essentially, the use of standard conditions has the effect of reversing the compensatory effect of respiration and thus to make more evident the eventual presence of a metabolic component in the pH unbalance.
Standard bicarbonate, a concept developed by Astrup and Siggaard Andersen in 1957, is the concentration of bicarbonate one measures when a sample of arterial blood is equilibrated under standard conditions. The rationale of this procedure is that of setting the concentration of one of the components of the major blood buffer (CO2) and determinining that of the other. The number one measures is not the same one would obtain in a standard hemogas analysis, given that part of the bicarbonate originally present in the blood may be lost as CO2 (if the P CO2 was higher than 40 mmHg) or part of the gas may be absorbed and converted to bicarbonate. The standard bicarbonate measures the metabolic component of the acid-base balance of the blood and corrects for respiratory compensation.
Base excess, again by Astrup and Siggaard Andersen, is the amount of alkali (or acid, in which case the resulting value is < zero) required to restore the normal pH of a blood sample equilibrated under standard conditions.
Anion gap, is the difference between the concentrations of (sodium + potassium) and (chloride + bicarbonate). The normal value is 10-15 mM. An increased anion gap may provide a gross indication of metabolic acidosis, i.e. of the presence of unmeasured negatively charged ions (e.g. lactic acid or acetoacetic acid).

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