Fluid, Electrolyte and Acid–Base Balance

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Fluid, Electrolyte and Acid–Base Balance

The realization that the enzyme systems and metabolic processes responsible for the maintenance of cellular function are dependent on an environment with stable electrolyte and hydrogen ion concentrations led Claude Bernard to describe the ‘milieu interieur’ over 100 years ago. Complex homeostatic mechanisms have evolved to maintain the constancy of this internal environment and thus prevent cellular dysfunction.

BASIC DEFINITIONS

Osmosis refers to the movement of solvent molecules across a membrane into a region in which there is a higher concentration of solute. This movement may be prevented by applying a pressure to the more concentrated solution – the effective osmotic pressure. This is a colligative property; the magnitude of effective osmotic pressure exerted by a solution depends on the number rather than the type of particles present.

The amounts of osmotically active particles present in solution are expressed in osmoles. One osmole of a substance is equal to its molecular weight in grams (1 mol) divided by the number of freely moving particles which each molecule liberates in solution. Thus, 180 g of glucose in 1 L of water represents a solution with a molar concentration of 1 mol L^–1 and an osmolarity of 1 osmol L^–1. Sodium chloride ionizes in solution and each ion represents an osmotically active particle. Assuming complete dissociation into Na⁺ and Cl^–, 58.5 g of NaCl dissolved in 1 L of water has a molar concentration of 1 mol L^–1 and an osmolarity of 2 osmol L^–1. In body fluids, solute concentrations are much lower (mmol L^–1) and dissociation is incomplete. Consequently, a solution of NaCl containing 1 mmol L^–1 contributes slightly less than 2 mosmol L^–1.

The term osmolality refers to the number of osmoles per unit of total weight of solvent, whereas osmolarity refers to the number of osmoles per litre of solvent. Osmolality (unlike osmolarity), is not affected by the volume of various solutes in solution. Confusion regarding the apparently interchangeable use of the terms osmolarity (measured in osmol L^–1) and osmolality (measured in osmol kg^–1) is caused by their numerical equivalence in body fluids; plasma osmolarity is 280–310 mosmol L^–1 and plasma osmolality is 280–310 mosmol kg^–1. This equivalence is explained by the almost negligible solute volume contained in biological fluids and the fact that most osmotically active particles are dissolved in water, which has a density of 1 (i.e. osmol L^–1 = osmol kg^–1). As the number of osmoles in plasma is estimated by measurement of the magnitude of freezing point depression, the more accurate term in clinical practice is osmolality.

Cations (principally Na⁺) and anions (Cl^– and ) are the major osmotically active particles in plasma. Glucose and urea make a smaller contribution. Plasma osmolality (P_OSM) may be estimated from the formula:

Osmolality is a chemical term and may be confused with the physiological term, tonicity. This term is used to describe the effective osmotic pressure of a solution relative to that of plasma. The critical difference between osmolality and tonicity is that all solutes contribute to osmolality, but only solutes that do not cross the cell membrane contribute to tonicity. Thus, tonicity expresses the osmolal activity of solutes restricted to the extracellular compartment, i.e. those which exert an osmotic force affecting the distribution of water between intracellular fluid (ICF) and extracellular fluid (ECF). As urea diffuses freely across cell membranes, it does not alter the distribution of water between these two body fluid compartments and does not contribute to tonicity. Other solutes that contribute to plasma osmolality but not tonicity include ethanol and methanol, both of which distribute rapidly throughout the total body water. In contrast, mannitol and sorbitol are restricted to the ECF and contribute to both osmolality and tonicity. The tonicity of plasma may be estimated from the formula:

COMPARTMENTAL DISTRIBUTION OF TOTAL BODY WATER

The volume of total body water (TBW) may be measured using radioactive dilution techniques involving either deuterium or tritium, both of which cross all membranes freely and equilibrate rapidly with hydrogen atoms in body water. Such measurements show that approximately 60% of lean body mass (LBM) is water in the average 70 kg male adult. As fat contains little water, females have proportionately less TBW (55%) relative to LBM. TBW decreases with age, decreasing to 45–50% in later life.

The distribution of TBW between the main body compartments is illustrated in Figure 12.1. One-third of TBW is contained in the extracellular fluid volume (ECFV) and two-thirds in the intracellular fluid volume (ICFV). The ECFV is subdivided further into the interstitial and intravascular compartments. In addition to the absolute volumes of each compartment, Figure 12.1 shows the relative size of each compartment compared with body weight.

FIGURE 12.1 Distribution of total body water in a 70 kg male, and related to body weight (BW).

SOLUTE COMPOSITION OF BODY FLUID COMPARTMENTS

Extracellular Fluid

The capillary endothelium behaves as a freely permeable membrane to water, cations, anions and many soluble substances such as glucose and urea (but not protein). As a result, the solute compositions of interstitial fluid and plasma are similar. Each contains sodium as the principal cation and chloride as the principal anion. Protein behaves as a non-diffusible anion and is present in a higher concentration in plasma. The concentration of Cl^– is slightly higher in interstitial fluid in order to maintain electrical neutrality (Donnan equilibrium).

Intracellular Fluid

This differs from ECF in that the principal cation is potassium and the principal anion is phosphate. In addition, there is a high protein content. In contrast to the capillary endothelium, the cell membrane is permeable selectively to different ions and freely permeable to water. Thus, equalization of osmotic forces occurs continuously and is achieved by the movement of water across the cell membrane. The osmolalities of ICF and ECF at equilibrium must be equal. Water moves rapidly between ICF and ECF to eliminate any induced osmolal gradient. This principle is fundamental to an understanding of fluid and electrolyte physiology.

Figure 12.2 shows the solute composition of the main body fluid compartments. Although the total concentration of intracellular ions exceeds that of extracellular ions, the numbers of osmotically active particles (and thus the osmolalities) are the same on each side of the cell membrane (290 mosmol kg^–1 of solution).

FIGURE 12.2 Principal solute composition of body fluid compartments. All concentrations are expressed in mmol L^–1.

WATER HOMEOSTASIS

Normal day-to-day fluctuations in TBW are small (< 0.2%) because of a fine balance between input, controlled by the thirst mechanisms, and output, controlled mainly by the renal–ADH (antidiuretic hormone) system.

The principal sources of body water are ingested fluid, water present in solid food and water produced as an end-product of metabolism. Intravenous fluids are another common source in hospital patients. Actual and potential outlets for water are classified conventionally as sensible and insensible losses. Insensible losses emanate from the skin and lungs; sensible losses occur mainly from the kidneys and gastrointestinal tract. Figure 12.3 depicts the daily water balance in a 70 kg adult in whom input and output balance. It should be noted that sources of potential loss are not evident in this diagram. For example, over 5 L of fluid are secreted daily into the gut in the form of saliva, bile, gastric juices and succus entericus, yet only 100 mL of fluid is present in faeces. This illustrates the potential that exists for significant fluid loss in the presence of disease.

FIGURE 12.3 Daily water balance. Input and output in millilitres.

PRACTICAL FLUID BALANCE

Calculation of the daily prescription of fluid is an arithmetic exercise to balance the input and output of water and electrolytes.

Table 12.1 shows the electrolyte contents of five intravenous solutions used commonly in the United Kingdom. These solutions are adequate for most clinical situations. Two self-evident but important generalizations may be made regarding solutions for intravenous infusion.

TABLE 12.1

Electrolyte Contents of Commonly Used Intravenous Fluids

Rule 1: All infused Na⁺ remains in the ECF; Na⁺ cannot gain access to the ICF because of the sodium pump. Thus, if saline 0.9% is infused, all Na⁺ remains in the ECF. As this is an isotonic solution, there is no change in ECF osmolality and therefore no water exchange occurs across the cell membrane. Thus, saline 0.9% expands ECFV only. However, if saline 0.45% is given, ECF osmolality decreases; this causes a shift of water from ECF to ICF. If saline 1.8% is administered, all Na⁺ remains in the ECF, its osmolality increases and water moves from ICF to ECF to maintain osmotic equality.

Rule 2: Water without sodium expands the TBW. After infusion of a solution of glucose 5%, the glucose enters cells and is metabolized. The infused water enters both ICF and ECF in proportion to their initial volumes.

Table 12.2 illustrates the results of infusion of 1 L of saline 0.9%, saline 0.45% or glucose 5% in a 70 kg adult.

TABLE 12.2

Compartmental Expansion Resulting from Infusion of 1 L of Saline 0.9%, Saline 0.45% or Glucose 5%

Assessment of daily fluid requirements may be allocated usefully into three processes:

normal maintenance needs

abnormal losses resulting from the underlying pathology

correction of pre-existing deficits.

Normal Maintenance Needs

Water. Regardless of the disease process, water and electrolyte losses occur in urine and as evaporative losses from skin and lungs. It is evident from Figure 12.3 that a normothermic 70 kg patient with a normal metabolic rate may lose 2500 mL of water per day. Allowing for a gain of 400 mL from water of metabolism, this hypothetical patient needs about 2000 mL day^–1 of water. As a rule of thumb, a volume of 30–35 mL kg^–1 day^–1 of water is a useful estimate for daily maintenance needs.

Sodium. The normal requirement is 1 mmol kg^–1 day^–1 (50–80 mmol day^–1) for adults.

Potassium. The normal requirement is 1 mmol kg^–1 day^–1 (50–80 mmol day^–1) for adults.

Thus, a 70 kg patient requires daily provision of 2000–2500 mL of water and approximately 70 mmol each of Na⁺ and K⁺. This could be administered as one of the following:

2000 mL of glucose 5% + 500 mL of saline 0.9%

2500 mL of glucose 4%/saline 0.18%; plus potassium as KCl, 1 g (13 mmol) added to each 500 mL of fluid.

SODIUM AND POTASSIUM

Sodium Balance

Daily ingestion amounts to 50–300 mmol. Losses in sweat and faeces are minimal (approximately 10 mmol day^–1) and the kidney makes final adjustments. Urine sodium excretion may be as little as 2 mmol day^–1 during salt restriction or may exceed 700 mmol day^–1 after salt loading. Sodium balance is related intimately to ECFV and water balance.

Disorders of Sodium/Water Balance

Hypernatraemia

Hypernatraemia is defined as a plasma sodium concentration of more than 150 mmol L^–1 and may result from pure water loss, hypotonic fluid loss or salt gain. In the first two conditions, ECFV is reduced, whereas salt gain is associated with an expanded ECFV. For this reason, the clinical assessment of volaemic status is important in the diagnosis and management of hypernatraemic states. The common causes of hypernatraemia are summarized in Table 12.3. The abnormality common to all hypernatraemic states is intracellular dehydration secondary to ECF hyperosmolality. Primary water loss resulting in hypernatraemia may occur during prolonged fever, hyperventilation or severe exercise in hot, dry climates. However, a more common cause is the renal water loss that occurs when there is a defect in either the production or release of ADH (cranial diabetes insipidus) or an abnormality in response to ADH (nephrogenic diabetes insipidus).

TABLE 12.3

Causes of Hypernatraemia

The administration of osmotic diuretics results temporarily in plasma hyperosmolality. An osmotic diuresis may occur also in hyperglycaemia. During an osmotic diuresis, the solute causing the diuresis (e.g. glucose, mannitol) constitutes a significant fraction of urine solute, and the sodium content of the urine becomes hypotonic relative to plasma sodium. Thus, osmotic diuretics cause hypotonic urine losses which may result in hypernatraemic dehydration.

Hypertonic dehydration may occur also in paediatric patients. Diarrhoea, vomiting and anorexia lead to loss of water in excess of solute (hypotonic loss). Concomitant fever, hyperventilation and the use of high-solute feeds may combine to exaggerate the problem. ECFV is maintained by movement of water from ICF to ECF to equalize osmolality, and clinical evidence of dehydration may not be apparent until 10–15% of body weight has been lost. Rehydration must be undertaken gradually to prevent the development of cerebral oedema.

Measurement of urine and plasma osmolalities and assessment of urine output help in the diagnosis of hypernatraemic, volume-depleted states. If urine output is low and urine osmolality exceeds 800 mosmol kg^–1, then both ADH secretion and the renal response to ADH are present. The most likely causes are extrarenal water loss (e.g. diarrhoea, vomiting or evaporation) or insufficient intake. High urine output and high urine osmolality suggest an osmotic diuresis. If urine osmolality is less than plasma osmolality, reduced ADH secretion or impairment of the renal response to ADH should be suspected; in both cases, urine output is high.

Usually, hypernatraemia caused by salt gain is iatrogenic in origin. It occurs when excessive amounts of hypertonic sodium bicarbonate are administered during resuscitation or when isotonic fluids are given to patients who have only insensible losses. Treatment comprises induction of a diuresis with a loop diuretic if renal function is normal; urine output is replaced in part with glucose 5%. Dialysis or haemofiltration is necessary in patients with renal dysfunction.

Consequences of Hypernatraemia: The major clinical manifestations of hypernatraemia involve the central nervous system. Severity depends on the rapidity with which hyperosmolality develops. Acute hypernatraemia is associated with a prompt osmotic shift of water from the intracellular compartment, causing a reduction in cell volume and water content of the brain. This results in increased permeability and even rupture of the capillaries in the brain and subarachnoid space. The patient may present with pyrexia (a manifestation of impaired thermoregulation), nausea, vomiting, convulsions, coma and virtually any type of focal neurological syndrome. The mortality and long-term morbidity of sustained hypernatraemia (Na⁺ > 160 mmol L^–1 for over 48 h) is high, irrespective of the underlying aetiology. In many cases, the development of hypernatraemia can be anticipated and prevented, e.g. cranial diabetes insipidus associated with head injury, but in situations where preventative strategies have failed, treatment should be instituted without delay.

Treatment of Hypernatraemia: The magnitude of the water deficit can be estimated from the measured plasma sodium concentration and calculated total body water:

Thus, in a 75 kg patient with a serum sodium of 170 mmol L^–1:

For hypernatraemic patients without volume depletion, 5% glucose is sufficient to correct the water deficit. However, the majority of hypernatraemic patients are frankly hypovolaemic and intravenous fluids should be prescribed to repair both the sodium and the water deficits. Regardless of the severity of the condition, isotonic saline is the initial treatment of choice in the volume-depleted, hypernatraemic patient, as even this fluid is relatively hypotonic in patients with severe hypernatraemia. When volume depletion has been corrected, further repair of any water deficit may be accomplished with hypotonic fluids. Fluid therapy should be prescribed with the intention of correcting hypernatraemia over a period of 48–72 h to prevent the onset of cerebral oedema.

Hyponatraemia

This is defined as a plasma sodium concentration of less than 135 mmol L^–1. Hyponatraemia is a common finding in hospital patients. It may occur as a result of water retention, sodium loss or both; consequently, it may be associated with an expanded, normal or contracted ECFV. As in hypernatraemia, the state of ECFV is important in determining the cause of the electrolyte imbalance.

As plasma osmolality decreases, an osmolal gradient is created across the cell membrane and results in movement of water into the ICF. The resulting expansion of brain cells is responsible for the symptomatology of hyponatraemia or ‘water intoxication’: nausea, vomiting, lethargy, weakness and obtundation. In severe cases (plasma Na⁺ < 115 mmol L^–1), seizures and coma may result.

A scheme depicting the causes of hyponatraemia is shown in Figure 12.4. True hyponatraemia must be distinguished from pseudohyponatraemia. Sodium ions are present only in plasma water, which constitutes 93% of normal plasma. In the laboratory, the concentration of sodium in plasma is measured in an aliquot of whole plasma and the concentration is expressed in terms of plasma volume (mmol L^–1 of whole plasma). If the percentage of water present in plasma is decreased, as in hyperlipidaemia or hyperproteinaemia, the amount of Na⁺ in each aliquot of plasma is also decreased, even if its concentration in plasma water is normal. A clue to this cause of hyponatraemia is the finding of a normal plasma osmolality. Pseudohyponatraemia is not encountered when plasma sodium concentration is measured by increasingly used ion-specific electrodes, because this method assesses directly the sodium concentration in the aqueous phase of plasma.

FIGURE 12.4 Causes of hyponatraemia.

True hyponatraemic states may be classified conveniently into depletional and dilutional types. Depletional hyponatraemia occurs when a deficit in TBW is associated with an even greater deficit of total body sodium. Assessment of volaemic status reveals hypovolaemia. Losses may be renal or extrarenal. Excessive renal loss of sodium occurs in Addison’s disease, diuretic administration, renal tubular acidosis and salt-losing nephropathies; usually, urine sodium concentration exceeds 20 mmol L^–1. Extrarenal losses occur usually from the gastrointestinal tract (e.g. diarrhoea, vomiting) or from sequestration into the ‘third space’ (e.g. peritonitis, surgery). Normal kidneys respond by conserving sodium and water to produce a urine that is hyperosmolal and low in sodium. In both situations, treatment should be directed at expanding the ECFV with saline 0.9%.

Dilutional hyponatraemic states may be associated with hypervolaemia and oedema or with normovolaemia. Again, assessment of volaemic status is important. If oedema is present, there is an excess of total body sodium with a proportionately greater excess of TBW. This is seen in congestive heart failure, cirrhosis and the nephrotic syndrome and is caused by secondary hyperaldosteronism. Treatment comprises salt and water restriction and spironolactone.

In normovolaemic hyponatraemia, there is a modest excess of TBW and a modest increase in ECFV associated with normal total body sodium. Pseudo-hyponatraemia is excluded by finding high protein or lipid levels and a normal plasma osmolality. True normovolaemic hyponatraemia is commonly iatrogenic in origin. The syndrome of inappropriate intravenous therapy (SIIVT) is caused usually by administration of intravenous fluids with a low sodium content to patients with isotonic losses.

A more chronic water overload may occur in patients with hypothyroidism and in conditions associated with an inappropriately elevated concentration of ADH. The syndrome of inappropriate ADH secretion (SIADH) is characterized by hyponatraemia, low plasma osmolality and an inappropriate antidiuresis, i.e. a urine osmolality higher than anticipated for the degree of hyponatraemia. It occurs in the presence of malignant tumours (e.g. lung, prostate, pancreas), which produce ADH-like substances, in neurological disorders (e.g. head injury, tumours, infections) and in some severe pneumonias. A number of drugs are associated with increased ADH secretion or potentiate the effects of ADH (Table 12.4). In patients with SIADH, the urine is concentrated in spite of hyponatraemia. Management comprises restriction of fluid intake to encourage a negative fluid balance. In severe or refractory cases, demeclocycline or lithium may result in improvement. Both drugs induce a state of functional diabetes insipidus and have been used effectively in SIADH if the primary disease cannot be treated.

TABLE 12.4

Drugs Associated with Antidiuresis and Hyponatraemia

Increased ADH Secretion

Hypnotics – barbiturates

Analgesics – opioids

Hypoglycaemics – chlorpropamide, tolbutamide

Anticonvulsants – carbamazepine

Miscellaneous – phenothiazines, tricyclics

Potentiation of ADH at Distal Tubule

Paracetamol

Indometacin

Chlorpropamide

Consequences of Hyponatraemia: Symptoms vary with the underlying aetiology, the magnitude of the reduction of plasma sodium and the rapidity with which the plasma sodium concentration decreases. Serious consequences involve the central nervous system and result from intracellular overhydration, cerebral oedema and raised intracranial pressure. Nausea, vomiting, delirium, convulsions and coma result.

Treatment of Hyponatraemia: Acute symptomatic hyponatraemia is a medical emergency and requires prompt intervention using hypertonic saline. The rapidity with which hyponatraemia should be corrected is the subject of controversy because of observations that rapid correction may cause central pontine myelinolysis, a disorder characterized by paralysis, coma and death. As a causal relationship between this syndrome and the rate of increase of plasma sodium has not been established and it is clear that there is a prohibitive mortality associated with inadequately treated water intoxication, rapid correction of the symptomatic hyponatraemic state is warranted. Sufficient sodium should be given to return the plasma concentration to 125 mmol L^–1 only and this should be administered over a period of no less than 12 h. The amount of sodium needed to cause the desired correction in the plasma sodium can be calculated as follows:

Hypertonic saline (3%) contains 514 mmol L^–1 of Na⁺ and administration poses the risk of pulmonary oedema, especially in oedematous patients, in whom renal dialysis is preferable.

Potassium Balance

The normal daily intake of potassium is 50–200 mmol. Minimal amounts are lost via the skin and faeces; the kidney is the primary regulator. However, the mechanisms for the retention of potassium are less efficient than those for sodium. In periods of K⁺ depletion, daily urinary excretion cannot decrease to less than 5–10 mmol. A considerable deficit of total body potassium occurs if intake is not restored. Hypokalaemia is a more common abnormality than hyperkalaemia.

Hypokalaemia

This is defined as a plasma potassium concentration of less than 3.5 mmol L^–1. Non-specific symptoms of hypokalaemia include anorexia and nausea, effects on skeletal and smooth muscle (muscle weakness, paralytic ileus) and abnormal cardiac conduction (delayed repolarization with ST-segment depression, reduced height of the T wave, increased height of the U wave and a widened QRS complex).

The causes of hypokalaemia are summarized in Table 12.5. Management includes diagnosis and treatment of the underlying disorder in addition to repletion of total body potassium stores. As a general rule, a reduction in plasma K⁺ concentration by 1 mmol L^–1 reflects a total body K⁺ deficit of approximately 100 mmol. Potassium supplements may be given orally or intravenously. The maximum infusion rate should not exceed 0.5 mmol kg^–1 h^–1 to allow equilibration with the intracellular compartment; much slower rates are generally used.

TABLE 12.5

Causes of Hypokalaemia

Cause	Comments
Reduced intake	Usually only contributory
Tissue redistribution	Insulin therapy, alkalaemia, β₂-adrenergic agonists, familial periodic paralysis, vitamin B₁₂ therapy
Increased Loss
Gastrointestinal (urine K⁺ < 20 mmol L^–1)	Diarrhoea, vomiting, fistulae, nasogastric suction, colonic villous adenoma
Renal	Diuretic therapy, primary or secondary hyperaldosteronism, malignant hypertension, renal artery stenosis (high renin), renal tubular acidosis, hypomagnesaemia, renal failure (diuretic phase)

The potassium salt used for replacement therapy is important. In most situations, and especially in the presence of alkalosis, potassium should be replaced as the chloride salt. Supplements are available also as the bicarbonate and phosphate salts.

Hyperkalaemia

This is defined as a plasma potassium concentration exceeding 5 mmol L^–1. Vague muscle weakness progressing to flaccid paralysis may occur. However, the major clinical feature of an increasing plasma potassium concentration is the characteristic sequence of ECG abnormalities. The earliest change is the development of tall, peaked T waves and a shortened QT interval, reflecting more rapid repolarization (6–7 mmol L^–1). As plasma K⁺ increases (8–10 mmol L^–1), abnormalities in depolarization become manifest as widened QRS complexes and widening, and eventually loss, of the P wave; the widened QRS complexes merge finally into the T waves (sine wave pattern). Plasma concentrations in excess of 10 mmol L^–1 are associated with ventricular fibrillation. The cardiac toxicity of K⁺ is enhanced by hypocalcaemia, hyponatraemia or acidaemia. The causes of hyperkalaemia are summarized in Table 12.6.

TABLE 12.6

Causes of Hyperkalaemia

Factitious (Pseudohyperkalaemia)

In vitro haemolysis

Thrombocytosis

Leucocytosis

Tourniquet

Exercise

Impaired Excretion

Renal failure

Acute or chronic hyperaldosteronism

Addison’s disease

K⁺-sparing diuretics

Indometacin

Tissue Redistribution

Tissue damage (burns, trauma)

Rhabdomyolysis

Tumour necrosis

Hyperkalaemic periodic paralysis

Massive intravascular haemolysis

Succinylcholine

Excessive Intake

Blood transfusion

Excessive i.v. administration

Immediate treatment is necessary if the plasma potassium concentration exceeds 7 mmol L^–1 or if there are any serious ECG abnormalities. Specific treatment may be achieved by four mechanisms:

chemical antagonism of the membrane effects

enhanced cellular uptake of K⁺

dilution of ECF

removal of K⁺ from the body.

Methods by which the plasma potassium concentration may be reduced are summarized in Table 12.7.

TABLE 12.7

Treatment of Hyperkalaemia

Calcium gluconate 10% i.v. (0.5 ml kg^–1 to maximum of 20 mL) given over 5 min. No change in plasma [K⁺]. Effect immediate but transient

Glucose 50 g (0.5–1.0 g kg^–1) plus insulin 20 units (0.3 unit kg^–1) as single i.v. bolus dose. Then infusion of glucose 20%, plus insulin 6–20 units h^–1 (depending on blood glucose)

Sodium bicarbonate 1.5–2.0 mmol kg^–1 i.v. over 5–10 min

Calcium resonium 15 g p.o. or 30 g p.r. 8-hourly

Peritoneal or haemodialysis

ACID–BASE BALANCE

Hydrogen ion homeostasis is a fundamental prerequisite to virtually all biochemical processes; hydrogen ion concentration [H⁺] significantly influences protein, including enzyme, structure and function and therefore nearly all biochemical pathways and many drug mechanisms. Unlike the majority of ions, [H⁺] is controlled at the nanomolar rather than millimolar level. Total body H⁺ turnover per day is in the order of 150 mmol, although most of this is ‘trapped’ within metabolic pathways (particularly ATP hydrolysis). The resultant acids may be considered as volatile (from metabolic CO₂ production) and non-volatile (from carbohydrate, fat and protein metabolism). While the lungs and kidneys play a primary role in [H⁺] homeostasis, the liver and gastrointestinal tract are also important particularly in relation to ammonium metabolism.

Due to the very low concentration of hydrogen ions in body fluids the pH notation was adopted for the sake of practicality. This system expresses [H⁺] on a logarithmic scale:

A more logical arithmetic convention which expresses [H⁺] in nmol L^–1 is gaining popularity. Table 12.8 compares values of [H⁺] expressed as pH and nmol L^–1 and reveals several disadvantages of the pH notation. The most obvious disadvantage is that it moves in the opposite direction to [H⁺]; a decrease in pH is associated with increased [H⁺] and vice versa. It is also apparent that the logarithmic scale distorts the quantitative estimate of change in [H⁺]; for example, twice as many hydrogen ions are required to reduce pH from 7.1 to 7.0 as are needed to reduce it from 7.4 to 7.3. The pH scale gives the false impression that there is relatively little difference in the sensitivity of biological systems to an equivalent increase or decrease in [H⁺]. However, when [H⁺] is expressed in nmol L^–1, it becomes apparent that tolerance is limited to a reduction in [H⁺] of only 24 nmol L^–1 from normal, but to an increase of up to 120 nmol L^–1. Nevertheless, the pH notation remains the most widely used system and is used in the remainder of this chapter.

TABLE 12.8

Comparison of Logarithmic and Arithmetic Methods of Expressing Hydrogen Ion Concentration in the Range of Blood [H⁺] Compatible with Life

Basic Definitions

An acid is a substance that dissociates in water to produce H⁺; a base is a substance that can accept H⁺. Strong acids dissociate completely in aqueous solution, whereas weak acids (e.g. carbonic acid, H₂CO₃) dissociate only partially. The conjugate base of an acid is its dissociated anionic product. For example, bicarbonate ion () is the conjugate base of carbonic acid:

A buffer is a combination of a weak acid and its conjugate base (usually as a salt) which acts to minimize any change in [H⁺] that would occur if a strong acid or base were added to it. Buffers in body fluids represent an important defence against [H⁺] change. The carbonic acid/bicarbonate system is an important buffer in blood and has historically been used as the principle determinant of physiological pH (this is in no small part due to the relationship this system has with the PaCO₂). However, it is important to appreciate the existence of other buffer systems such as plasma proteins, haemaglobin and phosphate. The pH of a buffer system may be determined from the Henderson–Hasselbalch equation, which, for the carbonic acid/bicarbonate system, relates pH, [H₂CO₃] and []:

where K = dissociation constant and pK = –log₁₀K.

This equation shows that [H⁺] in body fluids is a function of the ratio of base to acid. For the bicarbonate buffer system, pK is 6.1. As most of the carbonic acid pool exists as dissolved CO₂, the equation may be rewritten:

The value 0.225 represents the solubility coefficient of CO₂ in blood (mL kPa^–1). Normally, [] is 24 mmol L^–1 and PaCO₂ is 5.3 kPa. Thus:

Most acid–base disorders may be formulated in terms of the Henderson-Hasselbalch equation. The pH of plasma is kept remarkably constant at 7.36–7.44, i.e. a hydrogen ion concentration of 40 ± 5 nmol L^–1. This is achieved by:

regulation of H⁺ excretion and bicarbonate regeneration by the kidney

regulation of CO₂ by the alveolar ventilation of the lungs.

Cellular metabolism poses a constant threat to buffer systems by the production of volatile and non-volatile acids. Thus, the acid–base status of body fluids reflects the metabolism of both H⁺ and CO₂.

Acid–Base Disorders

The normal pH of body fluids is 7.36–7.44. Conventional acid–base nomenclature involves the following definitions:

acidosis – a process that causes acid to accumulate

acidaemia – this is present if pH < 7.36

alkalosis – a process that causes base to accumulate

alkalaemia – this is present if pH > 7.44.

Simple acid–base disorders are common in clinical practice and their successful management can usually be achieved by analysis of the carbonic acid/bicarbonate system as outlined above. In particular determination of pH, [] and PaCO₂, along with calculation of standard bicarbonate, base excess and anion gap (see below) will enable meaningful diagnosis and treatment. The first step involves diagnosis of the primary disorder; this is followed by an assessment of the extent and appropriateness of any compensation.

Primary acid–base disorders are either respiratory or metabolic. The disorder is respiratory if the primary disturbance involves CO₂, and metabolic if it involves . Thus, four potential primary disturbances exist (Table 12.9) and each may be identified by analysis of pH, [] and PaCO₂. Both pH and PaCO₂ are measured directly by the blood gas machine. [] is measured directly on the electrolyte profile but is derived in most blood gas machines. Other derived variables include standard bicarbonate and base excess. The standard bicarbonate is not the actual bicarbonate of the sample but an estimate of bicarbonate concentration after elimination of any abnormal respiratory contribution to [], i.e. an estimate of [] at a PaCO₂ of 5.3 kPa. The base excess (in alkalosis) or base deficit (in acidosis) is the amount of acid or base (in mmol) required to return the pH of 1 L of blood to normal at a PaCO₂ of 5.3 kPa; it is a measure of the magnitude of the metabolic component of the acid–base disorder.

TABLE 12.9

Compensatory Mechanisms in Acid–Base Disturbances

After the primary disorder has been identified, it is necessary to consider if it is acute or chronic and if any compensation has occurred. The body defends itself against changes in pH by compensatory mechanisms, which tend to return pH towards normal. Primary respiratory disorders are compensated by a metabolic mechanism and vice versa. For example, a primary respiratory acidosis is compensated for by renal retention of , whereas a primary metabolic acidosis is compensated for by hyperventilation and a decrease in PaCO₂. Thus, in each case, the acidaemia produced by the primary acidosis is reduced by a compensatory alkalosis. The response to a respiratory alkalosis is increased renal elimination of , and metabolic alkalosis results in hypoventilation and increased PaCO₂, pH being restored towards normal by the compensatory respiratory acidosis. In each case, the efficiency of compensatory mechanisms is limited; compensation is usually only partial and rarely complete. Overcompensation does not occur.

Metabolic Acidosis

The cardinal features of a metabolic acidosis are a decreased [], a low pH and an appropriately low PaCO₂. The extent of the acidaemia depends upon the nature, severity and duration of the initiating pathology in addition to the efficiency of compensatory mechanisms. The magnitude of the compensatory response is proportional to the decrease in [HCO₃]. The lower limit of the respiratory response is a PaCO₂ of 1.3 kPa. In a steady state:

predicted PaCO₂ = (0.2×observed bicarbonate) +1.1 (kPa)

If the observed PaCO₂ differs from the predicted value, then an independent respiratory disturbance is present.

In most instances, establishing the presence and the cause of a metabolic acidosis is straightforward. In difficult cases, an important clue to the nature of the abnormality is given by the measurement of the anion gap in plasma:

In reality, the numbers of cations and anions in plasma are the same and an anion gap exists because negatively charged proteins, together with phosphate, lactate and organic anions (which maintain electrical neutrality), are not measured. The normal anion gap is 12–18 mmol L^–1. In the critically ill population adjustments for hypoalbuminaemia (albumin itself being an anion) and hypophosphataemia should be made as follows:

Clinically, it is useful to divide the metabolic acidoses into those associated with a normal anion gap and those with an increased anion gap. The former are caused by loss of from the body and replacement with chloride. In acidoses associated with an increased anion gap, has been titrated by either endogenous, e.g. lactic acidosis, diabetic ketoacidosis, or exogenous acids (e.g. poisons), thus increasing the number of unmeasured plasma anions without altering the plasma chloride concentration (Table 12.10). Another useful concept is the osmolal/osmolar (depending on units used) gap:

TABLE 12.10

Types and Causes of Metabolic Acidosis

osmolal gap = measured osmolality – calculated osmolality

The concept is similar to the anion gap. A raised osmolal gap infers unrecognized/unmeasured osmotically active molecules within the plasma. A raised osmolal gap in conjunction with metabolic acidosis should immediately raise concern of methanol, ethylene glycol, paraldehyde or formaldehyde poisoning requiring urgent treatment. Other causes of raised osmolal gap in the absence of acidaemia include hyperglycaemia, hyperlipidaemias and paraproteinaemias.

Clinical Effects and Treatment: Metabolic acidosis results in widespread physiological disturbances, including reduced cardiac output, pulmonary hypertension, arrhythmias, Kussmaul respiration and hyperkalaemia; the severity of the disturbances is related to the extent of the acidaemia. Treatment should be directed initially at identifying and reversing the cause. If acidaemia is considered to be life-threatening (pH < 7.2, [] < 10 mmol L^–1), measures may be required to restore blood pH to normal. Overzealous use of sodium bicarbonate may lead to rapid correction of blood pH, with the risks of tetany and convulsions in the short term and volume overload and hypernatraemia in the longer term. The required quantity of bicarbonate should be calculated:

bicarbonate requirement (mmol) = body weight (kg) × base deficit (mmol L^–1) × 0.3

Administration of sodium bicarbonate should be followed by repeated measurements of plasma [] and pH. Sodium bicarbonate is available as isotonic (1.4%; 163 mmol L^–1) and hypertonic (8.4%; 1000 mmol L^–1) solutions. Slow infusion of the hypertonic solution is advisable to minimize adverse effects.

When considering the use of sodium bicarbonate in the context of metabolic acidaemia, it is important to realize that carbon dioxide is generated during the buffering process. This may result in a superimposed respiratory acidosis, especially in those patients with impaired ventilatory reserve or at the limit of compensation. It is also important to distinguish those acidoses associated with tissue hypoxia (e.g. cardiac arrest, septic shock) from those where tissue hypoxia is not a factor. It appears that therapy with sodium bicarbonate often exacerbates the acidosis if tissue hypoxia is present. For example, in patients with type A lactic acidosis, NaHCO₃ increases mixed venous PaCO₂, which rapidly crosses cell membranes resulting in an intracellular acidosis, particularly in cardiac and hepatic cells. Theoretically, this could result in decreased myocardial contractility and cardiac output and decreased lactate extraction by the liver, aggravating the lactic acidosis. Current guidelines for the management of cardiopulmonary arrest no longer recommend the routine use of sodium bicarbonate. However, if the acidosis is not associated with tissue hypoxaemia (e.g. uraemic acidosis) then the use of sodium bicarbonate results in a potentially beneficial increase in arterial pH.

Metabolic Alkalosis

The cardinal features of a metabolic alkalosis are an increased plasma [], a high pH and an appropriately raised PaCO₂. The compensatory response of hypoventilation is limited and not very effective. For diagnostic and therapeutic reasons, it is usual to subdivide metabolic alkalosis into the chloride-responsive and chloride-resistant varieties (Table 12.11). The differential diagnosis of metabolic alkalosis, and in particular the classification of patients on the basis of the urinary chloride concentration, is important because of the differences in treatment of the two groups. In chloride-responsive alkalosis, the administration of saline causes volume expansion and results in the excretion of excess bicarbonate; if potassium is required, it should be given as the chloride salt. In patients in whom volume administration is contraindicated, the use of acetazolamide results in renal loss of and an improvement in pH. H₂-receptor antagonists may be helpful if nasogastric suction is contributing to hydrogen ion loss.

TABLE 12.11

Types and Causes of Metabolic Alkalosis

Chloride-Responsive (urine chloride < 20 mmol L^–1)

Loss of acid

Vomiting

Nasogastric suction

Gastrocolic fistula

Chloride depletion

Diarrhoea

Diuretic abuse

Excessive alkali

NaHCO₃ administration

Antacid abuse

Chloride-Resistant (urine chloride > 20 mmol L^–1)

Primary or secondary hyperaldosteronism

Cushing’s syndrome

Severe hypokalaemia

Carbenoxolone

Severe alkalaemia with compensatory hypoventilation may result in seizures or CNS depression. In life-threatening metabolic alkalosis, rapid correction is necessary and may be achieved by administration of hydrogen ions in the form of dilute hydrochloric acid. Acid administration requires central vein cannulation, as peripheral infusion causes sclerosis of veins. Acid is given as 0.1 normal HCl in glucose 5% at a rate no greater than 0.2 mmol kg^–1 h^– 1.

Respiratory Acidosis

The cardinal features of a respiratory acidosis are a primary increase in PaCO₂, a low pH and an appropriate increase in plasma bicarbonate concentration. The extent of the acidaemia is proportional to the degree of hypercapnia. Buffering processes are activated rapidly in acute hypercapnia and may remove enough H⁺ from the extracellular fluid to result in a secondary increase in plasma [].

Usually, hypoxaemia and the manifestations of the underlying disease dominate the clinical picture, but hypercapnia per se may result in coma, raised intracranial pressure and a hyperdynamic cardiovascular system (tachycardia, vasodilatation, ventricular arrhythmias) resulting from release of catecholamines. There are many causes of respiratory acidosis, the most important of which are classified in Table 12.12. Treatment consists of reversing the underlying pathology if possible and mechanical ventilatory support if required.

TABLE 12.12

Causes of Respiratory Acidosis

Central Nervous System

Drug overdose

Trauma

Tumour

Degeneration or infection

Cerebrovascular accident

Cervical cord trauma

Peripheral Nervous System

Polyneuropathy

Myasthenia gravis

Poliomyelitis

Botulism

Tetanus

Organophosphorus poisoning

Primary Pulmonary Disease

Airway obstruction

Asthma

Laryngospasm

Chronic obstructive airways disease

Parenchymal disease

ARDS

Pneumonia

Severe pulmonary oedema

Chronic obstructive airways disease

Loss of mechanical integrity

Flail chest

Respiratory Alkalosis

The cardinal features of respiratory alkalosis are a primary decrease in PaCO₂ (alveolar ventilation in excess of metabolic needs), an increase in pH and an appropriate decrease in plasma bicarbonate concentration. Usually, hypocapnia indicates a disturbance of ventilatory control (in patients not receiving mechanical ventilation). As in respiratory acidosis, the manifestations of the underlying disease usually dominate the clinical picture. Acute hypocapnia results in cerebral vasoconstriction and reduced cerebral blood flow and may cause light-headedness, confusion and, in severe cases, seizures. Circumoral paraesthesia, hyperreflexia and tetany are common. Cardiovascular manifestations include tachycardia and ventricular arrhythmias secondary to the alkalaemia.

The causes of respiratory alkalosis are summarized in Table 12.13. Treatment comprises correction of the underlying cause and thus differential diagnosis is important.

TABLE 12.13

Causes of Respiratory Alkalosis

Supratentorial

Voluntary/hysterical hyperventilation

Pain, anxiety

Specific Conditions

CNS disease

Meningitis/encephalitis

Cerebrovascular accident

Tumour

Trauma

Respiratory disease