The [H+] of body fluids must be tightly regulated to ensure normal enzyme activity and cell function. Hydrogen ion is present in very low concentration in body fluids, approximately one-millionth the concentration of electrolytes such as sodium, potassium, chloride and bicarbonate. Deviations of systemic acidity can have adverse consequences. Hydrogen ion concentration less than 16 nanomoles/l or greater than 160 nanomoles/l is considered to be incompatible with life. [H+] is maintained at approximately 40 nanomoles/l despite the daily production of 50-100 mEq of hydrogen ions (fixed or nonvolatile acid) from the metabolism of proteins and phospholipids and 10,000-15,000 mmols of CO2 (volatile acid) from the metabolism of carbohydrates and fat. In addition, a number of toxins and pathophysiologic process result in the production of fixed acids (i.e., DKA, ethylene glycol). The arterial pH is maintained by the processes of 1) Chemical buffering by the extracellular and intracellular buffers, 2) control of the PCO2 by alterations in alveolar ventilation, and 3) control of plasma bicarbonate concentration by changes in renal H+ excretion.
Acid-base abnormalities are common in critically ill and injured patients. Successful management of patients with acid base disorders requires not only identification of an acid-base disturbance, but also the ability to determine the underlying mechanism for the disturbance. There are two different approaches described to assess changes in acid-base balance. The traditional approach based in the Henderson-Hasselbalch equation is more familiar than the newer quantitative strong ion approach. This presentation will describe the differences and similarities between the two methods and how each can be used in patient management.
Physiology of the Acid-Base and Buffer Systems
An acid is a hydrogen (H+) ion (i.e., proton) donor, and a base is a proton acceptor. Acids are typically represented by the notation HA, which signifies a hydrogen ion and any negatively charged particle. When placed in solution, HA dissociates into H+ (acid) and A- (base). Hydrogen ions are nonvolatile acids produced by normal metabolism of proteins and phospholipids, and are renally excreted. A base will combine with an acid to lower the amount of acid in solution, or to buffer the solution.
Carbon dioxide (CO2) is a volatile acid, or fat soluble gas, that can combine with water in the presence of carbonic anhydrase to form carbonic acid (H2CO3). Carbon dioxide is formed during normal carbohydrate and fat metabolism and is excreted via the respiratory system. These two sources of acid (H+ and CO2) are interrelated as is shown in the carbonic acid equation:
H+ + HCO3- <-> H2CO3 <-> H2O + CO2
This equation can go either direction depending upon availability of substrate on either side of the formula. The enzyme carbonic anhydrase catalyzes this reaction, and any cell containing carbonic anhydrase is capable of this reaction.
By definition, pH is the negative log of the hydrogen ion concentration. A gain of H+, or acid gain, results in a decrease in blood pH (acidemia), while loss of H+ results in increased pH (alkalemia). Acid can be gained systemically from abnormal renal elimination of a naturally occurring compound, or ingestion of an exogenous acid source. Changes in CO2 influence H+ concentration as evidenced by the carbonic acid equation. As CO2 is eliminated by increasing alveolar ventilation, carbonic acid dissociates to form more CO2, and H+ and HCO3- combine in turn to form carbonic acid. This effectively lowers H+ concentration and increases pH. Conversely, as CO2 increases from ventilation impairment, pH decreases.
A buffer is a compound that can accept protons (H+ ions) and minimize a change in pH. Buffer systems that exist in the body include nonbicarbonate buffers (proteins and phosphates), which are primarily intracellular, and bicarbonate (HCO3-), which is the primary extracellular buffer. Bicarbonate is an effective buffer because it exists in relatively large concentrations compared to other buffers, and it participates in the carbonic acid formula to produce CO2 gas that can be eliminated through ventilation. Therefore the HCO3- buffer system is considered an open system which can continue to buffer as long as the respiratory system is functional. When HCO3- is lost excessively from the urinary or gastrointestinal systems, CO2 and H2O combine to form carbonic acid, which dissociates to increase H+ and cause acidemia.
The traditional or descriptive approach is based on the Henderson-Hasselbalch equation:
pH = 6.1 + log ( [HCO3-]/ 0.03 PCO2)
The equation mathematically links the variables of pH, partial pressure of carbon dioxide (carbonic acid) and bicarbonate concentration. Since a predictable change in HCO3- will occur with gain or loss of H+ ions, HCO3- can be used to correctly identify acid base abnormalities originating from metabolic disorders. Acidemia or alkalemia resulting from a respiratory disorder should have altered PCO2. Respiratory acidosis results in increased PCO2, while respiratory alkalosis decreases PCO2. In metabolic acidosis, H+ increase will shift the carbonic acid equation to cause a decrease in HCO3-, and in metabolic alkalosis, H+ decrease will have the opposite effect on HCO3-. Blood gas analyzers typically measure pH and PCO2, and calculate HCO3- since the degree of change of HCO3- is predictable based on pH and PCO2. This equation can also be used to predict how compensatory mechanisms engage to reduce the degree of change of pH. When metabolic acidosis develops, the respiratory system is stimulated to increase respiratory rate to eliminate CO2 from the lungs and create respiratory alkalosis. Likewise, with a primary respiratory disorder, the opposite metabolic disorder will be generated. The body's response to a fixed acid and to a volatile acid differs. Hydrogen ions from a fixed acid load immediately titrate bicarbonate ions in the extracellular fluid and then titrate intracellular buffers. This response occurs within minutes. Alveolar ventilation is stimulated to decrease PCO2 and normalize the ratio of [HCO3-] and PCO2. This response begins immediately and is completed within hours. Finally, the kidney retitrates bicarbonate by increasing H+ excretion. The renal response begins within hours but requires 2-5 days to reach maximal effectiveness. In contrast, volatile acids (CO2) cannot be buffered by HCO3-. Hydrogen ions resulting from the dissociation of carbonic acid must be titrated by intracellular buffers. The compensatory renal response requires 2-5 days to achieve maximal effectiveness.
Base excess is a value in which the amount of base above or below the normal buffer base is calculated, taking into account expected change in HCO3- secondary to acute changes in PCO2. The general rule of thumb is that HCO3- concentration rises about 1 to 2 mEq/l for each acute 10-mm Hg increase in PaCO2 above 40 to a maximum increase of 4mEq/l, and that the HCO3- concentration falls 1 to 2 mEq for each acute 10-mm Hg decrease in PaCO2 below 40, to a maximum decrease of 6 mEq/l. Base excess is expressed in mEq/l. Some authors refer to a negative base excess as a base deficit. A simple acid base disorder is limited to the primary disorder and the appropriate compensatory response. A mixed disorder is one in which there are at least two separate abnormalities occurring simultaneously. Normal values at sea level for venous blood gas interpretation are pH 7.35-7.45, PCO2 40-45 mmHg, and HCO3- 19-24 mEq/l. Base excess should normally be between -5 and 5 mEq/l.
This traditional approach to acid-base evaluation is familiar and user-friendly, however it does not include the effect of changes in electrolytes and/or plasma proteins on acid-base balance and has been faulted for its inability to aid in the determination of the causative mechanism of an identified acid base disturbance. In some complex clinical situations, additional information can be obtained using the Stewart model of acid-base evaluation.
Strong Ion Approach
The basic principle of the quantitative or strong ion approach is that the pH of blood is determined by three mathematically independent determinates; the strong ion difference, the total weak acid concentration (Atot) and the PCO2. Electrolytes such as sodium, potassium, magnesium and chloride exist in body fluids completely ionized and are considered strong ions. Strong ion difference (SID) is the resulting net charge of all of the strong ions. This includes both the cations (Na+, K+, Ca2+, and Mg2+) and the anions (Cl- and lactate). Sodium is the only cation present at high enough concentration in body fluids that a change in its concentration is likely to have a substantial effect on SID. In healthy humans, the SID is +40mEq/l.1 The SID changes if the difference between the sum of the strong cations and sum of the strong anions changes. An increased in SID causes a metabolic alkalosis and a decreased SID is associated with a metabolic acidosis.
To preserve electroneutrality, there must be an equal and opposing charge to balance the positive SID and this force is comprised mostly of weak acids. These weak acids include plasma proteins (predominately albumin), phosphates and bicarbonate. Proteins and phosphate constitute the independent variable Atot, whereas HCO3- is a dependent variable. Decreased albumin can be associated with a mild metabolic alkalosis.
The pH of blood is directly influenced by the carbon dioxide concentration. Arterial PCO2 is inversely proportional to the alveolar ventilation and can be used to evaluate primary respiratory disorders. Alveolar hypoventilation results in respiratory acidosis and an elevation in PaCO2. Conversely, respiratory alkalosis and a decreased PaCO2 are seen with alveolar hyperventilation.
Several modifications of the strong ion approach have been suggested to simply its use in clinically situations. These will be discussed in the following presentation.
1. Gunnerson KJ. Clinical review: The meaning of acid-base abnormalities in the intensive care unit-epidemiology. Crit Care Oct 2005 Vol9 No 5,p508.