Vocal Cord Dysfunction (VCD ) or PARADOXICAL VOCAL CORD MOVEMENT (PVCM)

FROM BICARBONATE TO STRONG ION DIFFERENCE- INTRIGUING STORY OF BLOOD ACID-BASE ANALYSIS 

“Life is struggle, not against sin, not against money power…. but against Hydrogen ion.”

- H.L. Mencken, 1919

 

Since the understanding of similarity between fermentation of wine and respiration of animals, evolution of human physiology and measurement of carbon dioxide is coupled with studies of acids and bases. Alkalinity of blood was demonstrated by color indicators as early as eighteenth century by French chemist Hilaire Marin Rouelle, and one century later, its relation with gastric acid secretion was recognized by Henry Bence Jones. In 1831, William B. O’Shaughnessy, an Irish physician working in India, demonstrated that Cholera reduced the free alkali of blood.

But the discovery of relationship between blood alkalinity and carbon dioxide was contained within the mystery of diabetic coma.

 

1.   Story of Carbon Dioxide and Alkalinity of blood

 

In Nineteenth century clinicians were intrigued by the mystery surrounding two common diseases causing coma, diabetes and nephritis. It was presumed that some toxic metabolite led to decreased level of consciousness. 

 

In 1877, Friedrich Walter working on his thesis in Strasbourg, France, ingested rabbits, large amount of diluted hydrochloric acid to investigate the mechanism of disposal of metabolically generated acids in blood. He observed that carbon dioxide content of blood was reduced following acid ingestion, which was restored to normal with intravenous injection of sodium bicarbonate. Walter proposed that total carbon dioxide content in blood is useful measure of alkali content, degree to which alkalinity of blood is neutralized by non-volatile acid. Thus, Walter paved the way for studying acidosis by extracting and quantifying carbon dioxide from blood¹.

 

In 1884, German Physiologist, Oscar Minkowski, working in Bernhard Naunyn’s laboratory identified acid in diabetic coma patients and named it as 3-hydroxybutyric acid. He further substantiated Walter’s propositionthat carbon dioxide content of blood in comatose diabetic patients was much less than normal, and patients improved clinically after parenteral administration of sodium bicarbonate. Minkowski, thus, proposed that diabetic coma was the result of generalized acidosis rather than direct toxic effect of abnormal acid¹.

 

2.   Story of Electrolytes and Acid-Base

 

By 1878, both organic and inorganic components of blood were identified. While function of organic components like urea, uric acid, creatinine, bile pigments and hemoglobin had been recognized, that of inorganic components was still unknown. Inorganic components were measured in terms of potash, soda, lime, magnesia, chlorine and phosphoric acid, as it was assumed that they existed in same state in blood as in nature.

 

In 1884, Svante Arrhenius, a student of Chemistry and Mathematics at University of Upsala in Sweden, proposed theory of ionization, which assumed that when a salt like sodium chloride dissolves in water, it dissociates into individual ions bearing electrical charge. These ions are attracted towards electrodes with opposite charge. He further noticed that strong acids were better conductor of electricity than weak acids and linked it to degree of dissociation of acids. Arrhenius proposed that acids and bases are substances that add hydrogen ion (H⁺) and hydroxyl ion (OH^-) respectively, to water. For example, Hydrochloric acid, HCl, adds (H⁺) ions to water and sodium hydroxide, adds (OH^-) to the water. In the solution, acids and bases neutralize each other, because the H⁺ and OH^-) ions react to form water, leaving salts behind.

 

 HCl + NaOH = NaCl + H₂O

 

However, Arrhenius thesis which showed first link between chemistry and physics, was not accepted well and he was given doctorate with just passing marks, as current understanding of electrolytes assumed that isolated sodium and chloride particle could not exist in water. It was expected that sodium will react violently with water and chloride will produce poisonous gas chlorine.

 

Two decades later, discovery of subatomic particles finally explained Arrhenius’s theory of ionization. Studies by Joseph T. Thompson, Ernest Rutherford, Gilbert N. Lewis and Niels Bohr revealed that atoms consist of nuclei with orbiting electrons. Therefore, when sodium chloride dissociates in solution, sodium atom transfers its single electron in outer orbit to chloride atom, which is missing single electron in its outer orbit. This result in positive sodium ion and negative chloride ion.

 

Despite this, Arrhenius’s theory of ionization was not discussed in academic circles and standard textbook, until he received noble prize in 1903.

 

3.   Story of bicarbonate as measure of acid base balance

 

Henderson equation

 

In 1894, Lawrence Joseph Henderson, a student of physical chemistry in Harvard College in Cambridge, Massachusetts, was so impressed with Arrhenius’s ionization theory that he wrote an essay about it. Thirteen years later, observing the remarkable ability of blood to neutralize large amounts of acid, Henderson, now an instructor in biochemistry, investigated the relationship of bicarbonate to dissolved carbon dioxide gas, and its role as buffer of fixed acids. 

 

Based on his experiments, Henderson rewrote the laws of mass action for weak acids and their salts and published his classic paper, “The Theory of Neutrality Regulation in the Animal Organisms” and proposed famous Henderson’s equation in 1908. He explained that the chemical neutrality of body fluids is maintained by the presence of weak acids, which react with dissociated strong acids to form neutral salts and a slightly dissociated weak acids, thus minimizing changes in hydrogen ion concentration. These weak acids are bicarbonates, phosphates and proteins, and function as donor or acceptor of hydrogen ions according to circumstances. 

 

In the case of carbon dioxide and bicarbonate he defined the dissociation constant k as the hydrogen ion concentration at which half of the carbonic acid is dissociated: 

 

CO₂ + H2O = H₂CO₃ = [H⁺] + [HCO₃^-]

[H⁺] [HCO₃^-] = k × [H₂CO₃]

[H⁺] = k × [H₂CO₃] / [HCO₃^-]

 

k =  [H⁺] [HCO₃^-] / [H₂CO₃]

 

Assuming that all dissolved carbon dioxide was carbonic acid,  

 

K = [H⁺] [HCO₃^-] / [dCO₂],

 

where dCO₂ (dissolved CO2) includes both the CO2 gas and H2CO3.

 

In nutshell, Henderson described how the neutral reaction of blood is determined by carbonic acid (H₂CO₃) being in equilibrium with an excess of mineral bases over mineral acids. This excess of mineral bases over mineral acids was later defined as buffer base and strong ion difference. Henderson further proposed that hydrogen ion concentration in body is proportional to the ratio of carbonic acid to bicarbonate, placing bicarbonate at the centre of the acid-base homeostasis, and a criteria to determine acid-base status of body. 

 

[H⁺] ∝ [H₂ CO₃] / [HCO₃^-]

[H⁺] = 4 × 10^-8 or 0.000000004)

 

It was this insight that helped chemists and physiologists to realize that when acids are added to blood, the hydrogen ions react with blood bicarbonate, generating carbon dioxide gas, which is then excreted by the lungs, almost eliminating the added acid.

 

Henderson’s equation was initially difficult to interpret mathematically, both by clinicians and physiologist, who had recently started appreciating Arrhenius ionization concept.

 

4.   Story of pH and Henderson-Hasselbalch equation

 

In 1909, across Atlantic ocean, in Copenhagen Denmark, Soren P. L. Sorenson was studying the effect of ionic concentration in the analysis of proteins. Fed up with writing eight zero to describe hydrogen ion concentration, Sorenson suggested use of negative log of hydrogen ion concentration in terms of pH. Thus, the acidity of normal blood would be designated pH 7.4 instead writing 4×10^-8 N or 0.00000004 N. 

 

Sorenson further described Henderson’s neutrality systems as shock absorbers in trains and proposed the concept of buffer, which resisted moderate production of acids with little change in hydrogen ion concentration.

 

pH = -log [H+]

 

In 1917, in Germany, a physician and chemist Karl A. Hasselbalch, merged Sorenson’s pH equation and Henderson’s equilibrium equation into the immortal Henderson-Hasselbalch equation.

 

pH = pK_a + log [A^-] / [HA]

 

where K_a is the dissociation constant for HA, a weak acid, and square brackets enclose the concentrations of undissociated acid and the salt of that acid.

 

in terms of HCO₃^- and CO₂, the equation may be rewritten as,

 

pH = pK + log [HCO₃^-]/ [P_CO2]

pH = 6.1 + log [HCO3^-] / [P_CO2 × 0.03] 

 

By this time, it was understood that changes in the partial pressure of carbon dioxide directly lead to changes in a patient’s acid–base status and therefore, partial pressure of carbon dioxide (P_CO2) is a measure of respiratory acidosis or respiratory alkalosis. In 1917, acknowledging the interdependence of bicarbonate and partial pressure of carbon dioxide (P_CO2), Hasselbalch coined the term “compensation” to describe homeostasis in acid-base milieu brought about by kidneys and lungs in response to respiratory and non-respiratory (metabolic) acid-base abnormality. But, non-respiratory (metabolic) components of acid-base disorders was still not completely deciphered. Physiologist and clinicians desperately wanted an ideal parameter to identify and quantify metabolic acidosis, as it has clinical implication in treatment of diabetic coma and kidney disease.

 

5.   Story of the search for ideal parameter of metabolic acid-base disorder

Reduced pH

 

In 1916, Hasselbalch proposed reduced pH as a measure of the metabolic acid-base balance, by equilibrating a sample of blood at 38⁰C with a carbon dioxide tension (P_CO2) of 40 mmHg. However, this excellent idea was never adopted for widespread clinical use, probably because there was no easy means of measuring pH, especially in comparison to availability of the CO₂ content of blood, using the Van Slyke method described later.

 

Carbon dioxide content in blood 

 

Henderson’s equation emphasised the importance of electrolytes in acid-base disorders, it was difficult to measure electrolytes, thereby making it difficult to apply in patients to diagnose and evaluate these disorders¹.

But thousands of kilometres, across ocean, on man was going to solve this problem. Donald Van Slyke, an organic and analytical chemist, was invited by director of new Rockefeller hospital to help clinicians in studying various diseases. Initially reluctant to take this assignment, because of not being physician, Van Slyke accepted this challenge, and studied physiology text books and enhanced his knowledge about diabetes and other diseases¹.

 

With help of Wlater W. Plamer and Reginald Fitz, former associates of Henderson, Van Slyke deduced that if diabetic coma were due to accumulation of acetoacetic and 3-hydroxybutyric acids, these metabolites would react and lower the concentration of bicarbonates¹. He conducted an experiment, where sulfuric acid was injected into dogs, to mimic effects of diabetic ketoacidosis. Van Slyke observed that serum bicarbonate accounted for about 30% buffering of injected acid. He suggested resulting bicarbonate deficit as measure of metabolic acidosis, which could be therapeutic targeted to restore acid-base balance to a normal pH, in diabetic ketoacidosis.

 

Van Slyke conducted these experiments by adopting Walter’s idea of measuring CO₂ content of blood as surrogate of bicarbonate. CO₂ content of blood is comprised of bicarbonate, carbonic acid, carbon dioxide bound to haemoglobin and dissolved carbon dioxide. As bicarbonate constitutes 95% of all these, CO₂ content may be reliable estimate of blood bicarbonate. He used terms CO₂ combining power, alkali reserve and  base-bicarbonate for CO₂ content of blood. Later CO₂ content of blood was renamed as total CO₂ content (_ctCO₂). He also introduced terms alkali deficit and alkali excess as alternatives to metabolic acidosis and alkalosis, respectively.

 

He proposed to determine the HCO3^- at a pH of 7.4 and reintroduced the term alkali reserve for this quantity.

 

Van Slyke further explained that free carbonic acid is present in the body fluids in such concentration that it converts all bases not bound by other acids, into bicarbonate. The bicarbonate therefore, represents the excess of base, which is left after all the non-volatile acids (metabolic acids- organic and mineral) have been neutralized, and is available for the immediate neutralization of further acids. Therefore, bicarbonate measured at pH of 7.4 and PCO₂ 40 mmHg should be a physiological constant which he named the alkali reserve of body. He suggested that CO₂ content of blood, measured after equilibration at P_CO2 of 40 mmHg, may be used as a measure of metabolic acid-base balance. Four decades later, this was re-introduced as standard bicarbonate by Astrup. 

 

For measuring CO₂ content of blood, Van Slyke utilized his gastrometric technique that he had developed for measuring the amount of ammonia and carbon dioxide released by the action of urease on urea. In 1917, Van Slyke developed a glass instrument to measure CO₂ content of blood in patients of diabetic ketoacidosis. This glass instrument known as ‘Van Slyke apparatus’ was going to be used for next 50 years in almost every clinical laboratory as only practical method to estimate metabolic acidosis and to determine amount of bicarbonate to be given in DKA to correct acidosis until discovery of insulin by Banting in 1921.

 

However, impression that total CO2 content of blood is a measure of metabolic acid-base disturbance was flawed, as serum bicarbonate changes with changing CO₂ concentration with respiration. In other words, measurement of CO2 content of blood will not differentiate between respiratory and metabolic disorder, unless pH is determined. This presumption was going to present a historical error in next 30 years, which shaped the course of history of clinical and laboratory medicine.

 

Buffer Base- BB

 

When strong acid is added to blood, both carbonic (HCO₃^-) and non-carbonic (weak acids [A^-]- hemoglobin, albumin, phosphate) buffers concentrations will be reduced. Taking only HCO₃^- buffer into computation of acid base calculation, omits the effect of non-carbonic buffers. 

 

Hence, in 1948 Singer and Hastings introduced buffer base (BB) as sum of all buffer anions to quantify metabolic acid–base changes, independent of carbon dioxide. They embraced old definitions of acids and bases being anions and cations, respectively. Instead of utilizing Henderson-Hasselbalch equation, they employed the Gamble diagram with the two columns of cations and anions of equal height, illustrating the law of electro-neutrality. They defined buffer base as the sum of strong bases minus the sum of strong acids, which should be equal to carbonic (HCO₃^-) and non- carbonic (weak acids [A^-]- hemoglobin, albumin, phosphate) buffers. Thus, Buffer base is a measure of the concentration of both carbonic and non-carbonic buffers present in either plasma or blood. 

 

BB = Na⁺ – Cl^- = HCO₃ + A^-

 

Accumulation of strong acid or base is reflected stoichiometrically in a decrease or  increase of buffer base respectively. Singer and Hasting suggest that buffer base is not affected by P_CO, as compensatory change in bicarbonate concentration with change in P_CO2, is matched by a similar change in concentration of other buffer anions. 

 

To quantify buffer base, calculating left half of equation (measurement of all cations and anions) was tedious compared to calculating right half of equation (bicarbonate and other buffer anions- albumin, haemoglobin, phosphates) in that era of time. Singer and Hastings used this approach and constructed a nomogram which, among other calculations, allows calculation of plasma and whole blood buffer base from pH, pCO2, and haematocrit.

 

BB = HCO3 + A-

 

As buffer base may be altered by non-carbonic buffers especially albumin, and Singer and Hastings did not consider changes in albumin as acid-base disorders. Therefore, to eliminate its effect, they defined delta buffer base (ΔBB) as difference between actual buffer base and normal buffer base (measured at normal pH 7.4 and P_CO2 40 mmHg), as measure of a metabolic acid-base disturbance. However, it was realized later that change in non-carbonic buffer leads to similar but opposite change in carbonic buffer, to maintain electroneutrality. 

 

Using this approach they constructed a nomogram which, among other calculations, allows calculation of plasma and whole blood buffer base from pH, P_CO2, and haematocrit.

 

Delta buffer base was later redefined as base excess by Astrup and Siggard- Anderson.

 

ΔBB = BB (measured) – NBB (measured at normal pH and PCO2)

 

However, buffer base never gained acceptance because of several reasons, (1) it is altered by changes in non-carbonic buffers (albumin, hemoglobin, phosphate), (2) organic acids (lactic acid, ketoacids etc.) are titrated differentially in blood and interstitial compartments (3) it is influenced by chronic change in P_CO2. However, 60 years later its re-introduction by Stewart was going to create a revolution in the understanding of acid-base disorders.

 

Ignorance is not bliss- the polio epidemic

 

With the introduction of commercial pH glass electrode in 1932, It was possible to estimate both respiratory and metabolic components of acid base status by calculating P_CO2 after measuring pH with glass electrodes and carbon dioxide content (HCO₃^-) with the Van Slyke technique. But, it was very cumbersome and time consuming. In 1950s in Julius Comroe's laboratory at the University of Pennsylvania in America, blood gas results were screened through a committee headed by Arthur DuBois before being reported by telephone within a week and followed in several weeks by a written report. 

 

Therefore, in clinical practice, total CO₂ content in plasma as surrogate for HCO₃^-, was the only laboratory parameter available for assessment of a patient’s acid-base status. This limitation was going to create a clinical and organizational challenge which forced the transition of blood gas analysis from a physiologic laboratory exercise to a clinical necessity, and led to evolution of intensive care medicine. Also, this incidence expedited development of blood gas analyser and positive pressure mechanical ventilator.

 

In 1952, Denmark was struck by a devastating epidemic of poliomyelitis. Nearly 3000 patients were admitted in first four months of epidemic, mostly in Blegdam hospital, an Infectious disease hospital in Copenhagen. About 345 of these patients had bulbar poliomyelitis causing respiratory paralysis. During the first 3 weeks of the epidemic 27 of 31 patients with bulbar polio, died. Blood acid base analysis by Van Slyke method revealed high CO₂ content which was interpreted as metabolic alkalosis. Disappointed by high mortality, Professor H. C. A. Lassen, Head of the Medical Department, reluctantly called a young anaesthetist Bjørn Ibsen to solve the riddle. in those days, anaesthesia was an evolving speciality and anaesthetists were not highly regarded by other physicians. After examining patients and going through records of diseased patients, Ibsen realized that bulbar polio patients died because of respiratory acidosis (hypercapnia) resulting from respiratory muscle paralysis. This was confirmed by Poul Astrup, director of the clinical laboratory, who measured blood pH of patients with electrode borrowed from Radiometer AS, a company which had recently developed an electrode capable of measuring pH in small samples of blood.

 

But, what led to the misunderstanding of high CO₂ content of blood as metabolic alkalosis, and not respiratory acidosis. In research laboratories, respiratory disturbances of the acid-base equilibrium were well known, but in hospitals and hospital laboratories they were almost never encountered. Diseases of the lung causing significant carbon dioxide abnormalities did not get enough time to meet metabolic compensation, as patients used to die earlier because of complication. Negative pressure ventilators (tank ventilator) were not effective in supporting failing respiratory system and positive pressure ventilator were in infancy period of development. This ignorance of chronic metabolic compensation of respiratory acidosis in clinical experience, misled physicians of Blegdams hospital, to interpret high carbon dioxide content in plasma in patients of bulbar polio, as metabolic alkalosis.

 

Frustrated by the incidence, Poul Astrup studied Van Slyke’s work and realized that the titration line for CO₂ to pH was approximately a straight line in physiological range, if P_CO2 was plotted logarithmically. He developed an equilibration technique and designed an apparatus in which one could first measure the pH of a blood sample, and then bubble gas of known P_CO2 through the sample, and measure the pH again to draw a titration line. Blood samples were taken in three capillary tubes (approximately 50 µL in each tube). Two of these samples were transferred to the equilibration chambers and equilibrated with special gas mixtures of known high and low P_CO2 for 3-4 minutes. Third sample was not equilibrated with CO₂. The pH was measured for all three samples. A titration line was drawn based on the measured pH and known P_CO2 values in the two equilibrated samples. Now, P_CO2 of third sample could be read from the titration line with the measured pH of this sample.

Astrup emphasized that measuring PaO2, PaCO2, and pH should be the ultimate goal for determining how much one should artificially ventilate bulbar polio patients to correct respiratory acidosis. Ultimately, discovery by Clark of the oxygen electrode and by Stow and colleagues of the carbon dioxide electrode, with its improvement by Severinghaus and Bradley in the mid-1950s, set the scene for development by commercial interests of blood gas machines that could give accurate and rapid measurements of blood P02, PC02, and pH in routine hospital laboratories. These became increasingly available commercially in the early 1960s.The availability of methods for quickly measuring arterial blood gas values and pH enormously simplified the monitoring of patients who were undergoing mechanical ventilation. The improvement in managing patients receiving mechanical ventilation with arterial blood gas measurements, as opposed to clinical signs alone, was like driving out o f a dense fog into bright sunshine.  

Standard Bicarbonate- HCO₃^- (st)

 

In 1957 Astrup along with Jorgensen introduced standard bicarbonate as measure of metabolic acid base disorder, independent of P_CO2. They defined it as plasma bicarbonate in fully oxygenated blood equilibrated at 38⁰C to P_CO2 of 40 mmHg.

 

But, standard bicarbonate did not consider the buffer effect of non-carbonic buffers (albumin and hemoglobin), i.e. proteins which normally contributed to buffering with 14-16 negative charges (A^-) per liter.  Indeed, when strong acid is added to blood both carbonic (HCO₃^-) and non-carbonic buffer (A^-) concentration will be reduced. In open system, like healthy subject efficiently regulating P_CO2 through breathing, carbonic buffers have a predominant role (about 75-80%), however, non-carbonic buffers cannot be disregarded completely. Therefore, the difference between HCO₃^- (st) and the ideal normal bicarbonate value slightly underestimate the acid-base added to the system.

 

Base Excess- BE

 

One year later, still searching the ideal parameter for metabolic acid base status, independent of P_CO2, Astrup and Siggard-Anderson re-introduced Singer and Hasting’s delta buffer base as base excess (excess of actual buffer base above normal buffer base). Base excess was proposed as an ideal quantitative measure of metabolic disorder as it assesses all the extracellular buffers, carbonic (HCO₃) and non-carbonic (haemoglobin, albumin, phosphate) in the blood sample and was considered to be independent of carbon dioxide in in-vitro environment (blood gas machine).

 

BE = Delta BB = BB – NBB

 

Siggard-Anderson further explained base excess in terms of net titrable concentration of hydrogen ion. In any metabolic acid-base disorder (acidosis or alkalosis), amount of hydrogen ion added to or removed from the system may be determined by back titration of the system to the normal pH, by adding or removing hydrogen ion (HCl).

In blood or plasma, where P_CO2 is an independent variable (influences carbonic and non-carbonic buffers), the concentration of titrable hydrogen ion is determined by titrating blood to a pH of 7.40 at a P_CO2 of 40 mmHg and temperature of 37°C. This quantity with opposite sign is called the base excess, understanding that a negative value indicates a base deficit, which is equivalent with organic acid (lactic acid, ketoacids etc) excess. In view of the ambiguity of the words base in base excess, it was realized later by Siggard-Anderson that  (net) concentration of titratable hydrogen ion (∆ctH+) would have been preferable name than base excess. This would emphasize that the quantity refers to hydrogen ions, not cations or anions, and indicate that the quantity may be either positive or negative. 

 

Siggard-Andersen conducted in vitro studies that equilibrated blood with a  P_CO2 of 40 mmHg, effectively removing any respiratory abnormality. He then quantified the amount of strong, fully dissociated acid (hydrochloric acid HCl) or base (sodium hydroxide NaOH) in millimoles required to return a litre of blood to pH 7.40. This quantity is the base excess in millimoles per litre and is considered negative if NaOH is needed (acidosis) and positive if HCl is needed (alkalosis). 

Base excess was defined the amount of strong acid (in millimoles per liter) that needs to be added or removed in vitro to 1 liter of fully oxygenated whole blood, titrated to a Pco2 of 40 mmHg at temperature of 37oC, to return the pH to 7.40.

However, base excess was unable to prove itself an ideal parameter for metabolic acid-base status which was fiercely pointed out by two physicians from across Atlantic Ocean.

 

6.   Story of The Great Transatlantic Acid-Base Debate

 

By 1960, Astrup and Siggard-Andersen were recognized not only in Europe but across Atlantic Ocean, in America too. Their approach to acid-base analysis, centered upon base excess, became recognized as “The Copenhagen School”. 

 

Standard Base Excess or Base Excess in extracelluar fluid

Around same time in 1963, two internists with special interest in Nephrology, William Schwartz and Arnold Relman, at Tufts Medical School in Boston, Massachusetts, published a critique of the various parameters proposed for determination of acid-base balance in the New England Journal of Medicine. They specifically criticized base excess as a measure of the metabolic acid-base balance, independent of P_CO2. Principle reason for this criticism was the difference in base excess value, depending upon whether titration with P_CO2 has been done in vitro or in vivo. When P_CO2 is varied in vivo by CO₂ inhalation or hyperventilation, blood as well as whole extracellular fluid is equilibrated with the new P_CO2. When P_CO2 increases, pH tends to decrease more in the poorly buffered interstitial fluid than in the well-buffered blood containing hemoglobin as hemoglobin is the major buffer of CO_2. Hydrogen ions therefore, tend to diffuse from interstitial fluid into the blood, leading to further fall in base excess of blood, creating difference between in vivo and in vitro estimation of base excess. Therefore, base excess may not be the ideal index of metabolic component of acid base balance.

 

As in vitro base excess failed to represent in vivo base excess, Schwartz and Relman termed base excess  an artificially derived, physiologically unregulated parameter, and otherwise irrelevant to an “open” in vivo system. 

Another objection of Schwartz and Relman was that base excess is not independent of P_CO2 in chronic respiratory acidosis and resulting compensatory rise in base excess was being called metabolic alkalosis. They opined that it is a normal compensation for respiratory acidosis, and therefore, should not be regarded metabolic acid-base disorder. 

 

In nutshell, Schwartz and Relman rejected Base excess as an index of metabolic acid-base status because, (1) in vivo measurement din not represent in vitro status, (2) it was not independent of P_CO2 in chronic respiratory acidosis and (3) compensatory changes to acid-base disorders could be misinterpreted as abnormal. 

 

Siggard-Andersen dealt with first criticism by modifying the computation of base excess by assuming a model where blood volume is diluted with the interstitial fluid to eliminate the differential buffering capacity of blood and extracellular fluid. This was done by lowering normal hemoglobin concentration of 15 g/dL to 5 g/dL to represent hemoglobin deficient extracellular fluid, in computation of base excess, and was defined as in vivo base excess, standard base excess (SBE) or base excess of extracellular fluid (cBase(ecf).  

 

T. J. Morgon described four SBE-based bedside rules to assess compensation for acid-base disorders,

 

Metabolic Acidosis-

expected CO₂ = 40 + (1.0 × SBE)

Metabolic Alkalosis-

expected CO₂ = 40 + (0.6 × SBE)

Acute Respiratory Acidosis or Alkalosis-

An acute change in PaCO₂ will not change the Standard Base Excess.

expected CO₂ = 40 + (0 × SBE)

Chronic Respiratory Acidosis or Alkalosis-

expected SBE = 0.4 × (PaCO₂ - 40)

 

However, Schwartz and Relman believed that actual bicarbonate could well be used as the metabolic parameter. Therefore, in-vivo titration experiments were performed at Tufts university, around 1962-1963. A large number of healthy persons (exposed to recirculated CO₂), individuals with chronic respiratory acid-base abnormalities (chronically alkalotic pregnant women, chronic respiratory acidotic COPD patients) and sick patients with metabolic acid-base disorders were sampled, and the results were plotted on a graph. From this graph, based on relationship between H⁺, P_CO2 and HCO₃^-, six primary disorders of acid base imbalance were described, namely metabolic acidosis, metabolic alkalosis, acute respiratory acidosis, acute respiratory alkalosis, chronic respiratory acidosis and chronic respiratory alkalosis. To estimate compensatory changes in P_CO2 or HCO₃^- for change of other in individual primary disorders, six equations were proposed which were nicknamed as six rules of thumb. All of these equations were based on the averaged data conglomerated into point plots, where lines of best fit are drawn to represent the relationship between HCO₃^- and CO₂ under different circumstances.

 

The six rule of thumb describes the physiological compensation to acid–base changes to optimize acid–base homeostasis. Having allowed for expected physiological compensation, residual changes in carbon dioxide or bicarbonate are then seen as the mechanisms for changes in acid–base status, rather than disorder.

 

Schwartz and Relman approach utilizing relationship between P_CO2 and bicarbonate, for blood acid-base analysis, came to be known as P_CO2/HCO₃^- approach of Boston School.

 

Six rule of thumb for calculating compensation of acid-base disorders, borrowed from Brandis as The 1963 Schwartz and Relman article in NEJM is paywalled,

 

Metabolic Acidosis- (winter’s formula)

For complete compensation, expected PaCO₂ = (1.5 × HCO₃^-) + 8

Metabolic Alkalosis-

For complete compensation, expected PaCO₂ = (0.7 × HCO₃^-) + 20

Acute Respiratory Acidosis-

For every 10 mmHg increase in PaCO₂, the HCO₃^- will rise by 1 mmol/L

In other words, expected HCO₃ = 24 + ((PaCO₂-40) / 10)

Chronic Respiratory Acidosis-

For every 10 mmHg increase in PaCO₂, the HCO₃^- will rise by 4 mmol/L

In other words, expected HCO₃ = 24 + (4 × (PaCO₂-40) / 10)

Acute Respiratory Alkalosis-

For every 10 mmHg decrease in PaCO₂, the HCO₃^- will fall by 2 mmol/L

In other words, expected HCO₃ = 24 - (2 ×(40-PaCO₂) / 10)

Chronic Respiratory Alkalosis-

For every 10 mmHg decrease in PaCO₂, the HCO₃^- will fall by 5 mmol/L

In other words, expected HCO₃ = 24 - (5 ×(40-PaCO₂) / 10)

 

In 1977, Severinghaus suggested a truce between the Boston and the Copenhagen schools. Severinghaus proposed a modified Siggard-Andersen nomogram to permit one to estimate the chronic compensation of hypercapnia, according to the Boston school's work. Despite this attempt, both schools remained unreconciled and continue to differ.  

 

Copenhagen school or Boston school- which is better

 

Both Copenhagen and Boston schools are bicarbonate centered approach utilizing Henderson-Hasselbalch equation. There has been considerable discussion about the merits and demerits of the SBE compared with the P_CO2/HCO₃^- approach. In reality, there is little difference between the two; both equations and nomograms were derived from in vivo patient data and abstracted backward.  

 

Boston school is easier to use at bedside, if one memorizes rule of thumb, to diagnose acid-base disorder qualitatively. On the other hand, SBE is useful for diagnosis as well as quantification of metabolic acid–base disorders. But, SBE has its own limitations. First, it is a calculated value, that relies on several measured variables (pH, P_CO2, hemoglobin concentration) and does not takes into account albumin, which is an important non-carbonic buffer in critically ill patients. Second, SBE by itself does not provide information about the underlying mechanism. Therefore, while an abnormal SBE is a reliable marker of an acid-base derangement, a normal SBE may not exclude it. For example, a patient could have hyperchloremia and therefore, increased strong anions (Cl^-) and hypoalbuminemia causing decreased weak anions (A^-). If the increase in strong anions is equal to decrease in weak anions, HCO₃^- concentration will be normal. This patient’s blood acid-base analysis, will reveal normal acid-base parameters (pH 7.4, P_CO2  40 mm Hg, HCO₃ 24 mmol/L) and normal SBE of 0 mmol/L despite hyperchloremic acidosis. Similarly, in a patient with lactic acidosis and hypoalbuminemia, blood acid-base analysis may reveal normal parameters, falsely suggesting normal acid-base status.

 

Therefore, SBE may miss metabolic acid-base abnormality, if multiple pathologic processes, acting in opposite direction, are present. 

 

Another major shortcoming of both Boston and Copenhagen approaches is inability to identify mechanism of acid-base disorder.

 

 

7.   Story of anion gap     

In 1977, to overcome limitations of Copenhagen (SBE) and Boston (P_CO2/HCO₃^-) approaches, Emmette and Narins proposed anion gap (AG) based on the law of electroneutrality. As non-carbonic buffers (weak anions [A]) are not considered while assessing acid-base by blood gas machine, sum of the difference in charge of major extracellular ions (Na⁺, Cl^- and HCO₃^-) reveals a gap” of 14-16 mEq/L. This gap is equal to unmeasured non-carbonic buffer (albumin, phosphate). Emmette and Narin termed this gap as anion gap.

 

Na⁺ - Cl^-  = HCO₃^- + A^-

(Na⁺ - Cl^-) - HCO₃^- = A^- = Anion Gap

 

If a patient develops metabolic acidosis due to accumulation of unmeasured anions (organic acids- lactic acid, ketoacids etc.), bicarbonate being a major buffer for metabolic acids, decreases, leading to widening of anion gap to more than 16. Therefore, high anion gap is a measure accumulation of unmeasured anions (organic acids), which is defined as high anionic gap metabolic acidosis (HAGMA).

 

In presence of unmeasured anions (UA), above equation may be written as-

 

Na⁺ - Cl^- - UA = HCO₃^- + A^- 

(Na⁺ - Cl^-) - HCO₃^- = A^- + UA = high anion gap

If a patient has metabolic acidosis due to hyperchloremia, increase in chloride will balance out decrease in bicarbonate and therefore, anion gap will be normal. This was defined as normal anionic gap metabolic acidosis (NAGMA).

 

But, in a patient with metabolic acidosis due to accumulation of unmeasured anions (organic acids), presence of hypoalbuminemia, which is a common in critically ill patients, results in normal anion gap. Decrease in bicarbonate due to accumulation of organic acid is balanced by increase in bicarbonate due to hypoalbuminemia, forcing anion gap to remain normal.

 

Corrected anion gap- AG_cor

 

This deficiency in anion gap was corrected by Figge and colleagues who proposed corrected anion gap (AG_cor) by introducing a correcting factor for albumin.

 

Corrected anion gap (AG_cor) = calculated anion gap + 2.5 × (normal albumin g/dL − observed albumin g/dL)

However, major limitation of anion gap is the use of bicarbonate in the equation. An alteration in [HCO₃^-] concentration can occur as compensation of respiratory acid-base disorder. Also, like base excess, anion gap may underestimate the extent of the metabolic disturbance in presence of multiple pathologic processes, acting in opposite direction.  Anion gap is, nevertheless, useful in discriminating acidosis between unmeasured anions (organic acids) and hyperchloremic acidosis.

 

Both Copenhagen and Boston approaches are based upon Henderson-Hasselbalch equation considering bicarbonate as independent variable in acid-base determination. Together, they were named as standard theory of acid-base balance, to differentiate from a revolutionary theory that was going to be proposed in the eighth decade of last century.

 

The standard theory believes in Bronsted-Lowry definition of acid-base and follows four basic principles- (1) acid is a H+ donor and base is a H+ acceptor, (2) the quantity of H+ added to or removed from the blood determines the final pH, (3) plasma membranes may be permeable to H+, and thus intracellular as well as extracellular chemical reactions influence the pH, (4 ) an analysis of non-carbonic buffers (albumin, phosphate) is not necessary to describe acid- base balance.

 

According to the Henderson-Hasselbalch equation, the pH of serum depends upon only two variables, P_CO2 and HCO₃^-. Rapid elimination of CO₂ by respiration efficiently regulates the concentration of these two interdependent variables and determines hydrogen ion concentration, thereby, obviating the need to describe non-carbonic buffers mathematically.

 

However, standard theory fails to identify mechanism of acid-base disorders and may falsely show normal status in presence of multiple, interacting acid-base disorders.

 

8.   Story of Strong ion difference

In 1981, unsatisfied with standard theory, Peter Steward, a Canadian physiologist put forward a new theory, that came to be known as Stewart’s theory of acid-base analysis. 

 

Considering Arrhenius definition of acid-base, Stewart proposed six assumptions: (1) an acid is any species that raises the hydrogen ion concentration [H+] of a solution; (2 ) [H+] is dependent variable and therefore, the quantity of H+ added or removed from a physiologic system is not relevant to the final pH; (3 ) human plasma consists of fully dissociated ions called strong ions (sodium, potassium, Mg, Ca, chloride, and lactate), partially dissociated weak acids (non-carbonic buffers- albumin and phosphate), and carbonic buffer (bicarbonate); (4) an evaluation of weak acids (non-carbonic buffers- albumin and phosphate) is important to the description of acid- base balance; (5) the weak acids (non-carbonic buffers- albumin and phosphate) of plasma can be described as a pseudomonoprotic acid, HA; and (6 ) as plasma membranes may be permeable to strong ions, transport of strong ions across cell membranes may influence [H+]. 

 

In contrast to standard theory (bicarbonate-centered approaches), Stewart proposed that SID, total concentration of weak acids (A_TOT) and P_CO2 are three independent variables that determine acid-base balance in body and bicarbonate and hydrogen ion concentrations are dependent on combined effect of these three independent variables.

Further, utilizing physiochemical approach, Stewart suggested that it is possible to determine effect of strong ions, weak acids and carbon dioxide on water dissociation, and hence hydrogen ion concentration. 

 

Strong ions difference (SID)

 

Strong ion difference (SID) is charge difference between strong cations and strong anions. It is similar to buffer base (BB) introduced by Singer and Hastings in 1948. Strong ions are completely dissociated at physiological pH. Major strong ions in extracellular space are sodium (Na⁺) and chloride (Cl^-). Other strong ions are K⁺, Ca²⁺, Mg²⁺ and lactic acid. Organic acids (ketoacids, sulfate etc.) are also strong anion, and are depicted mathematically as unmeasured anions, as they are not routinely measured.

 

SID = ([Na⁺] +[K⁺] +[Ca²⁺] + [Mg²⁺]) - ([Cl^-] + [Lactic acid]) = 40-44 mEq/L

 

Strong ion difference (SID) is always positive and is balanced by an equal amount of buffer base, carbonic (HCO₃^-) and non-carbonic (weak acids [A_TOT]- albumin, phosphate). SID independently influences water dissociation via electrical neutrality (sum of all positive and negative charges should be equal to zero), and mass conservation (if all other factors such as P_CO2, albumin and phosphates are kept constant). Thus, an increase in SID will decrease hydrogen ion liberation from and water and cause alkalosis. A decrease in SID will increase hydrogen ion liberation and cause acidosis.

 

Total concentration of Weak Acids- [A_TOT]

 

[A_TOT] is sum of non-carbonic buffer anions and their corresponding acids. In practical terms, it is same as weak acids (A^-).

[A_TOT] = [HA] + [A−]

 

Total concentration of weak acids [A_TOT] are non-carbonate buffers like albumin and phosphates. Their degree of dissociation is related to temperature and pH. The independent effect of weak acids, symbolized as A_TOT, on acid base balance, depends on absolute quantity and dissociation equilibria. Failure to account for weak acids (A_TOT) limits the applicability of standard theory of acid base balance in critical ill patients. 

 

The principal weak acids routinely measured in plasma are albumin and phosphates of which albumin usually being more important quantitatively. However, In critically ill patients with hypoalbuminemia and renal failure with severe hyperphosphatemia, phosphate also becomes a major contributor. Total amount of weak acid (A_TOT) present is an independent contributor to acid-base status and has a reciprocal relationship with bicarbonate concentration. The bicarbonate concentration will decrease with an increase in the total weak acids and vice versa. 

A reduction in (A_TOT) leads to metabolic alkalosis and increase in weak acids leads to metabolic acidosis.

 

Carbon dioxide 

 

The major source of acid in body is carbon dioxide, produced as by-product of aerobic metabolism. Carbon dioxide is buffered principally by haemoglobin. Deoxyhemoglobin is strong base, and therefore better buffer than oxyhemoglobin. Carbon dioxide easily passes through cell membranes of erythrocytes where it combines with H₂O, catalysed by carbonic anhydrase to form H₂CO₃, which ionizes to hydrogen and bicarbonate. Hydrogen ions bind to histidine residues on deoxyhemoglobin while bicarbonate is actively pumped out of cells. Chloride moves inwards to maintain electroneutrality (chloride shift). Large increase in P_CO2 (respiratory acidosis) overwhelms this system, leading to a rapid, dramatic drop in pH.

 

With these assumptions on SID, weak acids and carbon dioxide, Stewart wrote six equations to determine [H⁺] concentration.

Water dissociation equilibrium: [H+]× [OH−] = K_W

Weak acid dissociation equilibrium: [H+]× [A−]=K_A× [HA] 

Conservation of mass for weak acids: [HA] + [A−] =[A_TOT] 

Bicarbonate ion formation equilibrium: [H+] × [HCO₃−] = K_C × P_CO2

Carbonate ion formation equilibrium: [H+] × [CO₃ ²−] = K₃ × [HCO₃−] 

Electric neutrality: [SID] +[H+] −[HCO3−] −[A−] −[CO₃ ²−] −[OH−] =0 

 

Although above equations look relatively simple, fourth order polynomials are required for resolution.

Solving the equation for [H⁺]:

 

[SID] + [H⁺] – K_C * Pc/[H⁺] - K_A * [A_TOT] / (K_A + [H⁺] – K₃ * K_C P_C / [H⁺]² – K_W/[H⁺] = 0

 

In simpler terms, [H+] is a function of SID, A_TOT, P_CO2 and a number of constants. All other variable notably [H⁺], [OH^-] and [HCO^3-] are dependent, and thus cannot independently influence acid base balance. As a result, it is possible to reduce all acid base abnormalities into a problem related to one or more of these three variables (SID, A_TOT, P_CO2).

 

In traditional theory, metabolic acid-base disorders are caused by the production or removal of H+. In Stewart’s theory, three independent variables SID, A_TOT, P_CO2 determine [H⁺] and cause acid-base disturbances. A change in pH may be brought about only by a change in one or more of these variables. The carbonate buffer [HCO₃^-] is dependant variable and as such do not and cannot regulate [H⁺]. 

 

Thus, acid-base disorders can be classified according to Stewart’s three independent variables. 

 

Acidosis results from decrease in SID or increase in  P_CO2 and A_TOT. Metabolic acidosis may be due decrease in SID resulting from accumulation of organic acids (lactic acid, ketoacids, sulfates, formic acid, salicylate etc.), loss of cations (Na⁺ loss in diarrhoea) and gain of anions (Cl^- gain in  renal tubular disorders- RTA or NS infusion) or increase in A_TOT from hyperalbuminemia and hyperphosphatemia.  

 

Alkalosis results from increase SID or decrease in P_CO2, and A_TOT. Metabolic alkalosis results from increased SID resulting from gain of cations (Na⁺ gain from hypertonic saline, dehydration), loss of anions (Cl^- loss in vomiting, diuretic use) or decrease in A_TOT from hypoalbuminemia.

 

In contrast to Singer and Hastings buffer base and standard theory, Stewart considered changes in [A_TOT] mainly albumin, as separate acid-base abnormality.

 

Thus, unlike base excess, Stewart’s three independent variables may quantify as well as explain mechanism of acid-base disorders. But, complex calculations make it difficult to apply clinically at bedside. To overcome this limitation, Fencle, Gilfix, Balsubramanayan and Story tried to simplify Stewart’s theory for practical applicability.

 

Story of strong ion gap (SIG)

 

Fencle along with Stewart proposed strong ion gap (SIG) as measure of unmeasured anions (organic acids like ketoacids etc). SIG is difference between apparent SID (SID_a) and effective SID (SID_e). 

 

Apparent SID (SID_a) is defined as sum of all strong ions; 

SID_a = [Na+] - ([Cl−]

 

Effective SID (SID_e) is defined as sum of bicarbonate (HCO_3-) and total concentration of weak acids [A_TOT]. In other words, SID_e is total buffer base, carbonic and non-carbonic.

 

SID_e = [HCO₃^-] +[charge on albumin] + [charge on Pi] (in mmol/L).

SID_e = HCO₃^- + A_TOT

SIG = SID_a – SID_e

       = ([Na+] - ([Cl−]) – ([HCO₃^-] - + [A_TOT])

[A-] = 2.8 [albumin g/dL] + 0.6 [phosphate mg/dL] 

 

The SIG is an estimate of unmeasured ions and mathematically similar to anion gap.

 

AG = [Na+] – [Cl^-] – [HCO₃^-]  

SIG = [Na+] - ([Cl−] – [HCO₃^-- + A^-]

SIG = |([Na+] - ([Cl−] – [HCO₃^-]) – [ A^-]

SIG = AG - [A-] 

 

But, unlike the anion gap, the SIG is normally close to zero. In Analogy to conventional interpretation of the anion gap, metabolic acidosis with a high SIG is due to unmeasured strong anions while metabolic acidosis with normal SIG (0 mEq/L) is usually due to hyperchloremia.  

 

Although accurate, the SIG is cumbersome and expensive, requiring measurement of multiple ions and albumin.

 

base deficit- excess gap- BDE gap

 

Gilfix, Balasubramanyan, and Story independently, proposed base deficit- excess gap (BDE gap) by combining base excess and stewart’s approaches. This amalgamation allows calculation of BE using base excess effect of strong ions (Na⁺, Cl^-, Lactate) and weak acids (albumin). Base deficit- excess gap is calculated as difference between estimated base excess (estimated by blood gas machine) and calculated base excess (sum of the base excess effect of Na⁺, Cl^-, Lactate and albumin). 

 

Base excess effect is calculated by difference between value of measured variable in meq/L and normal value of variable in meq/L.

 

The resulting base deficit-excess gap resembles strong ion gap (SIG) and anion gap (AG). Story and associates called this simplified Stewart approach.

 

Base excess effect of Strong ions (Na⁺, Cl^- and Lactate)

 

Base excess effect of Na and Cl

Na⁺ and Cl^- are major anions in extracellular space. Base excess effect of Na⁺ and Cl^- is the difference between measured SID (measured [Na⁺ - Cl^-]) and normal SID. Considering normal reference value of Na and chloride as 140 and 102 meq/L, normal SID would be 38. Therefore, 

 

Base excess effect of Na⁺, Cl^- in meq/L (BE_NaCl) = measured [Na⁺] – measured [Cl^-] – 38

 

Base excess effect of Lactate

Normal reference value of lactate is 1 meq/L, therefore base excess effect of lactate may be estimated as follows: 

Lactate base-excess effect in meq/L (BE_Lactate) = 1 − measured lactate. 

 

Base excess effect of Albumin

The principal weak acid in plasma is albumin. As albumin is measured in g/dl, Figge suggested a simple method to calculate the effective ionic concentration of albumin in meq/L as,

Ionic concentration of albumin in meq/L = 2.5 × albumin gm/dL

Ionic concentration of normal albumin level is (4.2 g/dL) = 2.5 × 4.2 g/dL = 10.5 meq/L. Base excess effect of albumin is calculated as difference between ionic concentration of measured albumin and ionic concentration of normal albumin value (4.2 g/dl)

Base excess effect of Albumin in meq/L (BE_Alb) = 2.5 × (4.2 − measured albumin gm/dL)

Therefore, for every 1.0 g/dL decrease in plasma albumin, the base excess will increase by 2.5 meq/L, making the patient more alkalotic. 

 

Calculated base excess

Summing up the base excess effect of strong ions (Na⁺, Cl^-, Lactate) and albumin gives calculated base excess.

Calculated Base-excess (BE_Cal) = BE_NaCl + BE_Lactate + BE_Alb

Calculated Base-excess (BE_Cal) = (Na⁺ - Cl^- – 38) + (1 – lactate) + 2.5 × (4.2 −albumin)   

 

Base deficit-excess gap (BDE gap)                                

Base deficit-excess gap (BDE gap) = estimated base excess (BE_est) – calculated base excess (BE_Cal)

In normal acid base status, estimated base excess would be equal to calculated base excess. In patient with unmeasured anions (organic acids like ketoacids, sulphates etc) estimated base excess would be more than calculated base excess, leading to positive base deficit-excess gap. 

 

Mathematically BDE gap is similar to strong ion gap (SIG) and anion gap (AG).

 

BE NaCl = ([Na+] – [Cl-]) − 38

BE Lactate = 1 – [lactate meq/L]

BEAlb = 2.5 (4.2 − albumin g/dL)

BEcalc = BENaCl + BELactate + BEAlb

BDE gap = BEest (measured by blood gas machine)− BEcalc = unmeasured anions (organic acids- ketoacids, sulphate etc.), normal BDE gap is zero

These calculations simplify the framework for “eyeballing” a chemistry series:

Normal Na+ = 140 mEq/L

For every 1 mEq/L increase in Na+ from 140, base excess increases by +1 (Na+ 150 = BE +10 = contraction alkalosis)

For every 1 mEq/L decrease in Na+ from 140, base excess decreases by −1 (Na+ 130 = BE −10 = dilutional acidosis)

Normal Cl- = 102 mEq/L

For every 1 mEq/L increase in Cl- from 102, base excess decreases by -1 (Cl- 110 = BDE − 8 = hyperchloremic acidosis)

For every 1 mEq/L decrease in Cl- from 102, base excess increases by +1 (Cl- 90 = BDE +12 = hypochloremic alkalosis)

Normal albumin = 4.2 g/dL

For every 1.0 g/dL decrement in albumin from 4.2, there is a 2.5 mEq/L increase in base excess (metabolic alkalosis)

 

Let’s conclude this story with evaluation of blood gas of a patient with Traditional and simplifies Steward approach separately-

 

Consider the following patient, 

Sodium 117, potassium 3.9, chloride 92, inorganic phosphate 0.6 mmol/L, albumin 0.6 gdL, lactate 1.0 mmol/L, pH 7.33, Pco2 30 mm Hg, HCO3 15, anion gap 13, base excess −10,

 

Standard theory interpretation-

Boston approach- Acidemia, compensated metabolic acidosis

Anion gap approach- Corrected AG = NAG + 2.5 × (4.2 −albumin) = 13+9 = 22

            Delta gap = (AG-12)- (24-HCO3) = (22-12) – (24-15) = 10-9= 1

              HAG metabolic acidosis with NAG metabolic alkalosis

              Acidemia, compensated metabolic acidosis- HAGMA + NAGMA, unmeasured anion +

Copenhagen approach- Acidemia, compensated metabolic acidosis, unmeasured anion= 10.

 

Stewart Theory interpretation

SIG approach- SID_a = 117-92- 1 = 24, SID_e = 15 + 1.44 + 0.36 = 16.8

                         SIG = 24 – 16.8 = 7.2 = unmeasured anion

BDE gap approach- BE_NaCl = 117 – 92 - 38 = -13, BE_Lac = 1 - 1 = 0, BE_Alb = 2.5 × (4.2 – 0.6) = 9,

BE_Cal = -13 + 0 + 9 = - 4

BDE gap = BE_Est - BE_Cal = -10 – (-4) = 6 = unmeasured anions

Acidemia, metabolic acidosis- unmeasured anions, hyponatremia (dilutional), metabolic alkalosis- hypochloraemia, hypoalbuminemia, hypophosphatemia.

 

Traditional theory does not reveal complete picture. Stewart approach reveals a much more complex situation. SID is reduced to 24 mEq/L, caused by free water excess (hyponatremia) and hypochloremia. High SIG and BDE gap suggest accumulation of unmeasured anions (organic acids- ketoacids, sulphate etc). BE calculation for albumin shows metabolic alkalosis due to hypoalbuminemia.

 

The degree of acidosis (pH 7.33 and BE -10) does not mirror this metabolic disturbance, because of alkalizing effect of hypoalbuminemia and hypophosphatemia. The corrected anion gap mirrors the change in SID, but this is grossly underestimated by the base excess.