How Should Acid-Base Disorders Be Diagnosed and Managed?

Arterial blood gas (ABG) analysis, a core component of critical care monitoring, provides immediate information on the status of the patient’s respiratory system and whether a state of acidosis or alkalosis is present. With a variety of empiric “rules” applied, the information contained in an ABG is often sufficient to allow one to identify the presence, cause, and progression of a disease. The diagnostic sensitivity of blood gas analysis is augmented when a serum chemistry panel and glucose, lactate, and ketone measurements are added. Unlike abnormalities of serum and urinary electrolytes, radiographic or electrocardiographic changes, there is no clear agreement between intensive care specialists regarding the optimal method of evaluating acid base balance.

Several different approaches to acid–base balance are in widespread use. These can be described as descriptive, based on changes in the Henderson-Hasselbalch equation; semiquantitative, based on calculations and nomograms; or quantitative, based on physical chemistry. The descriptive approach uses the interrelationship between partial pressure of carbon dioxide in arterial blood (Pa co ₂ ) and bicarbonate [HCO3−]
[ HC O 3 − ]
to detect and diagnose acid base abnormalities. An extension of this is the anion gap (AG). The semiquantitative approach includes the buffer base concept, the standardized base deficit–excess, and the base-deficit gap (BDG). The quantitative approach uses strong ion difference (SID) and total weak acid concentration (A _TOT ) and is quantified with the strong ion gap (SIG).

Scientific Background

In the early part of the twentieth century it was widely known that, in critical illness, the carbon dioxide (CO ₂) content of the blood fell. As early as 1831 O’Shaughnessy identified loss of “carbonate of soda” from the blood as a fundamental disturbance in patients dying of cholera. L. J. Henderson, in 1909, coined the term “acid base” balance. He was able to define this process in terms of carbonic acid equilibrium. This work was later refined by Hasselbalch in 1916. Their method described acid-base balance in terms of the hydration equation for CO ₂, the only clinical chemistry test available at that time.

CO2+H2O→H2CO3→H++HCO3−CO2+H2O→H2CO3→H++HCO3−
C O 2 + H 2 O → H 2 C O 3 → H + + HC O 3 −

pH=pKa+log[HCO3−]/[H2CO3]pH=pKa+log[HCO3−]/[H2CO3]
pH = pKa + log [ HC O 3 − ] / [ H 2 C O 3 ]

TotalCO2=[HCO3−]+[DissolvedCO2]+[CarbaminoCO2]+[H2CO3]TotalCO2=[HCO3−]+[DissolvedCO2]+[CarbaminoCO2]+[H2CO3]
Total C O 2 = [ HC O 3 − ] + [ Dissolved C O 2 ] + [ Carbamino C O 2 ] + [ H 2 C O 3 ]

≈PCO2×0.03mmolCO2/L/mmHg≈PCO2×0.03mmolCO2/L/mmHg
≈ PC O 2 × 0.03 mmol C O 2 / L / mm Hg

thus, substituting into the equation above:

pH=6.1+log[HCO3−]/PCO2×0.03:TheHenderson-HasselbalchequationpH=6.1+log[HCO3−]/PCO2×0.03:TheHenderson-Hasselbalchequation
pH = 6.1 + log [ HC O 3 − ] / PC O 2 × 0.03 : The Henderson-Hasselbalch equation

Svante Arrhenius (1859-1927) in 1903 established the foundations of acid–base chemistry. In an aqueous solution, an Arrhenius acid is any substance that delivers a hydrogen ion into the solution. A base is any substance that delivers a hydroxyl ion into the solution. Water is a highly ionizing amphiprotic solution, so substances with polar bonds will dissociate into their component part in it. Water may act as an acid or a base.

The degree of dissociation of substances in water determines whether they are strong acids or strong bases. Thus lactic acid, which has an ion dissociation constant (pKa) of 3.4, is completely dissociated at physiologic pH, and is a strong acid. Conversely, carbonic acid, which has a pKa of 6.4, is incompletely dissociated, and is a weak acid. Similarly, ions, such as sodium, potassium, and chloride, which do not easily bind to other molecules, are considered strong ions—they exist free in solution.

In any solution, the ion dissociation constant for water, Kw′, dictates that the relative ratio of [H ⁺] to [OH ⁻] must always be constant, and electrical neutrality must always hold. Consequently, strong cations, Na ⁺, K ⁺, Ca ²⁺, Mg ²⁺, will act as Arrhenius bases (they deliver hydroxide into the aqueous solution), and strong anions, Cl ⁻, LA ⁻, ketones, sulfate, and formate, will act as Arrhenius acids (they deliver hydrogen ions into the aqueous solution).

The Arrhenius theory was superseded in 1923 by Brønsted and Lowry. They defined acids as proton donors and bases as proton acceptors.

NH3+H2O↔NH4++OH−NH3+H2O↔NH4++OH−
N H 3 + H 2 O ↔ N H 4 + + O H −

In this situation, water is the proton donor, the Brønsted-Lowry acid, and ammonia the proton acceptor, the Brønsted-Lowry (BL) base. Conversely, consider the following reaction:

HCL + H ₂O → H ₃O ⁺+ Cl ⁻

In this reaction, hydrogen chloride acts as a Brønsted-Lowry (BL) acid and water as a BL base.

CO2+H2O↔H2CO3↔H++HCO3−CO2+H2O↔H2CO3↔H++HCO3−
C O 2 + H 2 O ↔ H 2 C O 3 ↔ H + + HC O 3 −

In this reaction, CO ₂is hydrated to carbonic acid, a BL acid that subsequently dissociates to hydrogen and bicarbonate.

Strong Ions

Strong ions are completely dissociated at physiologic pH. The most abundant strong ions in the extracellular space are sodium (Na ⁺) and chloride (Cl ⁻). Other important strong ions include K ⁺, SO42−
S O 4 2 −
, Mg ²⁺, and Ca ²⁺. Each applies a direct electrochemical and osmotic effect.

In the extracellular space the difference between the charge carried on strong cations and strong anions is calculated by

SID=([Na+]+[K+]+[Ca2+]+[Mg2+])−([Cl−]+[Other strong anions:A−])=40−44mEqSID=([Na+]+[K+]+[Ca2+]+[Mg2+])−([Cl−]+[Other strong anions:A−])=40−44mEq
SID = ( [ Na + ] + [ K + ] + [ Ca 2 + ] + [ Mg 2 + ] ) − ( [ Cl − ] + [ Other strong anions: A − ] ) = 40 − 44 mEq

This excess of positive charge, called the strong ion difference, by Peter Stewart, is always positive, and is balanced by an equal amount of “buffer base,” principally in the form of phosphate, albumin, and bicarbonate. SID independently influences water dissociation, determined by electrical neutrality and mass conservation. If all other factors (P co ₂, albumin, and phosphate) are kept constant, an increase in SID, due to a relative increase in the ratio of strong cations to strong anions, will decrease hydrogen ion availability causing alkalosis. A decrease in SID, due to a relative decrease in the relative ratio of strong cations to strong anions, results in greater accumulation of hydrogen ions causing acidosis.

The chief determinant of SID is the relationship between the relative concentrations of sodium, chloride, and free water in extracellular fluid (ECF). The normal ratio of sodium to chloride is approximately 1.4:1. Any process that reduces that ratio reduces SID and leads to acidosis (sodium loss, chloride gain, or free water gain). Any process that increases that ratio increases SID and leads to alkalosis (sodium gain, chloride loss, or free water gain).

Weak Acids

Albumin and phosphate are weak acids, in which the degree of dissociation is related to temperature and pH. Weak acids, represented by the symbol A _TOT, independently influence acid base balance, depending on absolute quantity and dissociation equilibria.

The principal limitation of traditional approaches to acid base balance has been the limited attention paid to changes in A _TOT. Although this may be valid in otherwise healthy patients, perioperative care and critical illness cause hypoalbuminemia due to crystalloid administration, hepatic reprioritization, and capillary leak. A reduction in serum albumin or phosphate leads to metabolic alkalosis. Hypophosphatemia is associated with malnutrition, refeeding, diuresis, and hemodilution. Hyperphosphatemia occurs in renal failure. Hyperphosphatemia leads to metabolic acidosis.

Carbon Dioxide

Aerobic metabolism results in the production of large quantities of CO ₂. CO ₂is hydrated by carbonic anhydrase in red cell erythrocytes to carbonic acid. This liberates the equivalent of 12,500 mEq of H ⁺per day. Hydrogen ions bind to histidine residues on deoxyhemoglobin, and bicarbonate is actively pumped out of the cell. CO ₂exists in four forms: CO ₂[denoted CO ₂(d)], carbonic acid (H ₂CO ₃), bicarbonate ions (HCO3−)
( HC O 3 − )
and carbonate ions CO32−
C O 3 2 −
. The principal mechanism of excretion is via alveolar ventilation, although some CO ₂is excreted from the kidney as bicarbonate as part of a sodium-chloride cotransporter.

Chronic respiratory acidosis is associated with increase in total body CO ₂content, reflected principally by an increase in serum bicarbonate. Mathematically ΔHCO3−=0.5ΔPaCO2.
Δ HC O 3 − = 0.5 Δ PaC O 2 .
It is important that this is not confused with “metabolic compensation for hypercarbia,” a relatively slow process that reduces SID by increased urinary chloride excretion.