Disturbances of acid-base homeostasis are common, particularly in the critical care setting as a result of cardiovascular, pulmonary, metabolic, and/or renal complications. Digression of [H⁺] outside the normal range (tables 3 & 4 ) can have marked effects on cell metabolism and membrane function. Severe deviations cause neural dysfunction and may lead to coma and death. The range of [H⁺] compatible with life is approximately 16 - 160 nanomoles/L (pH: 6.8 - 7.8).

Advances in medical technology have made the measurement of "arterial blood gases" including the acid-base variables (pH and P_CO2) simple and readily available. However, correct interpretation and use of these lab values in the diagnosis and treatment of acid-base disorders remain elusive. Poor understanding of acid-base chemistry and physiology stems largely from the confusing mathematical expressions (negative logarithms, inverted fractions, etc.) that are commonly used in this field. The problem is further complicated by the rote use of vague, derived parameters such as "base excess" and "standard bicarbonate". A good grasp of the fundamental principles of acid base chemistry and physiology is both necessary and sufficient to correctly diagnose and quantify acid-base disturbances.

1. An hydrogen ion (H⁺) = an hydrogen atom (H) minus its sole electron.
2. H⁺ is also called a proton, because that is all that remains when the H atom loses its electron.
3. An acid is, in most cases, a molecule or ion that tends to dissociate releasing protons into the solution. An acid may also be a molecule that binds (takes up) a hydroxide ion (OH^-). In other words, an acid is a substance that, when added to a solution, it increases [H⁺] relative to the [OH^-] in the solution.
   • Examples of acids: acetic, ß-OH-butyric, HCl, H₂CO₃, H₂PO₃^-, NH₄⁺
   • When a molecule releases a proton, it becomes a conjugate base (Table 1).
4. A base is a molecule that tends to pick up (bind) protons from the solution or release OH^- into the solution. In other words, a base is a substance that, when added to a solution, it increases the OH^- concentration in the solution relative to H⁺ concentration.
   • Examples of bases: HCO₃^-, NH₃ , HPO₃⁼, tromethamine (THAM^® or tris buffer)
5. H⁺ concentrations in biological fluids are extremely small, especially when compared to other ions. Whereas the plasma concentration of K⁺ may be 4 mmol/L, that of H⁺ is normally 0.00004 mmol/L or 0.00000004 moles/L. To overcome the difficulty in dealing with these small fractions, two methods were devised to express H⁺ concentration.:
   1. In 1909 Sorenson invented the "potential of hydrogen" or pH:

      pH = - log [H+] = log (1 / [H+] )

      
      
      where: [H⁺] = H⁺ activity in moles per liter.
      Using this convention, there is an inverse relationship between H⁺ concentration and pH. As a solution becomes more acid, i.e., as the [H⁺] increases, the pH decreases.
      
   2. In 1962 Campbell expressed [H⁺] in nanomoles/L; one nanomole = 10^-9 mole = 1/(1000,000,000) mole
   
   

   [H+], nmol/L = 109 - pH

6. Body fluids are normally significantly alkaline relative to a neutral solution (pH = 6.65) at the same temperature (37° C).

7. The pK_a
   The tendency of a weak acid (HB) to dissociate (HB ----> H⁺ + B^- ) depends in part on the [H⁺] (or the pH ) of the solution, and in part on its inherent strength, i.e., on its dissociation constant (K_a), which is given by the ratio [H⁺][B^-] / [HB]. The K_a values for weak biological acids are extremely small fractions. Therefore, they are commonly expressed as their negative logarithms (pK_a = - log K_a). Thus, the pK_a of an acid is the negative logarithm of its dissociation constant (K_a = [H⁺][B^-] / [HB]). The pK_a of an acid corresponds to the pH of the solution when the acid is 50% dissociated (i.e., when [B^-] = [HB] ). Strong acids dissociate to a greater extent (high K_a ) and have low pK_a values. The opposite is true for weak acids. Thus, the higher the pK_a the weaker is the acid and vice versa..
   

Acidemia: [H⁺] > 44 nmoles/L or a pH < 7.36.

Acidosis: a physiologic or pathophysiologic process that tends to cause acidemia.

Alkalemia: [H⁺] < 36 nmoles/L ( pH > 7.44).

Alkalosis: a physiologic or pathophysiologic process that tends to cause alkalemia.

1. Sulfuric acid (H₂SO₄), which is derived from the sulfur-containing amino acids such as methionine and cysteine/cystine (e.g., methionine glucose + urea + 2 H⁺ + SO₄⁼).
2. Phosphoric acid (H₂PO₄^-), which is derived from phospholipids, phosphoproteins, nucleic acids, etc.

Net fixed-acid production is highly dependent on diet; it is high with protein-rich diets. However, a round figure of one mEq/kg body wt/day (50 - 100 mEq per day) is a reasonable estimate for an average meat-eating adult. In addition to normal metabolism, there are pathological conditions that can result in abnormally high rates of fixed acid production:

• Keto-acidosis: Excessive lipolysis excess acetyl-CoA hepatic production of ketone bodies (ß-OH-butyric and acetoacetic acids). Diabetes mellitus, starvation, and severe exercise all can lead to ketosis.
• Lactic acidosis: Absolute or relative hypoxia promotes anaerobic glycolysis and increases lactic acid production. This occurs, for example, during severe exercise, or as a result of circulatory shock (hemorrhagic, cardiogenic, septic, etc). Excessive lactic acid production is often triggered by a relative decrease in tissue oxygen delivery or hypoxia brought about by circulatory and/or pulmonary problems. It may also be drug-induced (e.g., biguanide overdose).

1. Chemical buffering A buffer is a mixture of an undissociated weak acid (HB) and its conjugate (buffer) base (B^-). Although normally the protons produced in the course of cell metabolism (or added to body fluids from exogenous sources) are eventually excreted in the urine, the renal elimination does not take place immediately. Chemical buffering provides a transient solution to this problem. Protons are bound (neutralized) by buffer bases like HCO₃^-. This process minimizes [H⁺], but it consumes buffer bases.
   H⁺ + B^- ® HB.
   The body's buffer-base stores are eventually replenished at the same time that protons are excreted in the urine.
   By contrast, the undissociated acid (HB) component of the buffer pair guards against added base (OH^-):
   OH^- + HB ® B^- + H₂O.
   The buffer capacity of a solution (its the ability to maintain a given pH) is defined as the amount of H⁺ that can be added to, or removed from, a solution per 1 unit change in pH. It depends on the total concentration of the buffer ([HB] + [B^-]) and on its pK_a . The closer is the buffer pK_a to the pH of body fluid the higher is the buffering capacity.
   The main ICF buffers are proteins and organic phosphates, and the main ECF buffer is HCO₃^-. Hemoglobin is a major blood buffer, and so is plasma HCO₃^-.
   Hemoglobin plays a major role in the buffering and transport of CO₂ (see Figure below). It has a favorable buffer profile for it is normally present in relatively large amounts ( 14 - 17 grams per dL), and the pKa of the imidazole groups (of the histidine rediues) gives it a dynamic pKa that brackets the pH of body fluids (HbO₂ = 6.7 and HbH = 7.9).
   

   

   The HCO₃^- buffer system consists of HCO₃^- (the buffer base) and CO₂ which represents the conjugate acid (H₂CO₃). Unlike other buffer pairs, the HCO₃^-/ CO₂ represents an open system in view of the gaseous nature of its acid (CO₂). An additional major advantage of HCO₃^-/ CO₂ system is the fact that the two components are regulated by different organ-systems and somewhat independently - HCO₃^- by the kidney and CO₂ by the respiratory system. Due to the HCO₃^-/ CO₂ system in the ECF, the buffer response to added acid or base is essentially instantaneous albeit incomplete (the reactions involving bone and ICF buffers may take a few hours to complete).

   Approximately 45% of the total buffer response to an acid load is provided by extracellular bicarbonate (the CO₂ generated in this process is immediately eliminated in the expired gas). The remaining 55% is provided by intracellular proteins and phosphates and by bone buffers (carbonate and phosphate).

2. Ventilation
   As stated above, the respiratory system normally eliminates CO₂ as fast as it is produced. This is due to the fine control provided by the central and peripheral chemoreceptors. The central chemoreceptors are sensitive to minor changes in [CO₂] (and hence in [H⁺]). An increase of only 3-5 mm Hg in P_CO2 can lead to a doubling of the alveolar ventilation rate (V_A). Acid-base alterations can have marked effects on the ventilation rate. These effects are mediated mainly by the central chemoreceptors, but also by the peripheral chemoreceptors: ñ[H⁺] Þ ñ V_A Þ ò P_CO2 Þ Þ ò [CO₂] Þ ò [H⁺]

   Thus, whenever the [H⁺] is increased for non-respiratory reasons, the body compensates by increasing ventilation and lowering [CO₂]. This respiratory response is relatively rapid (minutes).

   

   Renal H⁺ Excretion And HCO₃^- Regeneration
   Normally, the renal tubule cells in almost all parts of the nephron generate and secrete protons (H⁺) into the tubule lumen. Protons are generated inside the tubule cell in the process of CO₂ hydration which is catalyzed by carbonic anhydrase (CA):
   CO₂ + H₂O ® H₂CO₃ ® H⁺ + HCO₃^-

   In the tubule lumen, the secreted proton may combine with HCO₃^-, ammonia (NH₃), or HPO₄⁼. Thus, proton secretion serves to reabsorb (save) the filtered HCO₃^-, and to excrete the fixed acid in the form of ammonium (NH₄⁺) and titratable acids (mainly H₂PO₄^-). Urinary ammonium excretion is usually twice as great as the excretion of titratable acids. Ammonium (NH₄⁺) is derived primarily from glutamine metabolism in the tubule cells. It is then secreted into the lumen of the proximal tubule, only to be reabsorbed by the thick ascending loop of Henle into the interstitium of the renal medulla. From the interstitium (where NH₄⁺ and NH₃ exist in equilibrium) NH₃ can diffuse into the lumen of the collecting ducts where it combines with the secreted protons to be excreted in the urine. In chronic acidosis, the capacity of the kidney to produce NH₃/ NH₄⁺ and excrete fixed acids is increased by as much as four fold.

   The buffer bases in the tubule fluid (mainly NH₃ and HPO₄⁼) permit the excretion of fixed acids without exceedingly lowering the tubule fluid pH. However, as these bases are exhausted toward the end of the nephron, the urine pH falls markedly. The final urine pH may be as low as 4.5.

   For each H⁺ excreted in the form of titratable acids (e.g., H₂PO₄^-) or (NH₄⁺) one HCO₃^- is added to the blood. This new HCO₃^- replenishes the body store of buffer base. Under steady state conditions, HCO₃^- is generated by the kidney at the same rate as it is consumed in the process of buffering the fixed acids produced by metabolism. Indeed, the role of the kidney in acid-base balance may be summarized as follows:

   "it maintains an adequate level of buffer bases in body fluids". When there is a deficit the kidneys generate new bicarbonate in the process of acid excretion. When there is an excess of bicarbonate (e.g., purely vegetarian diets, administration of citrate with blood transfusions, etc.) the healthy kidneys function like an overflow system - the excess is simply eliminated in the urine.

   The rate of proton secretion by the tubule cells is related to the degree of intracellular acidosis (­[H⁺]_i). Furthermore, as the acidosis becomes chronic the production and secretion of N₄H⁺ can increase up to four times the normal rate. Consequently, with normal renal function, fixed-acid excretion and HCO₃^- generation during acidosis are increased several fold. In chronic renal failure, the capacity of the remaining nephrons to excrete acid and generate HCO₃^- is increased several fold. However, the number of these functioning nephrons may not be sufficient to maintain acid-base balance, hence the development of metabolic acidosis in these patients.

The elements of the acid-base picture are the pH (or [H⁺]), P_CO2, and HCO₃^-. These are usually determined in a sample of arterial blood. In the Henderson-Hasselbalch equation, the normal pH of 7.4 corresponds to a [HCO₃^-] / [CO₂] ratio of 20.

pH = 6.1 + log ([HCO₃^-] / [CO₂])

pH = 6.1 + log { 24 / (0.03)(40)}

pH = 6.1 + log 20 = 6.1 + 1.3 = 7.4

Deviations of this ratio from its normal value of 20 result in acid-base disturbances. Such deviations could be due to a primary change in either the numerator ([HCO₃^-]) or in the denominator ([CO₂]) . A primary acid-base disturbance can either be an acidosis (­[H⁺]; ¯pH) or an alkalosis (¯[H⁺]; ­pH). Disturbances caused by a primary change in [HCO₃^-] are called metabolic or non-respiratory. Those caused by a primary change in [CO₂] are called respiratory. A primary acid-base disturbance is almost always accompanied by a secondary response (compensation) that is aimed at minimizing the effect of the primary disturbance on pH (i.e., minimizing pH). A primary metabolic disturbance is accompanied by a respiratory compensation, whereas a primary respiratory disturbance is accompanied by a metabolic (renal) compensation. It is also common for more than one primary disturbance to coexist in the same patient. There are four primary acid-base disturbances:

1. Non-respiratory (metabolic) acidosis, in which the decrease in pH is due to a decline in the level of buffer bases (reflected in ¯[HCO₃^-]). Metabolic acidosis can develop for a variety of reasons (see table ). The decline in [HCO₃^-] could be due to excessive metabolic acid production as in diabetic ketoacidosis and circulatory shock (lactic acidosis) or it could be due to the inability of the diseased kidney to reabsorb (save) the filtered HCO₃^- (e.g., dysfunction of the proximal tubule) or its inability to excrete the fixed acids produced by metabolism. The metabolic acidosis could also be due the loss of HCO₃^- via the gastrointestinal tract (e.g., diarrhea).
2. Non-respiratory (metabolic) alkalosis, in which the fall in [H⁺] (i.e., pH ) is due to an increased level of buffer bases (i.e., [HCO₃^-]). Many factors can elevate extracellular [HCO₃^-] including ECF-volume contraction, gastric H⁺ loss (vomiting or via nasogastric suction), and urinary H⁺ loss (loop or thiazide diuretics, hyperaldosteronism, etc)(see table ).
3. Respiratory acidosis, in which the decrease in pH is caused by a primary increase in CO₂ concentration (P_CO2> 44 mm Hg), which is often referred to as hypercapnia or CO₂ retention. It is always due to alveolar hypoventilation (¯V_A). The hypoventilation and the hypercapnia may be acute (e.g., choking or drug overdose) or chronic as in COPD. In the critical care setting, the most common causes are narcotic-induced respiratory center depression, and prolonged neuromuscular blockade. The renal compensation to hypercapnia is slow and incomplete. Also, hypercapnia is often accompanied by hyperkalemia ([K⁺]_o) as a result of K⁺ - H⁺ exchange between the ICF and the ECF. Removing the underlying cause is the first step in the treatment of hypercapnia. The P_CO2 should then be lowered gradually to its "normal" level. In the case of chronic hypercapnia this should be done over 2-3 days to avoid a sudden rise in CSF pH which can cause convulsions.
4. Respiratory alkalosis, in which the rise in pH is caused by a primary decrease in CO₂ concentration (P_CO2< 36 mm Hg) (hypocapnia). It is always due to alveolar hyperventilation (­V_A). . In the critical care setting, the most common causes is manual or mechanical hyperventilation. However, pain, fear, anxiety, and hysteria are also frequent causes of hyperventilation. The hyperventilation may also be due to a more serious cause such a neurological disorder (e.g., meningitis, encephalitis, trauma). Hyperventilation can be induced by salicylate overdose, fever and thyrotoxicosis. Also, hyperventilation can be induced by Gram-negative bacteremia, and by the hypoxemia associated with pneumonia, congestive heart failure, severe anemia, shock, etc.

* Bartter's is a rare syndrome characterized by hyperplasia of the juxtaglomerular cells of the JGA, hyper-reninemia, and hyperaldosteronism. In addition, there is a defect in ion transport in the thick ascending loop which results in Na loss and ECF volume depletion..

In all cases, the underlying cause of the disturbance should be identified and treated.

I. Normal Anion Gap
If arterial [HCO₃^-] < 8, it should be raised to 12 mmol/L. This should be done very slowly to avoid reducing intracellular pH further. The volume of distribution of HCO₃^- is equivalent to approximately 50% of Body weight.
Amount of bicarbonate required = 0.5 x Body Weight (12 - actual [HCO₃^-])

Thus if a patient weighs 65 kg and his actual [HCO₃^-] = 7 mmol/L , then:

Amount of bicarbonate required = 0.5 x 65 (12 - 7) 165 mmoles.

II. Wide Anion Gap
The first objective is to remove the underlying cause of acidosis if possible. Examples include improving cardiopulmonary function and tissue O₂ delivery in a patient with lactic acidosis, or providing insulin to a diabetic patient with ketoacidosis. Metabolic acidoses with wide anion gap rarely require the infusion of exogenous HCO₃ because the organic anions (lactate, ketones, etc) are eventually oxidized and the resulting CO₂ is converted to HCO₃^-.
Ketoacidosis

Regular Insulin: give bolus of 5 - 10 units, then 2 - 6 units / hr. After insulin, the infusion of an alkalinizing agent is rarely necessary.
If there is Na deficit (5 - 10 mmol/kg), infuse normal saline (150 mmol/L NaCl) until hemodynamically stable, then infuse half normal saline (75 mmol/L NaCl ). Supplement the saline with 20 - 40 mmoles of K⁺.

Treat underlying cause; improve tissue oxygenation (cardiac output, BP, P_O2, etc.).
Infuse a buffer base only if arterial [HCO₃^-] < 8 mmol/L. NaHCO₃ is no longer the buffer of choice in the treatment of lactic acidosis because of the danger of exacerbating the intracellular acidosis³. This may be avoided by using tromethamine (THAM) or an equimolar mixture of NaHCO₃ and Na₂CO₃ (called carbicarb).

• Treatment with ethanol may be started, based on strong suspicion, before diagnosis is confirmed.
• Oral treatment: Give a bolus of 120 mL of whiskey, followed by 60 mL /hr.
• IV treatment: infuse ethanol to achieve a blood ethanol level of 1 mg / mL (22 mmol/L). In chronic drinkers, tthis may be done by a loading dose of 600 mg/kg followed by 150 mg/kg/hr infusion.. In non-drinkers use 600 mg/kg load followed by 75 mg/kg/hr infusion.