Articles on Medical Diseases and Conditions

Category: Acid-Base and pH Measurements

Acid-Base and pH Measurements

  • Partial Pressure of Carbon Dioxide (PCO2) as a Measure of Respiratory Function

    In addition to its use in the Henderson-Hasselbalch equation, PCO2 provides information on pulmonary alveolar gas exchange (ventilation). If PCO2 is high, there is not a sufficient degree of alveolar ventilation. This may be due to primary lung disease (inability of the lungs to ventilate properly) or to some other reason. If PCO2 is low, there is alveolar hyperventilation, again either from primary or secondary etiology.

  • Summary of Acid-Base Changes

    To summarize plasma pH problems, in metabolic acidosis there is eventual HCO–3 deficit, leading to decreased plasma pH and decreased CO2 content (or CO2 combining power). In respiratory acidosis there is primary H2CO3 excess, which causes decreased plasma pH, but the CO2 content is increased due to renal attempts at compensation. In metabolic alkalosis there is eventual bicarbonate excess leading to increased plasma pH and increased CO2 content. In respiratory alkalosis there is primary carbonic acid deficit, which causes increased plasma pH, but the CO2 content is decreased due to renal attempts at compensation. The urine pH usually reflects the status of the plasma pH except in hypokalemic alkalosis, where there is acid urine pH despite plasma alkalosis.

    As noted, CO2 content or combining capacity essentially constitutes the numerator of the Henderson-Hasselbalch equation. PCO2 is essentially a measurement of the equation denominator and can be used in conjunction with pH to indicate acid-base changes. This is the system popularized by Astrup and Siggaard-Anderson. PCO2 follows the same direction as the CO2 content in classic acid-base syndromes. In metabolic acidosis, PCO2 is decreased, because acids other than H2CO3 accumulate, and CO2 is blown off by the lungs in attempts to decrease body fluid acidity. In metabolic alkalosis, PCO2 is increased if the lungs compensate by hypoventilation; in mild or acute cases, PCO2 may remain normal. In respiratory alkalosis, PCO2 is decreased because increased ventilation blows off more CO2. In respiratory acidosis, PCO2 is increased because of CO2 retention due to decreased ventilation.

  • Respiratory Alkalosis

    The other major subdivision of alkalosis is respiratory alkalosis, which occurs when the respiratory mechanism blows off more CO2 than it normally would due to respiratory center stimulation from some cause. The main conditions in which this happens are the hyperventilation syndrome caused by hysteria or anxiety, high fever, and direct stimulation of the respiratory center by drugs. Overdose of aspirin can cause respiratory alkalosis in the early stages; although later, after more of the aspirin is absorbed, a metabolic acidosis develops. In hyperventilation of whatever cause, respirations are increased and deeper, blowing off more CO2. This creates an H2CO3 deficit since it is being used up to replenish CO2 by the lung carbonic anhydrase enzymes. Therefore, the denominator of the Henderson-Hasselbalch equation is decreased, the 20:1 ratio is increased, and plasma pH increased. The CO2 content will decrease, because when H2CO3 is lost due to formation of CO2 in the lungs, HCO3– is converted to H2CO3 in the kidney to compensate secondarily for or to replace the decreasing plasma carbonic acid.

  • Metabolic Alkalosis

    Alkalosis may also be divided into metabolic and respiratory types. In metabolic alkalosis, three relatively common situations should be discussed.

    Alkali administration. Alkali administration most commonly occurs when sodium bicarbonate is taken in large quantities for the treatment of peptic ulcer symptoms. If this happens, excess HCO–3 is absorbed above the amount needed to neutralize stomach hydrochloric acid (HCl). The numerator of the Henderson-Hasselbalch equation is increased, the normal 20:1 ratio is increased, and the pH therefore rises. The CO2 content also rises because of the additional HCO–3. Lactate, citrate, or acetate in sufficient quantities may also produce alkalosis, since they are metabolized to HCO–3.

    Acid-losing alkalosis. Acid-losing alkalosis most frequently results from severe or protracted vomiting, as may occur with pyloric stenosis. Gastric HCl is lost in vomiting. Gastric HCl was originally produced by conversion of H2CO3 to HCO–3 and H+, mediated by carbonic anhydrase of the gastric mucosa. The HCO–3 is kept in the bloodstream, but the H+ is secreted into the gastric lumen as HCl. When HCl is lost through vomiting, the H+ component of HCl is also continually being lost. The CO2 content becomes increased, because the HCO–3 that is released when HCl is produced remains in the bloodstream and increases when additional HCl is formed to replace that which is being lost. Therefore, the 20:1 ratio is increased and the pH is increased. Since H2CO3 is decreased, as it is being continually used to produce more HCl, the lungs tend to retain CO2 to compensate. Therefore, PCO2 may actually increase, although not enough to prevent increase of the 20:1 ratio.

    Hypokalemic alkalosis. Hypokalemic alkalosis is most commonly due to excess potassium ion (K+) loss by the kidney, as might happen with overuse of certain diuretics that cause K+ as well as sodium ion (Na+) loss. Normally, most body K+ is intracellular, whereas most Na+ and H+ is extra-cellular. When excess K+ is lost in the urine, intracellular K+ diffuses out of the cells to replace some of that being lost from plasma; Na+ and H+ move into the cells to replace the K+ that has moved out. Thus, H+ is lost from extracellular fluid and plasma. A second mechanism depends on the fact that sodium is reabsorbed from the urine into the renal distal tubule cells by an active transport mechanism. This transport mechanism involves excretion of H+ and K+ into the urine to replace the reabsorbed Na+. In this exchange (or transport) system, H+ and K+ compete with each other. Therefore, if an intracellular deficit of K+ exists (in the tubule cells), more H+ is excreted into the urine to allow the reabsorption of the same quantity of sodium without losing as much K+. The result of renal H+ loss and extracellular H+ loss is an acid-losing type of alkalosis. Therefore, more H+ is manufactured by the kidney from H 2CO3 to replace lost extracellular fluid H+: more HCO–3 ions are thereby formed, and the numerator of the Henderson-Hasselbalch equation is increased. The denominator is eventually increased if the lungs attempt to compensate by increasing CO2 retention by slower breathing. However, respiratory compensation is frequently minimal or insignificant in hypokalemic alkalosis, so that the PCO2 frequently remains normal. Also, the urine pH is decreased because of the excess H+ being excreted in the urine. This is the opposite of what one would ordinarily expect, because in acidosis, the kidney usually attempts to compensate by excreting more H+(acid urine), whereas in alkalosis it normally tries to conserve H+ and thus produces a urine of higher pH (alkaline urine).

  • Respiratory Acidosis

    The second major category of acidosis is that called respiratory acidosis. This may be due to any condition that causes pulmonary CO2 retention. These conditions include the respiratory muscle paralysis of poliomyelitis, the respiratory brain center depression sometimes seen in encephalitis or with large doses of such drugs as morphine, primary lung disease (e.g., pulmonary fibrosis or severe emphysema) that destroys oxygen exchange ability, and sometimes heart diseases (e.g., chronic congestive heart failure). The basic problem is H2CO3 excess produced by CO2 retention. Thus, the denominator of the Henderson-Hasselbalch equation is increased, the normal 20:1 ratio is decreased, and the pH is decreased. The CO2 content is sometimes normal but is usually increased, because of kidney attempts to handle the excess CO2 by forming more HCO–3 and excreting more H+ ions.

  • Metabolic Acidosis

    This type of acidosis has at least three main causes.

    Acid-gaining acidosis. Hydrogen ions not included in the CO2 system are added to the blood. The common situations are:

    1. Direct administration, such as treatment with ammonium chloride, or the late effects of salicylate poisoning. Ammonium chloride (NH4Cl) releases H+ ions and Cl– ions as the liver utilizes this compound for NH3 to synthesize urea. Aspirin is acetylsalicylic acid, which in large quantities will eventually add enough H+ ions to cause acidosis, even though in the early stages there is respiratory alkalosis (to be discussed later).

    2. Excess metabolic acid formation is found in diabetic ketoacidosis, starvation, or severe dehydration. These conditions cause utilization of body protein and fat for energy instead of carbohydrate, with production of ketone bodies and various metabolic acids.

    The results of acid-gaining acidosis are a decrease in free HCO–3, which is used up trying to buffer the excess H+. Thus, the numerator of the Henderson-Hasselbalch equation is decreased, the normal 20:1 ratio is decreased, and the pH is, therefore, decreased. The CO2 content (CO2 combining power) is also decreased because the bicarbonate which it measures has been decreased as a primary response to the addition of excess acid.

    Base-losing acidosis. Base-losing acidosis is caused by severe intestinal diarrhea, especially if prolonged or in children. Diseases such as cholera or possibly ulcerative colitis or severe dysentery might cause this. The mechanism is direct loss of HCO–3 from the lumen of the small intestine. Normally, HCO–3 is secreted into the small intestine, so that the contents of the small intestine are alkaline in contrast to the acidity of the stomach. Most of the HCO–3 is reabsorbed; however, prolonged diarrhea or similar conditions could mechanically prevent intestinal reabsorption enough to cause significant HCO–3 loss in the feces. In addition, the H+ ions that were released from H2CO3 in the formation of HCO–3 by carbonic anhydrase are still present in the bloodstream and help decrease pH. However, the primary cause is the direct loss of HCO–3; the numerator of the Henderson-Hasselbalch equation is decreased, the 20:1 ratio is decreased, and the pH is decreased. Naturally, the CO2 content is also decreased.

    Renal acidosis. Renal acidosis occurs in kidney failure that produces the clinical syndrome of uremia. As mentioned previously, the kidney has the major responsibility for excreting large excesses of H+. In uremia, H+ from metabolic acids that normally would be excreted this way is retained in the bloodstream due to loss of renal tubular function. As in acid-gaining acidosis, the excess H+ must be buffered; therefore, part of the available body fluid HCO–3 is used up. This decreases the numerator of the Henderson-Hasselbalch equation, decreases the normal 20:1 ratio, and therefore decreases pH. Again, the CO2 content is decreased.

  • Clinical Disturbances of pH

    With this background, one may proceed to the various clinical disturbances of pH. These have been clinically termed “acidosis,” when the pH is decreased toward the acid side of normal and alkalosis, when the pH is elevated toward the alkaline side of normal. Acidosis, in turn, is usually subdivided into metabolic and respiratory etiology, and the same for alkalosis.

  • Acid-Base Test Specimens

    In the early days, acid-base studies were performed on venous blood. Venous specimens are nearly as accurate as arterial blood for pH and HCO3 (or PCO2) measurements if blood is obtained anaerobically from a motionless hand or arm before the tourniquet is released. Nevertheless, arterial specimens have mostly replaced venous ones because venous blood provides less accurate data in some conditions such as decreased tissue perfusion due to shock. Even more important, one can also obtain blood oxygen measurements (PO2) with arterial samples. Arterial blood is most often drawn from the radial artery in the wrist. Arterial puncture is little more difficult than venipuncture, and there is a small but definite risk of extravasation and hematoma formation that could compress the artery. Although glass syringes have some technical advantages over plastic syringes or tubes (such as a slightly smaller chance of specimen contamination with room air than when using plastic syringes), most hospitals use only plastic. It is officially recommended that the specimen tube or syringe should be placed in ice immediately for transport to the laboratory, both to prevent artifact from blood cell metabolism and to diminish gas exchange between the syringe and room air. The blood must be rewarmed before analysis. Actually, it is not absolutely necessary to ice the specimen in most cases if the analysis takes place less than 15 minutes after the specimen is obtained. Icing the specimen in plastic tubes can elevate PO2 values a little if they are already over 80 mm Hg (10.7 kPa). One investigator found that at 100 mm Hg (13.3 kPa), the false elevation averages 8 mm Hg (1.06 kPa). Also, icing in plastic tubes increases plasma viscosity over time and interferes with resuspension of the RBC, which affects hemoglobin assay in those instruments that calculate O2 content from PO2 and total hemoglobin. If mixing before assay is not thorough, hemoglobin values will be falsely decreased somewhat. In addition, if electrolytes are assayed on the arterial specimen, potassium may be falsely increased somewhat.

    Capillary blood specimens from heelstick are often used in newborn or neonatal patients because of their small blood vessels. Warming the heel produces a semiarterial (“arterialized”) specimen. However, PO2 is not reliable and PCO2 sometimes differs from umbilical artery specimens. The majority of reports do not indicate substantial differences in pH; however, one investigator found a mean decrease in PCO2 of 1.3 mm Hg (0.17 kPa), a mean pH increase of 0.02 units, and a mean decrease of 24.2 mm Hg (3.2 kPa) in PO2 from heelstick blood compared to simultaneously drawn umbilical artery blood.

    Heparin is the preferred anticoagulant for blood gas specimens. The usual method is to wash the syringe with a heparin solution and then expel the heparin (which leaves about 0.2 ml of heparin in the dead space of the syringe and needle). If too much heparin remains or the blood sample size is too small (usually when the sample is <3 ml), there is a disproportionate amount of heparin for the amount of blood. This frequently causes a significant decrease in PCO2 (and bicarbonate) and hemoglobin values, with a much smaller (often negligible) decrease in pH. These artifactual decreases in PCO2 are especially apt to occur when the sample is obtained from indwelling catheters flushed with heparin.

  • Carbon Dioxide of pH and Carbon Dioxide

    This section describes the laboratory tests used in pH abnormalities, which, incidentally, are often called “acid-base problems” because of the importance of the bicarbonate and carbonic acid changes involved.

    Carbon dioxide combining power. Venous blood is drawn aerobically with an ordinary syringe and the serum is then equilibrated to normal alveolar levels of 40 mm Hg by the technician blowing his or her own alveolar air into the specimen through a tube arrangement. This maneuver adjusts the amount of dissolved CO2 to the normal amounts found in normal arterial blood. Bicarbonate of the serum is then converted to CO2 by acid hydrolysis in a vacuum, and the released gas is measured. The released CO2 thus consists of the dissolved CO2 of the specimen already present plus the converted HCO–3 (and thus the denominator plus the numerator of the Henderson-Hasselbalch equation). Subtraction of the known amount of dissolved CO2 and H2CO3 in normal blood from this measurement gives what is essentially a value for serum HCO–3 alone (called the “combining power,” since HCO–3 combines with H+). The inaccuracy that may be caused by these manipulations should be obvious.

    Carbon dioxide content. Total CO2 content is determined from heparinized arterial or venous blood drawn anaerobically. This may be done in a vacuum tube or a syringe that is quickly capped. (Mineral oil is not satisfactory for sealing.) The blood is centrifuged and the plasma removed. At this point, all the CO2 present is still at the same CO2 tension or partial pressure of dissolved gas that the patient possessed. Next, the plasma is analyzed for CO2 by a method that converts HCO–3 and H2CO3 to the gas form. Thus CO2 content measures the sum of HCO–3, H2CO3, and dissolved CO2. Since the amount of dissolved CO2 and H2CO3 in blood is very small, normal values for CO2 content are quite close to those of the CO2 combining power (which measures only HCO–3). Since the specimen has been drawn and processed with little or no contact with outside air, the result is obviously more accurate than that obtained from the CO2 combining power.

    Serum bicarbonate. An order for serum CO2, serum HCO–3, or venous CO2 will usually result in serum being obtained from venous blood drawn aerobically and assayed for HCO–3 without equilibration. This technique is used in most automated equipment that assays “CO2” (really, bicarbonate) in addition to performing other tests, and is also popular in many laboratories using manual methods since the patients usually have other test orders that require serum. It is somewhat less accurate than CO2 combining power. The serum is frequently exposed to air for relatively long periods of time. Only relatively large changes in CO2 or HCO–3 will be detected. Underfilling of specimen collection tubes to only one third of capacity significantly decreases bicarbonate values.

    Partial pressure of carbon dioxide (PCO2).

    PCO2 is the partial pressure of CO2 gas in plasma or serum (in mm of Hg or in Torr); this is proportional to the amount of dissolved CO2 (concentration of CO2). Since most of the denominator of the Henderson-Hasselbalch equation represents dissolved CO2, and since PCO2 is proportional to the amount (concentration) of dissolved CO2, PCO2 is therefore proportional to the denominator of the Henderson-Hasselbalch equation and may be used as a substitute for the denominator. In practice, a small amount of whole blood (plasma can be used) collected anaerobically is analyzed for PCO2 by direct measurement using a PCO2 electrode. The HCO–3 may then be calculated (from the Henderson-Hasselbalch equation), or the PCO2 value itself may be used in conjunction with pH to differentiate acid-base abnormalities. PCO2 determined by electrode is without question the method of choice for acid-base problems. The pH is usually measured at the same time on the same specimen.

    pH measurement. pH determination originally involved measuring the difference in electric charge between two electrodes placed in a solution (e.g., plasma or whole blood). Current equipment uses a single direct-reading pH electrode, which makes pH determination very simple and reliable and enables pH determination to be a routine part of blood gas measurement. On the technical side, it should be noted that at room temperature, plasma pH decreases at the rate of about 0.015 pH unit every 30 minutes. Unless measurement is done within 30 minutes of drawing, the blood should be refrigerated immediately; it can then be kept up to 4 hours.

    There is not a universally accepted single reference range for arterial or venous acid-base parameters.

  • Blood pH: The Bicarbonate-Carbonic Acid System

    The term pH comes from the French puissance hydrogen, meaning the strength or power of hydrogen. The hydrogen ion concentration of blood expressed in gram molecular weights of hydrogen per liter (moles/L) is so much less than 1 (e.g., 0.0000001) that it is easier to communicate this information in terms of logarithms; thus, 0.0000001 becomes 1 x 10–7. The symbol pH represents an even greater simplification, because pH is defined as the negative logarithm of the hydrogen ion concentration (in the preceding example, 1 x 10–7 becomes 10–7, which then becomes 7.0). In this way, a relatively simple scale is substituted for very cumbersome tiny numbers. In the pH scale, therefore, a change in 1.0 pH unit means a tenfold change in hydrogen ion concentration (a change of pH from 7.0 to 6.0 represents a change from 10–7 to 10–6 moles/L).

    The normal pH of arterial blood is 7.4, with a normal range between 7.35 and 7.45. Blood pH must be maintained within relatively narrow limits, because a pH outside the range 6.8-7.8 is incompatible with life. Therefore, the hydrogen ion content is regulated by a series of buffers. A buffer is a substance that can bind hydrogen ions to a certain extent without inducing a marked change in pH. Among substances that act as buffers are hemoglobin, plasma protein, phosphates, and the bicarbonate-carbonic acid system. Bicarbonate is by far the body’s most important buffer substance; it is present in large quantities and can be controlled by the lungs and kidneys.

    A brief review of the bicarbonate-carbonic acid system recalls that carbon dioxide (CO2) in aqueous solutions exists in potential equilibrium with carbonic acid (CO2 + H2O = H2CO3). The enzyme carbonic anhydrase catalyzes this reaction toward attainment of equilibrium; otherwise the rate of reaction would be minimal. CO2 is produced by cellular metabolism and is released into the bloodstream. There, most of it diffuses into the red blood cells (RBCs), where carbonic anhydrase catalyzes its hydration to H2CO3. Carbonic acid readily dissociates into hydrogen ions (H+) and bicarbonate ions (HCO–3). Eventually only a small amount of dissolved CO2 and a much smaller amount of undissociated H2CO3 remain in the plasma. Therefore, the great bulk of the original CO2 (amounting to 75%) is carried in the blood as bicarbonate, with only about 5% still in solution (as dissolved CO2 or undissociated H2CO3) and about 20% coupled with hemoglobin as a carbamino compound or, to a much lesser extent, coupled with other buffers, such as plasma proteins.

    This situation can best be visualized by means of the Henderson-Hasselbalch equation. This equation states that    pH = pK + log(base/acid)  where pK is the dissociation constant (ability to release hydrogen ions) of the particular acid chosen, such as H2CO3. The derivation of this equation will be disregarded to concentrate on the clinically useful parts, the relationship of pH, base, and acid. If the bicarbonate-acid system is to be interpreted by means of the Henderson-Hasselbalch equation, then base/acid = HCO3- / H2CO3 since bicarbonate is the base and carbonic acid is the acid. Actually, most of the carbonic acid represented in the equation is dissolved CO2 which is present in plasma in an amount greater than 100 times the quantity of undissociated H2CO3. Therefore the formula should really be … but it is customary to let H2CO3 stand for the entire denominator. The next step is to note from the Henderson-Hasselbalch equation that pH is proportional to base/acid. This means that in the bicarbonate-carbonic acid system, pH is proportional to …  The kidney is the main regulator of HCO–3 production (the numerator of the equation) and the lungs primarily control CO2 excretion (the denominator).

    The kidney has several means of excreting hydrogen ions. One is the conversion of monohydrogen phosphate to dihydrogen phosphate (HPO42– to H2PO4–). Another is the formation of ammonia (NH3) in renal tubule cells by deamination of certain amino acids, such as glutamine. Ammonia then diffuses into the urine, where it combines with H+ to form ammonium ion (NH4+). A third mechanism is the one of most concern now, the production of HCO–3 in the renal tubule cells. These cells contain carbonic anhydrase, which catalyzes the formation of H2CO3 from CO2. The H2CO3 dissociates, leaving HCO–3 and H+. The H+ is excreted in the urine by the phosphate or ammonium pathways or combined with some other anion. The HCO–3 goes into the bloodstream where it forms part of the pH buffer system. Not only does HCO–3 assist in buffering H+ within body fluids, but HCO–3 filtered at the glomerulus into the urine can itself combine with urinary H+ (produced by the renal tubule cells from the H2CO3 cycle and excreted into the urine).

    The lungs, on the other hand, convert H2CO3, into CO2 and water with the aid of carbonic anhydrase and blow off the CO2. In this process, a mechanism for excreting H+ exists, because the HCO–3 in the plasma can combine with free H+ to form H2CO3, which can then be eliminated from the lungs in the form of CO2, as was just described. This process can probably handle excretion of most normal and mildly abnormal H+ quantities. However, when large excesses of free H+ are present in body fluids, the kidney plays a major role, because not only is H+ excreted directly in the urine, but HCO–3 is formed, which helps buffer additional amounts of H+ in the plasma.

    Going back to the Henderson-Hasselbalch equation, which is now modified to indicate that pH is proportional to HCO–3/H2CO3, it is easy to see that variations in either the numerator or denominator will change pH. If HCO–3 is increased without a corresponding increase in H2CO3, the ratio will be increased and the pH will rise. Conversely, if something happens to increase H2CO3 or dissolved CO2, the denominator will be increased, the ratio will be decreased, the pH will fall, and so on. Clinically, an increase in normal plasma pH is called “alkalosis” and a decrease is called “acidosis.” The normal ratio of HCO–3 to H2CO3 is 20:1.