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How to Interpret and Use Arterial Blood-Gas Data Interpretation and Application for the Non-Chemist Steven J. Barker, Ph.D, M.D.

Tucson, Arizona

Introduction The interpretation and application of arterial blood-gas (ABG) data is a task that anesthesiologists must often perform under difficult circumstances. The time is 3:00 AM; we are fatigued and distracted by multiple other simultaneous tasks; we need to take action on these ABG results now. In this setting, which bears similarities to piloting an aircraft on instruments in bad weather, it is useful to have a simple algorithm or “check-list,” both to ensure consistency and obtain a correct answer within a short time. The purpose of this talk is to develop such an algorithm and apply it to specific clinical examples, wherein we shall interpret both oxygenation and acid-base status, and then prescribe appropriate treatment. Case Example At 3:00 AM, you are presented with an 80 year-old, 60 kg man who has been found unconscious on the floor at home. He is brought to the operating room for emergency laparotomy, with a presumed diagnosis of ruptured appendix. His vital signs are BP: 90/60, HR: 120, RR: 26, SpO2: 96%. A pre-operative ABG reveals: pH = 7.20, PaCO2 = 25, PaO2 = 75, HCO3- = 8. Should you give sodium bicarbonate? How much? What else do you change? Oxygenation The first step in evaluating the arterial oxygen tension (PaO2) is to calculate the alveolar oxygen tension, using the alveolar gas equation: PAO2 = FIO2(PB – PH2O) – (1/RQ)PaCO2 . [1] For a normothermic (37oC) patient, breathing room air at sea level: P AO 2 = 0.21(760 – 47) – (1.25)40 = 99.73 mmHg . The alveolar PO2 at sea level is roughly 100 mmHg. Next, calculate the RATIO of arterial to alveolar oxygen tension: PaO2/ PAO2. Do NOT bother with the alveolar-arterial difference, sometimes erroneously referred to as the “gradient.” The normal value of this difference is a function of the FIO2, whereas the normal or healthy value of the ratio is roughly 0.85 at any FIO2. Thus a “healthy” PaO2 during normocarbia at sea level is PaO2 ~ 0.85 PAO2 = 85 mmHg. Now, just for fun, let’s go to the top of Mt. Everest, elevation 29,035 ft, barometric pressure PB = 247 mmHg. Here is how some humans have actually made it to the summit without supplemental oxygen: PAO2 = FIO2(PB – PH2O) – (1/RQ)PaCO2 = 0.21(247 – 47) – (1.25)7.5 = 32.6 mmHg Yes, your PaCO2 on top of Everest is 7.5 mmHg. Talk about hyperventilation! Furthermore, your predicted PaO2, using the arterial/alveolar ratio of 85%, will be 0.85 X 32.6 = 27.7 mmHg. In laboratory simulations of Everest conditions, this is almost exactly the resulting PaO2 value. If you are wondering why the PaCO2 is so low, try repeating this calculation for normocarbia. The treatment of hypoxemia is beyond the scope of this lecture, but in general the tools available in the operating room fall into three categories. 1. Ventilator adjustment: FIO2, PEEP, vent mode (reverse I/E, pressure, HFV, etc.). 2. Drugs: Diuretics, bronchodilators, PDE-5 inhibitors (Cialis is now used for HAPE). 3. Procedures: Chest tube, bronchoscopy, etc.

Refresher Course Lectures Anesthesiology 2012 © American Society of Anesthesiologists. All rights reserved. Note: This publication contains material copyrighted by others. Individual refresher course lectures are reprinted by ASA with permission. Reprinting or using individual refresher course lectures contained herein is strictly prohibited without permission from the authors/copyright holders.

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Acid-Base Balance At a concentration of 0.00004 mEq/L, hydrogen ion is one of the least plentiful electrolytes in the extracellular fluid (ECF). Nevertheless, because of its small size and high charge state, H+ is highly reactive and must be regulated within narrow limits, or we die. The hydrogen ion concentration, denoted [H+], varies between 0.1 Eq/L (gastric) and 0.000000008 Eq/L (duodenal) within the body. We commonly use a logarithmic scale to measure it: pH = -log[H+]. The rather wide concentration range above is thus converted to a pH range between 1.0 and 8.1. An H+ concentration in the ECF that is either too high (low pH: acidemia) or too low (high pH: alkalemia) will cause a variety of potentially fatal cardiovascular disturbances. These include decreased myocardial contractility, vasomotor instability, cardiac dysrythmias, and decreased enzyme function. The body uses several defenses against pH changes, but to understand these we must first review some basics of acid-base chemistry. A cids, Bases, Buffers Simply put, and acid is a chemical compound that can give up a free H+ ion in aqueous solution; a base is a compound that can accept a free H+ ion. This equilibrium is represented as: HA ↔ H+ + A- , [2] where HA is and acid, and A- is called the “conjugate base.” A “strong acid” is one that shifts this equilibrium to the right, producing large amounts of H+. A weak acid shifts the balance to the left, producing less free H+. Similar definitions apply for strong and weak bases. If an acid is strong, then its conjugate base is by definition weak, and vice versa. Since excess hydrogen ion (acidemia) is bad for our health, we would generally prefer weak acids to stronger ones. More on this concept later. The strength of the acid HA can be quantified by the ratio of concentrations of the two sides of equation 2: K = [H+][A-]/[HA] [3] In this and the following equations, square brackets denote concentrations. K is called the “dissociation constant” and equation 3 is referred to as the “law of mass action.” A large value of K implies a strong acid and weak conjugate base. Taking the logarithms of both sides of equation 3 and rearranging, we find: log K = log [H+] + log ([A-]/[HA]) -log [H+] = -log K + log ([A-]/[HA]) pH = pK + log ([A-]/[HA]) [4] In this last step we have inserted the definitions pH = -log [H+] and pK = -log K, yielding the familiar HendersonHasselbach Equation. This equation is simply a logarithmic form of the law of mass action, or definition of the dissociation constant. The most important acid-base equilibrium in the body is that of carbon dioxide and water: H2O + CO2 ↔ H2CO3 ↔ H+ + HCO3[5] The pK of this reaction is 6.1, making carbonic acid a rather weak acid. Henderson-Hasselbach (Eq. 4) for this reaction becomes: pH = pK + log {[HCO3-]/[H2CO3]} pH = 6.1 + log {[HCO3-]/(0.03 PaCO2)} ` [6] We don’t routinely measure blood carbonic acid concentration, so in the last step we have substituted the empirical formula: [H2CO3] = 0.03 PaCO2. Equation 6 states a simple and important fact about acid-base physiology: the pH is determined by the RATIO of the bicarbonate concentration to the PaCO2. We shall see that this is a key feature of the body’s defenses against pH changes. Inserting normal blood values into Eq. 6: pH = 6.1 + log {24/(0.03 X 40)} = 7.4 . It is always reassuring when a new formula gives the correct answer in a known situation. Body acids, buffers, and defenses Acids are formed in two ways in the human body. Respiratory acid is formed by the combination of carbon dioxide (product of aerobic metabolism) and water, as shown in Eq. 5. All other acids are called “metabolic acid.” The latter include lactic acid produced by anaerobic metabolism, as well as sulphuric, Refresher Course Lectures Anesthesiology 2012 © American Society of Anesthesiologists. All rights reserved. Note: This publication contains material copyrighted by others. Individual refresher course lectures are reprinted by ASA with permission. Reprinting or using individual refresher course lectures contained herein is strictly prohibited without permission from the authors/copyright holders.

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hydrochloric, and phosphoric acids. We have already noted that carbonic or respiratory acid is a weak acid. In fact, in the normal ECF equilibrium (Eq. 5) it takes 800,000 molecules of CO2 to produce 800 molecules of H2CO3, which in turn produce ONE free H+ ion. Insert Acids.jpg This brings us to the body’s first line of defense against acid “invasion”: a system of “buffers.” A buffer is something that effectively exchanges a strong acid for a weaker one, as in the following. H+ + Cl- + NaHCO3 ↔ H2CO3 + NaCl

[7]

As this reaction moves to the right, the sodium bicarbonate “trades” the very strong hydrochloric acid (HCl) for the very weak carbonic acid (H2CO3). A buffer solution thus consists of a weak acid (H2CO3) and a salt of the conjugate base (NaHCO3), and it uses this combination to trade strong acid for weak acid, thereby reducing the amount of free H+ released. Bicarbonate, the example we have just considered, is the most important buffer in the body. Other buffers include proteins, phosphate, and ammonia. Proteins are significant for two reasons: (1) they are very plentiful in the intracellular milieu (75% of total body buffering power), and (2) they can buffer both metabolic and respiratory acid, as shown below for the protein hemoglobin (Hb). HCl + KHb ↔ HHb + KCl [8a] H2CO3 + KHb ↔ HHb + KHCO3 [8b] There are two other front-line defenses against pH changes: respiratory and renal control. The lungs, driven by medullary chemoreceptors, regulate the PaCO2 level. The kidneys, by actively excreting H+ into the urine and reclaiming Na+ from the urine, effectively ‘pump’ bicarbonate from the urine back into the blood. The kidneys excrete a normal daily metabolic acid load of 50 mM at a urine pH of 6.0, and healthy kidneys can handle as much as ten times that amount by reducing the urine pH to 4.5. Looking again at equation 6, the lungs control the denominator and the kidneys the numerator of the ratio that determines pH. Evaluation of acid-base balance Given the above background on acids, bases, buffers, and the body’s pH controls, we are now ready to evaluate and treat the acid-base balance. The first step in this process is a plot of the plasma pH versus bicarbonate, shown in Figure 1.

Refresher Course Lectures Anesthesiology 2012 © American Society of Anesthesiologists. All rights reserved. Note: This publication contains material copyrighted by others. Individual refresher course lectures are reprinted by ASA with permission. Reprinting or using individual refresher course lectures contained herein is strictly prohibited without permission from the authors/copyright holders.

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Figure 1. The pH – Bicarbonate Diagram. Choosing a location on this graph fixes the values of pH and HCO3-, and Eq. 6 (Henderson-Hasselbach) then determines the corresponding value of PaCO2. The curves labeled 20, 30, 40, etc. are lines of constant PaCO2, called “isobars.” They are solutions of the Henderson-Hasselbach Equation. The normal acid base status of pH = 7.4, HCO3- = 24, PaCO2 = 40 corresponds to the center of the graph. Note the straight line passing through the center with a slight negative slope, labeled “10sl.” This “buffer line” is the path followed by a purely respiratory disturbance – either hypercarbia or hypocarbia with no metabolic disturbance at all. The slope of the buffer line represents the body’s chemical buffering power, and its units are called “Slykes,” named after the chemist Van Slyke. The normal slope is actually nearer to 12, but a value of 10 is easier for the calculations that follow and is close enough for our purposes. The two stippled regions along the buffer line, labeled “acute,” are the areas in which an acute, uncompensated, respiratory disturbance will lie. Take for example an acute respiratory acidosis – the stippled area to the left of the center of the graph. Referring again to Eq. 6, we see that the body’s compensation for this acute hypercarbia will be to increase the HCO3- level (the kidneys’ job) to bring the ratio back towards normal, thus correcting the pH. This “metabolic compensation” may take several days to complete, and it will move us from the “acute” stippled area upward along the appropriate isobar into the “chronic” stippled area in the upper-left quadrant. Thus a quick exam of the pH-bicarbonate diagram not only tells us that our patient has a respiratory acidosis, but also evaluates the metabolic compensation. The same reasoning applies to both acute and compensated respiratory alkalosis, shown in the lower right quadrant of the diagram. A pure metabolic acidosis, without compensation, would follow the PaCO2 isobar down into the lower left quadrant of Figure 1. (Ignore, for now, the arrow in the lower left quadrant. It will be used later.) Metabolic acidosis Refresher Course Lectures Anesthesiology 2012 © American Society of Anesthesiologists. All rights reserved. Note: This publication contains material copyrighted by others. Individual refresher course lectures are reprinted by ASA with permission. Reprinting or using individual refresher course lectures contained herein is strictly prohibited without permission from the authors/copyright holders.

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manifests as a decrease in HCO3-, so the body’s compensatory response is to decrease the denominator of the ratio in Eq. 6, that is, to lower the PaCO2. This compensatory hyperventilation takes us into the stippled area in the lower left quadrant of Figure 1. Similarly, the respiratory compensation for metabolic alkalosis is hypoventilation, as indicated by the stippled area of the upper right quadrant. Since hypoventilation will eventually cause hypoxemia, this compensation is less effective, as indicated by the slopes of the two stippled areas. Instead of referring the stippled areas of Figure 1, we can also use empirical “rules of thumb” to predict compensation. Remember that the body’s response to pH imbalance is to restore the ratio of HCO3-/PaCO2 toward its normal value of 24/40 = 0.6, because this ratio controls the pH (Eq. 6). The following rules are based on clinical data, and are only approximate predictors. 1. Metabolic compensation for a primary respiratory disturbance. a. Respiratory acidosis: For every 10 mmHg rise in PaCO2 above 40: i. Acute: HCO3- increases by 1 mEq/L. ii. Chronic (compensated): HCO3- increases by 4 mEq/L. b. Respiratory alkalosis: For every 10 mmHg fall in PaCO2 below 40: i. Acute: HCO3- decreases by 2 mEq/L. ii. Chronic (compensated): HCO3- decreases by 3 mEq/L. 2. Respiratory compensation for a primary metabolic disturbance. a. Metabolic acidosis with maximum compensation: PaCO2 = 1.5[HCO3-] + 8. b. Metabolic alkalosis with maximum compensation: PaCO2 = 0.7[HCO3-] + 20. One additional tool is required to quantify the metabolic disturbance: “base excess” or BE. The definition of BE is as follows. “Titrate the pH to 7.40 by varying only PaCO2. BE is the difference between the resulting HCO3- at the end of this titration and the normal value of 24 mEq/L.” It is easier to understand this definition by using an example on the pH-bicarbonate diagram, Figure 1. We are given the following arterial blood-gas results: pH = 7.1, PaCO2 = 32, HCO3- = 10, as indicated by the tail of the arrow in the lower left quadrant. To find BE, we titrate back to a pH of 7.4 by varying only PaCO2. This titration follows a line parallel to our “buffer line,” shown by the arrow shaft. Since the slope of this line is -10 Slyke, and the pH changes by 0.3 units (from 7.1 to 7.4), the bicarbonate will decrease by 3 mEq/L (0.3 times 10), reaching a new value of 7 (10 minus 3) at the head of the arrow. The BE is thus 7 – 24, or -14. A negative base excess is also called a “base deficit,” so our blood-gas shows a base deficit of 14. You can now calculate BE in your head if you know the pH and the bicarbonate. Should we treat this base deficit, and if so, how? This is our next topic. There are many possible causes of metabolic acidosis, but in the operating room it is most often the result of inadequate oxygen delivery to various organs and tissues. It is the direct result of anaerobic metabolism, and therefore manifests as a lactic acidosis. The “anion gap,” defined as [Na+] + [K+] – [Cl-] – [HCO3-], is above the normal range of 8-12. Like most physiological disorders, the ideal prescription is to “treat the underlying cause.” If the cause of the lactic acidosis is tissue hypoperfusion resulting from hypovolemic shock, correcting the patient’s volume status may be all that is needed to restore perfusion and rapidly correct the acidemia. On the other hand, a very low pH (7.1 would fall in that category) will depress cardiac function significantly, and the heart may not be able to generate sufficient cardiac output to restore organ perfusion while the pH remains so low. This is a judgment call that we must make on a case-by-case basis. If we decide that the pH must be at least partially corrected before the heart can do its job, here is how we treat it. Treatment of acid-base disorders The first step in treating a metabolic acidosis is to determine the “bicarbonate deficit.” That is, how much bicarbonate would be required to restore the extracellular fluid to its normal value of 24 Refresher Course Lectures Anesthesiology 2012 © American Society of Anesthesiologists. All rights reserved. Note: This publication contains material copyrighted by others. Individual refresher course lectures are reprinted by ASA with permission. Reprinting or using individual refresher course lectures contained herein is strictly prohibited without permission from the authors/copyright holders.

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mEq/L? This quantity is simply the base deficit times the weight in kg, times the ECF fraction. The ECF comprises the following fraction of body weight: ECF fraction =

0.5 (premature neonate) 0.4 (term neonate) 0.3 (2 year-old) 0.25 (older child) 0.2 (adult)

Thus, if the patient above with a BE of -14 is a 70 kg adult, the bicarbonate deficit will be 14 X 70 X 0.2, or 196 mEq. Treatment of metabolic acidosis with sodium bicarbonate (NaHCO3) is not a risk-free therapy. First, “bicarb” does not make metabolic acid disappear; it effectively converts it to respiratory acid (see Eq. 7). The additional CO 2 created by this exchange must be removed by increased ventilation. Furthermore, full strength sodium bicarbonate is a hyperosmolar solution (six times plasma osmolarity) and can thereby cause brain hemorrhage, particularly in children. Overtreatment can cause alkalemia, which is just as dangerous as acidemia. Insert Biocarbinate Therapy Dangers.jpg Therefore, the common recommendation is to give about half the amount of total bicarbonate deficit, increase ventilation to compensate for the additional CO2, and reevaluate in roughly 30 minutes. There are two approved alternatives to sodium bicarbonate: tris-hydroxy amino-methane (THAM) and “carbi-carb,” which is an equimolar mixture of sodium bicarbonate and sodium carbonate (Na2CO3). Both can buffer metabolic acid with little or no increase in CO2, but neither has been shown to have any outcomes advantage over conventional treatment. Evaluation/Treatment A lgorithm We conclude by encapsulating the above approach to acid-base evaluation into a simple four-step algorithm, which we shall apply to our original case example. Four Step Evaluation of Acid-Base Status: Move to chart or graphic 1. Evaluate the pH: if pH > 7.45, patient has alkalemia; if pH < 7.35, patient has acidemia. 2. Evaluate the PaCO2: if PaCO2 > 45, patient has respiratory acidosis; if PaCO2 < 35, patient has respiratory alkalosis. 3. Find the base excess BE, using the 10 Slyke approximation: if BE > +2, patient has metabolic alkalosis; if BE < -2, patient has metabolic acidosis. 4. Evaluate compensation: determine predicted maximum compensation by non-primary variable, using empirical rules or referring to stippled areas of Figure 1. Finally, let us use this algorithm to evaluate and treat our original patient, a 60 kg man with a blood gas analysis showing 7.20/25/75/8 (pH/ PaCO2/PaO2/HCO3-). 1. The patient is acidemic (pH = 7.20). 2. The patient has respiratory alkalosis (PaCO2 = 25). 3. Calculate BE (see Figure 1): titration to pH = 7.4 requires pH to change by 7.4 – 7.2 = 0.2; HCO3- therefore decreases by 0.2 X 10 = 2; new HCO3- = 8 – 2 = 6; BE = 6 – 24 = -18. 4. Evaluate compensation: Maximum predicted respiratory compensation for this metabolic acidosis is PaCO2 = 1.5[HCO3-] + 8 = (1.5 X 8) +8 = 20 mmHg. The patient’s actual PaCO2 is 25 mmHg. Given his physical status, this patient is nearly at maximum compensation. Treatment: The amount of bicarbonate required to fully correct the patient’s BE of -18 is: NaHCO3 = 60 kg X 0.2 X 18 mEq/L = 216 mEq. Refresher Course Lectures Anesthesiology 2012 © American Society of Anesthesiologists. All rights reserved. Note: This publication contains material copyrighted by others. Individual refresher course lectures are reprinted by ASA with permission. Reprinting or using individual refresher course lectures contained herein is strictly prohibited without permission from the authors/copyright holders.

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A reasonable therapy is to give half this amount, or two ampoules (55 mEq per ampoule), increase mechanical ventilation, and obtain another ABG measurement in 15-30 minutes. Conclusions Interpretation of arterial blood-gas data is straightforward, if we follow a prescribed algorithm such as the one presented in this lecture. The simplest interpretation of PaO2 comes from calculating the ratio of arterial to alveolar oxygen tension (using the alveolar gas equation) and comparing the result with the normal value of 0.85. To evaluate the acid-base status, simply follow the four step algorithm, which separates respiratory for metabolic disturbance and evaluates compensation. The pH-bicarbonate diagram (Figure 1) is an excellent tool for keeping all of these factors in perspective, and understanding the results of treatment. Finally, if the patient has a severe metabolic acidosis that may require treatment, calculate the bicarbonate deficit as shown above. I do not recommend an arbitrary BE threshold for treatment with sodium bicarbonate; that decision should take into account all aspects of the patient’s clinical status. However, once the decision to treat has been made, it is usually wise to give about half of the bicarbonate deficit as the initial dose, increase or monitor ventilation, and then reevaluate acid-base status a short time later. Of course there are exceptions to every rule – that is why you spent all those years in training. But these guidelines should serve you well as a starting point, and they are simple enough that even I can follow them at 3:00 AM. REFERENCES 1. 2. 3. 4. 5. 6. 7.

Guyton A, Hall J, Textbook of Medical Physiology, 11th Edition, Saunders, New York, 2005. Nunn J, Applied Respiratory Physiology, 6th Edition; Butterworth-Heinemann, London, 2005. Severinghaus JW, Astrup PB: History of blood gas analysis II. pH and acid-base balance measurements. J Clin Mon. 1985; 1(4):259-77. Stewart, PA: Modern quantitative acid-base chemistry. Can J Physiol Pharmacol 1983; 61:1444-1461. Adrogué HJ, Madias NE: Management of life-threatening acid-base disorders. N Engl J Med, 1998; 338:26–34, 107–111. Stewart PA: Stewart’s Textbook of Acid-Base, AcidBase.org, 2009. Tremper KK, Barker SJ” Blood Gas Analysis, in “Principles of Critical Care,” Hall JB, Schmidt GA, Wood LD (ed.); McGraw-Hill, New York, 1992.

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Refresher Course Lectures Anesthesiology 2012 © American Society of Anesthesiologists. All rights reserved. Note: This publication contains material copyrighted by others. Individual refresher course lectures are reprinted by ASA with permission. Reprinting or using individual refresher course lectures contained herein is strictly prohibited without permission from the authors/copyright holders.