e8 Abstract Fluid resuscitation is one of the oldest and most com mon interventions for critically ill children This chapter discusses the data informing selection of crystalloid and colloid fluids of[.]
e8 Abstract: Fluid resuscitation is one of the oldest and most common interventions for critically ill children This chapter discusses the data informing selection of crystalloid and colloid fluids of variable composition for different disease states and the impact of fluid volume overload post-resuscitation on morbidity and mortality In addition, this chapter presents the etiology, clinical manifestations, and management of the common electrolyte disturbances involving sodium, potassium, magnesium, calcium, and phosphorus Key words: fluid resuscitation, crystalloids, colloids, hypernatremia, hyponatremia, SIADH, hyperkalemia, hypomagnesemia, hypocalcemia, hypophosphatemia 72 Acid-Base Disorders MICHELLE C STARR AND SHINA MENON • Acid-base disorders are common in critical illness.1 Interpretation of these disorders requires an understanding of not only acids and bases but also knowledge of gas exchange and ventilation, dynamics of water and electrolyte movement at the cellular level, and renal-related mechanisms of hydrogen ion, electrolyte, and water excretion.2–4 In addition, to fully integrate the physiology with clinically relevant pathophysiology, one must also understand the possible dysfunctions, both structural and metabolic, that contribute to acid or base disturbances in the extracellular or intracellular spaces, such as sepsis, acute kidney injury, severe metabolic dysregulation (such as diabetic ketoacidosis [DKA] or inborn errors of metabolism), or toxic effects of certain compounds and medications.5 Most acid-base disturbances not directly cause harm and are self-limited after management of the underlying clinical condition Despite this, they are commonly considered a disease process, with significant time spent treating abnormal numbers in the intensive care unit (ICU) A majority of physicians experience difficulty in interpreting various acid-base disturbances.6,7 Given this, it is not a surprise to find easily available smartphone applications, podcasts, blood-gas calculators and interpreters, mnemonics, videos, tutorials, and primers for the interpretation of blood gases and acid-base status, with a variable level of quality and 882 • • Abnormalities in acid-base balance should be anticipated in all critically ill children due to their underlying disease or intensive care unit therapy and should be monitored closely In severe circulatory failure, such as in cardiac arrest or septic shock, differences between arterial and venous pH, partial pressure of carbon dioxide, and lactate commonly occur In the presence of tissue hypoperfusion, hypercapnia and acidemia are best detected in central venous blood Many available tools, equations, and corrections are available to assist in the interpretation of acid-base disorder Tools such as Henderson-Hasselbalch, bicarbonate, and standard base excess are the first line for assessing metabolic and respiratory acidbase disturbances For more complex situations, apparent strong ion difference, urinary anion gap, and corrected anion gap may provide additional information Acidosis is the most common derangement of acid-base homeostasis seen in the intensive care unit Metabolic acidosis • • • • PEARLS can represent a more severe disease presentation and portends a worse outcome The most common causes of metabolic acidosis in children with critical illness are sepsis-induced lactic acidosis, diabetic ketoacidosis, and renal insufficiency The most common causes of respiratory acidosis are severe status asthmaticus and acute respiratory distress syndrome Initial management of acid-base derangements in children with critical illness begins with stabilization of the patient, irrespective of the underlying cause of the acid-base disorder Depending on the severity of the derangement, the underlying cause may need to be aggressively identified and reversed Identification of the cause of the underlying disorder often comes after stabilization and treatment However, identification is essential in preventing worsening of the abnormality, treatment of the underlying disturbance, and may provide information to help determine prognosis accuracy.8–11 Contrary to their original purpose, some of these tools may provide easy interpretation and solutions to acid-base disorders without explaining the physiology.3,10,12 As disorders of acid-base homeostasis are common in critical illness, a thorough and practical understanding of the physiology and pathophysiology behind these disorders is crucial in the ICU Overview of Acid-Base Physiology Defining Acids and Bases Understanding of acid-base physiology was transformed at the end of the 19th century.13 Prior to this, Svante Arrhenius had outlined the theories of dissociation and ionization, and Ostwald described the ability of water to dissociate into hydrogen (H1) and hydroxide or hydroxyl (OH2) ions.14 Acid was defined by Arrhenius as any substance that increased the concentration of hydrogen ions when dissolved in water, and a base decreased the concentration of H1 ions.10–14 As water can itself be a source of hydrogen ions, it is not mandatory for a given substance to have hydrogen atoms.15 Later, work by Faraday and others showed that acid-base status in biological systems was partly determined by the concentration of electrolytes, most notably sodium and CHAPTER 72 Acid-Base Disorders chloride, which had been previously described as a base-forming cation and acid-forming anion, respectively.16 By the 1920s, Brønsted and Lowry, working independently, described an acid as a proton donor substance that contains hydrogen atoms and is able to release them.15,16 Thus, an acid (HA) may dissociate to donate a proton to the solution, forming the conjugate base (or anion) in a reversible manner according to the law of mass action: Ka [HA ] ↔ [H+] [ A– ] (Eq 72.1) In Eq 72.1, Ka (or Keq) is the equilibrium or dissociation constant, specific for every substance and influenced by the solution in which the reaction is taking place Based on this principle, a base is defined as a substance with the ability for binding free hydrogen ions.16,17 Examples of such bases include bicarbonate, proteins, and phosphates The Brønsted-Lowry definition also led to the concept of conjugate acid-base pairs.16–18 In the historical context of body fluid composition and acidbase balance, these two definitions work synergistically to explain the physiology of complex biological solutions The Arrhenius definitions recognize the principal role of sodium, chloride, and other strong ions, while Brønsted-Lowry’s definition highlights the role of the conjugate acid-base pairs, such as the bicarbonate/ carbonic acid system This definition gradually became the most popular among acid-base physiologists and the Arrhenius definitions became outdated, as many thought that this approach to acid-base balance did not include the central and direct role of hydrogen ions.15–17 As water supplies an inexhaustible reservoir of hydrogen (and hydroxyl) ions, acid-base balance must account for the properties of water.18–20 The hydrogen ion, H1, immediately protonates another water molecule to form H3O1, also known as the hydronium ion (or hydroxonium ion): H2 O H2 O ↔ H3O OH (Eq 72.2), with electrical charge and temperature as the main determinants for dissociation.5 According to the electroneutrality principle, the plasma concentration and activity of hydrogen ions depend on the plasma level of other ions, and variations in the ionic composition lead to a change in plasma pH.18,21,23 Normally, the extracellular concentration of hydrogen ions is extremely low, in the nanoequivalent (nEq) or nanomole (nmol) range The precision with which H1 is regulated highlights its critical impact on cellular functions The usual arterial blood [H1] is about 40 nEq/L (or nmol/L), or million times less than the serum sodium concentration.24 This value is usually expressed as pH units, which is the negative log10 of the hydrogen ion concentration Although the use of pH instead of concentration of H1 has been repeatedly challenged, the concept has survived, as it provides a measurement of much more than H1 concentration.25,26 A normal pH of 7.4 is equivalent to an H1 concentration in the blood of 40 nEq/L, at 37°C The relationship between pH and serum [H1] is nonlinear, but it is close to linear over the normal pH range of 7.35 to 7.45 (corresponding to 45–35 nEq/L of [H1]).24,26 The logarithmic scale of pH can be confusing and can obscure the magnitude of deviations from normal For example, a change in pH from 7.4 to 7.2, which may appear to be a decrease of only 0.2, is a 60% increase in [H1].21 The term acidosis is used to describe the underlying process that leads to increase in [H1], whether or not there is a change in blood pH Alkalosis, in contrast, is the process that tends to produce a decrease in [H1], with or without changes in blood pH Acidemia and alkalemia are the corresponding terms for situations with an associated change in blood pH.5,21 Acids, Bases, and Buffers Role of Water 883 (Eq 72.2) Therefore, acid-base abnormalities can also be regarded as an alteration of water dissociation For example, the presence of electrolytes, carbon dioxide (CO2), and other weak acids in plasma produces powerful electrochemical forces that influence the disassociation of water.21,22 Human plasma is an aqueous solution that must comply with certain physiologic principles, including electrical neutrality, mass action (which explains the dissociation equilibrium of partially dissociated substances), and the constancy of the ionic product for water (which explains that ionization of plasma water varies according to the plasma ionic composition).19,21,22 Principles of Electroneutrality In aqueous solutions, electrical neutrality must always be maintained—the sum of all positively charged ions (cations) must be equal to the sum of all negatively charged ions (anions) To maintain electroneutrality, the kidneys produce urine with varying concentration of ions depending on body requirements.18,21,22 Variations in the concentration of plasma ions (either anions or cations, organic acids or electrolytes) may predictably lead to a change in the plasma concentration of hydrogen ions by driving changes and adjustments in water ionization.20,21 In physiologic conditions, only a small percentage of water molecules disassociate into components Hydrogen ion concentration and activity have a strong influence on the function of almost all enzymatic systems of the body.13,27,28 Tight regulation of the H1 is mandatory for the body because of the high reactivity of H1 ions Given this, H1 ions undergo variable ionization with changes in pH.13,27,28 Cellular processes, enzymatic action, and transmembrane transport processes are all highly pH sensitive The body has a relatively large capacity to tolerate dramatic shifts in pH For example, from a pH of 7.4 (40 nEq/L or nmol/L) to a severe acidosis of pH 6.7 (equivalent to 175 nEq/L or nmol/L), the [H1] increases over fourfold, a change that could not be tolerated without intact buffering systems.13 Each acid or base has a characteristic tendency to release or to accept hydrogen ions in an aqueous solution according to the reversible reaction depicted in Eq 72.3.1 Ka [H+][ A– ] [HA ] (Eq 72.3) The larger the value of the equilibrium constant (Ka), the greater the tendency of the acid to dissociate and the stronger the acid.16,23 The Henderson-Hasselbalch equation is the logarithmic transformation of the equilibrium constant equation, in which pH is the 2log of [H1] and pKa is the 2log of Ka: pH pK a log [ A– ] ( Acid ) [HA ] (Eq 72.4) or pH pK a log [B ] [HB +] ( Base ) (Eq 72.5) 884 S E C T I O N V I I Pediatric Critical Care: Renal The pKa estimates the relative strength of an acid or base Strong acids have a low pKa (high Ka), and weak acids have high pKa (low Ka).23 An ideal buffer is a weak acid in equilibrium with its corresponding weak base, particularly when the pKa is close to physiologic pH.16,23,27 Buffers are a key defense mechanism against changes in systemic pH There are two main physiologic buffering systems: the bicarbonate/carbonic acid (HCO32/H2CO3) system, which acts in both the extracellular space and inside the erythrocytes; and the nonbicarbonate buffers, which include hemoglobin and oxyhemoglobin, phosphates, and plasma proteins These are naturally occurring weak acids and bases, which act by converting strong acids or bases to weak acids or bases, while minimizing the change in pH.27,29 Buffers are located in both the extracellular and intracellular fluids and in the bone The ability of each buffer system to protect against changes in pH is proportional to its pKa and its concentration The bicarbonate/carbonic acid system is physiologically the most important buffering system in the blood It is an open-buffer system given its interconversion with CO2, allowing effective buffering as Pco2 can be regulated by changes in alveolar ventilation: CO2 H2 O ↔ H2 CO3 ↔ H HCO3 (Eq 72.6) Any increase in [H1] (or drop in pH) will shift reaction to the left, increasing ventilation to eliminate the rise in CO2 This respiratory response begins within minutes, but it may not reach a steady state for 12 to 24 hours This respiratory response is particularly useful to cellular tissues, as it has a rapid effect on intracellular and extracellular pH CO2 crosses cell membranes easily; thus, changes in Pco2 affect intracellular pH rapidly and in a predictable direction The bicarbonate/carbonic acid buffering system has immense capacity to buffer, as the lungs can eliminate a vast amount of CO2 per day Additionally, the kidneys can regenerate bicarbonate, although their compensatory response is slower than that of the lungs The second buffering system involves cations such as ammonium and anions such as nonvolatile metabolic acids These acids are produced by the normal daily catabolic load or during incomplete catabolism of carbohydrates and fat (lactic acidosis, diabetic ketoacidosis, or, in cases with defective processing or excretion of certain metabolites, as occurs in some inborn errors of metabolism) The free hydrogen ions are neutralized by buffers; however, as these acids are not in equilibrium in the normal plasma, they must be metabolized, mainly in the liver, and then excreted by the kidneys.27,29 In addition, bone represents an important site of buffering of acid and base loads While it is difficult to measure the exact contribution of bone buffering, it has been estimated that as much as 40% of the buffering of an acute acid load takes place in bone.30 Tools for Interpreting Acid-Base Disorders Bicarbonate is a circulating anion in the plasma that contributes substantially to the negative charge and therefore plays a role in determining plasma pH Unlike other plasma ions, bicarbonate concentration is partially facilitated by the ventilatory activity of the lung in addition to the kidney’s handling of bicarbonate CO2 formed during cell metabolism diffuses across the plasma membrane into blood within the tissue capillary network and then is transported in a number of ways Most of it (90%–95%) diffuses into the red blood cells, where it is hydrated to bicarbonate (HCO32) by carbonic anhydrase II, generating hydrogen ions that bind to oxyhemoglobin.23 As a result, oxygen is released from oxyhemoglobin and leaves the erythrocyte, diffusing into tissues Approximately 5% of the CO2 remains as a gas in the aqueous phase of blood, measured as the partial pressure of CO2 (Pco2) An even smaller proportion is bound to plasma proteins The Henderson-Hasselbalch equation quantifies the relationship between CO2 and HCO32 Using this relationship, various tools were subsequently developed to quantify the metabolic component of acid-base disturbances (eBox 72.1) In an attempt to identify acid-base changes that are independent of CO2, Hastings and Singer developed the conceptual framework of a buffer base (BB), which is a summation of weak acids in plasma (HCO32 and all weak acid buffers).15,18,31,32 Other groups proposed a multitude of other Pco2-independent indicators of acid-base disturbances These approaches can detect metabolic disturbances, mostly the simple ones; however, if Pco2 varies, the interpretation may be more confusing or complex.15,18 One of these indicators is the standard bicarbonate, which is measured plasma HCO32 at a Paco2 of 40 mm Hg 32 Base excess or deficit (BE), developed by Siggard-Andersen, is another indicator as a measure of the deviation of BB from its normal value.32 Henderson-Hasselbalch Equation The Henderson-Hasselbalch equation expresses the relationship of the bicarbonate/carbonic acid buffering system to pH15,23,25,33: ( pH pK a log HCO3 H2 CO3 ) (Eq 72.7) Further refinement and understanding that most H2CO3 exists as dissolved CO2 is represented in the modified Henderson-Hasselbalch equation23,24,33: ( pH pK a log HCO3 0.03 Pco ) (Eq 72.8) where 0.03 is the solubility coefficient for CO2 in blood, and Pco2 represents the partial pressure of CO2 in blood The primary utility of the Henderson-Hasselbalch equation is that it allows quantification of the severity of the respiratory component in acid-base derangements When pH is altered as the result of changes in the increases or decreases of Pco2, the change is respiratory, as the primary alteration is associated with either hypoventilation or hyperventilation.24 When pH is modified by any other cause (e.g., changes in serum cations or anions or in nonvolatile acids), the alteration is considered metabolic in origin.15,21,24,33 Therefore, acid-base disorders can be classified using the Henderson-Hasselbalch equation to determine the primary type of acid being increased or decreased For example, in respiratory acidosis, the increase in the Pco2 quantifies the derangement even when there are mixed disorders The relationship between Pco2 and HCO32 can provide a clinical guide for discovering the metabolic origin of a derangement in simple acidosis cases, mainly those associated with the presence of organic acids However, this can be complicated by the fact that Pco2, H2CO3, and HCO32 are all interlinked (Eq 72.7 and 72.8), so that when bicarbonate increases, Pco2 increases, and vice versa.15,18,21,23,25,33 It is also important to note that the Henderson-Hasselbalch equation does not provide information about acids other than CO2/ H2CO3, which means that only the HCO32/CO2 buffering system is assessed and that the influence of other buffering systems and mineral anions and cations is overlooked.15,18,33 884.e1 • eBox 72.1 Acid-Base Parameters • pH: Log10 H1 Measured directly by electrode Normal 7.35–7.45 • Pco2: Partial pressure of gaseous CO2 Measured directly by electrode Usually expressed in mm Hg Normal value: 35–45 mm Hg (sea level, arterial blood) • HCO32: Bicarbonate concentration A calculated parameter, derived from pH and Pco2 values using a nomogram or the Henderson-Hasselbalch equation, or equal to the difference between serum total CO2 (CO2TOT) and the dissolved CO2 (Pco2 0.03) Normal value: 22–28 mEq/L • Base excess (BE or BE/D): Also known as base excess/deficit Defined as the amount of strong base (negative base excess, or “base deficit”) or strong acid (positive base excess) in mmol/L that would be needed to restore a pH of 7.4 to a liter of whole blood equilibrated at Pco2 40 mm Hg It is calculated using a nomogram • Standard base excess (SBE): An improvement of base excess to allow equilibration across the entire extracellular fluid space with the goal of increasing accuracy at variable Pco2 values The new equation assumes hemoglobin g/dL • CO2TOT or CO2: Total CO2 or CO2 concentration A serum chemistry measured value About 95% exists as HCO32 and 4%–5% as dissolved gaseous CO2 Remaining forms are negligible Because the difference between CO2TOT and HCO32 is about mEq/L at physiologic pH, for clinical purposes they are taken almost as equivalents • Anion gap (AG): A calculated value that, taking advantage of the electroneutrality principle (cations anions), indicates the presence of unmeasured anions (mostly organic acids) AG (Na1 K1) (CI2 HCO32) Normal values are 16 mEq/L (if K1 is included) or 12 mEq/L (without K1, with variations of 2–4 mEq/L This may be influenced by the way the values of some of the parameters are measured; thus, it is always better to consult each institution’s own expected normal AG • Corrected anion gap (AGcorr): AG is corrected for the patient’s albumin concentration as follows: AGCORRECTED AGOBSERVED 2.5 [normal albumin (g/dL)] [observed albumin (g/dL)], considering normal albumin from 3.2–4.5 g/dL If there is a concern for elevated lactate, lactate correction should also be considered, as follows: AGCORRLACT Albumin-corrected AG [lactate2 (mmol/L)] • Serum osmol gap: The difference between measured and calculated serum osmolality Serum osmolality can be calculated as follows: SOSM (2 Na1) (Glu/18) (BUN/2.8) The “gap” between the measured and calculated values should be within 6; an elevation is suggestive of an unmeasured solute • Urinary anion gap (uGap): Urinary equivalent of serum AG uGap estimates unmeasured urinary ions It has also been designated by some as the urine strong ion difference (uSID), a functionally correct name, as the most important unmeasured cation (Ur1) is ammonium (NH41) In physiologically normal conditions, uSID must equal the plasma SIDAPP, or 38–42 mEq/L ... increasing ventilation to eliminate the rise in CO2 This respiratory response begins within minutes, but it may not reach a steady state for 12 to 24 hours This respiratory response is particularly useful... carbonic acid system This definition gradually became the most popular among acid-base physiologists and the Arrhenius definitions became outdated, as many thought that this approach to acid-base... Understanding of acid-base physiology was transformed at the end of the 19th century.13 Prior to this, Svante Arrhenius had outlined the theories of dissociation and ionization, and Ostwald described