Perioperative Acid-Base Balance

Key Points

▪
The presence of a significant acid-base abnormality often signals a sinister underlying problem.
▪
All acid-base abnormalities result from alterations in the dissociation of water.
▪
Only three factors independently affect acid-base balance—the arterial partial pressure of carbon dioxide (PaCO2), the strong ion difference (SID), and the total concentration of weak acids (A _TOT ).
▪
Respiratory acidosis and alkalosis are caused by hypercarbia and hypocarbia, respectively.
▪
Metabolic acidosis is caused by decreased SID or increased A _TOT . Decreased SID results from accumulation of metabolic anions (shock, ketoacidosis, and renal failure), hyperchloremia, and free water excess. Increased A _TOT results from hyperphosphatemia.
▪
Metabolic alkalosis is caused by increased SID or decreased A _TOT . SID increases due to sodium gain, chloride loss, or free water deficit. A _TOT decreases in hypoalbuminemia and hypophosphatemia. This is particularly common in critical illness.
▪
Most acid-base disorders are treated by reversal of the cause.

Acknowledgment

The editors, publisher, and Dr. Patrick Neligan would like to thank Dr. Clifford S. Deutschman for his contribution to previous editions of this chapter.

Introduction—Why Is Acid-Base Balance Important?

Blood gas and pH analyses have been the most robust laboratory and bedside tools for identifying and monitoring critical illness since the advent of modern medicine. There is good reason to believe that understanding acid-base chemistry will continue to be important for physicians for decades to come.

The human body is composed principally of water, partitioned into intracellular and extracellular compartments. The electrolyte composition of each of these spaces is tightly controlled to maintain homeostasis. Alterations in the relative concentrations of electrolytes and of carbon dioxide (CO ₂ ) impact the tendency of water to auto-ionize into its component parts: hydrogen and hydroxyl ions. Alterations in the water, gas, and electrolyte composition of the fluid compartments manifests as changes in the chemistry profile of body water and acid-base balance.

The hydrogen ion concentration is conventionally measured as pH (literally the “power of hydrogen”), the negative logarithmic value of its concentration. Deviations in extracellular pH away from the resting value of 7.4 have long been associated with acute and critical illness. Such deviations are known as “acid-base abnormalities.” All acid-base abnormalities result from changes in the local concentration of strong ions, weak acids, and CO ₂ .

This chapter first looks at the basic science behind acid-base abnormalities. Subsequently, we explore the detection and treatment of acid-base conundrums, with specific reference to perioperative medicine and critical care.

What Are Acids and Bases?

The concept of acids and bases is relatively new in medicine, and arose with the development of laboratory science in the early part of the 20th century. However, as early as 1831, O’Shaughnessy identified loss of “carbonate of soda” from the blood as a fundamental disturbance in patients dying of cholera. This led directly to the development of crystalloid replacement therapy for hypovolemic shock. In 1909, L.J. Henderson coined the term “acid-base” balance. He was able to define this process in terms of carbonic acid equilibrium. Henderson’s work was later refined by Hasselbalch in 1916. Their method described acid-base balance in terms of the hydration equation for CO ₂ :

{CO}_{2} + H_{2} O \to H_{2} {CO}_{3} \to H^{+} + {HCO}_{3}^{-}

${CO}_{2} + H_{2} O \to H_{2} {CO}_{3} \to H^{+} + {HCO}_{3}^{-}$

{pH = pK}_{a} + log [{HCO}_{3}^{-}] / [H_{2} {CO}_{3}]

${pH = pK}_{a} + log [{HCO}_{3}^{-}] / [H_{2} {CO}_{3}]$

[{Total CO}_{2}] = [{HCO}_{3}^{-}] + [{Dissolved CO}_{2}] + [{Carbamino CO}_{2}] + [H_{2} {CO}_{3}]

$[{Total CO}_{2}] = [{HCO}_{3}^{-}] + [{Dissolved CO}_{2}] + [{Carbamino CO}_{2}] + [H_{2} {CO}_{3}]$

\approx P_{{CO}_{2}} \times 0 {. 03 mmol CO}_{2} /L/mm Hg

$\approx P_{{CO}_{2}} \times 0 {. 03 mmol CO}_{2} /L/mm Hg$

So, substituting into the foregoing equation:

pH = 6 {. 1 + log [HCO}_{3}^{-}] / {PCO}_{2} \times 0.03

$pH = 6 {. 1 + log [HCO}_{3}^{-}] / {PCO}_{2} \times 0.03$

This is the Henderson-Hasselbalch equation.

The introduction of this concept into clinical practice became possible with the development of volumetric CO ₂ analysis by Van Slyke and others in 1919. This led to 60 years of research and interest in CO ₂ and its derivative, bicarbonate, as the principle agents that impact acid-base chemistry, despite the knowledge in the 1920s of the importance of chloride in acid-base equilibria.

An understanding of acid-base chemistry in the human body requires familiarity with physical chemistry. Because the body contains a large quantity of water, the physical properties of water have enormous implication for maintenance of homeostasis. Water is a simple triatomic molecule. Its chemical formula is H ₂ O and its structural formula is H-O-H. The charge distribution of each covalent bond is unequal, and the molecule has a polar conformation and an H-O-H bond angle of 105 degrees. Water molecules attract and form hydrogen bonds with one another. Consequently, water has a high surface tension, a low vapor pressure, a high specific heat capacity, a high heat of vaporization, and a high boiling point.

Water molecules are in continual motion. Occasionally a collision generates sufficient energy to transfer a proton from one water molecule to another. Thus, water is always slightly dissociated into a negatively charged hydroxyl (OH ⁻ ) ion and a positively charged hydronium (H ₃ O ⁺ ) ion. Conventionally, this self-ionization of water is written as follows:

[H_{2} O] \leftrightarrow H^{+} {+ OH}^{-}

$[H_{2} O] \leftrightarrow H^{+} {+ OH}^{-}$

The symbol H ⁺ is convenient because, although protons dissociating from water have many aliases (such as H ₃ O ⁺ and H ₉ O ₄ ⁺ ), most physicians and chemists refer to them as hydrogen ions.

The self-ionization of water is miniscule. In pure water at 25°C, the [H ⁺ ] and [OH ⁻ ] are 1.0 × 10 ⁻⁷ mmol/L. The tendency for water to dissociate into its component parts is represented by the expression

K_{eq} H_{2} O = {[H}^{+} {] [OH}^{-}]

$K_{eq} H_{2} O = {[H}^{+} {] [OH}^{-}]$

The molarity of water is extremely high—55.5 M (“there is a lot of water in water”). As the concentration of water and the K _eq are constants, the ion-product dissociation constant (pKa) for water can be expressed as follows:

K_{eq} H_{2} {O = K}_{eq} (55.5) = K_{w}^{'} = {[H}^{+} {] [OH}^{-}]

$K_{eq} H_{2} {O = K}_{eq} (55.5) = K_{w}^{'} = {[H}^{+} {] [OH}^{-}]$

The implication is that the product of the concentrations of hydroxyl and hydrogen is constant, so when there is an increase in the concentration of hydrogen ions, there is a concomitant decrease in the concentration of hydroxyl ions, and vice versa.

Pure water is considered neutral because the relative concentrations of hydrogen and hydroxyl are equal at 1.0 × 10 ⁻⁷ mmol/L. A solution is considered acidic if the concentration of hydrogen ions exceeds that of hydroxyl ions ([H ⁺ ] > 1.0 × 10 ⁻⁷ mmol/L, [OH ⁻ ] < 1.0 × 10 ⁻⁷ mmol/L). A solution is considered alkaline if the hydroxyl ion concentration exceeds the hydrogen ion concentration.

In 1903, Svante Arrhenius (1859-1927) 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. Because of its high dielectric constant, water is a highly ionizing solution. Therefore, substances with polar bonds will dissociate into their component parts (dissolve) in water. Brønsted and Lowry (BL) independently advanced this concept using slightly different terminology: an acid is a proton donor, a base a proton acceptor. Water itself is amphoteric and can act as either an acid or as a base. Thus, when hydrochloric acid (HCl) is dissolved in water, chloride, the acid, donates a proton to water, the base. Similarly, when potassium hydroxide (KOH) is dissolved in water, potassium, the base, receives a hydrogen ion from water, the acid or proton donor.

The degree of dissociation of substances in water determines whether they are strong acids or strong bases. Lactic acid, which has an ion 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 (Na ⁺ ), potassium (K ⁺ ), and chloride (Cl ⁻ ), which do not easily bind other molecules, are considered strong ions— they exist free in solution. As each Na ⁺ delivers a hydroxyl moiety into extracellular fluid (ECF), it is functionally a base, as are all cations. As each Cl ⁻ delivers a hydrogen moiety into ECF, it is functionally an acid, as are all anions. The hydrogen and hydroxyl ions delivered in this way bind to one another, forming water molecules, and relatively few free hydrogen or hydroxyl ions remain free in solution.

{HCL + H}_{2} O \to H_{3} O^{+} + {Cl}^{-}

${HCL + H}_{2} O \to H_{3} O^{+} + {Cl}^{-}$

In this reaction, hydrogen chloride acts as a BL acid and water as a BL base.

{NaOH + H}_{2} O \to H_{2} {O + OH}^{-} + {Na}^{+}

${NaOH + H}_{2} O \to H_{2} {O + OH}^{-} + {Na}^{+}$

In this reaction, water acts as a BL acid and sodium as a BL base.

{OH}^{-} + {Na}^{+} + H_{3} O^{+} + {Cl}^{-} = {Na}^{+} + {Cl}^{-} + H_{2} O

${OH}^{-} + {Na}^{+} + H_{3} O^{+} + {Cl}^{-} = {Na}^{+} + {Cl}^{-} + H_{2} O$

Because of electrical neutrality, the hydrogen and hydroxyl ions delivered by chloride and sodium become water.

In summary, all acid-base reactions in the human body relate to the presence of charged particles within an aqueous environment. In the following section, how the components of ECF influence the acid-base status of the body, as measured by clinicians, is discussed. This is followed by an explanation of different acid-base abnormalities, and the tools used to identify them. These approaches are neither distinct from one another nor scientifically incompatible.

What Determines the Acidity or Alkalinity of a Solution?

Because all acid-base reactions are based on the principles of physical chemistry, three simple rules must be followed:

1.
Electrical neutrality : in aqueous solutions, in any compartment, the sum of all the positive charged ions must equal the sum of all the negative charged ions.
2.
Dissociation equilibria : the dissociation equilibria of all incompletely dissociated substances, as derived from the law of mass action, must be satisfied at all times.
3.
Mass conservation : the amount of a substance in a given compartment remains constant unless it is added, removed, generated, or destroyed. Therefore, the total concentration of an incompletely dissociated substance is the sum of concentrations of its dissociated and undissociated forms.

To determine the acid-base status of a fluid, all substances to which these rules could be applied must be accounted for. Fundamentally this involves counting up all the charges (and hence hydrogen and hydroxyl ions) delivered by strong cations (bases) and strong anions (acids), weak acid buffers, and CO ₂ . A discussion of key groups follows.

Strong Ions

The first group of ions, the strong ions, dissociate completely. The most abundant strong ions in the extracellular space are Na ⁺ and Cl ⁻ . Other important strong ions include K ⁺ , sulfate (SO ₄ ²⁻ ), magnesium (Mg ²⁺ ), and calcium (Ca ²⁺ ). These are occasionally referred to as “mineral” acids or bases because they cannot be metabolized. Organic acids are generated from metabolism and accumulate when there is metabolic dysfunction—such as kidney failure, splanchnic hypoperfusion, or hormonal deficiency.

In a solution containing strong ions, created, for example, using specified concentrations of sodium hydroxide (NaOH) and HCL, the hydrogen ion concentration can be calculated by solving for electric neutrality:

({Na}^{+} {- Cl}^{-}) + {(H}^{+} - {OH}^{-}) = 0

$({Na}^{+} {- Cl}^{-}) + {(H}^{+} - {OH}^{-}) = 0$

This creates two separate simultaneous equations :

H^{+} = \sqrt{K_{W}^{'} + \frac{{([{Na}^{+}] - [{Cl}^{-}])}^{2}}{4}} - \frac{({[Na]}^{+} - [{Cl}^{-}])}{2}

$H^{+} = \sqrt{K_{W}^{'} + \frac{{([{Na}^{+}] - [{Cl}^{-}])}^{2}}{4}} - \frac{({[Na]}^{+} - [{Cl}^{-}])}{2}$

and

{OH}^{-} = \sqrt{K_{W}^{'} + \frac{{([{Na}^{+}] - [{Cl}^{-}])}^{2}}{4}} + \frac{([{Na}^{+}] - [{Cl}^{-}])}{2}

${OH}^{-} = \sqrt{K_{W}^{'} + \frac{{([{Na}^{+}] - [{Cl}^{-}])}^{2}}{4}} + \frac{([{Na}^{+}] - [{Cl}^{-}])}{2}$

These equations tell us that hydrogen and hydroxyl concentrations are determined by the Kw′ (water pKa), and the difference in charge between Na ⁺ and Cl ⁻ . Because the former is constant, in this system ([Na ⁺ ] − [Cl ⁻ ]) must determine [H ⁺ ] and [OH ⁻ ]. Given that the concentration of both Na+ and Cl ⁻ are known, the net positive charge minus net negative charge can be quantified. It is the strong ion difference (SID). Logically, in any solution, the sum total of the charges imparted by strong cations minus the charges from strong anions will represent the SID. SID independently influences hydrogen ion concentration ( Fig. 48.1 ). In human ECF, SID is always positive.

\begin{matrix} SID = ([{Na}^{+}] + [K^{+}] + [{Ca}^{2 +}] + [{Mg}^{2 +}]) \\ - ([{Cl}^{-}] + [A^{-}]) = 40 - 44 mEq / L \end{matrix}

$\begin{matrix} SID = ([{Na}^{+}] + [K^{+}] + [{Ca}^{2 +}] + [{Mg}^{2 +}]) \\ - ([{Cl}^{-}] + [A^{-}]) = 40 - 44 mEq / L \end{matrix}$

Fig. 48.1, Effect of Changes in the Strong Ion Difference ( SID ) on Hydrogen and Hydroxyl Ion Concentration.

Hydroxyl ions almost always exceed hydrogen ions quantitatively in solution. The relationship between SID and [H ⁺ ] is nonlinear in these conditions. Any change in SID will change both [H ⁺ ] and [OH ⁻ ] concentrations. Because of the Kw′, this relationship is inverse: as [H ⁺ ] increases, [OH ⁻ ] decreases (see Fig. 48.1 ). SID is an independent variable, and [H ⁺ ] and [OH ⁻ ] are dependent, meaning that the addition of hydrogen ions alone (without strong corresponding anions) cannot influence the pH of the solution.

Weak Acid “Buffer” Solutions

The degree of water dissociation, and thus the hydrogen ion concentration, is also influenced by charges derived from weak acids. These acids are partially dissociated compounds whose degree of dissociation is determined by the prevailing temperature and pH. The predominant molecules in this group are albumin and phosphate. Stewart used the term A _TOT to represent the total concentration of weak anions or acids that influence acid-base balance.

The acid HA will only partly dissociate, represented by the equilibrium

[H^{+}] \times [A^{-}] = K_{A} \times [HA]

$[H^{+}] \times [A^{-}] = K_{A} \times [HA]$

K _A is the weak acid pKa. If we assume that HA and A ⁻ play no further part in this reaction (the law of mass conservation), the amount of A ⁻ present in the solution must equal the amount initially present, so:

[H A] + [A^{-}] = [A_{T O T}]

$[H A] + [A^{-}] = [A_{T O T}]$

where [A _TOT ] is the total weak acid concentration.

In order to calculate the effect of weak acid dissociation on [H ⁺ ], we must take into account water dissociation and electrical neutrality:

[H^{+}] + [OH^{-}] = K_{w}^{'} (water dissociation) .

$[H^{+}] + [OH^{-}] = K_{w}^{'} (water dissociation) .$

[SID] + {[H}^{+}] - {[A}^{-}] - [{OH}^{-}] = 0 (electrical neutrality) .

$[SID] + {[H}^{+}] - {[A}^{-}] - [{OH}^{-}] = 0 (electrical neutrality) .$

These four simultaneous equations determine the [H ⁺ ] in this solution containing strong ions and weak acids. SID and A _TOT are independent variables. Kw′ and KA are constants. Consequently, the other variables [HA], [H ⁺ ], [OH ⁻ ], and [A ⁻ ] must adjust to satisfy the foregoing equations. They are dependent variables.

Carbon Dioxide

Along with strong ions and weak bases, ECF contains CO ₂ . The concentration of CO ₂ in ECF is determined by tissue production and alveolar ventilation. CO ₂ , when in solution, exists in four forms: CO ₂ (denoted CO ₂ [d]), carbonic acid (H ₂ CO ₃ ), bicarbonate ions (HCO ₃ ⁻ ), and carbonate ions CO ₃ ²⁻ .

The concentration of CO ₂ (d) is determined by the solubility coefficient of CO ₂ (SCO ₂ ), which depends on body temperature, PCO ₂ , and other factors. Several equilibrium equations can be derived from the hydration of CO ₂ :

[{CO}_{2} (d)] = {[Sco}_{2}] \times {Pco}_{2}

$[{CO}_{2} (d)] = {[Sco}_{2}] \times {Pco}_{2}$

The tendency for CO ₂ to hydrate to H ₂ CO ₃ , and subsequently dissociate to H ⁺ and HCO ₃ ⁻ , is reflected by the equation:

[{CO}_{2} (d)] \times {[OH}^{-} {] = K}_{1} \times [{HCO}_{3}^{-}]

$[{CO}_{2} (d)] \times {[OH}^{-} {] = K}_{1} \times [{HCO}_{3}^{-}]$

These equations can be combined together and with water equilibrium to the following:

[H^{+}] \times {[HCO}_{3}^{-}] = K_{c} \times {pCO}_{2}

$[H^{+}] \times {[HCO}_{3}^{-}] = K_{c} \times {pCO}_{2}$

HCO ₃ ⁻ also dissociates to release hydrogen ions and carbonate, the equilibrium reaction which is represented by this equation:

[H^{+}] \times {[CO}_{3}^{2 -}] = K_{3} {[HCO}_{3}^{-}]

$[H^{+}] \times {[CO}_{3}^{2 -}] = K_{3} {[HCO}_{3}^{-}]$

Factors Independently Influencing Water Dissociation

Now that we have looked at the different factors that may influence the concentration of hydrogen ions in a solution—strong ions, weak acids, and CO ₂ —we can combine the derived equations to a solve for [H ⁺ ]:

1.
Water dissociation equilibrium: $[H^{+}] \times [{OH}^{-}] = K_{W}^{'}$ $[H^{+}] \times [{OH}^{-}] = K_{W}^{'}$
2.
Weak acid dissociation equilibrium: $[H^{+}] \times {[A}^{-}] = K_{A} \times [HA]$ $[H^{+}] \times {[A}^{-}] = K_{A} \times [HA]$
3.
Conservation of mass for weak acids: $[HA] {+ [A}^{-} {] = [A}_{TOT}]$ $[HA] {+ [A}^{-} {] = [A}_{TOT}]$
4.
HCO ₃ ⁻ formation equilibrium: $[H^{+}] \times {[HCO}_{3}^{-}] = K_{C} \times {pCO}_{2}$ $[H^{+}] \times {[HCO}_{3}^{-}] = K_{C} \times {pCO}_{2}$
5.
Carbonate ion formation equilibrium: $[H^{+}] \times [{CO}_{3}^{2 -}] = K_{3} \times [{HCO}_{3}^{-}]$ $[H^{+}] \times [{CO}_{3}^{2 -}] = K_{3} \times [{HCO}_{3}^{-}]$
6.
Electrical neutrality: $\begin{matrix} [SID] + [H^{+}] - [{HCO}_{3}^{-}] - [A^{-}] - \\ [{CO}_{3}^{2 -}] - [{OH}^{-}] = 0 \end{matrix}$ $\begin{matrix} [SID] + [H^{+}] - [{HCO}_{3}^{-}] - [A^{-}] - \\ [{CO}_{3}^{2 -}] - [{OH}^{-}] = 0 \end{matrix}$

There are six independent simultaneous equations, and just six unknown, dependent variables determined by them: [HA], [A ⁻ ], [HCO ₃ ⁻ ], [CO ₃ ²⁻ ], [OH ⁻ ], and [H ⁺ ]. There are three known independent variables: [SID], [A _TOT ], and PCO ₂

Solving the equations for [H ⁺ ]:

\begin{matrix} [SID] + [H^{+}] - K_{C} \times P_{C} / [H^{+}] - K_{A} \times [A_{TOT}] / \\ (K_{A} + [H^{+}]) - K_{3} \times K_{C} P_{C} / {[H^{+}]}^{2} - K_{W}^{'} / [H^{+}] = 0 \end{matrix}

$\begin{matrix} [SID] + [H^{+}] - K_{C} \times P_{C} / [H^{+}] - K_{A} \times [A_{TOT}] / \\ (K_{A} + [H^{+}]) - K_{3} \times K_{C} P_{C} / {[H^{+}]}^{2} - K_{W}^{'} / [H^{+}] = 0 \end{matrix}$

In other words, [H ⁺ ] is a function of SID, A _TOT , PCO ₂ , and a number of constants. All other variables, most notably [H ⁺ ], [OH ⁻ ], and [HCO ₃ ⁻ ] are dependent and cannot independently influence acid-base balance.

Acid-Base Abnormalities

The value of the physical chemistry (“Stewart”) approach is that it allows us to use a simple model for understanding acid-base disturbances, as all abnormalities can be explained in terms of SID, A _TOT , or PCO ₂ . Traditionally, acid-base disturbances have been classified as resulting from either alterations in arterial carbon dioxide (PaCO ₂ ) tension (respiratory acidosis or alkalosis) or alterations in blood chemistry (metabolic acidosis or alkalosis). This remains a useful classification, although it must be remembered that respiratory or metabolic abnormalities rarely occur independent of one another.

Respiratory Acid-Base Abnormalities

Respiratory Alkalosis

The normal PaCO ₂ is 40 mm Hg (5.3 kPa). Respiratory alkalosis occurs when there is an acute decrease in PaCO ₂ as a result of hyperventilation. Acute respiratory alkalosis is characterized by a pH > 7.45, a low PaCO ₂ , and a low HCO ₃ ⁻ . A simple rule of thumb for this reaction is as follows:

Acute respiratory alkalosis:

{ΔHCO}_{3}^{-} = 0.2 {ΔPaCO}_{2}

${ΔHCO}_{3}^{-} = 0.2 {ΔPaCO}_{2}$

The hyperventilating patient will present with symptoms and signs of vasoconstriction: lightheadedness, visual disturbances, dizziness, and hypocalcemia from increased binding of calcium to albumin. The hypocalcemia is caused by an increase in the available negative charge on albumin in alkaline states. Acute hypocalcemia is associated with paresthesia and tetany. In anesthesia practice, patients may hyperventilate preoperatively as a consequence of anxiety, and postoperatively secondary to pain, agitation, or a full bladder. More commonly, hyperventilation results from poor mechanical ventilation strategy and may result in significant systemic and, in particular, cerebral vasoconstriction. Therapeutic hyperventilation is no longer used to treat intracranial hypertension because of significant concerns regarding cerebral hypoperfusion and ischemia.

Respiratory Acidosis

Respiratory acidosis occurs when there is an acute rise in PaCO ₂ usually associated with respiratory failure. This may result from problems with

▪
Central ventilator control—such as toxicity from anesthetic agents, benzodiazepines, or opioids, stroke, spinal cord injury;
▪
Peripheral ventilator control—such as myasthenia gravis, poliomyelitis, polymyopathy, or neuromuscular blockade;
▪
Ventilation-perfusion mismatch—associated with pneumothorax, pleural effusion, atelectasis, pneumonia, or pulmonary edema.

Clinically, patients have signs of CO ₂ retention: cyanosis, vasodilatation, and narcosis.

Respiratory acidosis causes a rapid increase in [H ⁺ ]. Compensation for hypercarbia is slow, requiring increased urinary excretion of Cl ⁻ , and pH falls rapidly. A concomitant increase in the plasma HCO ₃ ⁻ occurs and reflects a higher total CO ₂ load. In the mid-1960s, Brackett, Cohen, and Schwartz elegantly described the relative changes in HCO ₃ ⁻ in response to acute and chronic elevations in PaCO ₂ , providing us with extremely useful “rules of thumb”:

In acute hypercarbia, the bicarbonate concentration of plasma rises slowly

▪
An increase in PaCO ₂ by 10 mm Hg (1.3 kPa) results in an increase in HCO ₃ ⁻ by 1 mmol/L (1 mEq/L)

A patient returning to the intensive care unit from the operating room who has been under ventilation for several hours may be hypercapnic, for example, with a PaCO ₂ of 80 mm Hg (10.5 kPa). Using this rule, the expected HCO ₃ will be 28 mmol/L ( Table 48.1 ).

Table 48.1

Changes in Bicarbonate HCO ₃ ⁻ and PaCO ₂ in Acute and Chronic Hypercarbia

Acute hypercarbia
Measured Bicarbonate	Expected PaCO ₂ in Acute Respiratory Failure
HCO ₃ ⁻ mEq/L or mmol/L	PaCO ₂ mm Hg	PaCO ₂ kPa
24	40	5.3
25	50	6.6
26	60	8.0
27	70	9.3
28	80	10.5
29	90	11.8
30	100	13

Chronic hypercarbia
Measured Bicarbonate	Expected PaCO ₂ in Chronic Respiratory Failure
HCO ₃ ⁻ mEq/L or mmol/L	PaCO ₂ mm Hg	PaCO ₂ kPa
24	40	5.3
27	50	6.6
30	60	8.0
33	70	9.3
36	80	10.5
39	90	11.8
42	100	13

In chronic respiratory failure, the total CO ₂ load in the body increases substantially, reflected by relatively high levels of plasma HCO3 ⁻ . There is a concomitant fall in plasma chloride, reflecting compensation for elevated levels of carbonic acid.

In chronic respiratory acidosis:

▪
$Δ H^{+} (nEq/L) = 0.8 (Δ {PCO}_{2})$ $Δ H^{+} (nEq/L) = 0.8 (Δ {PCO}_{2})$
▪
An increase in PaCO ₂ by 10 mm Hg (1.3 kPa) will increase plasma [HCO ₃ ⁻ ] by 3 mmol/L (3 mEq/L)

A patient returning to the intensive care unit from the operating room, ventilated to a PaCO ₂ of 40 mm Hg (5.3 kPa) but with high total CO ₂ , secondary to chronic respiratory failure (for example, from COPD), may fail to liberate from mechanical ventilation, due to acute metabolic alkalosis. Using the “rule” the patient’s preoperative total CO ₂ of 33 mEq/L (mmol/L) will indicate a baseline PaCO ₂ of 70 mm Hg (9.3 kPa) for that patient (see Table 48.1 ).

For anesthesiologists, who routinely administer mechanical ventilation to patients with various levels of CO ₂ retention, it is important to consider goals for both arterial and end tidal carbon dioxide (etCO ₂ ) in the operating room. This should also include consideration of the impact of drugs, such as opioids and benzodiazepines, on respiratory function and on the effect of intravenous fluids on overall acid-base dynamics. For example, both isotonic saline (0.9%) solution and Ringer’s Lactate (or Hartmann’s) solution will cause progressive increase in extracellular chloride in a dose-dependent manner. This does not occur with truly balanced solutions such as Plasmalyte-148 or Normosol-R. In theory, this may impact postoperative CO ₂ -HCO ₃ ⁻ homeostasis and respiratory function.

In critical illness, in particular in patients with acute respiratory distress syndrome (ARDS), there is general consensus among intensivists that aggressive mechanical ventilation to normalize pH and PaCO ₂ is more harmful than “permissive” hypercapnia. Data support that such an approach—a “lung protective” strategy— is well tolerated, and may indeed be beneficial.

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