Chapter 8. Acid-Base Balance

The reader understands the basic concepts of the regulation of the acid-base status of the body.

• Defines acids, bases, and buffers.
• Lists the buffer systems available in the human body.
• Describes the interrelationships of the pH, the
  of the blood, and the plasma bicarbonate concentration, and states the Henderson-Hasselbalch equation.
• States the normal ranges of arterial pH,
  , and bicarbonate concentration, and defines alkalosis and acidosis.
• Lists the potential causes of respiratory acidosis and alkalosis and metabolic acidosis and alkalosis.
• Discusses the respiratory and renal mechanisms that help compensate for acidosis and alkalosis.
• Evaluates blood gas data to determine a subject’s acid-base status.
• Classifies and explains the causes of tissue hypoxia.

The maintenance of a relatively constant internal environment is one of the major physiologic functions of the organ systems of the body. Body temperature, fluid volume and osmolarity, and electrolytes—including acids and bases—are normally carefully regulated. A thorough knowledge of the mechanisms that control these variables is essential to clinical practice.

The respiratory system is intimately involved in the maintenance of the balance of acids and bases in the body. This chapter will introduce the major concepts of acid-base balance, particularly with respect to the respiratory system; a more detailed study of this important subject is strongly encouraged.

Although there are several ways to define acids and bases, the most useful physiologically is to define an acid as a substance that can donate a hydrogen ion (a proton) to another substance and a base as a substance that can accept a hydrogen ion from another substance. A strong acid is a substance that is completely or almost completely dissociated into a hydrogen ion and its corresponding or conjugate base in dilute aqueous solution; a weak acid is only slightly ionized in aqueous solution. In general, a strong acid has a weak conjugate base and a weak acid has a strong conjugate base. The strength of an acid or a base should not be confused with its concentration.

A buffer is a mixture of substances in aqueous solution (usually a combination of a weak acid and its conjugate base) that can resist changes in hydrogen ion concentration when strong acids or bases are added. That is, the changes in hydrogen ion concentration that occur when a strong acid or base is added to a buffer system are much smaller than those that would occur if the same amount of acid or base were added to pure water or another nonbuffer solution.

The Quantification of Acidity

The acidity of a solution is determined by the activity of the hydrogen ions in the solution. The hydrogen ion activity, which is denoted by the symbol αH+, is closely related to the concentration of hydrogen ions ([H+]) in a solution. In extremely dilute solutions, the hydrogen ion activity is equal to the hydrogen ion concentration; in highly concentrated solutions, the activity is less than the concentration. The hydrogen ion concentration of the blood is low enough that the hydrogen ion activity may be considered to be equal to the hydrogen ion concentration.

The hydrogen ion activity of pure water is about 1.0 × 10−7 mol/L. By convention, solutions with hydrogen ion activities above 10−7 mol/L are considered to be acid; those with hydrogen ion activities below 10−7 are considered to be alkaline. The range of hydrogen ion concentrations or activities in the body is normally from about 10−1 for gastric acid to about 10−8 for the most alkaline pancreatic secretion. This wide range of hydrogen ion activities led to the use of the somewhat more convenient pH scale. The pH of a solution is the negative logarithm of its hydrogen ion activity. With the exception of the highly concentrated gastric acid, in most instances in the body the hydrogen ion activity is about equal to the hydrogen ion concentration. Therefore,

pH = −log (αH +)

or pH = −log [H+]

Thus, the pH of gastric acid is on the order of 1; the pH of the alkaline pancreatic secretion may be as high as 8.

The pH of arterial blood is normally close to 7.40, with a normal range considered to be about 7.35 to 7.45. An arterial pH less than 7.35 is considered acidemia; an arterial pH greater than 7.45 is considered alkalemia. The underlying condition characterized by hydrogen ion retention or by loss of bicarbonate or other bases is referred to as acidosis; the underlying condition characterized by hydrogen ion loss or retention of base is referred to as alkalosis. Under pathologic conditions the extremes of arterial blood pH have been noted to range as high as 7.8 and as low as 6.9. These correspond to hydrogen ion concentrations as follows [hydrogen ion concentrations are expressed as nanomoles (10−9 mol/L) for convenience]:

Note that the pH scale is “inverted” by the negative sign and is also logarithmic. Therefore, an increase in pH from 7.40 to 7.70 represents a decrease in hydrogen ion concentration. The increase of only 0.3 pH units indicates that hydrogen ion concentration is cut in half because the logarithm of 2 in base 10 is 0.3.

The Importance of Body pH Regulation

Hydrogen ions are the most reactive cations in body fluids, and they interact with negatively charged regions of other molecules, such as those of body proteins. Interactions of hydrogen ions with negatively charged functional groups of proteins can lead to marked changes in protein structural conformations with resulting alterations in the behavior of the proteins. An example of this was already seen in Chapter 7, where hemoglobin was noted to combine with less oxygen at a lower pH (the Bohr effect). Alterations in the structural conformations and charges of protein enzymes can affect their activities, with resulting alterations in the functions of body tissues. The absorption and efficacy of drugs administered by the physician may also be affected by the pH. Extreme changes in the hydrogen ion concentration of the body can result in loss of organ system function and structural integrity; under acute conditions, arterial pHs above approximately 7.80 or below 6.9 are not compatible with life.

Sources of Acids in the Body

Under normal circumstances, cellular metabolism is the main source of acids in the body. These acids are the waste products of substances ingested as foodstuffs. The greatest source of hydrogen ions is the carbon dioxide produced as one of the end products of the oxidation of glucose and fatty acids during aerobic metabolism. The hydration of carbon dioxide results in the formation of a hydrogen ion and a bicarbonate ion, as discussed in Chapter 7. This is reversed in the pulmonary capillaries, and CO2 then diffuses through the alveolar-capillary barrier into the alveoli, from which it is removed by alveolar ventilation. Carbonic acid is therefore said to be a volatile acid because it can be converted into a gas and then removed from an open system like the body. Very great amounts of carbon dioxide can be removed from the lungs by alveolar ventilation: Under normal circumstances, about 15,000 to 25,000 mmol of carbon dioxide is removed via the lungs daily.

A much smaller quantity of fixed or nonvolatile acids is also normally produced during the course of the metabolism of foodstuffs. The fixed acids produced by the body include sulfuric acid, which originates from the oxidation of sulfur-containing amino acids such as cysteine; phosphoric acid from the oxidation of phospholipids and phosphoproteins; hydrochloric acid, which is produced during the conversion of ingested ammonium chloride to urea and by other reactions; and lactic acid from the anaerobic metabolism of glucose. Lactic acid is sometimes converted to carbon dioxide, and so it is not always a fixed acid. Other fixed acids may be ingested accidentally or formed in abnormally large quantities by disease processes, such as the acetoacetic and butyric acid formed during diabetic ketoacidosis. About 70 mEq of fixed acids is normally produced and removed from the body each day (about 1 mEq/kg/body weight/day); the range is 50 to 100 mEq. A vegetarian diet may produce significantly less fixed acid and may even result in no net production of fixed acids. The removal of fixed acids is accomplished mainly by the kidneys, as will be discussed later in this chapter. Some may also be removed via the gastrointestinal tract. Fixed acids normally represent only about 0.2% of the total body acid production.

The body contains a variety of substances that can act as buffers in the physiologic pH range. These include bicarbonate, phosphate, and proteins in the blood, the interstitial fluid, and inside cells. One way to express the ability of a substance to act as a buffer is its buffer value. The buffer value of a solution is the amount of hydrogen ions in milliequivalents per liter that can be added to or removed from the solution with a resultant change of 1 pH unit. Another way is to determine the substance’s titration curve.

The pK of the acid is another important consideration. An acid, HA, can dissociate into a hydrogen ion, H+, and its base, A−:

According to the law of mass action, the relationship between the undissociated acid and the proton and the base at equilibrium can be expressed as the following ratio:

That is, the product of the concentrations of hydrogen ion and base divided by the concentration of the acid is equal to a constant K, the dissociation constant. This can be rearranged to

After taking the logarithm of both sides,

After multiplying both sides by −1,

This is the general form of the Henderson-Hasselbalch equation.

The buffer value or buffering capacity of a buffer pair is greatest at or near the pK of the weak acid. Note that when the concentrations of HA and A− are equal, the pH of a solution is equal to its pK.

As already stated, the human body contains a number of buffers and buffer pairs. The isohydric principle states that all the buffer pairs in a homogeneous solution are in equilibrium with the same hydrogen ion concentration. For this reason, all the buffer pairs in the plasma behave similarly, with the relative concentrations of their undissociated acids and their bases determined by their respective pKs.

An implication of the isohydric principle is that the detailed analysis of a single buffer pair, like the bicarbonate buffer system, can reveal a great deal about the chemistry of all the plasma buffers.

Buffers of the Blood

The main buffers of the blood are bicarbonate, phosphate, and proteins.

Bicarbonate

The bicarbonate buffer system consists of the buffer pair of the weak acid, carbonic acid, and its conjugate base, bicarbonate. As already stated, in the body:

The ability of the bicarbonate system to function as a buffer of fixed acids in the body is largely due to the ability of the lungs to remove carbon dioxide from the body. In a closed system bicarbonate would not be nearly as effective.

At a temperature of 37°C about 0.03 mmol of carbon dioxide per mm Hg of

where [

where pK′ is the pK of the

The pK′ of this system at physiologic pHs and at 37°C is 6.1. Therefore, at an arterial pH of 7.40 and an arterial

Therefore, the arterial plasma bicarbonate concentration is about 24 mmol/L (the normal range is 23–28 mmol/L) because the logarithm of 20 is equal to 1.3.

Note that the term total CO2 refers to the dissolved carbon dioxide (including carbonic acid) plus the carbon dioxide present as bicarbonate.

A useful way to display the interrelationships among the variables of pH,

As can be seen from Figure 8–1, pH is on the abscissa of the pH-bicarbonate diagram, and the plasma bicarbonate concentration in millimoles per liter is on the ordinate. The hydrogen ion concentration corresponding to the pH on the X-axis is shown at the top of the graph. Note that the hydrogen ion concentration is expressed in nanomoles; the bicarbonate concentration on the Y-axis is expressed in millimoles. For each value of pH and bicarbonate ion concentration, there is a single corresponding

The buffer value of bicarbonate in a closed system without the presence of hemoglobin is only about −5.4 mmol/L/pH unit. That is, the bicarbonate concentration increases only 5.4 mmol/L as enough acid in the form of carbon dioxide is added to the bicarbonate system to lower the pH by 1 unit. The bicarbonate buffer system is therefore a poor buffer for carbonic acid. The presence of hemoglobin makes blood a much better buffer, as can be seen in Figure 8–2.

The figure shows that increasing the hemoglobin concentration in this in vitro experiment makes the buffering curve steeper. That is, the bicarbonate concentration increases more at greater hemoglobin concentrations as carbonic acid (in the form of CO2) is added to the blood. The increase in bicarbonate concentration is greater with more hemoglobin because as carbon dioxide is added to the blood, the hydrogen ions formed by the dissociation of carbonic acid are buffered by hemoglobin (as will be discussed shortly). Most of the bicarbonate ions formed by this dissociation can therefore move into the plasma. The buffer value of plasma in the presence of hemoglobin is thus 4 to 5 times that of plasma separated from erythrocytes. Therefore, the slope of the normal in vivo buffer line shown in Figures 8–1 and 8–3 is mainly determined by the nonbicarbonate buffers present in the body.

Phosphate

The phosphate buffer system mainly consists of the buffer pair of the dihydrogen phosphate (

The pK of the acid form is 6.8, so that in pHs ranging near 7.0, the acid form can readily donate a proton and the base form can accept a proton. Many organic phosphates found in the body also have pKs within ±0.5 pH units of 7.0, and these compounds can also function as buffers under physiologic conditions. These organic phosphates include such compounds as glucose-1-phosphate and adenosine triphosphate.

Proteins

Although several potential buffering groups are found on proteins, only 1 large group has pKs in the pH range encountered in the blood. These are the imidazole groups in the histidine residues of the peptide chains. The pKs of the various histidine residues on the different plasma proteins range from about 5.5 to about 8.5, thus providing a broad spectrum of buffer pairs. The protein present in the greatest quantity in the blood is hemoglobin. Thirty-six of the 540 amino acid residues in hemoglobin are histidine, with pKs ranging from 7 to 8; the N-terminal valine residues also have a pK of about 7.8. As already noted, deoxyhemoglobin is a weaker acid than is oxyhemoglobin. That is, the pK of an imidazole group of one of the histidine residues in deoxygenated hemoglobin is greater than it is in the oxyhemoglobin state. Thus, as oxygen leaves hemoglobin in the tissue capillaries, the imidazole group removes hydrogen ions from the erythrocyte interior, allowing more carbon dioxide to be transported as bicarbonate. This process is reversed in the lungs.

Buffers of the Interstitial Fluid

The bicarbonate buffer system is the major buffer found in the interstitial fluid, including the lymph. The phosphate buffer pair is also found in the interstitial fluid. The volume of the interstitial compartment is much larger than that of the plasma, and so the interstitial fluid may play an important role in buffering.

Bone

The extracellular portion of bone contains very large deposits of calcium and phosphate salts, mainly in the form of hydroxyapatite. Although bone growth in a child causes a net production of hydrogen ions, in an otherwise healthy adult, where bone growth and resorption are in a steady state, bone salts can buffer hydrogen ions in chronic acidosis. Chronic buffering of hydrogen ions by the bone salts may therefore lead to demineralization of bone.

Intracellular Buffering

The intracellular proteins and organic phosphates of most cells can function to buffer both fixed acids and carbonic acid. Again, this is largely a function of the histidine groups on the proteins and phosphate groups on such compounds as ATP (adenosine triphosphate) and glucose-1-phosphate. Of course, buffering by the hemoglobin in erythrocytes is intracellular buffering.

Acid-base disorders can be divided into 4 major categories: respiratory acidosis, respiratory alkalosis, metabolic acidosis, and metabolic alkalosis. These primary acid-base disorders may occur singly (“simple”) or in combination (“mixed”) or may be altered by compensatory mechanisms.

Respiratory Acidosis

The arterial

In respiratory acidosis, the ratio of bicarbonate to CO2 decreases. Yet, as can be seen at point C in Figure 8–3, in uncompensated primary (simple) respiratory acidosis, the absolute plasma bicarbonate concentration does increase somewhat because of the buffering of some of the hydrogen ions liberated by the dissociation of carbonic acid by nonbicarbonate buffers.

Any impairment of alveolar ventilation can cause respiratory acidosis. As shown in Table 8–1, depression of the respiratory centers in the medulla (see Chapter 9) by anesthetic agents, narcotics, hypoxia, central nervous system disease or trauma, or even greatly elevated

Respiratory Alkalosis

The causes of respiratory alkalosis include anything leading to hyperventilation. As shown in Table 8–2, hyperventilation syndrome, a psychological dysfunction of unknown cause, results in chronic or recurrent episodes of hyperventilation and respiratory alkalosis. Drugs, hormones (such as progesterone), toxic substances, central nervous system diseases or disorders, bacteremias, fever, overventilation by mechanical ventilators (or the physician), or ascent to high altitude may all result in respiratory alkalosis. Students are often surprised to see acute asthma included in a list of problems that can cause respiratory alkalosis. Asthma is an episodic obstructive disease and it is reasonable to assume that it would cause CO2 retention and therefore respiratory acidosis during attacks. That is true in very severe asthma attacks, but most asthma attacks result in hypocapnia and respiratory alkalosis. As the asthma attack occurs, bronchial smooth muscle spasm and mucus secretion obstruct the ventilation to some alveoli. Although some hypoxic pulmonary vasoconstriction may occur, it is not sufficient to divert all of the mixed venous blood flow away from these poorly ventilated alveoli. That results in a right-to-left shunt or shunt-like state, which would therefore be expected to cause the arterial

Irritant receptors in the airways are stimulated by the mucus and by chemical mediators released during the attack. Hypoxia caused by the shunt stimulates the arterial chemoreceptors; the patient also has the feeling of dyspnea (many asthma attacks have an emotional component). All of these things cause increased breathing and therefore increased alveolar ventilation.

Increasing ventilation will get more CO2 out of the blood perfusing ventilated alveoli (and therefore out of the body) but it will not get much oxygen into alveoli supplied by obstructed airways, nor will it get much more oxygen into the blood of the unobstructed alveoli because of the shape of the oxyhemoglobin dissociation curve. Remember that the hemoglobin is already 97.4% saturated with oxygen and not much more will dissolve in the plasma. Therefore, during the attack the patient has hypoxemia, hypocapnia, and respiratory alkalosis. It is only when the attack is so severe that the patient can’t do the additional work of breathing that hypercapnia and respiratory acidosis occur.

Metabolic Acidosis

As shown in Table 8–3, ingestion of methyl alcohol or salicylates can cause metabolic acidosis by increasing the fixed acids in the blood. (Salicylate poisoning—eg, aspirin overdose—causes both metabolic acidosis and later respiratory alkalosis.) Diarrhea can cause very great bicarbonate losses, resulting in metabolic acidosis. Renal dysfunctions can lead to an inability to excrete hydrogen ions, as well as an inability to reabsorb bicarbonate ions, as will be discussed in the next section. True “metabolic” acidosis may be caused by an accumulation of lactic acid in severe hypoxemia or shock and by diabetic ketoacidosis.

Metabolic Alkalosis

As shown in Table 8–4, loss of gastric juice by vomiting results in a loss of hydrogen ions and may cause metabolic alkalosis. Excessive ingestion of bicarbonate or other bases (eg, stomach antacids) or over infusion of bicarbonate by the physician may cause metabolic alkalosis. In addition, diuretic therapy, treatment with steroids (or the overproduction of endogenous steroids), and conditions leading to severe potassium depletion may also cause metabolic alkalosis.

Uncompensated primary acid-base disturbances, such as those indicated by points B, C, D, and G on Figure 8–3, are seldom seen because respiratory and renal compensatory mechanisms are called into play to offset these disturbances. The 2 main compensatory mechanisms are functions of the respiratory and renal systems.

Respiratory Compensatory Mechanisms

The respiratory system can compensate for metabolic acidosis or alkalosis by altering alveolar ventilation. As discussed in Chapter 3, if carbon dioxide production is constant, the alveolar

The respiratory compensation for metabolic alkalosis is to decrease alveolar ventilation, thus raising

Under most circumstances the cause of respiratory acidosis or alkalosis is a dysfunction in the ventilatory control mechanism or the breathing apparatus itself. Compensation for acidosis or alkalosis in these conditions must therefore come from outside the respiratory system. The respiratory compensatory mechanism can operate very rapidly (within minutes) to partially correct metabolic acidosis or alkalosis.

Renal Compensatory Mechanisms

The kidneys can compensate for respiratory acidosis and metabolic acidosis of nonrenal origin by excreting fixed acids and by retaining filtered bicarbonate, thus increasing net acid excretion. The kidneys can also compensate for respiratory alkalosis or metabolic alkalosis of nonrenal origin by decreasing hydrogen ion excretion and by increasing the excretion of bicarbonate, thus decreasing net acid excretion.

Renal Mechanisms in Acidosis

The renal tubular cells secrete hydrogen ions into the tubular fluid. This is mainly accomplished by the generation of hydrogen ions and bicarbonate ions within the cell by the dissociation of carbonic acid. The carbonic acid is formed by the hydration of carbon dioxide via the carbonic anhydrase reaction. The carbon dioxide may be metabolically produced by the tubular cell itself or may be carried dissolved into the tubular fluid after production elsewhere in the body. The hydrogen ion generated by this process is actively secreted into the tubular lumen, and the new bicarbonate ion is “reabsorbed” into the peritubular capillary. Sodium ions in the tubular fluid are exchanged for the hydrogen ions secreted into the tubular fluid via the apical sodium-hydrogen exchanger to maintain electrical neutrality. There are also proton pumps and proton exchange pumps located on the apical membranes of specific nephron segments. Following metabolism of glutamine by the proximal tubule cells, ammonia or ammonium ions are secreted into the tubular fluid and bicarbonate ions enter the capillaries. The hydrogen ion secreted into the tubular lumen that is buffered by tubular phosphate, sulphate, ammonia, or other urinary buffers represents the excretion of acid and the synthesis of new bicarbonate. The increased urinary excretion of acid contributes to the renal compensation for acidosis.

Normally, the kidneys synthesize about 70 mEq of new bicarbonate to buffer nonvolatile acid production by the body daily. The urinary excretion of ammonium can be increased 10-fold during severe acidosis. Net acid excretion can increase during acidosis to the extent that the urine can be acidified to a pH of 4.0 to 5.0. This is about 800 times as acidic as normal plasma.

Renal Mechanisms in Alkalosis

In alkalotic states, the kidney decreases the secretion of hydrogen ions and decreases the reabsorption of filtered bicarbonate resulting in an increase in bicarbonate excretion. The enhanced renal excretion of bicarbonate returns the pH and bicarbonate concentration toward normal.

Time Course of Renal Mechanisms

Renal compensatory mechanisms for acid-base disturbances operate much more slowly than respiratory compensatory mechanisms because the tubular epithelial cells must increase the synthesis of enzymes and membrane transporters. The renal compensatory responses to sustained respiratory acidosis or alkalosis may take 3 to 6 days.

Summary of Renal and Respiratory Contributions to Acid-Base Balance

The kidneys help regulate acid-base balance by altering the excretion of fixed acids and the retention of the filtered bicarbonate; the respiratory system helps regulate body acid-base balance by adjusting alveolar ventilation to alter alveolar

Samples of arterial blood are usually analyzed clinically to determine the “arterial blood gases”: the arterial

Table 8–5 summarizes the changes in pHa,

A simple approach to interpreting a blood gas set is to first look at the pH to determine whether the predominant problem is acidosis or alkalosis. (Note that an acidemia could represent more than 1 cause of acidosis, an acidosis with some compensation, or even an acidosis and a separate underlying alkalosis. Similarly, an alkalemia could represent more than 1 cause of alkalosis, an alkalosis with some compensation, or even an alkalosis and a separate underlying acidosis.) After looking at the pH, look at the arterial

If the pH is high and the

Base Excess

Calculation of the base excess or base deficit may be very useful in determining the therapeutic measures to be administered to a patient. The base excess or base deficit is the number of milliequivalents of acid or base needed to titrate 1 L of blood to pH 7.4 at 37°C if the

The base deficit can be used to estimate how much sodium bicarbonate (in mEq) should be given to a patient by multiplying the base deficit (in mEq/L) times the patient’s estimated extracellular fluid (ECF) space (in liters), which is the distribution space for the bicarbonate. The ECF is usually estimated to be 0.3 times the lean body mass in kilograms.

Anion Gap

Calculation of the anion gap can be helpful in determining the cause of a patient’s metabolic acidosis. It is determined by subtracting the sum of a patient’s plasma chloride and bicarbonate concentrations (in mEq/L) from his or her plasma sodium concentration:

The anion gap is normally 12 ± 4 mEq/L.

The sum of all of the plasma cations must equal the sum of all of the plasma anions, so the anion gap exists only because all of the plasma cations and anions are not measured when standard blood chemistry is done. Sodium, chloride, and bicarbonate concentrations are almost always reported. The normal anion gap is a result of the presence of more unmeasured anions than unmeasured cations in normal blood.

The anion gap is therefore the difference between the unmeasured anions and the unmeasured cations.

The negative charges on the plasma proteins probably make up most of the normal anion gap because the total charges of the other plasma cations (K+, Ca2+, Mg2+) are approximately equal to the total charges of the other anions (

Thus, metabolic acidosis with an abnormally great anion gap (ie, greater than 16 mEq/L) would probably be caused by lactic acidosis or ketoacidosis; ingestion of organic anions such as salicylate, methanol, and ethylene glycol; or renal retention of anions such as sulfate, phosphate, and urate.

Thus far only 2 of the 3 variables referred to as the arterial blood gases, the arterial

Hypoxic Hypoxia

Hypoxic hypoxia refers to conditions in which the arterial

Low Alveolar

Conditions causing low alveolar

Hypoventilation can be caused by depression or injury of the respiratory centers in the brain (discussed in Chapter 9); interference with the nerves supplying the respiratory muscles, as in spinal cord injury; neuromuscular junction diseases such as myasthenia gravis; and altered mechanics of the lung or chest wall, as in noncompliant lungs due to sarcoidosis, reduced chest wall mobility because of kyphoscoliosis or obesity, and airway obstruction. Ascent to high altitudes causes alveolar hypoxia because of the reduced total barometric pressure encountered above sea level. Reduced

Diffusion Impairment

Alveolar-capillary diffusion is discussed in greater detail in Chapter 6. Conditions such as interstitial fibrosis and interstitial or alveolar edema can lead to low arterial

Shunts

True right-to-left shunts, such as anatomic shunts and absolute intrapulmonary shunts, can cause decreased arterial

Ventilation-Perfusion Mismatch

Alveolar-capillary units with low ventilation-perfusion (

Note that diffusion impairment, shunts, and

Anemic Hypoxia

Anemic hypoxia is caused by a decrease in the amount of functioning hemoglobin, which can be a result of decreased hemoglobin or erythrocyte production, the production of abnormal hemoglobin or red blood cells, pathologic destruction of erythrocytes, or interference with the chemical combination of oxygen and hemoglobin. Carbon monoxide poisoning, for example, results from the greater affinity of hemoglobin for carbon monoxide than for oxygen. Methemoglobinemia is a condition in which the iron in hemoglobin has been altered from the Fe2+ to the Fe3+ form, which does not combine with oxygen.

Anemic hypoxia results in a decreased oxygen content when both alveolar and arterial

Hypoperfusion Hypoxia

Hypoperfusion hypoxia (sometimes called stagnant hypoxia) results from low blood flow. This can occur either locally, in a particular vascular bed, or systemically, in the case of a low cardiac output. The alveolar

Histotoxic Hypoxia

Histotoxic hypoxia refers to a poisoning of the cellular machinery that uses oxygen to produce energy. Cyanide, for example, binds to cytochrome oxidase in the respiratory chain and effectively blocks oxidative phosphorylation. Alveolar

Other Causes of Hypoxia

Tissue edema or fibrosis may result in impaired diffusion of oxygen from the blood to the tissues. It is also conceivable that the delivery of oxygen to a tissue is completely normal, but the tissue’s metabolic demands still exceed the supply and tissue hypoxia could result. This is known as overutilization hypoxia.

The Effects of Hypoxia

Hypoxia can result in reversible tissue injury or even tissue death. The outcome of a hypoxic episode depends on whether the tissue hypoxia is generalized or localized, on how severe the hypoxia is, on the rate of development of the hypoxia (see Chapter 11), and on the duration of the hypoxia. Different cell types have different susceptibilities to hypoxia; unfortunately, brain cells and heart cells are the most susceptible.

From Raff H, Levitzky MG, eds. Medical Physiology: A Systems Approach. New York: McGraw-Hill; 2011:383–384.

A 15-year-old adolescent entered the emergency department with dyspnea (shortness of breath), a feeling of chest tightness, coughing, wheezing, and anxiety. His nail beds and lips showed some cyanosis.

He has had frequent episodes of dyspnea and wheezing for several years, especially in the spring and has been diagnosed with asthma. Pulmonary function tests done at the time of his diagnosis showed lower than predicted FEV1 (forced expiratory volume in the first second), FVC (forced vital capacity), FEV1/FVC, and PEF (peak expiratory flow). Inhaling a bronchodilator improved all of these.

An arterial blood gas was obtained to help determine the severity of the episode. The arterial

Asthma is an episodic obstructive disease and, as discussed in the text, it is reasonable to assume that it would cause CO2 retention and respiratory acidosis during attacks. That is true in very severe asthma attacks, but most asthma attacks result in hypocapnia and respiratory alkalosis. As the asthma attack occurs, bronchial smooth muscle spasm and mucus secretion obstruct the ventilation to some alveoli. Hypoxic pulmonary vasoconstriction may occur, but it is not sufficient to divert all of the mixed venous blood flow away from these poorly ventilated alveoli, resulting in a right-to-left shunt or shunt-like state (Chapter 5), which would be expected to cause the arterial

Increasing ventilation will get more CO2 out of the blood perfusing ventilated alveoli (and therefore out of the body) but it will not get much oxygen into alveoli supplied by obstructed airways, nor will it get much more oxygen into the blood of the unobstructed alveoli because of the shape of the oxyhemoglobin dissociation curve. Therefore, during the attack the patient has hypoxemia, hypocapnia, and respiratory alkalosis.

Acute treatment of asthma is aimed at dilating the airways with a bronchodilator, such as a β2 adrenergic agonist and relieving the hypoxemia with oxygen. Mechanical ventilation may be used in more severe cases. Chronic treatment includes bronchodilators such as β2 adrenergic agonists; anticholinergics to block parasympathetically mediated constriction and mucus production; antileukotriene drugs and corticosteroids to prevent inflammation; and inhibition of mast cells to prevent them from releasing cytokines.