27. ACID–BASE BALANCE

27. ACID–BASE BALANCE

•  various ions play different roles that help maintain homeostasis.

• A major homeostatic challenge is keeping the H+ concentration (pH) of body fluids at an appropriate level. 
• This task—the maintenance of acid–base balance—is of critical importance to normal cellular function. 

For example, 

• the three-dimensional shape of all body proteins, which enables them to perform specific functions, is very sensitive to pH changes. 

When the diet contains a large amount of protein, as is typical in North America,

•  cellular metabolism produces more acids than bases, 
• which tends to acidify the blood. 

• In a healthy person, several mechanisms help maintain the pH of systemic arterial blood between 7.35 and 7.45. 
• (A pH of 7.4 corresponds to a H concentration of 0.00004 mEq/liter 40 nEq /liter.) 

Because metabolic reactions often produce a huge excess of H+ , 

• the lack of any mechanism for the disposal of H + would cause H+ level in body fluids to rise quickly to a lethal level. 

• Homeostasis of H+ concentration within a narrow range is thus essential to survival. 
• The removal of H + from body fluids and its subsequent elimination from the body depend on the following three major mechanisms: 

1. Buffer systems. 

• Buffers act quickly to temporarily bind H+ ,
• removing the highly reactive, excess H + from solution. 
• Buffers thus raise pH of body fluids but do not remove H+ from the body.

2. Exhalation of carbon dioxide. 

• By increasing the rate and depth of breathing, more carbon dioxide can be exhaled. 
• Within minutes this reduces the level of carbonic acid in blood, 
• which raises the blood pH (reduces blood H + level). 

3. Kidney excretion of H+ . 

• The slowest mechanism, but the only way to eliminate acids other than carbonic acid, is through their excretion in urine.

The Actions of Buffer Systems

 Most buffer systems in the body consist of 

1. a weak acid and 
2. the salt of that acid, 

• which functions as a weak base. 

• Buffers prevent rapid, drastic changes in the pH of body fluids by converting strong acids and bases into weak acids and weak bases within fractions of a second. 

• Strong acids lower pH more than weak acids because strong acids release H+  more readily
• and thus contribute more free hydrogen ions. 

• Similarly, strong bases raise pH more than weak ones. 

• The principal buffer systems of the body fluids are 

1. the protein buffer system, 
2. the carbonic acid–bicarbonate buffer system, 
3. and the phosphate buffer system.

Protein Buffer System

• The protein buffer system is the most abundant buffer in intracellular fluid and blood plasma. 

For example, 

• the protein hemoglobin is an especially good buffer within red blood cells, and
• albumin is the main protein buffer in blood plasma. 
• Proteins are composed of amino acids, organic molecules that contain at least

1. one carboxyl group (−COOH) and 
2. at least one amino group (−NH2); 

• These groups are the functional components of the protein buffer system. 

• The free carboxyl group at one end of a protein acts like an acid by releasing H+  when pH rises; 

it dissociates as follows:

                 

              R

              |

         NH₂ -C-­ COOH  

              |

              H 

      

              ↓

             R

             |

       NH₂ ­- C - COO− + H+

             |

             H    

• The H+ is then able to react with any excess  OH−  in the solution to form water. 
• The free amino group at the other end of a protein can act as a base by combining with H+  when pH falls, as follows: 

     

              R

              |

         NH₂ -C-­ COOH + H+ 

              |

              H 

      

              ↓

              R

              |

        +NH3 ­- C - COOH

              |

              

              H 

• So proteins can buffer both acids and bases. 

• In addition to the terminal carboxyl and amino groups, side chains that can buffer H+  are present on seven of the 20 amino acids.

•  the protein hemoglobin is an important buffer of H+ in red blood cells . 

As blood flows through the systemic capillaries,

• carbon dioxide (CO₂) passes from tissue cells into red blood cells, 
• where it combines with water (H₂O) 
• to form carbonic acid H₂CO₃
• Once formed, H₂CO₃ dissociates into H+ and HCO₃−

At the same time that CO2 is entering red blood cells, 

• oxyhemoglobin (Hb - O2) is giving up its oxygen to tissue cells. 
• Reduced hemoglobin (deoxyhemoglobin) picks up most of the H+ . 
• For this reason, reduced hemoglobin usually is written as Hb - H.

The following reactions summarize these relations:

(H₂O)    +    CO₂     →   H₂CO₃
Water       Carbon            Carbonic acid
                dioxide 
         (entering RBCs)


      H₂CO₃       →  H +    +    HCO3−

Carbonic            Hydrogen     Bicarbonate
    acid                      ion              ion

Hb–O2        +   H+      → Hb–H       +     O2
Oxy                Hydrogen       Reduced        Oxygen
hemoglobin     ion                 hemo            (released 
(in RBCs)       (from             globin              to
                        carbonic                              tissue 
                           acid)                                cells)

Carbonic Acid–Bicarbonate Buffer System

• The carbonic acid–bicarbonate buffer system is based on the bicarbonate ion (HCO3− ), 
• which can act as a weak base, 

• and carbonic acid (H₂CO₃), 
• which can act as a weak acid.  

• HCO3−  is a significant anion in both intracellular and extracellular fluids. 
• Because the kidneys also synthesize new HCO3−  and reabsorb filtered HCO3− , this important buffer is not lost in the urine. 

• If there is an excess of H+ , the HCO3−  can function as a weak base and remove the excess H as follows:

     H+          +    HCO3−        →   H₂CO₃
Hydrogen         Bicarbonate          Carbonic acid
    ion                     ion
                       (weak base)

• Then,  H₂CO₃ dissociates into water and carbon dioxide, 
• and the CO2 is exhaled from the lungs.

Conversely, if there is a shortage of H+, 

• the  H₂CO₃ can function as a weak acid and 
• provide H+ as follows:

        H₂CO₃         →        H+       +       HCO3− 

Carbonic acid      Hydrogen       Bicarbonate 
(weak acid)             ion                    ion


At a pH of 7.4,

•  HCO3−  concentration is about 24 mEq/liter
• and H₂CO₃ concentration is about 1.2 mmol/liter, 
• so bicarbonate ions outnumber carbonic acid molecules by 20 to 1. 

• Because CO2 and H₂O combine to form H₂CO₃, 
• this buffer system cannot protect against pH changes due to respiratory problems in which there is an excess or shortage of CO2.

Phosphate Buffer System

• The phosphate buffer system acts via a mechanism similar to the one for the carbonic acid–bicarbonate buffer system. 
•  The components of the phosphate buffer system are the ions 

1. dihydrogen phosphate H₂PO₄⁻ and 
2. monohydrogen phosphate (HPO₄²⁻ ). 

phosphates 

• are major anions in intracellu-lar fluid 
• and minor ones in extracellular fluids. 

The dihydrogen phosphate ion 

• acts as a weak acid 
• and is capable of buffering strong bases such as OH , as follows:

          OH⁻    +    H₂PO₄⁻       →     H₂O       +   HPO₄²⁻

Hydroxide      Dihydrogen      Water      Mono         
     ion               phosphate                     hydrogen
                                                              phosphate
(strong base)   (weak acid)                  (weak base)


The monohydrogen phosphate ion is capable of buffering the H+ released by a strong acid such as hydrochloric acid (HCl) by acting as a weak base:

        H+     +   HPO₄²⁻                           →    H₂PO₄⁻

Hydrogen         Monohydrogen            Dihydrogen
ion                       phosphate                  phosphate
(strong acid)      (weak base)                (weak acid)

Because the concentration of phosphates is highest in intracellular fluid,

•  the phosphate buffer system is an important regulator of pH in the cytosol. 
• It also acts to a smaller degree in extracellular fluids and
•  buffers acids in urine. 
• H₂PO₄ ²⁻ is formed when excess H+ in the kidney tubule fluid combines with HPO₄²⁻. 
• The H+  that becomes part of the H₂PO₄⁻ passes into the urine. 
• This reaction is one way the kidneys help maintain blood pH by excreting H in the urine.

Exhalation of Carbon Dioxide

• The simple act of breathing also plays an important role in maintaining the pH of body fluids. 
• An increase in the carbon dioxide (CO2) concentration in body fluids increases H+ concentration
• and thus lowers the pH (makes body fluids more acidic).
• Because H₂CO₃  can be eliminated by exhaling CO2 , it is called a volatile acid. 

Conversely, 

• a decrease in the CO2 concentration of body fluids raises the pH (makes body fluids more alkaline). 
• This chemical interaction is illustrated by the following reversible reactions:

  CO2    +  H₂O   ⇋  H₂CO₃  ⇋    H+    +   HCO₃−


Carbon    Water    Carbonic    Hydrogen        Bi 
dioxide                     acid             ion        carbonate
                                                                      ion

• Changes in the rate and depth of breathing can alter the pH of body fluids within a couple of minutes. 

• With increased ventilation, more  CO2 is exhaled. 

When CO2  levels decrease, 

• the reaction is driven to the left , 

• H+  concentration falls,
• and blood pH increases. 

• Doubling the ventilation increases pH by about 0.23 units, from 7.4 to 7.63. 

• If ventilation is slower than normal, less carbon dioxide is exhaled. 

When CO2 levels increase, 

• the reaction is driven to the right ,
•  the H+ concentration increases,
•  and blood pH decreases. 

• Reducing ventilation to one-quarter of normal lowers the pH by 0.4 units, from 7.4 to 7.0. 
• These examples show the powerful effect of alterations in breathing on the pH of body fluids.

• The pH of body fluids and the rate and depth of breathing interact via a negative feedback loop. 
• When the blood acidity increases, the decrease in pH (increase in concentration of H+ ) is detected by  

1. central chemoreceptors in the medulla oblongata 
2. and peripheral chemoreceptors in the aortic and carotid bodies, 

• both of which stimulate the inspiratory area in the medulla oblongata. 
• As a result, the diaphragm and other respiratory muscles contract more forcefully and frequently, so more CO2 is exhaled. 

• As less H2CO3 forms and fewer H+  are present, blood pH increases. 
• When the response brings blood pH (H+  concentration) back to normal, there is a return to acid–base homeostasis. 

• The same negative feedback loop operates if the blood level of CO2 increases. 
• Ventilation increases, 
• which removes more CO2, reducing the H concentration and
• increasing the blood’s pH.

By contrast,

•  if the pH of the blood increases, 
• the respiratory center is inhibited and
•  the rate and depth of breathing decreases.

• A decrease in the CO2 concentration of the blood has the same effect. 
• When breathing decreases, CO2 accumulates in the blood so its H concentration increases. 

Kidney Excretion of H+

• Metabolic reactions produce nonvolatile acids such as sulfuric acid at a rate of about 1 mEq of H+per day for every kilogram of body mass. 
• The only way to eliminate this huge acid load is to excrete H+ in the urine.

•  Given the magnitude of these contributions to acid–base balance, it’s not surprising that renal failure can quickly cause death.

• cells in both the proximal convoluted tubules (PCT) and the collecting ducts of the kidneys secrete hydrogen ions into the tubular fluid. 
• In the PCT, Na+ /H+ antiporters secrete H+ as they reabsorb Na+. 

• Even more important for regulation of pH of body fluids, however, are the intercalated cells of the collecting duct.

• The apical membranes of some intercalated cells include proton pumps (H ATPases) that secrete H+  into the tubular fluid. 
• Intercalated cells can secrete H+  against a concentration gradient so effectively that urine can be up to 1000 times(3 pH units) more acidic than blood.  
• HCO3⁻ produced by dissociation of H₂CO₃ inside intercalated cells crosses the basolateral membrane by means of Cl − /HCO3−  antiporters 
• and then diffuses into peritubular capillaries . 

• The HCO3⁻ that enters the blood in this way is new (not filtered). 
• For this reason, blood leaving the kidney in the renal vein may have a higher HCO3− concentration than blood entering the kidney in the renal artery.

• Interestingly, a second type of intercalated cell has proton pumps in its basolateral membrane and Cl− / HCO3− antiporters in its apical membrane. 
• These intercalated cells secrete HCO3−
• and reabsorb H+ . 

• Thus, the two types of intercalated cells help maintain the pH of body fluids in two ways—

1. by excreting excess H+  when pH of body fluids is too low 
2. and by excreting excess HCO3− when pH is too high.

• Some H+  secreted into the tubular fluid of the collecting duct are buffered, 
• but not by HCO3− , most of which has been filtered and reabsorbed. 

• Two other buffers combine with H+  in the collecting duct . 

• The most plentiful buffer in the tubular fluid of the collecting duct is HPO₄²⁻ (monohydrogen phosphate ion). 
• In addition, a small amount of NH3 (ammonia) also is present. 

H+  combines 

• with HPO₄²⁻ to form  H₂PO₄⁻  (dihydrogen  phosphate ion) 
• and with NH3 to form NH4 + (ammonium ion). 

• Because these ions cannot diffuse back into tubule cells, they are excreted in the urine.

Mechanisms That Maintain pH of Body Fluids
MECHANISM	COMMENTS
Buffer systems	Most consist of a weak acid and the salt of that acid, which functions as a weak base.  They prevent drastic changes in body fluid pH.
Proteins	The most abundant buffers in body cells and blood.  Hemoglobin inside red blood cells is a good buffer.
Carbonic acid– bicarbonate	Important regulator of blood pH.  The most abundant buffers in extracellular fluid (ECF).
Phosphates	Important buffers in intracellular fluid and in urine.
Exhalation of CO2	With increased exhalation of CO2, pH rises (fewer H ).  With decreased exhalation of CO2, pH falls (more H ).
Kidneys	Renal tubules secrete H into the urine and reabsorb HCO3− so it is not lost in the urine.

Acid–Base Imbalances

• The normal pH range of systemic arterial blood is between 7.35( 45 nEq of H /liter) and 7.45 ( 35 nEq of H /liter).

Acidosis (or acidemia)

• is a condition in which blood pH is below 7.35; 

alkalosis (or alkalemia) 

• is a condition in which blood pH is higher than 7.45.

The major physiological effect of acidosis is 

• depression of the central nervous system through depression of synaptic transmission.

• If the systemic arterial blood pH falls below 7, 
• depression of the nervous system is so severe 
• that the individual becomes disoriented, then comatose, and may die. 

• Patients with severe acidosis usually die while in a coma. 

A major physiological effect of alkalosis, by contrast, is 

• overexcitability in both the central nervous system and peripheral nerves. 

• Neurons conduct impulses repetitively, even when not stimulated by normal stimuli;

the results are 

• nervousness,
•  muscle spasms, 
• and even convulsions and death.

A change in blood pH that leads to acidosis or alkalosis may be countered by compensation, 

• the physiological response to an acid–base imbalance that acts to normalize arterial blood pH.

• Compensation may be either 

1. complete, if pH indeed is brought within the normal range, 
2. or partial, if systemic arterial blood pH is still lower than 7.35 or higher than 7.45. 

If a person has altered blood pH due to metabolic causes, 

• hyperventilation or hypoventilation can help bring blood pH back toward the normal range; this form of compensation,
•  termed respiratory compensation,
• occurs within minutes 
• and reaches its maximum within hours. 

• If, however, a person has altered blood pH due to respiratory causes, 
• then renal compensation—changes in secretion of H+ and reabsorption of HCO3⁻ by the kidney tubules—
• can help reverse the change. 
• Renal compensation may begin in minutes,
• but it takes days to reach maximum effectiveness.

• note that both respiratory acidosis and respiratory alkalosis are disorders resulting from changes in the partial pressure of CO2 (PCO2) in systemic arterial blood (normal range is 35–45 mmHg). 

• By contrast, both metabolic acidosis and metabolic alkalosis are disorders resulting from changes in HCO3⁻ concentration 
• (normal range is 22–26 mEq/liter in systemic arterial blood).

Respiratory Acidosis

• The hallmark of respiratory acidosis is an abnormally high PCO2 in systemic arterial blood—above 45 mmHg. 
• Inadequate exhalation of CO2 causes the blood pH to drop. 

• Any condition that decreases the movement of CO2 from the blood to the alveoli of the lungs to the atmosphere causes a buildup of  CO2, H2CO3, and H+  . 
• Such conditions include

1. emphysema, 
2. pulmonary edema, 
3. injury to the respiratory center of the medulla oblongata, 
4. airway obstruction, 
5. or disorders of the muscles involved in breathing. 

• If the respiratory problem is not too severe, the kidneys can help raise the blood pH into the normal range by 

1. increasing excretion of H+  
2.  and reabsorption of HCO3 ⁻(renal compensation). 

The goal in treatment of respiratory acidosis 

• is to increase the exhalation of CO2,
•  as, for instance, by providing ventilation therapy. 

In addition, 

• intravenous administration of HCO3− may be helpful. 

Respiratory Alkalosis

• In respiratory alkalosis, systemic arterial blood PCO2 falls below 35 mmHg. 
• The cause of the drop in PCO2 and the resulting increase in pH is hyperventilation,
• which occurs in conditions that stimulate the inspiratory area in the brain stem. 

• Such conditions include 

1. oxygen deficiency due to high altitude or pulmonary disease, 
2. cerebrovascular accident (stroke), 
3. or severe anxiety. 

• Again, renal compensation may bring blood pH into the normal range if the kidneys are able to decrease excretion of H+ and reabsorption of HCO3⁻

Treatment of respiratory alkalosis 

• is aimed at increasing the level of CO2 in the body. 
• One simple treatment is to have the person inhale and exhale into a paper bag for a short period; 
• as a result, the person inhales air containing a higher-than-normal concentration of CO2.

Metabolic Acidosis

• In metabolic acidosis,  the systemic arterial blood HCO3 ⁻ level drops below 22 mEq/liter. 
• Such a decline in this important buffer causes the blood pH to decrease. 

Three situations may lower the blood level of HCO3−


(1) actual loss of HCO3−, 

• such as may occur with severe diarrhea or renal dysfunction;

(2) accumulation of an acid other than carbonic acid, 

• as may occur in ketosis ; or 

(3) failure of the kidneys to excrete H from metabolism of dietary proteins. 

• If the problem is not too severe, hyperventilation can help bring blood pH into the normal range (respiratory compensation).

Treatment of metabolic acidosis 

• consists of administering intravenous solutions of sodium bicarbonate 
• and correcting the cause of the acidosis.

Metabolic Alkalosis

• In metabolic alkalosis, the systemic arterial blood HCO3− concentration is above 26 mEq/liter. 

• A nonrespiratory loss of acid or excessive intake of alkaline drugs causes the blood pH to increase above 7.45. 

1. Excessive vomiting of gastric contents, which results in a substantial loss of hydrochloric acid, is probably the most frequent cause of metabolic alkalosis. Other causes include 
2. gastric suctioning, 
3. use of certain diuretics, 
4. endocrine disorders, 
5. excessive intake of alkaline drugs (antacids), and
6. severe dehydration. 

• Respiratory compensation through hypoventilation may bring blood pH into the normal range. 

Treatment of metabolic alkalosis 

• consists of giving fluid solutions to correct Cl ⁻, K+ , and other electrolyte deficiencies 
• plus correcting the cause of alkalosis.

Summary of Acidosis and Alkalosis
CONDITION	DEFINITION	COMMON                CAUSES	COMPENSATORY          MECHANISM
Respiratory acidosis	Increased PCO2 (above 45 mm Hg), and decreased pH (below 7.35 if there is no compensation.	Hypoventilation due to emphysema)   pulmonary edema,  trauma to respiratory center,  airway obstructions, or dysfunction of muscles of respiration.	Renal: increased excretion of H+; increased reabsorption of HCO3−.  If compensation is complete, pH will be within the normal range  but PCO2 will be high.
Respiratory alkalosis	Decreased PCO2 (below 35 mmHg)  and increased pH (above 7.45) if there is no compensation	Hyperventilation due to  oxygen deficiency,  pulmonary disease,  cerebrovascular accident (CVA),  or severe anxiety.	Renal:   decreased excretion of H+;  decreased reabsorption of HCO3−. If compensation is complete, Ph will be within the normal range  but PCO2 will be low.
Metabolic acidosis	Decreased HCO3−_ (below22 mEq/liter)  and decreased pH (below 7.35) if there is no compensation.	Loss of bicarbonate ions due to   diarrhea,  accumulation of acid (ketosis),  renal dysfunction.	Respiratory:  hyperventilation, which increases loss of CO2.  If compensation is complete, pH will be within the normal range  but HCO3−will be low
Metabolic alkalosis	Increased HCO3−(above26 mEq/liter)  and increased pH (above 7.45) if there is no compensation.	Loss of acid due to  vomiting,  gastric suctioning,  or use of certain diuretics;  excessive intake of alkaline drugs	Respiratory:  hypoventilation, which slows loss of CO2.  If compensation is complete,    pH will be within the normal range   but HCO3−will be high

CLINICAL CONNECTION 

Diagnosis of Acid–Base Imbalances

• One can often pinpoint the cause of an acid–base imbalance by careful evaluation of three factors in a sample of systemic arterial blood: 

1. pH,
2. concentration of HCO3⁻ and
3. PCO2.

• These three blood chemistry values are examined in the following four-step sequence:

1. Note whether the pH is high (alkalosis) or low (acidosis).

2. Then decide which value—PCO2 or HCO3− —is out of the normal range 

• and could be the cause of the pH change. 

For example, 

• elevated pH could be caused by low PCO2
•  or high HCO3−

.
3. If the cause is a change in PCO2, the problem is respiratory; 

• if the cause is a change in HCO3−, the problem is metabolic.

4. Now look at the value that doesn’t correspond with the observed pH change. 

• If it is within its normal range, there is no compensation. 
• If it is outside the normal range, compensation is occurring and partially correcting the pH imbalance.