Sodium and Potassium Imbalances
Extracellular fluids usually have high sodium ion concentrations, and intracellular fluid usually has high potassium ion concentration. The renal regulation of sodium is closely related to that of potassium because active reabsorption of sodium (under the influence of aldosterone) is accompanied by secretion (and excretion) of potassium. Thus, it is not surprising that conditions that alter sodium ion balance also affect potassium ion balance.
Such disorders can be summarized as follows:
1. Low sodium concentration (hyponatremia). Possible causes of hyponatremia include prolonged sweating, vomiting, or diarrhea; renal disease in which sodium is inadequately reabsorbed; adrenal cortex disorders in which aldosterone secretion is insufficient to promote the reabsorption of sodium (Addison disease); and drinking too much water.
Possible effects of hyponatremia include the development of extracellular fluid that is hypotonic and promotes the movement of water into the cells by osmosis. This is accompanied by the symptoms of water intoxication described in Clinical Application 21.1.
2. High sodium concentration (hypernatremia). Possible causes of hypernatremia include excessive water loss by evaporation and diffusion, as may occur during high fever, or increased water loss accompanying diabetes insipidus, in one form of which ADH secretion is insufficient to maintain water conservation by the renal tubules and collecting ducts. Possible effects of hypernatremia include disturbances of the central nervous system, such as confusion, stupor, and coma.
3. Low potassium concentration (hypokalemia). Possible causes of hypokalemia include excessive release of aldosterone by the adrenal cortex (Cushing syndrome), which increases renal excretion of potassium; use of diuretic drugs that promote potassium excretion; kidney disease; and prolonged vomiting or diarrhea. Possible effects of hypokalemia include muscular weakness or paralysis, respiratory difficulty, and severe cardiac disturbances, such as atrial or ventricular arrhythmias.
4. High potassium concentration (hyperkalemia). Possible causes of hyperkalemia include renal disease, which decreases potassium excretion; use of drugs that promote renal conservation of potassium; insufficient secretion of aldosterone by the adrenal cortex (Addison disease); or a shift of potassium from the intracellular fluid to the extracellular fluid, a change that accompanies an increase in plasma hydrogen ion concentration (acidosis). Possible effects of hyperkalemia include paralysis of the skeletal muscles and severe cardiac disturbances, such as cardiac arrest. ■
Acids that ionize more completely (release more H+) are strong acids, and those that ionize less completely are weak acids. For example, the hydrochloric acid (HCl) of gastric juice is a strong acid and dissociates completely to release a lot of H+, but the carbonic acid (H2CO3) produced when carbon dioxide reacts with water is weak and dissociates less completely to release less H+.
Bases release ions, such as hydroxyl ions (OH-), which can combine with hydrogen ions, thereby lowering their concentration. Thus, sodium hydroxide (NaOH), which releases hydroxyl ions and sodium bicarbonate (NaHCO3), which releases bicarbonate ions
(HCO3-), are bases. Strong bases dissociate to release more OH- or its equivalent than do weak bases. Often, the negative ions themselves are called bases. For example, HCO3- acting as a base combines with H+ from the strong acid HCl to form the weak acid carbonic acid (H2CO3).
Either an acid shift or an alkaline (basic) shift in the body fluids could threaten the internal environment. However, normal metabolic reactions generally produce more acid than base. These reactions include cellular metabolism of
Some of the metabolic processes that provide hydrogen ions.
Some of the metabolic processes that provide hydrogen ions.
glucose, fatty acids, and amino acids. Consequently, the maintenance of acid-base balance usually entails elimination of acid. This is accomplished in three ways:
1. Acid-base buffer systems
2. Respiratory excretion of carbon dioxide
3. Renal excretion of hydrogen ions
H Explain why the regulation of hydrogen ion concentration is so important.
What are the major sources of hydrogen ions in the body?
Acid-base buffer systems occur in all of the body fluids and involve chemicals that combine with acids or bases when they are in excess. Recall that buffers are substances that stabilize the pH of a solution, despite the addition of an acid or a base. More specifically, the chemical components of a buffer system can combine with strong acids to convert them into weak acids. Likewise, these buffers can combine with strong bases to convert them into weak bases. Such activity helps minimize pH changes in the body fluids. The three most important buffer systems in body fluids are the bicarbonate buffer system, the phosphate buffer system, and the protein buffer system.
In the following discussion, associated anions and cations have been omitted for clarity. For example, the weak base sodium bicarbonate (NaHCO3) is represented by bicarbonate (HCO3-). Sodium is also the cation associated with the phosphate ions.
1. Bicarbonate buffer system. In the bicarbonate buffer system, which is present in both intracellular and extracellular fluids, the bicarbonate ion (HCO3-) acts as a weak base, and carbonic acid (H2CO3) acts as a weak acid. In the presence of excess hydrogen ions, bicarbonate ions combine with hydrogen ions to form carbonic acid, thus minimizing any increase in the hydrogen ion concentration of body fluids:
On the other hand, if conditions are basic or alkaline, carbonic acid dissociates to release bicarbonate ion and hydrogen ion:
It is important to remember that, even though this reaction releases bicarbonate ion, it is the increase of free hydrogen ions at equilibrium that is important in minimizing the shift toward a more alkaline pH.
2. Phosphate buffer system. The phosphate buffer system is also present in both intracellular and extracellular fluids. However it is particularly important in the control of hydrogen ion concentration in the intracellular fluid and in renal tubular fluid and urine. This buffer system consists of two phosphate ions, dihydrogen phosphate (H2PO4-) and monohydrogen phosphate (HPO42-).
In the presence of excess hydrogen ions, monohydrogen phosphate ions act as a weak base, combining with hydrogen ions to form dihydrogen phosphate, thus minimizing any increase in the hydrogen ion concentration of the body fluids.
On the other hand, if conditions are basic or alkaline, dihydrogen phosphate, acting as a weak acid, dissociates to release hydrogen ion:
3. Protein buffer system. The protein acid-base buffer system consists of the plasma proteins, such as albumins, and certain proteins within the cells, including the hemoglobin of red blood cells.
As described in chapter 2 (p. 54), proteins are chains of amino acids. Some of these amino acids have freely exposed groups of atoms, called carboxyl groups. If the H+ concentration drops, a carboxyl group (—COOH) can become ionized, releasing a hydrogen ion, thus resisting the pH change:
Notice that this is a reversible reaction. In the presence of excess hydrogen ions, the —COO-portions of the protein molecules accept hydrogen ions and become —COOH groups again. This action decreases the number of free hydrogen ions in the body fluids and again minimizes the pH change.
Some of the amino acids within a protein molecule also contain freely exposed amino groups (—NH2). If the H+ concentration rises, these amino groups can accept hydrogen ions in another reversible reaction:
In the presence of excess hydroxyl ions (OH-), the —NH3+ groups within protein molecules give up hydrogen ions and become —NH2 groups again. These hydrogen ions then combine with hydroxyl ions to form water molecules. Once again, pH change is minimized. Thus, protein molecules can function as acids by releasing hydrogen ions under alkaline conditions or as bases by accepting hydrogen ions under acid conditions. This special property allows protein molecules to operate as an acid-base buffer system.
Hemoglobin is an especially important protein that buffers hydrogen ions. As explained in chapter 19 (p. 811), carbon dioxide, produced by cellular oxidation of glucose, diffuses through the capillary wall and enters the plasma and then the red blood cells. The red cells contain an enzyme, carbonic anhydrase, that speeds the reaction between carbon dioxide and water, producing carbonic acid:
The carbonic acid quickly dissociates, releasing hydrogen ions and bicarbonate ions:
In the peripheral tissues, where CO2 is generated, oxygen is used in the metabolism of glucose. As a result, hemoglobin has given up much of its oxygen and is in the form of deoxyhemoglobin. In this form, hemoglobin can bind hydrogen ions, thus acting as a buffer to minimize the pH change that would otherwise occur.
The above two reactions can be written as a single reversible reaction:
Thus, in the peripheral tissues, where CO2 levels are high, the reaction equilibrium shifts to the right, generating H+, which is buffered by hemoglobin, and HCO3-, which becomes a plasma electrolyte. In the lungs, where oxygen levels are high, hemoglobin is no longer a good buffer and it releases its H+. However, the released H+ combines with plasma HCO3, shifting the reaction equilibrium to the left, generating carbonic acid, which quickly dissociates to form CO2 and water. The water is added to the body fluids, and the CO2 is exhaled. Because of this relationship to CO2, carbonic acid is sometimes called a volatile acid.
Individual amino acids in body fluids can also function as acid-base buffers by accepting or releasing hydrogen ions. This is possible because every amino acid has an amino group (—NH2) and a carboxyl group (—COOH).
To summarize, acid-base buffer systems take up hydrogen ions when body fluids are becoming more acidic and give up hydrogen ions when the fluids are becoming more basic (alkaline). Buffer systems convert stronger acids into weaker acids or convert stronger bases into weaker bases, as table 21.3 summarizes.
In addition to minimizing pH fluctuations, acid-base buffer systems in body fluids buffer each other. Consequently, whenever the hydrogen ion concentration begins to change, the chemical balances within all of the buffer systems change too, resisting the drift in pH.
Converts a strong
acid into a
Was this article helpful?