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24.3 Acid–Base Balance
    Expected Learning Outcomes

    When you have completed this section, you should be able to

    1. define buffer and write chemical equations for the bicarbonate, phosphate, and protein buffer systems;

    2. discuss the relationship between pulmonary ventilation, pH of the extracellular fluids, and the bicarbonate buffer system;

    3. explain how the kidneys secrete hydrogen ions and how these ions are buffered in the tubular fluid;

    4. identify some types and causes of acidosis and alkalosis, and describe the effects of these pH imbalances; and

    5. explain how the respiratory and urinary systems correct acidosis and alkalosis, and compare their effectiveness and limitations.

As we saw in chapter 2, metabolism depends on the functioning of enzymes, and enzymes are very sensitive to pH. Slight deviations from the normal pH can shut down metabolic pathways as well as alter the structure and function of other macromolecules. Consequently, acid–base balance is one of the most important aspects of homeostasis.

The blood and tissue fluid normally have a pH of 7.35 to 7.45. Such a narrow range of variation is remarkable considering that our metabolism constantly produces acid: lactic acid from anaerobic fermentation, phosphoric acids from nucleic acid catabolism, fatty acids and ketones from fat catabolism, and carbonic acid from carbon dioxide. Here we examine mechanisms for resisting these challenges and maintaining acid–base balance.

    Apply What You Know

In the systemic circulation, arterial blood has a mean pH of 7.40 and venous blood has a mean of 7.35. What do you think causes this difference?

Acids, Bases, and Buffers

The pH of a solution is determined solely by its hydrogen ions (H+). An acid is any chemical that releases H+ in solution. A strong acid such as hydrochloric acid (HCl) ionizes freely, gives up most of its hydrogen ions, and can markedly lower the pH of a solution. A weak acid such as carbonic acid (H2CO3) ionizes only slightly and keeps most hydrogen in a chemically bound form that does not affect pH. A base is any chemical that accepts H+. A strong base such as the hydroxide ion (OH-) has a strong tendency to bind H+ and raise the pH, whereas a weak base such as the bicarbonate ion (HCO3-) binds less of the available H+ and has less effect on pH.

A buffer, broadly speaking, is any mechanism that resists changes in pH by converting a strong acid or base to a weak one. The body has both physiological and chemical buffers. A physiological buffer is a system—namely, the respiratory or urinary system—that stabilizes pH by controlling the body's output of acids, bases, or CO2. Of all buffer systems, the urinary system buffers the greatest quantity of acid or base, but it requires several hours to days to exert an effect. The respiratory system exerts an effect within a few minutes but cannot alter the pH as much as the urinary system can.

A chemical buffer is a substance that binds H+ and removes it from solution as its concentration begins to rise, or releases H+ into solution as its concentration falls. Chemical buffers can restore normal pH within a fraction of a second. They function as mixtures called buffer systems composed of a weak acid and a weak base. The three major chemical buffer systems of the body are the bicarbonate, phosphate, and protein systems.

The amount of acid or base that can be neutralized by a chemical buffer system depends on two factors: the concentration of the buffers and the pH of their working environment. Each system has an optimum pH at which it functions best; its effectiveness is greatly reduced if the pH of its environment deviates too far from this. The relevance of these factors will become apparent as you study the following buffer systems.

The Bicarbonate Buffer System

The Bicarbonate Buffer System   The bicarbonate buffer system is a solution of carbonic acid and bicarbonate ions. Carbonic acid (H2CO3) forms by the hydration of carbon dioxide and then dissociates into bicarbonate (HCO3-) and H+:

This is a reversible reaction. When it proceeds to the right, carbonic acid acts as a weak acid by releasing H+ and lowering pH. When the reaction proceeds to the left, bicarbonate acts as a weak base by binding H+, removing the ions from solution, and raising pH.

At a pH of 7.4, the bicarbonate system would not ordinarily have a particularly strong buffering capacity outside of the body. This is too far from its optimum pH of 6.1. If a strong acid was added to a beaker of carbonic acid–bicarbonate solution at pH 7.4, the preceding reaction would shift only slightly to the left. Much surplus H+ would remain and the pH would be substantially lower. In the body, by contrast, the bicarbonate system works quite well because the lungs and kidneys constantly remove CO2 and prevent an equilibrium from being reached. This keeps the reaction moving to the left, and more H+ is neutralized. Conversely, if there is a need to lower the pH, the kidneys excrete HCO3-, keep this reaction moving to the right, and elevate the H+ concentration of the ECF. Thus, you can see that the physiological and chemical buffers of the body function together in maintaining acid–base balance.

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The Phosphate Buffer System

The Phosphate Buffer System   The phosphate buffer system is a solution of HPO42- and H2PO4-. It works in much the same way as the bicarbonate system. The following reaction can proceed to the right to liberate H+ and lower pH, or it can proceed to the left to bind H+ and raise pH:

The optimal pH for this system is 6.8, closer to the actual pH of the ECF (7.4). Thus, the phosphate buffer system has a stronger buffering effect than an equal amount of bicarbonate buffer. However, phosphates are much less concentrated in the ECF than bicarbonate, so they are less important in buffering the ECF. They are more important in the renal tubules and ICF, where not only are they more concentrated, but the pH is lower and closer to their functional optimum. In the ICF, the constant production of metabolic acids creates pH values ranging from 4.5 to 7.4, probably averaging 7.0. The reason for the low pH in the renal tubules is discussed later.

The Protein Buffer System

The Protein Buffer System   Proteins are more concentrated than either bicarbonate or phosphate buffers, especially in the ICF. The protein buffer system accounts for about three-quarters of all chemical buffering in the body fluids. The buffering ability of proteins is due to certain side groups of their amino acid residues. Some have carboxyl (-COOH) side groups, which release H+ when pH begins to rise and thus lower pH:

Others have amino (-NH2) side groups, which bind H+ when pH falls too low, thus raising pH toward normal:

    Apply What You Know

What protein do you think is the most important buffer in blood plasma? In erythrocytes?

Respiratory Control of pH

The equation for the bicarbonate buffer system shows that the addition of CO2 to the body fluids raises H+ concentration and lowers pH, while the removal of CO2 has the opposite effects. This is the basis for the strong buffering capacity of the respiratory system. Indeed, this system can neutralize two or three times as much acid as the chemical buffers can.

Carbon dioxide is constantly produced by aerobic metabolism and is normally eliminated by the lungs at an equivalent rate. As explained in chapter 22, rising CO2 concentration and falling pH stimulate peripheral and central chemoreceptors, which stimulate an increase in pulmonary ventilation. This expels excess CO2 and thus reduces H+ concentration. The free H+ becomes part of the water molecules produced by this reaction:

Conversely, a drop in H+ concentration raises pH and reduces pulmonary ventilation. This allows metabolic CO2 to accumulate in the ECF faster than it is expelled, thus lowering pH to normal.

These are classic negative feedback mechanisms that result in acid–base homeostasis. Respiratory control of pH has some limitations, however, that are discussed later under acid–base imbalances.

Renal Control of pH

The kidneys can neutralize more acid or base than either the respiratory system or the chemical buffers. The essence of this mechanism is that the renal tubules secrete H+ into the tubular fluid, where most of it binds to bicarbonate, ammonia, and phosphate buffers. Bound and free H+ are then excreted in the urine. Thus the kidneys, in contrast to the lungs, actually expel H+ from the body. The other buffer systems only reduce its concentration by binding it to another chemical.

Figure 24.10 shows how the renal tubule secretes and neutralizes H+. The hydrogen ions are colored so you can trace them from the blood (step 1) to the tubular fluid (step 6). Notice that it is not a simple matter of transporting free H+ across the tubule cells; rather, the H+ travels in the form of carbonic acid and water molecules.

FIGURE 24.10
Secretion and Neutralization of Hydrogen Ions in the Kidneys.The colored hydrogen symbols allow you to trace hydrogen from H+ in the blood to H2O in the urine.

  If the pH of the tubular fluid went down, how would its Na+ concentration change?

The tubular secretion of H+ takes place at step 6, where the ion is pumped out of the tubule cell into the tubular fluid. This can happen only if there is a steep enough concentration gradient between a high H+ concentration within the cell and a lower concentration in the tubular fluid. If the pH of the tubular fluid drops any lower than 4.5, H+ concentration in the fluid is so high that tubular secretion ceases. Thus, pH 4.5 is the limiting pH for tubular secretion. This has added significance later in our discussion.

In a person with normal acid–base balance, all bicarbonate ions (HCO3-) in the tubular fluid are consumed by neutralizing H+; thus there is no HCO3- in the urine. Bicarbonate ions are filtered by the glomerulus, gradually disappear from the tubular fluid, and appear in the peritubular capillary blood. It appears as if HCO3- is reabsorbed by the renal tubules, but this is not the case; indeed, the renal tubules are incapable of reabsorbing HCO3- directly. The cells of the proximal convoluted tubule, however, have carbonic anhydrase (CAH) on their brush borders facing the lumen. This breaks down the H2CO3 in the tubular fluid to CO2 + H2O (step 10). It is the CO2 that is reabsorbed, not the bicarbonate. For every CO2 reabsorbed, however, a new bicarbonate ion is formed in the tubule cell and released into the blood (step 5). The effect is the same as if the tubule cells had reabsorbed bicarbonate itself.

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Note that for every bicarbonate ion that enters the peritubular capillaries, a sodium ion does too. Thus, Na+ reabsorption by the renal tubules is part of the process of neutralizing acid. The more acid the kidneys excrete, the less sodium the urine contains.

The tubules secrete somewhat more H+ than the available bicarbonate can neutralize. The urine therefore contains a slight excess of free H+, which gives it a pH of about 5 to 6. Yet if all of the excess H+ secreted by the tubules remained in this free ionic form, the pH of the tubular fluid would drop far below the limiting pH of 4.5, and H+ secretion would stop. This must be prevented, and there are additional buffers in the tubular fluid to do so.

The glomerular filtrate contains Na2HPO4 (dibasic sodium phosphate), which reacts with some of the H+ (fig. 24.11). A hydrogen ion replaces one of the sodium ions in the buffer, forming NaH2PO4 (monobasic sodium phosphate). This is passed in the urine, and the displaced Na+ is transported into the tubule cell and from there to the bloodstream.

FIGURE 24.11
Acid Buffering in the Urine.Reactions in the tubule cells are the same as in figure 24.10 but are simplified in this diagram. The essential differences are the buffering mechanisms shown in the tubular fluid. Hydrogen symbols are colored to allow tracing them from carbonic acid to the urine.

In addition, tubule cells catabolize certain amino acids and release ammonia (NH3) as a product (fig. 24.11). Ammonia diffuses into the tubular fluid, where it acts as a base to neutralize acid. It reacts with H+ and Cl- (the most abundant anion in the glomerular filtrate) to form ammonium chloride (NH4Cl), which is passed in the urine.

Since there is so much chloride in the tubular fluid, you might ask why H+ is not simply excreted as hydrochloric acid (HCl). Why involve ammonia? The reason is that HCl is a strong acid—it dissociates almost completely, so most of its hydrogen would be in the form of free H+. The pH of the tubular fluid would drop below the limiting pH and prevent excretion of more acid. Ammonium chloride, by contrast, is a weak acid—most of its hydrogen remains bound to it and does not lower the pH of the tubular fluid.

Disorders of Acid–Base Balance

Figure 24.12 represents acid–base balance with an instructive metaphor to show its dependence on the bicarbonate buffer system. At a normal pH of 7.4, the ECF has a 20:1 ratio of HCO3- to H2CO3. Excess hydrogen ions convert HCO3- to H2CO3 and tip the balance to a lower pH. A pH below 7.35 is considered to be a state of acidosis. On the other hand, a H+ deficiency causes H2CO3 to dissociate into H+ and HCO3-, thus tipping the balance to a higher pH. A pH above 7.45 is a state of alkalosis. Either of these imbalances has potentially fatal effects. A person cannot live more than a few hours if the blood pH is below 7.0 or above 7.7, and a pH below 6.8 or above 8.0 is quickly fatal.

FIGURE 24.12
The Relationship of Bicarbonate–Carbonic Acid Ratio to pH.At a normal pH of 7.40, there is a 20:1 ratio of bicarbonate ions (HCO3-) to carbonic acid (H2CO3) in the blood plasma. An excess of HCO3- tips the balance toward alkalosis, whereas an excess of H2CO3 tips it toward acidosis.
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In acidosis, H+ diffuses down its concentration gradient into cells, and to maintain electrical balance, K+ diffuses out (fig. 24.13a). The H+ is buffered by intracellular proteins, so this exchange results in a net loss of cations from the cell. This makes the resting membrane potential more negative than usual (hyperpolarized) and makes nerve and muscle cells more difficult to stimulate. This is why acidosis depresses the central nervous system and causes such symptoms as confusion, disorientation, and coma.

FIGURE 24.13
The Relationship Between Acid–Base Imbalances and Potassium Imbalances.(a) In acidosis, H+ diffuses into the cells and drives out K+, elevating the K+ concentration of the ECF. (b) In alkalosis, H+ diffuses out of the cells and K+ diffuses in to replace it, lowering the K+ concentration of the ECF.

  How would you change part (a) to show the effect of hyperkalemia on the pH of the ECF?

In alkalosis, the extracellular H+ concentration is low. Hydrogen ions diffuse out of the cells and K+ diffuses in to replace them (fig. 24.13b). The net gain in positive intracellular charges shifts the membrane potential closer to firing level and makes the nervous system hyperexcitable. Neurons fire spontaneously and overstimulate skeletal muscles, causing muscle spasms, tetanus, convulsions, or respiratory paralysis.

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Acid–base imbalances fall into two categories, respiratory and metabolic (table 24.2). Respiratory acidosis occurs when the rate of alveolar ventilation fails to keep pace with the body's rate of CO2 production. Carbon dioxide accumulates in the ECF and lowers its pH. This occurs in such conditions as emphysema, in which there is a severe reduction in the number of functional alveoli. Respiratory alkalosis results from hyperventilation, in which CO2 is eliminated faster than it is produced.

TABLE 24.2
Some Causes of Acidosis and Alkalosis




Hypoventilation, apnea, or respiratory arrest; asthma; emphysema; cystic fibrosis; chronic bronchitis; narcotic overdose

Hyperventilation due to pain or emotions such as anxiety; oxygen deficiency (as at high elevation)


Excess production of organic acids as in diabetes mellitus and starvation; long-term anaerobic fermentation; hyperkalemia; chronic diarrhea; excessive alcohol consumption; drugs such as aspirin and laxatives

Rare but can result from chronic vomiting; overuse of bicarbonates (antacids); aldosterone hypersecretion

Metabolic acidosis can result from increased production of organic acids, such as lactic acid in anaerobic fermentation and ketone bodies in alcoholism and diabetes mellitus. It can also result from the excessive ingestion of acidic drugs such as aspirin or from the loss of base due to chronic diarrhea or overuse of laxatives. Dying persons also typically exhibit acidosis. Metabolic alkalosis is rare but can result from overuse of bicarbonates (such as oral antacids and intravenous bicarbonate solutions) or from the loss of stomach acid by chronic vomiting.

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Compensation for Acid–Base Imbalances

In compensated acidosis or alkalosis, either the kidneys compensate for pH imbalances of respiratory origin, or the respiratory system compensates for pH imbalances of metabolic origin. Uncompensated acidosis or alkalosis is a pH imbalance that the body cannot correct without clinical intervention.

In respiratory compensation, changes in pulmonary ventilation correct the pH of the body fluids by expelling or retaining CO2. If there is a CO2 excess (hypercapnia), pulmonary ventilation increases to expel CO2 and bring the blood pH back up to normal. If there is a CO2 deficiency (hypocapnia), ventilation is reduced to allow CO2 to accumulate in the blood and lower the pH to normal.

This is very effective in correcting pH imbalances due to abnormal Pco2 but not very effective in correcting other causes of acidosis and alkalosis. In diabetic acidosis, for example, the lungs cannot reduce the concentration of ketone bodies in the blood, although one can somewhat compensate for the H+ that ketones release by increasing pulmonary ventilation and exhausting extra CO2. The respiratory system can adjust a blood pH of 7.0 back to 7.2 or 7.3 but not all the way back to the normal 7.4. Although the respiratory system has a very powerful buffering effect, its ability to stabilize pH is therefore limited.

Renal compensation is an adjustment of pH by changing the rate of H+ secretion by the renal tubules. The kidneys are slower to respond to pH imbalances but better at restoring a fully normal pH. Urine usually has a pH of 5 to 6, but in acidosis it may fall as low as 4.5 because of excess H+, whereas in alkalosis it may rise as high as 8.2 because of excess HCO3-. The kidneys cannot act quickly enough to compensate for short-term pH imbalances, such as the acidosis that might result from an asthmatic attack lasting an hour or two, or the alkalosis resulting from a brief episode of emotional hyperventilation. They are effective, however, at compensating for pH imbalances that last for a few days or longer.

In acidosis, the renal tubules increase the rate of H+ secretion. The extra H+ in the tubular fluid must be buffered; otherwise, the fluid pH could exceed the limiting pH and H+ secretion would stop. Therefore, in acidosis, the renal tubules secrete more ammonia to buffer the added H+, and the amount of ammonium chloride in the urine may rise to 7 to 10 times normal.

    Apply What You Know

Suppose you measured the pH and ammonium chloride concentration of urine from a person with emphysema and urine from a healthy individual. How would you expect the two to differ, and why?

In alkalosis, the bicarbonate concentration and pH of the urine are elevated. This is partly because there is more HCO3- in the blood and glomerular filtrate and partly because there is not enough H+ in the tubular fluid to neutralize all the HCO3- in the filtrate.

Acid–Base Imbalances in Relation to Electrolyte and Water Imbalances

The foregoing discussion once again stresses a point made early in this chapter—we cannot understand or treat imbalances of water, electrolyte, or acid–base balance in isolation from each other, because each of these frequently affects the other two. Table 24.3 itemizes and explains a few of these interactions. This is by no means a complete list of how fluid, electrolytes, and pH affect each other, but it does demonstrate their interdependence. Note that many of these relationships are reciprocal—for example, acidosis can cause hyperkalemia, and conversely, hyperkalemia can cause acidosis.

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TABLE 24.3
Some Relationships Among Fluid, Electrolyte, and Acid–Base Imbalances


Potential Effect




H+ diffuses into cells and displaces K+ (see fig. 24.13a). As K+ leaves the ICF, K+ concentration in the ECF rises.



Opposite from the above; high K + concentration in the ECF causes less K + to diffuse out of the cells than normally. H+ diffuses out to compensate, and this lowers the extracellular pH.



H+ diffuses from ICF to ECF. More K+ remains in the ICF to compensate for the H+ loss, causing a drop in ECF K+ concentration (see fig. 24.13b).



Opposite from the above; low K + concentration in the ECF causes K+ to diffuse out of cells. H+ diffuses in to replace K+, lowering the H+ concentration of the ECF and raising its pH.



More Cl- is excreted as NH4Cl to buffer the excess acid in the renal tubules, leaving less Cl- in the ECF.



More Cl- is reabsorbed from the renal tubules, so ingested Cl- accumulates in the ECF rather than being excreted.



More H+ is retained in the blood to balance the excess Cl-, causing hyperchloremic acidosis.



More Na+ is reabsorbed by the kidney. Na+ reabsorption is coupled to H+ secretion (see fig. 24.10), so more H+ is secreted and pH of the ECF rises.



Less Na + is reabsorbed, so less H + is secreted into the renal tubules. H+ retained in the ECF causes acidosis.



Acidosis causes more Ca2+ to bind to plasma protein and citrate ions, lowering the concentration of free, ionized Ca2+ and causing symptoms of hypocalcemia.



Alkalosis causes more Ca2+ to dissociate from plasma protein and citrate ions, raising the concentration of free Ca2+.

Before You Go On

Answer the following questions to test your understanding of the preceding section:

  1. Write two chemical equations that show how the bicarbonate buffer system compensates for acidosis and alkalosis and two equations that show how the phosphate buffer system compensates for these imbalances.

  2. Why are phosphate buffers more effective in the cytoplasm than in the blood plasma?

  3. Renal tubules cannot reabsorb HCO3-; yet HCO3- concentration in the tubular fluid falls while in the blood plasma it rises. Explain this apparent contradiction.

  4. In acidosis, the renal tubules secrete more ammonia. Why?

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Clinical Application
Fluid Replacement Therapy

One of the most significant problems in the treatment of seriously ill patients is the restoration and maintenance of proper fluid volume, composition, and distribution among the fluid compartments. Fluids may be administered to replenish total body water, restore blood volume and pressure, shift water from one fluid compartment to another, or restore and maintain electrolyte and acid–base balance.

Drinking water is the simplest method of fluid replacement, but it does not replace electrolytes. Heat exhaustion can occur when you lose water and salt in the sweat and replace the fluid by drinking plain water. Broths, juices, and sports drinks replace water, carbohydrates, and electrolytes.

Patients who cannot take fluids by mouth must be treated by alternative routes. Some fluids can be given by enema and absorbed through the colon. All routes of fluid administration other than the digestive tract are called parenteral6 routes. The most common of these is the intravenous (I.V.) route, but for various reasons, including inability to find a suitable vein, fluids are sometimes given by subcutaneous (sub-Q), intramuscular (I.M.), or other parenteral routes. Many kinds of sterile solutions are available to meet the fluid replacement needs of different patients.

In cases of extensive blood loss, there may not be time to type and cross-match blood for a transfusion. The more urgent need is to replenish blood volume and pressure. Normal saline (isotonic, 0.9% NaCl) is a relatively quick and simple way to raise blood volume while maintaining normal osmolarity, but it has significant shortcomings. It takes three to five times as much saline as whole blood to rebuild normal volume because much of the saline escapes the circulation into the interstitial fluid compartment or is excreted by the kidneys. In addition, normal saline can induce hypernatremia and hyperchloremia, because the body excretes the water but retains much of the NaCl. Hyperchloremia can, in turn, produce acidosis. Normal saline also lacks potassium, magnesium, and calcium. Indeed, it dilutes those electrolytes that are already present and creates a risk of cardiac arrest from hypocalcemia. Saline also dilutes plasma albumin and RBCs, creating still greater risks for patients who have suffered extensive blood loss. Nevertheless, the emergency maintenance of blood volume sometimes takes temporary precedence over these other considerations.

Fluid therapy is also used to correct pH imbalances. Acidosis may be treated with Ringer's lactate solution, which includes sodium to rebuild ECF volume, potassium to rebuild ICF volume, lactate to balance the cations, and enough glucose to make the solution isotonic. Alkalosis can be treated with potassium chloride. This must be administered very carefully, because potassium ions can cause painful venous spasms, and even a small potassium excess can cause cardiac arrest. High-potassium solutions should never be given to patients in renal failure or whose renal status is unknown, because in the absence of renal excretion of potassium, they can bring on lethal hyperkalemia. Ringer's lactate or potassium chloride also must be administered very cautiously, with close monitoring of blood pH, to avoid causing a pH imbalance opposite the one that was meant to be corrected. Too much Ringer's lactate causes alkalosis and too much KCl causes acidosis.

Plasma volume expanders are hypertonic solutions or colloids that are retained in the bloodstream and draw interstitial water into it by osmosis. They include albumin, sucrose, mannitol, and dextran. Plasma expanders are also used to combat hypotonic hydration by drawing water out of swollen cells, averting such problems as seizures and coma. A plasma expander can draw several liters of water out of the intracellular compartment within a few minutes.

Patients who cannot eat are often given isotonic 5% dextrose (glucose). A fasting patient loses as much as 70 to 85 g of protein per day from the tissues as protein is broken down to fuel the metabolism. Giving 100 to 150 g of I.V. glucose per day reduces this by half and is said to have a protein-sparing effect. More than glucose is needed in some cases—for example, if a patient has not eaten for several days and cannot be fed by nasogastric tube (due to lesions of the digestive tract, for example) or if large amounts of nutrients are needed for tissue repair following severe trauma, burns, or infections. In total parenteral nutrition (TPN), or hyperalimentation,7 a patient is provided with complete I.V. nutritional support, including a protein hydrolysate (amino acid mixture), vitamins, electrolytes, 20% to 25% glucose, and on alternate days, a fat emulsion.

The water from parenteral solutions is normally excreted by the kidneys. If the patient has renal insufficiency, however, excretion may not keep pace with intake, and there is a risk of hypotonic hydration. Intravenous fluids are usually given slowly, by I.V. drip, to avoid abrupt changes or overcompensation for the patient's condition. In addition to pH, the patient's heart rate, blood pressure, hematocrit, and plasma electrolyte concentrations are monitored, and the patient is examined periodically for respiratory sounds indicating pulmonary edema.

The delicacy of fluid replacement therapy underscores the close relationships among fluids, electrolytes, and pH. It is dangerous to manipulate any one of these variables without close attention to the others. Parenteral fluid therapy is usually used for persons who are seriously ill. Their homeostatic mechanisms are already compromised and leave less room for error than in a healthy person.