Physiology of Body Fluids

Objectives

Upon completion of this chapter, the student should be able to answer the following questions :

•

How do body fluid compartments differ with respect to their volumes and their ionic compositions?
•

What are the driving forces responsible for movement of water across cell membranes and the capillary wall?
•

How do the volumes of the intracellular and extracellular fluid compartments change under various pathophysiologic conditions?

In addition, the student should be able to define and understand the following properties of physiologically important solutions and fluids :

•

Molarity and equivalence
•

Osmotic pressure
•

Osmolarity and osmolality
•

Oncotic pressure
•

Tonicity

Key Terms

Steady-state balance

Positive balance

Negative balance

Molarity

Equivalence

Osmosis

Osmotic pressure

van’t Hoff’s law

Osmolarity

Osmolality

Tonicity (isotonic, hypotonic, and hypertonic)

Isosmotic

Permeable

Impermeable

Effective osmole

Ineffective osmole

Reflection coefficient

Osmotic coefficient

Oncotic pressure

Specific gravity

Total body water

Intracellular fluid (ICF)

Extracellular fluid (ECF)

Interstitial fluid

Plasma

Third space

Ascites

Capillary wall

Starling forces

Capillary filtration coefficient (K _f)

Aquaporin (AQP)

One of the major functions of the kidneys is to maintain the volume and composition of the body’s fluids constant despite wide variation in the daily intake of water and solutes. In this chapter, the concept of steady-state balance is introduced. Also, the volume and composition of the body’s fluids are discussed to provide a background for the study of the kidneys as regulatory organs. Some of the basic principles, terminology, and concepts related to the properties of solutes in solution also are reviewed.

Concept of Steady-State Balance

The human body is an “open system,” meaning that substances are added to the body each day, and similarly substances are lost from the body each day. The amounts added to or lost from the body can vary widely depending on the environment, access to food and water, disease processes, and even cultural norms. In such an open system, the volume and composition of the body fluids is maintained through the process of steady-state balance .

The concept of steady-state balance can be illustrated by considering a river on which a dam is built to create a manmade lake. Each day, water enters the lake from the various streams and rivers that feed it. In addition, water is added by underground springs, rain, and snow. At the same time water is lost through the spillways of the dam and by the process of evaporation. For the level of the lake to remain constant (i.e., steady-state balance), the rate at which water is added, regardless of source, must be exactly matched by the amount of water lost, again by whichever route. Because the addition of water is not easily controlled, nor can the loss by evaporation be controlled, the only way to maintain the level of the lake constant is to regulate the amount that is lost through the spillways.

To understand steady-state balance as it applies to the human body, the following key concepts are important.

1.

There must be a set-point from which deviations can be monitored (e.g., the level of the lake in the previous example or setting the temperature in a room by adjusting the thermostat).
2.

The sensors that monitor deviations from the set-point must generate effector signals that can lead to changes in either input or output, or both, to maintain the desired set-point (e.g., electrical signals to adjust the spillway in the dam analogy or electrical signals sent to either the furnace or air conditioner to maintain the proper room temperature).
3.

Effector organs must respond in an appropriate way to the effector signals generated by the set-point monitor (e.g., the spillway gates must operate and the furnace or air conditioner must turn on and off as appropriate).
4.

The sensitivity of the system (i.e., how much of a deviation away from the set-point is tolerated) depends on several factors, including the nature of the sensor (i.e., how much of a deviation from the set-point is needed for the sensor to detect the deviation), the time necessary for generation of the effector signals, and how rapidly the effector mechanisms respond to the effector signals.

It is important to recognize that deviations from steady-state balance do occur. When input is greater than output, a state of positive balance exists. When input is less than output, a state of negative balance exists. Although transient periods of imbalance can be tolerated, prolonged states of extreme positive or negative balance are generally incompatible with life.

Physicochemical Properties of Electrolyte Solutions

Molarity and Equivalence

The amount of a substance dissolved in a solution (i.e., its concentration) is expressed in terms of either molarity or equivalence . Molarity is the amount of a substance relative to its molecular weight. For example, glucose has a molecular weight of 180 g/mol. If 1 L of water contains 1 g of glucose, the molarity of this glucose solution would be determined as:

1g/L180g/mol=0.0056mol/Lor5.6mmol/L1g/L180g/mol=0.0056mol/Lor5.6mmol/L
1g/L180g/mol=0.0056mol/Lor5.6mmol/L

For uncharged molecules, such as glucose and urea, concentrations in the body fluids are usually expressed in terms of molarity. ^aBecause many of the substances of biologic interest are present at very low concentrations, units are more commonly expressed in the millimolar range (mmol/L).

a The units used to express the concentrations of substances in various body fluids differ among laboratories. The system of international units (SI) is used in most countries and in most scientific and medical journals in the United States. Despite this convention, traditional units are still widely used. For urea and glucose, the traditional units of concentration are milligrams per deciliter (mg/dL, or 100 mL), whereas the SI units are millimole per liter (mmol/L). Similarly, electrolyte concentrations are traditionally expressed as milliequivalent per liter (mEq/L), whereas the SI units are mmol/L (see Appendix B ).

The concentration of solutes, which normally dissociate into more than one particle when dissolved in solution (e.g., sodium chloride [NaCl]), is usually expressed in terms of equivalence. Equivalence refers to the stoichiometry of the interaction between cation and anion and is determined by the valence of these ions. For example, consider a 1 L solution containing 9 g of NaCl (molecular weight = 58.4 g/mol). The molarity of this solution is 154 mmol/L. Because NaCl dissociates into Na ⁺ and Cl ⁻ ions, and assuming complete dissociation, this solution contains 154 mmol/L of Na ⁺ and 154 mmol/L of Cl ⁻ . Because the valence of these ions is 1, these concentrations also can be expressed as milliequivalents of the ion per liter (i.e., 154 mEq/L for Na ⁺ and Cl ⁻ , respectively).

For univalent ions such as Na ⁺ and Cl ⁻ , concentrations expressed in terms of molarity and equivalence are identical. However, this is not true for ions having valences greater than 1. Accordingly, the concentration of Ca ⁺⁺ (molecular weight = 40.1 g/mol and valence = 2) in a 1 L solution containing 0.1 g of this ion could be expressed as:

0.1g/L40.1g/mol=2.5mmol/L2.5mmol/L×2Eq/mol=5mEq/L0.1g/L40.1g/mol=2.5mmol/L2.5mmol/L×2Eq/mol=5mEq/L
0.1g/L40.1g/mol=2.5mmol/L2.5mmol/L×2Eq/mol=5mEq/L

Although some exceptions exist, it is customary to express concentrations of ions in milliequivalents per liter.

Osmosis and Osmotic Pressure

The movement of water across cell membranes occurs by the process of osmosis . The driving force for this movement is the osmotic pressure difference across the cell membrane. Fig. 1.1 illustrates the concept of osmosis and the measurement of the osmotic pressure of a solution.

Osmotic pressure is determined solely by the number of solute particles in the solution. It is not dependent on factors such as the size of the solute particles, their mass, or their chemical nature (e.g., valence). Osmotic pressure (π), measured in atmospheres (atm), is calculated by van’t Hoff’s law as:

?=nCRTπ=nCRT
π=nCRT

where n is the number of dissociable particles per molecule, C is total solute concentration, R is gas constant, and T is temperature in degrees Kelvin (°K).

For a molecule that does not dissociate in water, such as glucose or urea, a solution containing 1 mmol/L of these solutes at 37°C can exert an osmotic pressure of 2.54 × 10− ² atm as calculated by Eq. (1.3) using the following values: n is 1, C is 0.001 mol/L, R is 0.082 atm L/mol, and T is 310°K.

Because 1 atm equals 760 mm Hg at sea level, π for this solution also can be expressed as 19.3 mm Hg. Alternatively, osmotic pressure is expressed in terms of osmolarity (see the following discussion). Thus a solution containing 1 mmol/L of solute particles exerts an osmotic pressure of 1 milliosmole per liter (1 mOsm/L).

For substances that dissociate in a solution, n of Eq. (1.3) has a value other than 1. For example, a 150 mmol/L solution of NaCl has an osmolarity of 300 mOsm/L because each molecule of NaCl dissociates into a Na ⁺ and a Cl ⁻ ion (i.e., n = 2). If dissociation of a substance into its component ions is not complete, n is not an integer. Accordingly, osmolarity for any solution can be calculated as:

Osmolarity=Concentration×NumberofdissociableparticlesmOsm/L=mmol/L×numberofparticles/molOsmolarity=Concentration×NumberofdissociableparticlesmOsm/L=mmol/L×numberofparticles/mol
Osmolarity=Concentration×NumberofdissociableparticlesmOsm/L=mmol/L×numberofparticles/mol

Osmolarity and Osmolality

Osmolarity and osmolality are often confused and incorrectly interchanged. Osmolarity refers to the number of solute particles per 1 L of solvent, whereas osmolality is the number of solute particles in 1 kg of solvent. For dilute solutions, the difference between osmolarity and osmolality is insignificant. Measurements of osmolarity are temperature dependent, because the volume of solvent varies with temperature (i.e., the volume is larger at higher temperatures). In contrast, osmolality, which is based on the mass of the solvent, is temperature independent. For this reason, osmolality is the preferred term for biologic systems and is used throughout this and subsequent chapters. Osmolality has the units of Osm/kg H ₂ O. Because of the dilute nature of physiologic solutions and because water is the solvent, osmolalities are expressed as milliosmoles per kilogram of water (mOsm/kg H ₂ O).

Table 1.1 shows the relationships among molecular weight, equivalence, and osmoles for a number of physiologically significant solutes.

TABLE 1.1

Units of Measurement for Physiologically Significant Substances

Substance	Atomic/Molecular Weight	Equivalents/Mol	Osmoles/Mol
Na ⁺	23.0	1	1
K ⁺	39.1	1	1
Cl ⁻	35.4	1	1
HCO ⁻₃	61.0	1	1
Ca ⁺⁺	40.1	2	1
Phosphate (P _i)	95.0	3	1
NH ₄⁺	18.0	1	1
NaCl	58.4	2 ∗	2 †
CaCl ₂	111	4 ‡	3
Glucose	180	—	1
Urea	60	—	1

CaCl ₂, Calcium chloride;

HC?−3HCO3−
HC O 3 −
, bicarbonate; K ⁺, potassium;

??+4NH4+
N H 4 +
, ammonium.

∗ One equivalent each from Na ⁺and Cl ⁻.

† NaCl does not dissociate completely in solution. The actual Osm/mol volume is 1.88. However, for simplicity, a value of 2 often is used.

‡ Ca ⁺⁺contributes two equivalents, as do the Cl ⁻ions.

Tonicity

The tonicity of a solution is related to its effect on the volume of a cell. Solutions that do not change the volume of a cell are said to be isotonic . A hypotonic solution causes a cell to swell, whereas a hypertonic solution causes a cell to shrink. Although it is related to osmolality, tonicity also takes into consideration the ability of the solute to cross the cell membrane.

Consider two solutions: a 300 mmol/L solution of sucrose and a 300 mmol/L solution of urea. Both solutions have an osmolality of 300 mOsm/kg H ₂ O and therefore are said to be isosmotic (i.e., they have the same osmolality). When red blood cells (which, for the purpose of this illustration, also have an intracellular fluid osmolality of 300 mOsm/kg H ₂ O) are placed in the two solutions, those in the sucrose solution maintain their normal volume, but those placed in urea swell and eventually burst. Thus the sucrose solution is isotonic and the urea solution is hypotonic. The differential effect of these solutions on red blood cell volume is related to the permeability of the plasma membrane to sucrose and urea. The red blood cell membrane contains uniporters for urea (see Chapter 4 ). Thus urea easily crosses the cell membrane (i.e., the membrane is permeable to urea), driven by the concentration gradient (i.e., extracellular [urea] > intracellular [urea]). In contrast, the red blood cell membrane does not contain sucrose transporters and sucrose cannot enter the cell (i.e., the membrane is impermeable to sucrose).

To exert an osmotic pressure across a membrane, a solute must not cross the membrane. Because the red blood cell membrane is impermeable to sucrose, it exerts an osmotic pressure equal and opposite to the osmotic pressure generated by the contents of the red blood cell (in this case, 300 mOsm/kg H ₂ O). In contrast, urea is readily able to cross the red blood cell membrane and it cannot exert an osmotic pressure to balance that generated by the intracellular solutes of the red blood cell. Consequently, sucrose is termed an effective osmole and urea is termed an ineffective osmole .

To account for the effect of a solute’s membrane permeability on osmotic pressure, it is necessary to rewrite Eq. (1.3) as:

?=?(nCRT)π=σ(nCRT)
π=σ(nCRT)

where σ is the reflection coefficient or osmotic coefficient and is a measure of the relative ability of the solute to cross a cell membrane.

For a solute that can freely cross the cell membrane (such as urea in this example), σ = 0 and no effective osmotic pressure is exerted. Thus urea is an ineffective osmole for red blood cells. In contrast, σ = 1 for a solute that cannot cross the cell membrane (e.g., sucrose). Such a substance is said to be an effective osmole. Many solutes are neither completely able nor completely unable to cross cell membranes (i.e., 0 < σ < 1) and generate an osmotic pressure that is only a fraction of what is expected from the molecules’ concentration in solution.

Oncotic Pressure

Oncotic pressure is the osmotic pressure generated by large molecules (especially proteins) in solution. As illustrated in Fig. 1.2 , the magnitude of the osmotic pressure generated by a solution of protein does not conform to van’t Hoff’s law. The cause of this anomalous relationship between protein concentration and osmotic pressure is not completely understood but appears to be related to the size and shape of the protein molecule. For example, the correlation to van’t Hoff’s law is more precise with small, globular proteins than with larger protein molecules.

The oncotic pressure exerted by proteins in human plasma has a normal value of approximately 26 to 28 mm Hg. Although this pressure appears to be small when considered in terms of osmotic pressure (28 mm Hg ≈ 1.4 mOsm/kg H ₂ O), it is an important force involved in fluid movement across capillaries (details of this topic are presented in the following section on fluid exchange between body fluid compartments).

Specific Gravity

The total solute concentration in a solution also can be measured as specific gravity . Specific gravity is defined as the weight of a volume of solution divided by the weight of an equal volume of distilled water. Thus the specific gravity of distilled water is 1. Because biologic fluids contain a number of different substances, their specific gravities are greater than 1. For example, normal human plasma has a specific gravity in the range of 1.008 to 1.010.

In the Clinic

The specific gravity of urine is sometimes measured in clinical settings and used to assess the concentrating ability of the kidney, which can be altered by diseases that alter the ability of the posterior pituitary gland to secrete arginine vasopressin (see Chapter 5 for details). The specific gravity of urine varies in proportion to its osmolality. However, because specific gravity depends on both the number of solute particles and their weight, the relationship between specific gravity and osmolality is not always predictable. For example, patients who have been injected with radiocontrast dye (molecular weight > 500 g/mol) for radiographic studies can have high values of urine-specific gravity (1.040–1.050), even though the urine osmolality is similar to that of plasma (e.g., 300 mOsm/kg H ₂ O).