Concentration of Solutions

The term concentration refers to the amount of solute dissolved into the solvent. Concentration can be described either qualitatively or quantitatively. Calling something a dilute solution is an example of a qualitative description. Stating that a specific container holds 50 ml of 0.4 molar solution of sodium hydroxide (NaOH) is a quantitative description (Figure 12-1, A). Saturated solutions occur when the solvent has dissociated the maximal amount of solute into itself. Additional solute added to a saturated solution does not dissociate into solution but remains at the bottom of the container (see Figure 12-1, B). Solute particles precipitate into the solid state at the same rate at which other particles go into solution. This equilibrium characterizes a saturated solution.

A solution is characterized as being supersaturated when the solvent contains more solute than a saturated solution at the same temperature and pressure. If a saturated solution is heated, the solute equilibrium is upset, and more solute goes into solution. If undissolved solute is removed and the solution is cooled gently, there is an excess of dissolved solute (see Figure 12-1, C). The excess solute of supersaturated solutions may be precipitated out if the solution is disturbed or if a “seed crystal” is introduced.

Starling Forces

Starling was a nineteenth-century British physiologist who studied fluid transport across membranes. His hypothesis states that the fluid movement secondary to filtration across the wall of a capillary depends on both the hydrostatic and the oncotic pressure gradients across the capillary.² The driving force for fluid filtration across the wall of the capillary is determined by four separate pressures: hydraulic (hydrostatic) and colloid osmotic pressure both within the vessel and in the tissue space.³ This process can be described mathematically using the following equation:

< ?xml:namespace prefix = "mml" />Jv=Lp [Pc−Pi−s (pc−pi)]

Where:

Osmotic Pressure of Solutions

Most of the solutions of physiologic importance in the body are dilute. Solutes in dilute solution show many of the properties of gases. This behavior results from the relatively large distances between the molecules in dilute solutions. The most important physiologic characteristic of solutions is their ability to exert pressure.

Osmotic pressure (oncotic pressure)⁴ is the force produced by solvent particles under certain conditions. A membrane that permits passage of solvent molecules but not solute is called a semipermeable membrane. If such a membrane divides a solution into two compartments, molecules of solvent can pass through it from one side to the other (Figure 12-2, A). The number of solvent molecules passing (or diffusing) in one direction must equal the number of solute molecules passing in the opposite direction. An equal ratio of solute to solvent particles (i.e., the concentration of the solution) is maintained on both sides of the membrane. A capillary wall is an example of a semipermeable membrane.^5,6

If a solution is placed on one side of a semipermeable membrane and pure solvent is placed on the other, solvent molecules move through the membrane into the solution. The force driving solvent molecules through the membrane is called osmotic pressure. Osmotic pressure tries to redistribute solvent molecules so that the same concentration exists on both sides of the membrane. Osmotic pressure may be measured by connecting a manometer to the expanding column of the solution (see Figure 12-2, B and C).

Osmotic pressure can also be visualized as an attractive force of solute particles in a concentrated solution. If 100 ml of a 50% solution is placed on one side of a membrane and 100 ml of a 30% solution is placed on the other side, solvent molecules move from the dilute to the concentrated side (see Figure 12-2, D and E). The particles in the concentrated solution attract solvent molecules from the dilute solution until equilibrium occurs. Equilibrium exists when the concentrations (i.e., ratio of solute to solvent) in the two compartments are equal (40% in Figure 12-2).

Osmolality is defined as the ratio of solute to solvent. In physiology, the solvent is water.1,5,7 Osmotic pressure depends on the number of particles in solution but not on their charge or identity. A 2% solution has twice the osmotic pressure of a 1% solution under similar pressures. For a given amount of solute, osmotic pressure is inversely proportional to the volume of solvent. Most cell walls are semipermeable membranes. Through osmotic pressure, water is distributed throughout the body within certain physiologic ranges. Tonicity describes how much osmotic pressure is exerted by a solution. Average body cellular fluid has a tonicity equal to a 0.9% solution of sodium chloride (NaCl; sometimes referred to as physiologic saline). Solutions with similar tonicity are called isotonic. Solutions with more tonicity are hypertonic, and solutions with less tonicity are hypotonic. Most cells reside in a hypotonic environment in which the concentration of water (solute) is lower inside the cell than in the surroundings. Water flows into the cell causing it to expand until the cell membrane restricts further expansion. Pressure increases inside the cell to counteract osmotic pressure. This pressure is called turgor, and it is what prevents more water from entering the cell. The equilibrium that develops allows the cell to maintain a gradient across the cell membrane. Some cells have selective permeability, allowing passage not only of water but also of specific solutes. Through these mechanisms, nutrients and physiologic solutions are distributed throughout the body.

In electrochemical terms, there are three basic types of physiologic solutions. Depending on the solute, solutions are ionic (electrovalent), polar covalent, or nonpolar covalent (Table 12-1). In ionic and polar covalent solutions, some of the solute ionizes into separate particles known as ions. A solution in which this dissociation occurs is called an electrolyte solution (Figure 12-3). If an electrode is placed in such a solution, positive ions migrate to the negative pole of the electrode. These ions are called cations. Negative ions migrate to the positive pole of the electrode; they are called anions. In nonpolar covalent solutions, molecules of solute remain intact and do not carry electrical charges; these solutions are referred to as nonelectrolytes. These nonelectrolytes are not attracted to either the positive or the negative pole of an electrode (hence the designation nonpolar). All three types of solutions coexist in the body. These solutions also serve as the media in which colloids and simple suspensions are dispersed. Gases such as O₂ and CO₂ are nonpolar molecules (along with N₂) and do not dissolve very well in water, which is a polar solvent.

Quantifying Solute Content and Activity

The amount of solute in a solution can be quantified in two ways: (1) by actual weight (grams or milligrams) and (2) by chemical combining power. The weight of a solute is easy to measure and specify. However, it does not indicate chemical combining power. The sodium ion (Na⁺) has a gram ionic weight of 23. The bicarbonate ion (HCO₃⁻) has a gram ionic weight of 61. Because the gram atomic weight of every substance has 6.023 × 10²³ particles, these ions have the same chemical combining power in solution. The number of chemically reactive units is usually more meaningful than their weight.

Equivalent Weights

In medicine, it is customary to refer to physiologic substances in terms of chemical combining power. The measure commonly used is equivalent weight. Equivalent weights are amounts of substances that have equal chemical combining power. For example, if chemical A reacts with chemical B, by definition, 1 equivalent weight of A reacts with exactly 1 equivalent weight of B. No excess reactants of A or B remain.

Two magnitudes of equivalent weights are used to calculate chemical combining power: gram equivalent weight (gEq) and milligram equivalent weight, or milliequivalent (mEq). One milliequivalent (1 mEq) is of 1 gEq.

Gram Equivalent Weight Values

A gEq of a substance is calculated as its gram molecular (formula) weight divided by its valence. Valence refers to the number of electrons that need to be added or removed to make the substance electrically neutral. The valence signs (+ or −) are disregarded.

gEq=Gram molecular weightValence

The gEq of sodium (Na⁺), with a valence of 1, equals its gram atomic weight of 23 g. The gEq of calcium (Ca⁺⁺) is its atomic weight (i.e., 40) divided by 2, or 20 g. The gEq of ferric iron (Fe⁺⁺⁺) is its atomic weight (i.e., 55.8) divided by 3, or approximately 18.6 g.

For radicals such as sulfate (SO₄²⁻), the formula for sulfuric acid (H₂SO₄) shows that one sulfate group combines with two atoms of hydrogen. Half (0.5) of a mole of sulfate is equivalent to 1 mole of hydrogen atoms. The gEq of SO₄²⁻ is half its gram formula weight, or 48 g. If an element has more than one valence, the valence must be specified or must be apparent from the observed chemical combining properties.

Gram Equivalent Weight of an Acid

The gEq of an acid is the weight of the acid (in grams) that contains 1 mole of replaceable hydrogen. The gEq of an acid may be calculated by dividing its gram formula weight by the number of hydrogen atoms in its formula, as shown in the following reaction:

HCl+Na+→NaCl+H+

The single H⁺ of hydrochloric acid (HCl) is replaced by Na⁺. In 1 mole of HCl, there is 1 mole of replaceable hydrogen. By definition, the gEq of HCl must be the same as its gram formula weight, or 36.5 g. The two hydrogen atoms of sulfuric acid (H₂SO₄) are both considered to be replaceable. In 1 mole of sulfuric acid, there are 2 moles of replaceable hydrogen, and the gEq of H₂SO₄ is half its gram formula weight, or 48 g.

Acids in which hydrogen atoms are not completely replaceable are exceptions to the rule. In some acids, H⁺ replacement varies according to specific reactions. Carbonic acid (H₂CO₃) and phosphoric acid (H₃PO₄) are examples of such exceptions. Their equivalent weights are determined by the conditions of their chemical reactions.

For example, H₂CO₃ has two hydrogen atoms. In physiologic reactions, only one is considered replaceable:

H2CO3+Na+→NaHCO3+H+

Only one hydrogen atom is released; the other remains bound. In 1 mole of carbonic acid, there is only 1 mole of replaceable hydrogen. The gEq of carbonic acid is the same as its gram formula weight, or 61 g.

Gram Equivalent Weight of a Base

The equivalent weight of a base is its weight (grams) containing 1 mole of replaceable hydroxyl (OH⁻) ions. Similar to acids, the gEq of bases is calculated by dividing gram formula weight by the number of OH⁻ groups in its formula.

Conversion of Gram Weight to Equivalent Weight

To determine the number of gEqs in a substance, the gram weight is divided by its calculated equivalent weight, as shown in the following example:

58.5g NaClgEq58.5g=1gEq

29.25g NaClgEq58.5=0.5gEq

Milligram Equivalent Weights

The concentrations of most chemicals in the body are quite small. The term milligram equivalent weight (milliequivalent) is preferred for expressing these minute values; 1 mEq is simply 0.001 gEq:

mEq=gEq1000

The normal concentration of potassium (K⁺) in plasma ranges from 0.0035 to 0.005 gEq/L. These values may be converted to milliequivalents by multiplying by a factor of 1000. The normal concentration of K⁺ in the plasma would be expressed as ranging from 3.5 to 5.0 mEq/L.

Solute Content by Weight

The measurement of many electrolytes is based on actual weight rather than on milliequivalents. This weight is often expressed as milligrams per 100 ml of blood or body fluid. The units for this measurement are abbreviated as mg% (mg percent) or mg/dl (milligrams per deciliter). This text uses the modern designation mg/dl. Some substances present in blood or body fluid are present in extremely small amounts and are expressed in micrograms ( of a milligram) per deciliter (µg/dl or mcg/dl).

Values stated in mg/dl may be converted into their corresponding equivalent weights and reported as mEq/L. Conversion between mEq/L and mg/dl may be calculated as follows:

mEq/L=mg/dl×10Equivalent weight (1)

(1)

mEq/L=mEq/L×Equivalent weight10 (2)

(2)

To convert a serum Na⁺ value of 322 mg/dl to mEq/L, the equation is used as follows:

mEq/L=mg/dl×10Equivalent weight=322×1023=140 mEq/L

In clinical practice, electrolyte replacement is common when a laboratory test identifies a significant deficiency. The electrolyte content of intravenous solutions is usually stated in milligrams per deciliter or in mEq per liter. Lactated Ringer’s solution is one such infusion used for electrolyte replacement (Table 12-2).

Calculating Solute Content

In addition to gEq, mEq, mg/dl, and µg/dl (mcg/dl), several other methods of calculating solute content exist. These common chemical standards are used to compute solute content and dilution of solutions.

Quantitative Classification of Solutions

The amount of solute in a solution may be quantified by the following six methods:

Dilution Calculations

Dilute solutions are made from a stock preparation. Preparation of medications often involves dilution. Dilution calculations are based on the weight-per-unit volume principle (the aforementioned W/V solution method).

Diluting a solution increases its volume without changing the amount of solute it contains, and this reduces the concentration of the solution. The amount of solute in a solution can be expressed as volume times concentration. For example, 50 ml of a 10% solution (10 g/dl) contains 50 × 0.1, or 5 g. In the dilution of a solution, initial volume (V₁) multiplied by initial concentration (C₁) equals final volume multiplied by final concentration. This can be expressed as follows:

V1C1=V2C2

This equation is sometimes referred to as the dilution equation. Whenever three of the variables are known, the fourth can be calculated as in the following examples:

1. Diluting 10 ml of a 2% (0.02) solution to a concentration of 0.5% (0.005) requires finding the new volume (V₂) by rearranging the dilution equation as follows:

V2=V1C1C2

V2=10ml×0.020.005

V2=40ml

2. If 50 ml of water is added to 150 ml of a 3% (0.03) solution, the new concentration is calculated by rearranging the dilution equation to find C₂ as follows:

C2=V1C1V2

C2=150ml×0.02(50ml+150ml)

C2=0.0225 (2.25%)

3. To dilute 50 ml of a 0.33 normal (N) solution to a 0.1N concentration, concentration is given as normality, but it can be used similar to a percentage. The new volume (V₂) can be calculated by rearranging the dilution equation as follows:

V2=V1C1C2

V2=50 ml×0.330.1

V2=165 ml

In the last example, the volume needed to produce a 0.1N solution would be 165 ml − 50 ml (the original volume), or 115 ml. In other words, 115 ml of solvent would have to be added to the original 50 ml of 0.33N solution to produce the desired concentration. The added solvent is called the diluent because it dilutes the original concentration to a lower concentration.

The term acid refers to either compounds that can donate [H⁺] (Brönsted-Lowry acid) or any compound that accepts an electron pair (Lewis acid). Although these two theories of acids differ in which is being transferred, both theories attempt to describe how reactive groups perform within an aqueous solution8,9:

NH4Cl+NaOH→NH3+NaCl+HOH

In this reaction, sodium and chloride ions are not involved in the proton transfer. The equation can be rewritten ionically as follows to show the acidity of the ammonium ion:

NH4++OH−→NH3+HOH

The ammonium ion donates a hydrogen ion (proton) to the reaction. The H⁺ combines with the hydroxide ion (OH⁻), and this converts the former into ammonia gas and the latter into water.

Mini Clini

Methacholine Dilution

The dilution equation (V₁C₁ = V₂C₂) is commonly used to calculate volumes or concentrations of medications when a specific dosage needs to be administered to a patient. If three of the variables are known, the fourth can be determined.

 Problem

Methacholine is a drug used to challenge the airways of patients suspected to have asthma. In healthy subjects, only higher doses of methacholine cause bronchospasm. In asthmatics, very low doses can precipitate a 20% decrease in the forced expiratory volume in 1 second (FEV₁). The methacholine challenge test begins with a low dose and increases the concentration (either doubling or quadrupling) until the patient has a significant change in FEV₁ or the highest dose has been given. Methacholine is supplied in vials that contain 100 mg of the active substance to which 6.25 ml of diluent (saline) can be added to produce a concentration of 16 mg/ml.* This is the highest dosage that is administered to the patient. How can you make serial dilutions of the drug so that five different dosages are available and each one is four times more concentrated than the previous dose?

Solution

Starting with a 16 mg/ml stock solution of methacholine, how much diluent needs to be added to 3 ml of the stock to make a 4 mg/ml dose (one-fourth of the original concentration)?

Using the dilution equation:

C1V1=C2V2

(16)(3.0)=(4)V2

484=V2

12=V1

Because there was 3 ml of the stock solution to begin with, the amount of diluent to add is the difference between 12 (V₂) and 3, or 9 ml. Adding 9 ml of diluent to the original 3 ml of stock (16 mg/ml) provides 12 ml of methacholine with a concentration of 4 mg/ml, exactly one-fourth of the highest dose. Additional dilutions can be prepared using 3 ml of solution according to the following table:

Start With	Add Diluent	To Make
3 ml of 4 mg/ml	9 ml	1 mg/ml
3 ml of 1 mg/ml	9 ml	0.25 mg/ml
3 ml of 0.25 mg/ml	9 ml	0.0625 mg/ml

Each of these dilutions uses the same proportions used in the first dilution as determined by the dilution equation. Methacholine is administered by nebulizer to the patient starting with the lowest concentration (0.0625 mg/ml) and increasing until a change in FEV₁ is observed. (See Chapter 19 for additional information on pulmonary function testing.)

*Only trained individuals should prepare and label solutions of methacholine.

Acids With Single Ionizable Hydrogen

Simple compounds such as HCl ionize into one cation and one anion:

HCl→H++Cl−

Acids With Multiple Ionizable Hydrogens

The H⁺ ions in an acid may become available in stages. The degree of ionization increases as an electrolyte solution becomes more dilute. Concentrated sulfuric acid ionizes only one of its two hydrogen atoms per molecule, as follows:

H2SO4→H++HSO4−

With further dilution, second-stage ionization occurs:

H2SO4→H++H++SO4−

Bases

A base is a compound that yields hydroxyl ions (OH⁻) when placed into aqueous solution. A substance capable of inactivating acids is also considered a base. These compounds, called hydroxides, consist of a metal that is ionically bound to a hydroxide ion or ions. The hydroxide may also be bound to an ammonium cation (NH₄⁺). An example of this type of base is sodium hydroxide (NaOH). The Brönsted-Lowry definition of a base is any compound that accepts a proton; bases are paired with acids that donate the proton, and these are called conjugate pairs. This definition includes substances other than hydroxides, such as ammonia, carbonates, and certain proteins.

Hydroxide Bases

In aqueous solution, the following are typical dissociations of hydroxide bases:

Na+OH→Na++OH−

K+OH→K++OH−

Ca++(OH−)2→Ca+++2(OH−)

Inactivation of an acid is part of the definition of a base. This inactivation is accomplished by OH⁻ reacting with H⁺ to form water:

NaOH+HCl→NaCl+HOH

Nonhydroxide Bases

Ammonia and carbonates are examples of nonhydroxide bases. Proteins, with their amino groups, also can serve as nonhydroxide bases.

Ammonia

Ammonia qualifies as a base because it reacts with water to yield OH⁻:

NH3+HOH→NH4++OH−

and neutralizes H⁺ directly:

NH3+H+→NH4+

In both instances, NH₃ accepts a proton to become NH₄⁺. Ammonia plays an important role in renal excretion of acid (see Chapter 13).

Carbonates

The carbonate ion, CO₃²⁻, can react with water in the following way to produce OH⁻:

Na2CO3⇔2Na++CO32− (1)

(1)

CO32−+HOH⇔HCO3−+OH− (2)

(2)

In this reaction, the carbonate ion accepts a proton from water, becoming the bicarbonate ion. It simultaneously produces a hydroxide ion. The carbonate ion also can react directly with H⁺ to inactivate it:

CO32−+H+⇔HCO3−

Protein Bases

Proteins are composed of amino acids bound together by peptide links. Physiologic reactions in the body occur in a mildly alkaline environment. This environment allows proteins to act as H⁺ receptors, or bases. Cellular and blood proteins acting as bases are transcribed as prot⁻.

The imidazole group of the amino acid histidine is an example of an H⁺ acceptor on a protein molecule (Figure 12-4). The ability of proteins to accept hydrogen ions limits H⁺ activity in solution, which is called buffering. The buffering effect of hemoglobin is produced by imidazole groups in the protein. Each hemoglobin molecule contains 38 histidine residues. Each oxygen-carrying component (heme group) of hemoglobin is attached to a histidine residue. The ability of hemoglobin to accept (i.e., buffer) H⁺ ions depends on its oxygenation state. Deoxygenated (reduced) hemoglobin is a stronger base (i.e., a better H⁺ acceptor) than oxygenated hemoglobin. This difference partially accounts for the ability of reduced hemoglobin to buffer more acid than oxygenated hemoglobin can (see Chapter 13). Plasma proteins also act as buffers, although with less buffering power than hemoglobin, which contains more histidine.

Designation of Acidity and Alkalinity

Pure water can be used as a reference point for determining acidity or alkalinity. The concentration of both H⁺ and OH⁻ in pure water is 10⁻⁷ mol/L. A solution that has a greater H⁺ concentration or lower OH⁻ concentration than water acts as an acid. A solution that has a lower H⁺ concentration or a greater OH⁻ concentration than water is alkaline, or basic.

The H⁺ concentration [H⁺] of pure water has been adopted as the standard for comparing reactions of other solutions. Electrochemical techniques are used to measure the [H⁺] of unknown solutions. Acidity or alkalinity is determined by variation of the [H⁺] greater than or less than 1 × 10⁻⁷. For example, a solution with a [H⁺] of 89.2 × 10⁻⁴ has a higher [H⁺] than water and is acidic. A solution with a [H⁺] of 3.6 × 10⁻⁸ has fewer hydrogen ions than water and is by definition alkaline. Two related techniques are used for expressing the acidity or alkalinity of solutions using the [H⁺] of water (i.e., 10⁻⁷) as a neutral factor: (1) the [H⁺] in nanomoles per liter and (2) the logarithmic pH scale.

pH Scale

The pH scale is used to describe the concentration of H⁺, ([H⁺]), (i.e., Brönsted-Lowry acid) in a solution. Rather than expressing the [H⁺] as a very small number or in nanomoles, it is more convenient to describe it in terms of the inverse logarithm of the nanomolar [H⁺]. pH is defined as:

pH=−log⁡[H+]

pH is always represented as a positive number and is derived by converting the value for [H⁺] to a negative exponent of 10 and calculating its logarithm. The [H⁺] of water is 1 × 10⁻⁷ mol/L. Because the negative logarithm of 1 × 10⁻⁷ is 7, the pH of water is 7.

Using this scheme, in a solution with a pH of 7.00, the [H⁺] is the same as would be seen in pure water, so by definition this is called “neutral.” As the pH decreases to less than 7.00, the solution is termed more acidic, and when the pH increases to greater than 7.00, the solution is considered to be basic. With a whole number change in pH (i.e., pH decreasing from 7.00 to 6.00), the [H⁺] is a factor of 10 less. With a pH increase from 7.00 to 8.00, the [H⁺] is 10 times greater (Figure 12-5).

All fluids in the body are aqueous in origin. pK is the inverse logarithm of the dissociation constant for each solute. A pH of 7.00 is equivalent to a [H⁺] of 100 nmol. A pH of 8.00 is equivalent to a [H⁺] concentration of 10 nmol. Similarly, a change in pH of 0.3 unit equals a twofold change in [H⁺].

The law of mass action states that acids and bases freely dissociate and r-associate in a solution at a constant rate relative to the structure of the acid and the temperature of the system.¹⁰ Using the Henderson-Hasselbalch equation, which describes the ratio of [H⁺] to base, we can calculate expected pH (see Chapter 13).

pH=pK+log⁡[H+]

Where:

pH = Inverse log value [H⁺]

pK = Inverse log of dissociation constant of solution

log [H⁺] = Logarithm of [H⁺], which can be expressed as the ratio between conjugate acid and total acid concentration

Applying these concepts in an example pertinent to clinical medicine yields the following:

[H+]=4.0×10−8mol/L

pH=−log⁡(4.0×10−8)=−log⁡4.0+−log⁡10−8=−log⁡4.0+log⁡108=−0.602+8=7.40

In this example, the [H⁺] in arterial blood of a healthy adult is approximately 4.0 × 10⁻⁸ mol/L, or 40 nmol/L.

 Rule of Thumb

The pH scale is logarithmic. pH is a positive number representing the negative log of the hydrogen ion concentration [H⁺] of a solution. To visualize changes in acidity or alkalinity, the following two rules are helpful:

1. A pH change of 0.3 unit equals a 2-fold change in [H⁺].

2. A pH change of 1 unit equals a 10-fold change in [H⁺].

For example, if a patient’s blood pH decreased from 7.40 (normal) to 7.10, the [H⁺] concentration would be twice as high. If a patient’s urine pH decreased from 7.00 to 6.00, the [H⁺] would have increased by 10 times.

Distribution

Total body water is divided into the following two major compartments: (1) intracellular (“within the cells”) and (2) extracellular (“outside the cells”). Intracellular water accounts for approximately two-thirds of the total body water, and extracellular water accounts for the remaining one-third. Extracellular water is found in three subcompartments: (1) intravascular water (plasma), (2) interstitial water, and (3) transcellular fluid. Intravascular water constitutes approximately 5% of the body weight. Interstitial water is water in the tissues between the cells. It constitutes approximately 15% of the body weight. The proportion of transcellular fluid is quite small in proportion to plasma and interstitial fluid. Interstitial fluid is a matrix—a collagen/gel substance that allows the interstitium to provide structural support during times of extracellular volume depletion.¹¹ Examples of transcellular fluid include cerebrospinal fluid, digestive juices, and mucus. Transcellular fluid can become an important third space in some pathologic conditions, such as ascites (excess fluid in the peritoneal cavity) or pleural effusion (fluid collection in the pleural space).

Composition

The concentration of ionic solutes in intracellular and extracellular fluids differs significantly. Sodium (Na⁺), chloride (Cl⁻), and bicarbonate (HCO₃⁻) are predominantly extracellular electrolytes. Potassium (K⁺), magnesium (Mg⁺⁺), phosphate (PO₄³⁻), sulfate (SO₄²⁻), and protein constitute the main intracellular electrolytes. Although protein does not dissociate ionically, it can create hydrogen and other weak bonds and distribute net extra charge within its molecule. Intravascular and interstitial fluids have similar electrolyte compositions. However, plasma contains substantially more protein than interstitial fluid. Proteins, chiefly albumin, account for the high osmotic pressure of plasma. Osmotic pressure is an important determinant of fluid distribution between vascular and interstitial compartments.

Regulation

Movement of certain ions and proteins between body compartments is restricted. However, water diffuses freely. Control of total body water occurs through regulation of water intake (thirst) and water excretion (urine production, insensible loss, and stool water). The kidneys are mainly responsible for water excretion. If water intake is low, the kidneys reduce urine volume. Solutes in the urine can be concentrated up to four times the concentration of solutes in the plasma. If water intake is high, the kidneys can excrete large volumes of dilute urine.

The kidneys maintain the volume and composition of body fluids via two related mechanisms. First, filtration and reabsorption of sodium adjust urinary sodium excretion to match changes in dietary intake. Second, water excretion is regulated by osmoreceptors which are located in the hypothalamus and modulate secretion of antidiuretic hormone (ADH, also known as vasopressin).7,12,13 These receptors are exceptionally sensitive; in vivo studies have shown that a single neuron can respond to either an osmotic or a nonosmotic baroreceptor simulus.¹³ These mechanisms allow the kidneys to maintain the volume and concentration of body fluid despite variations in salt and water intake. Analysis of the urine (urinalysis) often provides diagnostic clues in disorders of body fluid volume.

Transport Between Compartments

Homeostasis depends largely on the total volume of body fluids and on fluid transport between body compartments. The first stage of homeostasis is fluid exchange between systemic capillaries and interstitial fluid via passive diffusion. Capillary walls are permeable to crystalline electrolytes. This allows equilibrium between the two extracellular compartments to occur quickly. Except for the large protein molecules, plasma can also move through capillary walls into the tissue spaces. Because water and small molecules can cross the capillary membranes, they produce little or no osmotic effect.

Movement of fluid and solutes from capillary blood to interstitial spaces is enhanced by the difference in hydrostatic pressure between compartments. Hydrostatic pressure difference depends on blood pressure, blood volume, and the vertical distance of the capillary from the heart (i.e., the effects of gravity). Hydrostatic pressure tends to cause fluid to leak out of capillaries into the interstitial spaces.

Osmotic pressure differences between interstitial and intravascular compartments oppose hydrostatic pressure; that is, osmotic pressure tends to keep fluid in the capillaries. Proteins with molecular weights greater than about 70,000 in colloidal suspension in the plasma cause this difference in osmotic pressure. Proteins such as albumin are too large to pass through the pores of the capillary. Instead, these proteins remain in the intravascular compartment and exert osmotic pressure, which draws water and small solute molecules back into the capillaries; this is called plasma colloid osmotic pressure (oncotic pressure). Because these large proteins are negatively charged, they attract (but do not bind) an equivalent amount of cations to the intravascular compartment. These cations have the effect of increasing osmotic pressure within the capillary (Donnan effect).

In a typical capillary, blood pressure is approximately 30 mm Hg at the arterial end and approximately 20 mm Hg at the venous end (Figure 12-6). Colloid osmotic pressure of the intravascular fluid remains constant at approximately 25 mm Hg. Hydrostatic pressure along the capillary continually decreases. At the arterial end, hydrostatic pressure normally exceeds osmotic pressure, and water flows out of the vascular space into the interstitial space. At the venous end, colloidal osmotic pressure exceeds hydrostatic forces. Water is pulled back into the vascular compartment.

The outflow of water and electrolytes from the capillary at the arterial end is not completely balanced by the return on the venous end. Slightly more water diffuses out than is reabsorbed. This slight outward excess is balanced by fluid return through the lymphatic circulation (see Chapter 8). Fluid return via lymphatic channels also depends on pressure differences. The pressure in the interstitial space is determined by the volume of interstitial fluid and its electrolyte content. Interstitial fluid moves from a region of higher pressure (interstitial space) to a region of lower pressure (lymphatic channels). This lymph fluid moves into larger lymphatic spaces, where the pressure is continuously decreasing.

These relationships may be expressed by the Starling equilibrium equation:

Qf=K1(Pch−Pih)−K2(Pco−Pio)

Where:

Three examples of the forces in this equation are fluid return from gravity-dependent areas of the body, fluid exchange in the lung, and tissue edema.

Because of hydrostatic effects, capillary pressure in the feet can reach 100 mm Hg when an individual is standing. Reabsorption of tissue fluid can be accomplished, although hydrostatic pressure greatly exceeds colloidal osmotic pressure. Three factors favor reabsorption under these circumstances:

However, when an imbalance results from changes in the basic pressures (e.g., arterial hypertension), edema tends to occur in the dependent limbs.

The lungs present a different situation. In systemic tissues, a constant exchange of interstitial fluid is essential. In the lungs, the alveoli must be kept relatively dry. Otherwise, interstitial fluid in the alveolar-capillary spaces would impede the diffusion of gas. Colloid osmotic pressure in pulmonary blood vessels is the same as it is in the systemic circulation. To minimize interstitial fluid in the alveolar-capillary region, the hydrostatic pressure difference must be kept low. The pulmonary circulation is a low-pressure system. The mean pulmonary vascular pressures are approximately one-sixth of those in the systemic circulation. Colloid osmotic pressure exceeds hydrostatic forces across the entire length of the pulmonary capillaries in healthy individuals. The alveoli are relatively free of excess interstitial water.

If hydrostatic pressure increases in the pulmonary circulation, this balance can be upset. This causes fluid movement into the alveolar-capillary spaces. Excess fluid in the interstitial space is called edema. In the lungs, edema caused by increased hydrostatic pressure often is a result of backpressure from a failing left ventricle (e.g., in congestive heart failure).

Edema can be caused by other factors. The Starling equilibrium equation given earlier shows that edema can be caused by a decrease in colloid osmotic pressure or an increase in capillary permeability. If albumin is depleted in the blood, the balance of forces is upset, favoring increased movement of fluid into the interstitium. Likewise, an increase in capillary permeability results in more fluid leaving the capillaries. Increased capillary permeability is a major factor in certain types of acute lung injuries (see Chapter 27).¹⁵

Electrolytes

Electrolytes in the various body fluids are not passive solutes. Electrolytes maintain the internal environment while making possible essential chemical and physiologic events. There are seven major electrolytes: sodium, chloride, bicarbonate, potassium, calcium, magnesium, and phosphorus (phosphate).