Solutions, Body Fluids, and Electrolytes

Solutions, Body Fluids, and Electrolytes

Daniel F. Fisher

In healthy individuals, body water and various chemicals are regulated to maintain an environment in which biochemical processes can continue. Imbalances in the amount or concentration of chemicals in the body occur in many diseases. The nature and importance of body fluids and electrolytes require an understanding of physiologic chemistry. This chapter provides the reader with the background knowledge needed to understand body chemistry.

Solutions, Colloids, and Suspensions

Definition of a Solution

The body is based on liquid water chemistry and the interaction of various substances either dissolved or suspended within the fluid. Water itself is a polar covalent molecule and is referred to in chemistry as a universal solvent. Water is the primary component of any liquid within the body and has a great influence on the behavior of other materials as they are introduced. These substances and particles combine with water in the following three ways: as (1) colloids, (2) suspensions, or (3) solutions.

A solution is a stable mixture of two or more substances in a single phase that cannot be separated using a centrifuge. One substance is evenly distributed between the molecules of the other. The substance that dissolves is called the solute. The medium in which it dissolves is called the solvent. Gases, liquids, and solids all can dissolve to become solutes. The process of dissolving involves breaking the (relatively weak) bonds between the solute-solute molecules and the solvent-solvent molecules. These intermolecular forces must be broken before a new solute-solvent bond can be formed. A solute dissolves in a solvent if the solute-solvent forces of attraction are great enough to overcome the solute-solute and solvent-solvent forces of attraction. If the solute-solvent force is less than the solute-solute or solvent-solvent force, the solute does not dissolve. When all three sets of forces are approximately equal, the two substances typically are soluble in each other. The electrical properties of the solvent molecules determine how soluble a substance is for a particular solvent. Polar solvents, such as water, dissolve other polar covalent bonds; nonpolar solvents dissolve nonpolar solutes: “Like dissolves like.”

Colloids (sometimes called dispersions or gels) consist of large molecules that attract and hold water (hydrophilic: “water loving”). These molecules are uniformly distributed throughout the dispersion, and they tend not to settle. The protoplasm inside cells is a common example of a colloid. Physiologically, colloids provide very little free water to the patient’s system, and care should be taken not to create a hypotonic environment.1

Suspensions are composed of large particles that float in a liquid. Suspensions can be physically separated by centrifugation and do not possess the same interactions between solvent and solute that are found in a true solution. Red blood cells in plasma are an example of a suspension. Dispersion of suspended particles depends on physical agitation. Particles settle because of gravity when the suspension is motionless.

The ease with which a solute dissolves in a solvent is its solubility, which is influenced by the following five factors:

The effects of temperature and pressure on the solubility of gases are important. More gas dissolves in a liquid at lower temperatures. As the temperature of a liquid increases, gas dissolved in that liquid comes out of solution. Henry’s law describes the effect of pressure on solubility of a gas in a liquid. At a given temperature, the volume of a gas that dissolves in a liquid is proportional to the solubility coefficient of the gas and the partial pressure of gas to which the liquid is exposed. Oxygen (O2) and carbon dioxide (CO2) transport can change significantly with changes in body temperature or atmospheric pressure (see Chapter 6).

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:

Jv=Lp [PcPi(pcpi)]



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.

Mini Clini

Sputum Induction and Hypertonic Saline


Sputum induction is usually performed by having the patient inhale a sterile hypertonic saline solution. Isotonic saline is approximately 0.9% (i.e., normal saline); concentrations greater than 0.9% are considered hypertonic. In clinical practice, concentrations of 3% to 10% have been used. The exact mechanism by which hypertonic saline increases the sputum volume has not been completely elucidated. However, when the particles of hypertonic saline are deposited in the airway, osmotic pressure is assumed to play a key role. When hypertonic saline comes into contact with the respiratory mucosa, water moves from the cells lining the airway into the sol-gel matrix that lines the airways, increasing its volume. The combination of increased volume of respiratory secretions with irritation of the epithelial cells themselves promotes reflex coughing. The volume of sputum and the rate of clearance from the lungs seem to depend on the osmolarity of the inhaled aerosol. Exposure of mast cells normally present in the airways to hypertonic aerosols results in the release of mediators (e.g., histamine) and bronchospasm. These effects may be related to the stimulation of the cough reflex. For the same reason, hypertonic saline is also sometimes used for bronchial challenge testing.

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 image 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:



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:



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.

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 (image 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)

image (1)

mEq/L=mEq/L×Equivalent weight10 (2)

image (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).

Quantitative Classification of Solutions

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

1. Ratio solution. The amount of solute to solvent is expressed as a proportion (e.g., 1 : 100). Ratio solutions are sometimes used in describing concentrations of drugs.

2. Weight-per-volume solution (W/V). The W/V solution is commonly used for solids dissolved in liquids. It is defined as weight of solute per volume of solution. This method is sometimes erroneously described as a percent solution. W/V solutions are commonly expressed in grams of solute per 100 ml of solution. For example, 5 g of glucose dissolved in 100 ml of solution is properly called a 5% solution, according to the W/V scheme. A liquid dissolved in a liquid is measured as volumes of solute to volumes of solution.

3. Percent solution. A percent solution is weight of solute per weight of solution. For example, 5 g of glucose dissolved in 95 g of water is a true percent solution. The glucose is 5% of the total solution weight of 100 g.

4. Molal solution. A molal solution contains 1 mole of solute per kilogram of solvent, or 1 mmol/g solvent. The concentration of a molal solution is independent of temperature.

5. Molar solution. A molar solution has 1 mole of solute per liter of solution, or 1 mmol/ml of solution. Solute is measured into a container, and solvent is added to produce the solution volume desired.

6. Normal solution. A normal solution

Jun 12, 2016 | Posted by in RESPIRATORY | Comments Off on Solutions, Body Fluids, and Electrolytes
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