Why is too much sodium or cholesterol harmful? Why does an iron deficiency cause anemia and an iodine deficiency cause a goiter? Why does a pH imbalance make some drugs less effective? Why do some pregnant women suffer convulsions after several days of vomiting? How can radiation cause cancer as well as cure it? None of these questions can be answered, nor would the rest of this book be intelligible, without understanding the chemistry of life. A little knowledge of chemistry can help you choose a healthy diet, use medications more wisely, avoid worthless health fads and frauds, and explain treatments and procedures to your patients or clients. Thus, we begin our study of the human body with basic chemistry, the simplest level of the body’s structural organization. We will progress from general chemistry to biochemistry, study of the molecules that compose living organisms——especially those unique to living things, such as carbohydrates, fats, proteins, and nucleic acids. Most people have at least heard of these——it is common knowledge that we need proteins, fats, carbohydrates, vitamins, and minerals in our diet, and we should avoid consuming too much saturated fat and cholesterol. But most people, too, have only a vague concept of what these molecules are, much less how they function in the body. Such knowledge is very helpful in matters of personal fitness and patient education, and is essential to the comprehension of the rest of this book.
A chemical element is the simplest form of matter to have unique chemical properties. Water, for example, has unique properties, but it can be broken down into two elements, hydrogen and oxygen, that have unique chemical properties of their own. If we carry this process any further, however, we find that hydrogen and oxygen are made of protons, neutrons, and electrons—and none of these are unique. A proton of gold is identical to a proton of oxygen. Hydrogen and oxygen are the simplest chemically unique components of water and are thus elements. Each element is identified by an atomic number, the number of protons in its nucleus. The atomic number of carbon is 6 and that of oxygen is 8, for example. The periodic table of the elements (see appendix A) arranges the elements in order by their atomic numbers. The elements are represented by one- or two-letter symbols, usually based on their English names: C for carbon, Mg for magnesium, Cl for chlorine, and so forth. A few symbols are based on Latin names, such as K for potassium (kalium), Na for sodium (natrium), and Fe for iron (ferrum). There are 91 naturally occurring elements on earth, 24 of which play normal physiological roles in humans. Table 2.1 groups these 24 according to their abundance in the body. Six of them account for 98.5% of the body’s weight: oxygen, carbon, hydrogen, nitrogen, calcium, and phosphorus. The next 0.8% consists of another 6 elements: sulfur, potassium, sodium, chlorine, magnesium, and iron. The remaining 12 account for 0.7% of body weight, and no one of them accounts for more than 0.02%; thus they are known as trace elements. Despite their minute quantities, trace elements play vital roles in physiology. Other elements without natural physiological roles can contaminate the body and severely disrupt its functions, as in heavy metal poisoning with lead or mercury. Several of these elements are classified as minerals— inorganic elements that are extracted from the soil by plants and passed up the food chain to humans and other organisms. Minerals constitute about 4% of the human body by weight. Nearly three-quarters of this is Ca and P; the rest is mainly Cl, Mg, K, Na, and S. Minerals contribute significantly to body structure. The bones and teeth consist partly of crystals of calcium, phosphate, magnesium, fluoride, and sulfate ions. Many proteins include sulfur, and phosphorus is a major component of nucleic acids, ATP, and cell membranes. Minerals also enable enzymes and other organic molecules to function. Iodine is a component of thyroid hormone; iron is a component of hemoglobin; and some enzymes function only when manganese, zinc, copper, or other minerals are bound to them. The electrolytes needed for nerve and muscle function are mineral salts. The biological roles of minerals are discussed in more magnitude in the later chapters
In the fifth century B.C.E., the Greek philosopher Democritus reasoned that we can cut matter such as a gold nugget into smaller and smaller pieces, but there must ultimately be particles so small that nothing could cut them. He called these imaginary particles atoms1 (“indivisible”). Atoms were only a philosophical concept until 1803, when English chemist John Dalton began to develop an atomic theory based on experimental evidence. In 1913, Danish physicist Niels Bohr proposed a model of atomic structure similar to planets orbiting the sun (figs. 2.1 and 2.2). Although this planetary model is too simple to account for many of the properties of atoms, it remains useful for elementary purposes.At the center of an atom is the nucleus, composed of protons and neutrons. Protons (p) have a single positive charge and neutrons (n0 ) have no charge. Each proton or neutron weighs approximately 1 atomic mass unit (amu), defined as one-twelfth the mass of an atom of carbon-12. The atomic mass of an element is approximately equal to its total number of protons and neutrons. Around the nucleus are one or more concentric clouds of electrons (e), tiny particles with a single negative charge and very low mass. It takes 1,836 electrons to equal 1 amu, so for most purposes we can disregard their mass. A person who weighs 64 kg (140 lb) contains less than 24 g (1 oz) of electrons. This hardly means that we can ignore electrons, however. They determine the chemical properties of an atom, thereby governing what molecules can exist and what chemical reactions can occur. The number of electrons equals the number of protons, so their charges cancel each other and an atom is electrically neutral. Electrons swarm about the nucleus in concentric regions called electron shells (energy levels). The more energy an electron has, the farther away from the nucleus its orbit lies. Each shell holds a limited number of electrons (see fig. 2.1). The one closest to the nucleus holds a maximum of 2 electrons, the second one holds a maximum of 8, and the third holds a maximum of 18. The outermost shell never holds more than 8 electrons, but a shell can acquire more electrons after another one, farther out, begins to fill. Thus, the third shell will hold 18 electrons only in atoms with four or more shells. The elements known to date have up to seven electron shells, but those ordinarily involved in human physiology do not exceed four. The electrons of the outermost shell, called valence electrons, determine the chemical bonding properties of an atom. An atom tends to bond with other atoms that will fill its outer shell and produce a stable number of valence electrons. A hydrogen atom, with only one electron shell and one electron (fig. 2.2), tends to react with other atoms that provide another electron and fill this shell with a stable number of two electrons. All other atoms react in ways that produce eight electrons in the valence shell. This tendency is called the octet rule (rule of eights)
Dalton believed that every atom of an element was identical, but we now know that all elements have varieties called isotopes,2 which differ from each other only in number of neutrons and therefore in atomic mass. Most hydrogen atoms, for example, have only one proton; this isotope is symbolized 1 H. Hydrogen has two other isotopes: deuterium ( 2 H) with one proton and one neutron, and tritium ( 3 H) with one proton and two neutrons (fig. 2.2). Over 99% of carbon atoms have an atomic mass of 12 (6p, 6n0 ) and are called carbon-12 (12C), but a small percentage of carbon atoms are 13C, with seven neutrons, and 14C, with eight. All isotopes of a given element behave the same chemically. Deuterium (2 H), for example, reacts with oxygen the same way 1 H does to produce water. The atomic weight of an element accounts for the fact that an element is a mixture of isotopes. If all carbon were 12C, the atomic weight of carbon would be the same as its atomic mass, 12.000. But since a sample of carbon also contains small amounts of the heavier isotopes 13C and 14C, the atomic weight is slightly higher, 12.011. Although different isotopes of an element exhibit identical chemical behavior, they differ in physical behavior. Many of them are unstable and decay (break down) to more stable isotopes by giving off radiation. Unstable isotopes are therefore called radioisotopes, and the process of decay is called radioactivity (see insight 2.1). Every element has at least one radioisotope. Oxygen, for example, has three stable isotopes and five radioisotopes. All of us contain radioisotopes such as 14C and 40K—that is, we are all mildly radioactive!
In 1896, French scientist Henri Becquerel (1852–1908) discovered that uranium darkened photographic plates through several thick layers of paper. Marie Curie (1867–1934) and Pierre Curie (1859–1906), her husband, discovered that polonium and radium did likewise. Marie Curie coined the term radioactivity for the emission of energy by these elements. Becquerel and the Curies shared a Nobel Prize in 1903 for this discovery. Marie Curie (fig 2.3) was not only the first woman in the world to receive a Nobel Prize but also the first woman in France even to receive a Ph.D. She received a second Nobel Prize in 1911 for inventing radiation therapy for breast and uterine cancer. Curie crusaded to train women for careers in science, and in World War I, she and her daughter, Irène JoliotCurie (1897–1956), trained physicians in the use of X-ray machines. In the wake of such discoveries, radium was regarded as a wonder drug. Unaware of its danger, people drank radium tonics and flocked to health spas to bathe in radium-enriched waters. Marie herself suffered extensive damage to her hands from handling radioactive minerals and died of radiation poisoning at age 67. The following year, Irène and her husband, Frédéric Joliot (1900–1958), were awarded a Nobel Prize for work in artificial radioactivity and synthetic radioisotopes. Apparently also a martyr to her science, Irène died of leukemia, possibly induced by radiation exposure.
Many forms of radiation, such as light and radio waves, have low energy and are harmless. High-energy radiation, however, ejects electrons from atoms, converting atoms to ions; thus it is called ionizing radiation. It destroys molecules and produces dangerous free radicals and ions in human tissues. Examples of ionizing radiation include ultraviolet rays, X rays, and three kinds of radiation produced by nuclear decay: alpha () particles, beta () particles, and gamma () rays. An particle is composed of two protons and two neutrons (equivalent to a helium nucleus), and a particle is a free electron. Alpha particles are too large to penetrate the skin, and particles can penetrate only a few millimeters. They are relatively harmless when emitted by sources outside the body, but they are very dangerous when emitted by radioisotopes that have gotten into the body. Strontium-90 (90Sr), for example, has been released by nuclear accidents and the atmospheric testing of nuclear weapons. It settles onto pastures and contaminates cow’s milk. In the body, it behaves chemically like calcium, becoming incorporated into the bones, where it emits particles for years. Uranium and plutonium emit electromagnetic rays, which have high energy and penetrating power. Gamma rays are very dangerous even when emitted by sources outside the body. Each radioisotope has a characteristic physical halflife, the time required for 50% of its atoms to decay to a more stable state. One gram of 90Sr, for example, would be half gone in 28 years. In 56 years, there would still be 0.25 g left, in 84 years 0.125 g, and so forth. Many radioisotopes are much longer-lived. The half-life of 40K, for example, is 1.3 billion years. Nuclear power plants produce hundreds of radioisotopes that will be intensely radioactive for at least 10,000 years—longer than the life of any disposal container yet conceived. The biological half-life of a radioisotope is the time required for half of it to disappear from the body. This is a function of both physical decay and physiological clearance from the body. Cesium-137, for example, has a physical half-life of 30 years but a biological half-life of only 17 days. Chemically, it behaves like potassium; it is quite mobile and rapidly excreted by the kidneys. There are several ways to measure the intensity of ionizing radiation, the amount absorbed by the body, and its biological effects. To understand the units of measurement requires a grounding in physics beyond the scope of this book, but the standard international (SI) unit of radiation exposure is the sievert3 (Sv), which takes into account the type and intensity of radiation and its biological effect. Doses of 5 Sv or more are usually fatal. The average American receives about 3.6 millisieverts (mSv) per year in background radiation from natural sources and another 0.6 mSv from artificial sources. The most significant natural source is radon, a gas that is produced by the decay of uranium in the earth and that may accumulate in buildings to unhealthy levels. Artificial sources include medical X rays, radiation therapy, and consumer products such as color televisions, smoke detectors, and luminous watch dials. Such voluntary exposure must be considered from the standpoint of its risk-to-benefit ratio. The benefits of a smoke detector or mammogram far outweigh the risk from the low levels of radiation involved. Radiation therapists and radiologists face a greater risk than their patients, however, and astronauts and airline flight crews receive more than average exposure. U.S. federal standards set a limit of 50 mSv/year as acceptable occupational exposure to ionizing radiation.
Ions, Electrolytes, and Free Radicalsion are charged particles with unequal numbers of protons and electrons. Elements with one to three valence
electrons tend to give them up, and those with four to
seven electrons tend to gain more. If an atom of the first
kind is exposed to an atom of the second, electrons may
transfer from one to the other and turn both of them into
ions. This process is called ionization. The particle that
gains electrons acquires a negative charge and is called an
anion (AN-eye-on). The one that loses electrons acquires a
positive charge (because it then has a surplus of protons)
and is called a cation (CAT-eye-on).
Consider, for example, what happens when sodium
and chlorine meet (fig. 2.4). Sodium has three electron
shells with a total of 11 electrons: 2 in the first shell, 8
in the second, and 1 in the third. If it gives up the electron in the third shell, its second shell becomes the
valence shell and has the stable configuration of 8 electrons. Chlorine has 17 electrons: 2 in the first shell, 8 in the second, and 7 in the third. If it can gain one more
electron, it can fill the third shell with 8 electrons and
become stable. Sodium and chlorine seem “made for
each other”—one needs to lose an electron and the other
needs to gain one. This is just what they do. When they
interact, an electron transfers from sodium to chlorine.
Now, sodium has 11 protons in its nucleus but only 10
electrons. This imbalance gives it a positive charge, so
we symbolize the sodium ion Na. Chlorine has been
changed to the chloride ion with a surplus negative
charge, symbolized Cl.
Some elements exist in two or more ionized forms.
Iron, for example, has ferrous (Fe2) and ferric (Fe3) ions.
Note that some ions have a single positive or negative
charge, while others have charges of 2 or 3 because
they gain or lose more than one electron. The charge on an
ion is called its valence. Ions are not always single atoms
that have become charged; some are groups of atoms—phosphate (PO4
3) and bicarbonate (HCO3
) ions, for
example.
Ions with opposite charges are attracted to each other
and tend to follow each other through the body. Thus,
when Na is excreted in the urine, Cl tends to follow it.
The attraction of cations and anions to each other is
important in maintaining the excitability of muscle and
nerve cells, as we shall see in chapters 11 and 12.
Electrolytes are salts that ionize in water and form
solutions capable of conducting electricity (table 2.2). We
can detect electrical activity of the muscles, heart, and
brain with electrodes on the skin because electrolytes in
the body fluids conduct electrical currents from these
organs to the skin surface. Electrolytes are important for
their chemical reactivity (as when calcium phosphate
becomes incorporated into bone), osmotic effects (influence on water content and distribution in the body), and
electrical effects (which are essential to nerve and muscle
function). Electrolyte balance is one of the most important
considerations in patient care. Electrolyte imbalances
have effects ranging from muscle cramps and brittle bones
to coma and cardiac arrest.
Free radicals are chemical particles with an odd
number of electrons. For example, oxygen normally exists
as a stable molecule composed of two oxygen atoms, O2;
but if an additional electron is added, it becomes a free
radical called the superoxide anion, O2
•. Free radicals
are represented with a dot to symbolize the odd electron.
Free radicals are produced by some normal metabolic reactions of the body (such as the ATP-producing
oxidation reactions in mitochondria, and a reaction that
some white blood cells use to kill bacteria), by radiation
(such as ultraviolet radiation and X rays), and by chemicals (such as carbon tetrachloride, a cleaning solvent, and
nitrites, present as preservatives in some wine, meat, and
other foods). They are short-lived and combine quickly
with molecules such as fats, proteins, and DNA, converting them into free radicals and triggering chain reactions
that destroy still more molecules. Among the damages
caused by free radicals are some forms of cancer and
myocardial infarction, the death of heart tissue. One theory of aging is that it results in part from lifelong cellular
damage by free radicals.
Because free radicals are so common and destructive, we have multiple mechanisms for neutralizing
them. An antioxidant is a chemical that neutralizes free
radicals. The body produces an enzyme called superoxide dismutase (SOD), for example, that converts superoxide into oxygen and hydrogen peroxide. Selenium,
vitamin E (-tocopherol), vitamin C (ascorbic acid), and
carotenoids (such as -carotene) are some antioxidants
obtained from the diet. Dietary deficiencies of antioxidants have been associated with increased incidence of
heart attacks, sterility, muscular dystrophy, and other
disorders.
Molecules are chemical particles composed of two or more
atoms united by a covalent chemical bond (the sharing of
electrons). The atoms may be identical, as in nitrogen (N2),
or different, as in glucose (C6H12O6). Molecules composed of
two or more different elements are called compounds. Oxygen (O2) and carbon dioxide (CO2) are both molecules
because both consist of at least two atoms, but only CO2 is a
compound, because it has atoms of two different elements.
Molecules can be represented by molecular formulae, as shown here, that identify their constituent elements
and show how many atoms of each are present. Molecules
with identical molecular formulae but different arrangements of their atoms are called isomers4 of each other. For
example, both ethanol (grain alcohol) and ethyl ether have
the molecular formula C2H6O, but they are certainly not
interchangeable! To show the difference between them,
we use structural formulae that show the location of each
atom (fig. 2.5).
The molecular weight (MW) of a compound is the sum
of the atomic weights of its atoms. Rounding the atomic
mass units (amu) to whole numbers, we can calculate the
approximate MW of glucose (C6H12O6), for example, as
6 C atoms X 12 amu each = 72 amu
12 H atoms X 1 amu each = 12 amu
6 O atoms X 16 amu each = 96 amu
Molecular weight (MW) = 180 amu
Molecular weight is needed to compute some measures of
concentration, as we shall see later.
A molecule is held together, and molecules are
attracted to each other, by forces called chemical bonds.
The three bonds of greatest physiological interest are ionic
bonds, covalent bonds, and hydrogen bonds (table 2.3).
An ionic bond is the attraction of a cation to an
anion. Sodium (Na) and chloride (Cl) ions, for example, are attracted to each other and form the compound
sodium chloride (NaCl), common table salt. Ionic compounds can be composed of more than two ions. Calcium
has two valence electrons. It can become stable by donating one electron to one chlorine atom and the other electron to another chlorine, thus producing a calcium ion(Ca2) and two chloride ions. The result is calcium chloride, CaCl2. Ionic bonds are weak and easily dissociate
(break up) in the presence of something more attractive,
such as water. The ionic bonds of NaCl break down easily
as salt dissolves in water, because both Na and Cl are
more attracted to water molecules than they are to each
other.
Do you think ionic bonds are common in the human body? Explain your answer
Covalent bonds form by the sharing of electrons. For example, two hydrogen atoms share valence electrons to form a hydrogen molecule, H2 (fig. 2.6a). The two electrons, one donated by each atom, swarm around both nuclei in a dumbbell-shaped cloud. A single covalent bond is the sharing of a single pair of electrons. It is symbolized by a single line between atomic symbols, for example H H. A double covalent bond is the sharing of two pairs of electrons. In carbon dioxide, for example, a central carbon atom shares two electron pairs with each oxygen atom. Such bonds are symbolized by two lines, for example O=C=O e around each nucleus, they form a nonpolar covalent bond (fig. 2.7a), the strongest of all chemical bonds. Carbon atoms bond to each other with nonpolar covalent bonds. If shared electrons spend significantly more time orbiting one nucleus than they do the other, they lend their negative charge to the region where they spend the most time, and they form a polar covalent bond (fig. 2.7b). When hydrogen bonds with oxygen, for example, the electrons are more attracted to the oxygen nucleus and orbit it more than they do the hydrogen. This makes the oxygen region of the molecule slightly negative and the hydrogen regions slightly positive. The Greek delta ( ) is used to symbolize a charge less than that of one electron or proton. A slightly negative region of a molecule is represented and a slightly positive region is represented . A hydrogen bond is a weak attraction between a slightly positive hydrogen atom in one molecule and a slightly negative oxygen or nitrogen atom in another. Water molecules, for example, are weakly attracted to each other by hydrogen bonds (fig. 2.8). Hydrogen bonds also form between different regions of the same molecule, especially in very large molecules such as proteins and DNA. They cause such molecules to fold or coil into precise three-dimensional shapes. Hydrogen bonds are represented by dotted or broken lines between atoms: =C=O· · ·H N . Hydrogen bonds are the weakest of all the bond types we have considered, but they are enormously important to physiology
Bond Type Definition and Remarks
Ionic bond Relatively weak attraction between an anion
and a cation. Easily disrupted in water, as
when salt dissolves.
Covalent bond Sharing of one or more pairs of electrons
between nuclei.
Single covalent Sharing of one electron pair.
Double covalent Sharing of two electron pairs. Often occurs
between carbon atoms, between carbon and
oxygen, and between carbon and nitrogen
.
Nonpolar covalent Covalent bond in which electrons are equally
attracted to both nuclei. May be single or
double. Strongest type of chemical bond.
Polar covalent Covalent bond in which electrons are more
attracted to one nucleus than to the other,
resulting in slightly positive and negative
regions in one molecule. May be single or
double.
Hydrogen bond Weak attraction between polarized molecules
or between polarized regions of the same
molecule. Important in the three-dimensional
folding and coiling of large molecules. Weakest
of all bonds; easily disrupted by temperature
and pH changes.
Objectives
When you have completed this section, you should be able to
• define mixture and distinguish between mixtures and
compounds;
• describe the biologically important properties of water;
• show how three kinds of mixtures differ from each other;
• discuss some ways in which the concentration of a solution
can be expressed, and explain why different expressions of
concentration are used for different purposes; and
• define acid and base and interpret the pH scale.
Our body fluids are complex mixtures of chemicals. A mixture consists of substances that are physically blended but not chemically combined. Each substance retains its own chemical properties. To contrast a mixture with a compound, consider sodium chloride again. Sodium is a lightweight metal that bursts into flame if exposed to water, and chlorine is a yellow-green poisonous gas that was used for chemical warfare in World War I. When these elements chemically react, they form common table salt. Clearly, the compound has properties much different from the properties of its elements. But if you were to put a little salt on your watermelon, the watermelon would taste salty and sweet because the sugar of the melon and the salt you added would merely form a mixture in which each compound retained its individual properties.
Most mixtures in our bodies consist of chemicals dissolved or suspended in water. Water constitutes 50% to
75% of your body weight, depending on age, sex, fat content, and other factors. Its structure, simple as it is, has
profound biological effects. Two aspects of its structure are
particularly important: (1) its atoms are joined by polar
covalent bonds, and (2) the molecule is V-shaped, with a
105° bond angle (fig. 2.9a). This makes the molecule as a
whole polar, with a slight negative charge (
) on the oxygen and a slight positive charge (
) on each hydrogen.
Like little magnets, water molecules are attracted to each
other by hydrogen bonds (see fig. 2.8). This gives water a
set of properties that account for its ability to support life:
solvency, cohesion, adhesion, chemical reactivity, and
thermal stability.
Solvency is the ability to dissolve other chemicals.
Water is sometimes called the universal solvent because it
dissolves a broader range of substances than any other liquid. Substances that dissolve in water, such as sugar, are
said to be hydrophilic (HY-dro-FILL-ic); the relatively few substances that do not, such as fats, are hydrophobic (HYdro-FOE-bic). Virtually all metabolic reactions depend on
the solvency of water. Biological molecules must be dissolved in water to move freely, come together, and react.
The solvency of water also makes it the body’s primary
means of transporting substances from place to place.
To be soluble in water, a molecule must be charged so
that its charges can interact with those of water. When NaCl
is dropped into water, for example, the ionic bonds between
Na and Cl are overpowered by the attraction of each ion
to water molecules. Water molecules form a cluster, or
hydration sphere, around each sodium ion with the O
pole of each water molecule facing the sodium ion. They
also form a hydration sphere around each chloride ion,
with the H
poles facing it. This isolates the sodium ions
from the chloride ions and keeps them dissolved (fig. 2.9b).
Adhesion is the tendency of one substance to cling to
another, whereas cohesion is the tendency of molecules of
the same substance to cling to each other. Water adheres to
the body’s tissues and forms a lubricating film on membranes such as the pleura and pericardium. This helps
reduce friction as the lungs and heart contract and expand
and rub against these membranes. Water also is a very
cohesive liquid because of its hydrogen bonds. This is why,
when you spill water on the floor, it forms a puddle and
evaporates slowly. By contrast, if you spill a nonpolar substance such as liquid nitrogen, it dances about and evaporates in seconds, like a drop of water in a hot dry skillet. This is because nitrogen molecules have no attraction for
each other, so the little bit of heat provided by the floor is
enough to disperse them into the air. The cohesion of water
is especially evident at its surface, where it forms an elastic layer called the surface film held together by a force
called surface tension. This force causes water to hang in
drops from a leaky faucet and travel in rivulets down a
window.
The chemical reactivity of water is its ability to participate in chemical reactions. Not only does water ionize
many other chemicals such as acids and salts, but water
itself ionizes into H and OH. These ions can be incorporated into other molecules, or released from them, in the
course of chemical reactions such as hydrolysis and dehydration synthesis, described later in this chapter.
The thermal stability of water helps to stabilize the
internal temperature of the body. It results from the high
heat capacity of water—the amount of heat required to
raise the temperature of 1 g of a substance by 1°C. The
base unit of heat is the calorie7 (cal)—1 cal is the amount
of heat that raises the temperature of 1 g of water 1°C. The
same amount of heat would raise the temperature of a
nonpolar substance such as nitrogen about four times as
much. The difference stems from the presence or absence
of hydrogen bonding. To increase in temperature, the
molecules of a substance must move around more
actively. The hydrogen bonds of water molecules inhibit
their movement, so water can absorb a given amount of heat without changing temperature (molecular motion)
as much.
The high heat capacity of water also makes it a very
effective coolant. When it changes from a liquid to a
vapor, water carries a large amount of heat with it. One
milliliter of perspiration evaporating from the skin
removes about 500 calories of heat from the body. This
effect is very apparent when you are sweaty and stand in
front of a fan
Think About It
Why are heat and temperature not the same thing?
Mixtures of other substances in water can be classified as
solutions, colloids, and suspensions.
A solution consists of particles of matter called the
solute mixed with a more abundant substance (usually
water) called the solvent. The solute can be a gas, solid, or
liquid—as in a solution of oxygen, sodium chloride, or
alcohol in water, respectively. Solutions are defined by the
following properties:
• The solute particles are under 1 nanometer (nm) in
size. The solute and solvent therefore cannot be
visually distinguished from each other, even with a
microscope.
• Such small particles do not scatter light noticeably, so
solutions are usually transparent
• The solute particles will pass through most selectively
permeable membranes, such as dialysis tubing and
cell membranes.
• The solute does not separate from the solvent when
the solution is allowed to stand.
The most common colloid in the body is protein,
such as the albumin in blood plasma. Many colloids can
change from liquid to gel states—gelatin desserts, agar culture media, and the fluids within and between our cells,
for example. Colloids are defined by the following physical properties:
• The colloidal particles range from 1 to 100 nm in size.
• Particles this large scatter light, so colloids are usually
cloudy
• The particles are too large to pass through most
selectively permeable membranes.
• The particles are still small enough, however, to
remain permanently mixed with the solvent when the
mixture stands
The blood cells in our plasma exemplify a suspension. Suspensions are defined by the following properties:
• The suspended particles exceed 100 nm in size.
• Such large particles render suspensions cloudy or
opaque.
• The particles are too large to penetrate selectively
permeable membranes.
• The particles are too heavy to remain permanently
suspended, so suspensions separate on standing. If
allowed to stand, blood cells settle to the bottom of a
tube,
An emulsion is a suspension of one liquid in another, such
as oil and vinegar salad dressing. The fat in breast milk is
an emulsion, as are medications such as Kaopectate and
milk of magnesia.
A single mixture can fit into more than one of these
categories. Blood is a perfect example—it is a solution of
sodium chloride, a colloid of protein, and a suspension of cells. Milk is a solution of calcium, a colloid of protein,
and an emulsion of fat
Solutions are often described in terms of their concentration—how much solute is present in a given volume of solution. Concentration is expressed in different ways for different purposes, some of which are explained here. You may find the table of symbols and measures inside the back cover to be helpful as you study this section.
A simple way to express concentration is the weight of solute in a given volume of solution. For example, intravenous (I.V.) saline typically contains 8.5 grams of NaCl per liter of solution (8.5 g/L). For many biological purposes, however, we deal with smaller quantities such as milligrams per deciliter (mg/dL; 1 dL = 100 mL). For example, a typical serum cholesterol concentration may be 200 mg/dL, also expressed 200 mg/100 mL or 200 milligram-percent (mg-%).
Percentage concentrations are also simple to compute, but it is necessary to specify whether the percentage refers to the weight or the volume of solute in a given volume of solution. For example, if we begin with 5 g of dextrose (an isomer of glucose) and add enough water to make 100 mL of solution, the resulting concentration will be 5% weight per volume (w/v). A common intravenous fluid is D5W, which stands for 5% w/v dextrose in distilled water. If the solute is a liquid, such as ethanol, percentages refer to volume of solute per volume of solution. Thus, 70 mL of ethanol diluted with water to 100 mL of solution produces 70% volume per volume (70% v/v) ethanol.
Percent concentrations are easy to prepare, but that unit of measurement is inadequate for many purposes. The physiological effect of a chemical depends on how many molecules of it are present in a given volume, not the weight of the chemical. Five percent glucose, for example, contains almost twice as many glucose molecules as the same volume of 5% sucrose (fig. 2.11a). Each solution contains 50 g of sugar per liter, but glucose has a molecular weight (MW) of 180 and sucrose has a MW of 342. Since each molecule of glucose is lighter, 50 g of glucose contains more molecules than 50 g of sucrose. To produce solutions with a known number of molecules per volume, we must factor in the molecular weight. If we know the MW and weigh out that many grams of the substance, we have a quantity known as its gram molecular weight, or 1 mole. One mole of glucose is 180 g and 1 mole of sucrose is 342 g. Each quantity contains the same number of molecules of the respective sugar—a number known as Avogadro’s9 number, 6.023 1023. Such a large number is hard to imagine. If each molecule were the size of a pea, 6.023 1023 molecules would cover 60 earthsized planets 3 m (10 ft) deep! Molarity (M) is the number of moles of solute per liter of solution. A one-molar(1.0 M) solution of glucose contains 180 g/L, and 1.0 M solution of sucrose contains 342 g/L. Both have the same number of solute molecules in a given volume (fig. 2.11b). Body fluids and laboratory solutions usually are less concentrated than 1 M, so biologists and clinicians more often work with millimolar (mM) and micromolar ( M) concentrations—103 and 106 M, respectively
Electrolytes are important for their chemical, physical
(osmotic), and electrical effects on the body. Their electrical effects, which determine such things as nerve, heart,
and muscle actions, depend not only on their concentration but also on their electrical charge. A calcium ion
(Ca2) has twice the electrical effect of a sodium ion (Na),
for example, because it carries twice the charge. When we
measure electrolyte concentrations, we must therefore
take the charges into account.
One equivalent (Eq) of an electrolyte is the amount
that would electrically neutralize 1 mole of hydrogen ions
(H) or hydroxide ions (OH). For example, 1 mole (58.4
g) of NaCl yields 1 mole, or 1 Eq, of Na in solution. Thus,
an NaCl solution of 58.4 g/L contains 1 equivalent of Na
per liter (1 Eq/L). One mole (98 g) of sulfuric acid (H2SO4)
yields 2 moles of positive charges (H). Thus, 98 g of sulfuric acid per liter would be a solution of 2 Eq/L.
The electrolytes in our body fluids have concentrations less than 1 Eq/L, so we more often express their concentrations in milliequivalents per liter (mEq/L). If you
know the millimolar concentration of an electrolyte, you can easily convert this to mEq/L by multiplying it by the
valence of the ion:
1 mM Na+ = 1 mEq/L
1 mM Ca2+ = 2 mEq/L
1 mM Fe3+ = 3 mEq/L
Acids, Bases, and pH
Most people have some sense of what acids and bases are. Advertisements are full of references to excess stomach acid and pH-balanced shampoo. We know that drain cleaner (a strong base) and battery acid can cause serious chemical burns. But what exactly do “acidic” and “basic” mean, and how can they be quantified? An acid is any proton donor, a molecule that releases a proton (H+) in water. A base is a proton acceptor. Since hydroxide ions (OH-) accept H+, many bases are substances that release hydroxide ions—sodium hydroxide (NaOH), for example. A base does not have to be a hydroxide donor, however. Ammonia (NH3) is also a base. It does not release hydroxide ions, but it readily accepts hydrogen ions to become the ammonium ion (NH4+ ). Acidity is expressed in terms of pH, a measure derived from the molarity of H. Molarity is represented by square brackets, so the molarity of H+ is symbolized [H+]. pH is the negative logarithm of hydrogen ion molarity—that is, pH = log [H+]. In pure water, 1 in 10 million molecules ionizes into hydrogen and hydroxide ions: H2O ↔ H+ OH-. Pure water has a neutral pH because it contains equal amounts of H+ and OH-. Since 1 in 10 million molecules ionize, the molarity of H and the pH of water are [H+] = 0.0000001 molar = 10^-7 M log [H+] = -7 pH = -log [H+] = 7 The pH scale was invented in 1909 by Danish biochemist and brewer Sören Sörensen to measure the acidity of beer. The scale extends from 0.0 to 14.0. A solution with a pH of 7.0 is neutral; solutions with pH below 7 are acidic; and solutions with pH above 7 are basic (alkaline). The lower the pH value, the more hydrogen ions a solution has and the more acidic it is. Since the pH scale is logarithmic, a change of one whole number on the scale represents a 10-fold change in H concentration. In other words, a solution with a pH of 4 is 10 times as acidic as one with a pH of 5 and 100 times as acidic as one with a pH of 6. Slight disturbances of pH can seriously disrupt physiological functions and alter drug actions so it is important that the body carefully control its pH. Blood, for example, normally has a pH ranging from 7.35 to 7.45. Deviations from this range cause tremors, fainting, paralysis, or even death. Chemical solutions that resist changes in pH are called buffers. Buffers and pH regulation are laterly discussed
The pH of our body fluids has a direct bearing on how we react to drugs. Depending on pH, drugs such as aspirin, phenobarbital, and penicillin can exist in charged (ionized) or uncharged forms. Whether a drug is charged or not can determine whether it will pass through cell membranes. When aspirin is in the acidic environment of the stomach, for example, it is uncharged and passes easily through the stomach lining into the bloodstream. Here it encounters a basic pH, whereupon it ionizes. In this state, it is unable to pass back through the membrane, so it accumulates in the blood. This effect, called ion trapping or pH partitioning, can be controlled to help clear poisons from the body. The pH of the urine, for example, can be manipulated so that poisons become trapped there and thus rapidly excreted from the body
• define energy and work, and describe some types of energy;
• understand how chemical reactions are symbolized by
chemical equations;
• list and define the fundamental types of chemical reactions;
• identify the factors that govern the speed and direction of a
reaction;
• define metabolism and its two subdivisions;
• define oxidation and reduction and relate these to changes in
the energy content of a molecule.
Energy is the capacity to do work. To do work means to move something, whether it is a muscle or a molecule. Some examples of physiological work are breaking chemical bonds, building molecules, pumping blood, and contracting skeletal muscles. All of the body’s activities are forms of work. Energy is broadly classified as potential or kinetic energy. Potential energy is energy contained in an object because of its position or internal state, but which is not doing work at the time. Kinetic energy is energy of motion, energy that is doing work. It is observed in skeletomuscular movements, the flow of ions into a cell, and vibration of the eardrum, for example. The water behind a dam has potential energy because of its position. Let the water flow through, and it exhibits kinetic energy that can be tapped for generating electricity. Like water behind a dam, ions concentrated on one side of a cell membrane have potential energy that can be released by opening gates in the membrane. As the ions flow through the gates, their kinetic energy can be tapped to create a nerve signal or make the heart beat. Within the two broad categories of potential and kinetic energy, there are several forms of energy relevant to human physiology. Chemical energy is potential energy stored in the bonds of molecules. Chemical reactions release this energy and make it available for physiological work. Heat is the kinetic energy of molecular motion. The temperature of a substance is a measure of rate of this motion, and adding heat to a substance increases this rate. Electromagnetic energy is the kinetic energy of moving “packets” of radiation called photons. The most familiar form of electromagnetic energy is light. Electrical energy has both potential and kinetic forms. It is potential energy when charged particles have accumulated at a point such as a battery terminal or on one side of a cell membrane; it becomes kinetic energy when these particles begin to move and create an electrical current—for example, when electrons move through your household wiring or sodium ions move through a cell membrane.
A chemical reaction is a process in which a covalent or ionic bond is formed or broken. The course of a chemical reaction is symbolized by a chemical equation that typically shows the reactants on the left, the products on the right, and an arrow pointing from the reactants to the products. For example, consider this common occurrence: If you open a bottle of wine and let it stand for several days, it turns sour. Wine “turns to vinegar” because oxygen gets into the bottle and reacts with ethanol to produce acetic acid and water. Acetic acid gives the tart flavor to vinegar and spoiled wine. The equation for this reaction is CH3CH2OH O2 CH3COOH H2O Ethanol Oxygen → Acetic acid Water Ethanol and oxygen are the reactants, and acetic acid and water are the products of this reaction. Not all reactions are shown with the arrow pointing from left to right. In complex biochemical equations, reaction chains are often written vertically or even in circles. Chemical reactions can be classified as decomposition, synthesis, or exchange reactions. In decomposition reactions, a large molecule breaks down into two or more smaller ones (fig. 2.13a); symbolically, AB → A B. When you eat a potato, for example, digestive enzymes decompose its starch into thousands of glucose molecules, and most cells further decompose glucose to water and carbon dioxide. Starch, a very large molecule, ultimately yields about 36,000 molecules of H2O and CO2. Synthesis reactions are just the opposite—two or more small molecules combine to form a larger one; symbolically, A B → AB (fig. 2.13b). When the body synthesizes proteins, for example, it combines several hundred amino acids into one protein molecule. In exchange reactions, two molecules exchange atoms or groups of atoms; AB CD → AC BD (fig. 2.13c). For example, when stomach acid (HCl) enters the small intestine, the pancreas secretes sodium bicarbonate (NaHCO3) to neutralize it. The reaction between the two is NaHCO3 HCl → NaCl H2CO3. We could say the sodium atom has exchanged its bicarbonate group ( HCO3) for a chlorine atom. Reversible reactions can go in either direction under different circumstances and are represented with doubleheaded arrows. For example, carbon dioxide combines with water to produce carbonic acid, which in turn decomposes into bicarbonate ions and hydrogen ions: CO2 H2O ↔ H2CO3 ↔ HCO3 H Carbon Water Carbonic Bicarbonate Hydrogen dioxide acid ion ion This reaction appears in this book more often than any other, especially as we discuss respiratory, urinary, and digestive physiology. The direction in which a reversible reaction goes is determined by the relative abundance of substances on each side of the equation. If there is a surplus of CO2, this reaction proceeds to the right and produces bicarbonate and hydrogen ions. If bicarbonate and hydrogen ions are present in excess, the reaction proceeds to the left and generates CO2 and H2O. Reversible reactions follow the law of mass action: they proceed from the side with the greater quantity of reactants to the side with the lesser quantity. This law will help to explain processes discussed in later chapters, such as why hemoglobin binds oxygen in the lungs yet releases it to muscle tissue. In the absence of upsetting influences, reversible reactions exist in a state of equilibrium, in which the ratio of products to reactants is stable. The carbonic acid reaction, for example, normally maintains a 20:1 ratio of bicarbonate ions to carbonic acid molecules. This equilibrium can be upset, however, by a surplus of hydrogen ions, which drive the reaction to the left, or adding carbon dioxide and driving it to the right.
The basis for chemical reactions is molecular motion and
collisions. All molecules are in constant motion, and reactions occur when mutually reactive molecules collide with sufficient force and the right orientation. The rate of
a reaction depends on the nature of the reactants and on
the frequency and force of these collisions. Some factors
that affect reaction rates are:
• Concentration. Reaction rates increase when the
reactants are more concentrated. This is because the
molecules are more crowded and collide more
frequently.
• Temperature. Reaction rate increases as the
temperature rises. This is because heat causes
molecules to move more rapidly and collide with
greater force and frequency.
• Catalysts (CAT-uh-lists). These are substances that
temporarily bind to reactants, hold them in a favorable
position to react with each other, and may change the
shapes of reactants in ways that make them more
likely to react. By reducing the element of chance in
molecular collisions, a catalyst speeds up a reaction. It
then releases the products and is available to repeat
the process with more reactants. The catalyst itself is
not permanently consumed or changed by the
reaction. The most important biological catalysts are
enzymes, discussed later in this chapter.
All the chemical reactions in the body are collectively called metabolism. Metabolism has two divisions—catabolism and anabolism. Catabolism (ca-TAB-oh-lizm) consists of energy-releasing decomposition reactions. Such reactions break covalent bonds, produce smaller molecules from larger ones, and release energy that can be used for other physiological work. Energy-releasing reactions are called exergonic11 reactions. If you hold a beaker of water in your hand and pour sulfuric acid into it, for example, the beaker will get so hot you may have to put it down. If you break down energy-storage molecules to run a race, you too will get hot. In both cases, the heat signifies that exergonic reactions are occurring. Anabolism (ah-NAB-oh-lizm) consists of energystoring synthesis reactions, such as the production of protein or fat. Reactions that require an energy input, such as these, are called endergonic13 reactions. Anabolism is driven by the energy that catabolism releases, so endergonic and exergonic processes, anabolism and catabolism, are inseparably linked. Oxidation is any chemical reaction in which a molecule gives up electrons and releases energy. A molecule is oxidized by this process, and whatever molecule takes the electrons from it is an oxidizing agent (electron acceptor). The term oxidation stems from the fact that oxygen is often involved as the electron acceptor. Thus, we can sometimes recognize an oxidation reaction from the fact that oxygen has been added to a molecule. The rusting of iron, for example, is a slow oxidation process in which oxygen is added to iron to form iron oxide (Fe2O3). Many oxidation reactions, however, do not involve oxygen at all. For example, when yeast ferments glucose to alcohol, no oxygen is required; indeed, the alcohol contains less oxygen than the sugar originally did, but it is more oxidized than the sugar: C6H12O6 → 2 CH3CH2OH + 2 CO2 Glucose Ethanol Carbon dioxide Reduction is a chemical reaction in which a molecule gains electrons and energy. When a molecule accepts electrons, it is said to be reduced; a molecule that donates electrons to another is therefore called a reducing agent (electron donor). The oxidation of one molecule is always accompanied by the reduction of another, so these electron transfers are known as oxidation-reduction (redox) reactions. It is not necessary that only electrons be transferred in a redox reaction. Often, the electrons are transferred in the form of hydrogen atoms. The fact that a proton (the hydrogen nucleus) is also transferred is immaterial to whether we consider a reaction oxidation or reduction. Table 2.5 summarizes these energy transfer reactions. We can symbolize oxidation and reduction as follows, letting A and B symbolize arbitrary molecules and e represent one or more electrons: Ae B → A Be High-energy Low-energy Low-energy High-energy reduced oxidized oxidized reduced state state state state Ae is a reducing agent because it reduces B, and B is an oxidizing agent because it oxidizes Ae.
When you have completed this section, you should be able to
• explain why carbon is especially well suited to serve as the
structural foundation of many biological molecules;
• identify some common functional groups of organic
molecules from their formulae;
• discuss the relevance of polymers to biology and explain how
they are formed and broken by dehydration synthesis and
hydrolysis;
• discuss the types and functions of carbohydrates;
• discuss the types and functions of lipids;
• discuss protein structure and function;
• explain how enzymes function;
• describe the structure, production, and function of ATP;
• identify other nucleotide types and their functions;
• identify the principal types of nucleic acids.
Organic chemistry is the study of compounds of carbon.
By 1900, biochemists had classified the organic molecules
of life into four primary categories: carbohydrates, lipids,
Exergonic Reactions Reactions in which there is a net release of energy. The products have less total free energy than the reactants did.
Oxidation An exergonic reaction in which electrons are removed from a reactant. Electrons may be removed one or two at a time
and may be removed in the form of hydrogen atoms (H or H2). The product is then said to be oxidized.
Decomposition A reaction such as digestion and cell respiration, in which larger molecules are broken down into smaller ones.
Catabolism The sum of all decomposition reactions in the body.
Endergonic Reactions Reactions in which there is a net input of energy. The products have more total free energy than the reactants did.
Reduction An endergonic reaction in which electrons are donated to a reactant. The product is then said to be reduced.
Synthesis A reaction such as protein and glycogen synthesis, in which two or more smaller molecules are combined into a larger one.
Anabolism The sum of all synthesis reactions in the body
proteins, and nucleic acids. We examine the first three in
this chapter but describe the details of nucleic acids,
which are concerned with genetics, in chapter 4.
Carbon is an especially versatile atom that serves as
the basis of a wide variety of structures. It has four valence
electrons, so it bonds with other atoms that can provide it
with four more to complete its valence shell. Carbon atoms
readily bond with each other and can form long chains,
branched molecules, and rings—an enormous variety of
carbon backbones for organic molecules. Carbon also
forms covalent bonds with hydrogen, oxygen, nitrogen,
sulfur, and other elements.
Carbon backbones carry a variety of functional
groups—small clusters of atoms that determine many of
the properties of an organic molecule. For example,
organic acids bear a carboxyl (car-BOC-sil) group, and
ATP is named for its three phosphate groups. Other common functional groups include hydroxyl, methyl, and
amino groups
Since carbon can form long chains, some organic molecules are gigantic macromolecules with molecular weights that range from the thousands (as in starch and proteins) to the millions (as in DNA). Most macromolecules are polymers14—molecules made of a repetitive series of identical or similar subunits called monomers (MON-ohmurs). Starch, for example, is a polymer of about 3,000 glucose monomers. In starch, the monomers are identical, while in other polymers they have a basic structural similarity but differ in detail. DNA, for example, is made of 4 different kinds of monomers (nucleotides), and proteins are made of 20 kinds (amino acids). The joining of monomers to form a polymer is called polymerization. Living cells achieve this by means of a reaction called dehydration synthesis (condensation) (fig. 2.15a). A hydroxyl ( OH) group is removed from one monomer and a hydrogen ( H) from another, producing water as a by-product. The two monomers become joined by a covalent bond, forming a dimer. This is repeated for each monomer added to the chain, potentially leading to a chain long enough to be considered a polymer. The opposite of dehydration synthesis is hydrolysis15 (fig. 2.15b). In hydrolysis, a water molecule ionizes into OH and H. A covalent bond linking one monomer to another is broken, the OH is added to one monomer, and the H is added to the other one. All digestion consists of hydrolysis reactions.
A carbohydrate16 is a hydrophilic organic molecule with the general formula (CH2O)n, where n represents the number of carbon atoms. In glucose, for example, n 6 and the formula is C6H12O6. As the generic formula shows, carbohydrates have a 2:1 ratio of hydrogen to oxygen. The names of individual carbohydrates are often built on the word root sacchar- or the suffix -ose, both of which mean “sugar” or “sweet.” The most familiar carbohydrates are sugars and starches.The simplest carbohydrates are called monosaccharides17 (MON-oh-SAC-uh-rides), or simple sugars. The three of primary importance are glucose, fructose, and galactose, all with the molecular formula C6H12O6; they are isomers of each other (fig. 2.16). We obtain these sugars mainly by the digestion of more complex carbohydrates. Glucose is the “blood sugar” that provides energy to most of our cells. Two other monosaccharides, ribose and deoxyribose, are important components of DNA and RNA. Disaccharides are sugars composed of two monosaccharides. The three of greatest importance are sucrose (made of glucose fructose), lactose (glucose galactose), and maltose (glucose glucose) Sucrose is produced by sugarcane and sugar beets and used as common table sugar. Lactose is milk sugar. Maltose is a product of starch digestion and is present in a few foods such as germinating wheat and malt beverages. Polysaccharides (POL-ee-SAC-uh-rides) are long chains of glucose. Some polysaccharides have molecular weights of 500,000 or more (compared to 180 for a single glucose). Three polysaccharides of interest to human physiology are glycogen, starch, and cellulose. Animals, including ourselves, make glycogen, while starch and cellulose are plant products. Glycogen18 is an energy-storage polysaccharide made by cells of the liver, muscles, uterus, and vagina. It is a long branched glucose polymer (fig. 2.18). The liver produces glycogen after a meal, when the blood glucose level is high, and then breaks it down between meals to maintain blood glucose levels when there is no food intake. Muscle stores glycogen for its own energy needs, and the uterus uses it in pregnancy to nourish the embryo. Starch is the corresponding energy-storage polysaccharide of plants. They store it when sunlight and nutrients are available and draw from it when photosynthesis is not possible (for example, at night and in winter, when a plant has shed its leaves). Starch is the only significant digestible dietary polysaccharide. Cellulose is a structural polysaccharide that gives strength to the cell walls of plants. It is the principal component of wood, cotton, and paper. It consists of a few thousand glucose monomers joined together, with every other monomer “upside down” relative to the next. (The —CH2OH groups all face in the same direction in glycogen and starch, but alternate between facing up and down in cellulose.) Cellulose is the most abundant organic compound on earth and it is a common component of the diets of humans and other animals—yet we have no enzymes to digest it and thus derive no energy or nutrition from it. Nevertheless, it is important as dietary “fiber,” “bulk,” or “roughage.” It swells with water in the digestive tract and helps move other materials through the intestine. Carbohydrates are, above all, a source of energy that can be quickly mobilized. All digested carbohydrate is ultimately converted to glucose, and glucose is oxidized to make ATP, a high-energy compound discussed later. But carbohydrates have other functions as well (table 2.6). They are often conjugated19 with (covalently bound to) proteins and lipids. Many of the lipid and protein molecules at the external surface of the cell membrane have chains of up to 12 sugars attached to them, thus forming glycolipids and glycoproteins, respectively. Among other functions, glycoproteins are a major component of mucus, which traps particles in the respiratory system, resists infection, and protects the digestive tract from its own acid and enzymes.Proteoglycans (once called mucopolysaccharides) are macromolecules in which the carbohydrate component is dominant and a peptide or protein forms a smaller component. Proteoglycans form gels that help hold cells and tissues together, form a gelatinous filler in the umbilical cord and eye, lubricate the joints of the skeletal system, and account for the tough rubbery texture of cartilage. Their structure and functions are further considered in chapter 5. When discussing conjugated macromolecules it is convenient to refer to each chemically different component as a moiety20 (MOY-eh-tee). Proteoglycans have a protein moiety and a carbohydrate moiety, for example.
A lipid is a hydrophobic organic molecule, usually composed only of carbon, hydrogen, and oxygen, with a high ratio of hydrogen to oxygen. A fat called tristearin (triSTEE-uh-rin), for example, has the molecular formula C57H110O6—more than 18 hydrogens for every oxygen. Lipids are less oxidized than carbohydrates, and thus have more calories per gram. Beyond these criteria, it is difficult to generalize about lipids; they are much more variable in structure than the other macromolecules we are considering. We consider the five primary types of lipids in humans—fatty acids, triglycerides, phospholipids, eicosanoids, and steroids (table 2.7). A fatty acid is a chain of usually 4 to 24 carbon atoms with a carboxyl group at one end and a methyl group at the other. Fatty acids and the fats made from them are classified as saturated or unsaturated. A saturated fatty acid such as palmitic acid has as much hydrogen as it can carry. No more could be added without exceeding four covalent bonds per carbon atom; thus it is “saturated” with hydrogen. In unsaturated fatty acids such as linoleic acid, however, some carbon atoms are joined by double covalent bonds (fig. 2.19). Each of these could potentially share one pair of electrons with another hydrogen atom instead of the adjacent carbon, so hydrogen could be added to this molecule. Polyunsaturated fatty acids are those with many C C bonds. Most fatty acids can be synthesized by the human body, but a few, called essential fatty acids, must be obtained from the diet because we cannot synthesize them (see chapter 26). A triglyceride (try-GLISS-ur-ide) is a molecule consisting of three fatty acids covalently bonded to a threecarbon alcohol called glycerol; triglycerides are more correctly, although less widely, also known as triacylglycerols. Each bond between a fatty acid and glycerol is formed by dehydration synthesis (see fig. 2.19). Once joined to glycerol, a fatty acid can no longer donate a proton to solution and is therefore no longer an acid. For this reason, triglycerides are also called neutral fats. Triglycerides are broken down by hydrolysis reactions, which split each of these bonds apart by the addition of water. Triglycerides that are liquid at room temperature are also called oils, but the difference between a fat and oil is fairly arbitrary. Coconut oil, for example, is solid at room temperature. Animal fats are usually made of saturated fatty acids, so they are called saturated fats. They are solid at room or body temperature. Most plant triglycerides are polyunsaturated fats, which generally remain liquid at room temperature. Examples include peanut, olive, corn, and linseed oils. Saturated fats contribute more to cardiovascular disease than unsaturated fats, and for this reason it is healthier to cook with vegetable oils than with lard or bacon fat.The primary function of fat is energy storage, but when concentrated in adipose tissue, it also provides thermal insulation and acts as a shock-absorbing cushion for vital organs (see chapter 5). Phospholipids are similar to neutral fats except that, in place of one fatty acid, they have a phosphate group which, in turn, is linked to other functional groups. Lecithin is a common phospholipid in which the phosphate is bonded to a nitrogenous group called choline (COleen) (fig. 2.20). Phospholipids have a dual nature. The two fatty acid “tails” of the molecule are hydrophobic, but the phosphate “head” is hydrophilic. Thus, phospholipids are said to be amphiphilic21 (AM-fih-FIL-ic). Together, the head and the two tails of a phospholipid give it a shape like a clothespin. The most important function of phospholipids is to serve as the structural foundation of cell membranes (see chapter 3). Eicosanoids22 (eye-CO-sah-noyds) are 20-carbon compounds derived from a fatty acid called arachidonic (ah-RACK-ih-DON-ic) acid. They function primarily as hormonelike chemical signals between cells. The most functionally diverse eicosanoids are the prostaglandins, in which five of the carbon atoms are arranged in a ring (fig. 2.21). They were originally found in the secretions of bovine prostate glands, hence their name, but they are now known to be produced in almost all tissues. They play a variety of signaling roles in inflammation, blood clotting, hormone action, labor contractions, control of blood vessel diameter, and other processes (see chapter 17). A steroid is a lipid with 17 of its carbon atoms arranged in four rings (fig. 2.22). Cholesterol is the “parent” steroid from which the other steroids are synthesized. The others include cortisol, progesterone, estrogens, testosterone, and bile acids. These differ from each other in the location of C C bonds within the rings and in the functional groups attached to the rings. Cholesterol is synthesized only by animals (especially in liver cells) and is not present in vegetable oils or other plant products. The average adult contains over 200 g (half a pound) of cholesterol. Cholesterol has a bad reputation as a factor in cardiovascular disease (see insight 2.3), and it is true that hereditary and dietary factors can elevate blood cholesterol to dangerously high levels. Nevertheless, cholesterol is a natural product of the body. Only about 15% of our cholesterol comes from the diet; the other 85% is internally synthesized. In addition to being the precursor of the other steroids, cholesterol is an important component of cell membranes and is required for proper nervous system function.
There is only one kind of cholesterol, and it does far more good than harm. When the popular press refers to “good” and “bad” cholesterol, it is actually referring to droplets in the blood called lipoproteins, which are a complex of cholesterol, fat, phospholipids, and protein. So-called bad cholesterol refers to low-density lipoprotein (LDL), which has a high ratio of lipid to protein and contributes to cardiovascular disease. So-called good cholesterol refers to high-density lipoprotein (HDL), which has a lower ratio of lipid to protein and may help to prevent cardiovascular disease. Even when food products are advertised as cholesterol-free, they may be high in saturated fat, which stimulates the body to produce more cholesterol. Palmitic acid seems to be the greatest culprit in stimulating elevated cholesterol levels, while linoleic acid has a cholesterol-lowering effect. Both are shown in figure 2.19. Cardiovascular disease is further discussed at the end of chapter 19, and LDLs and HDLs are more fully explained in
The word protein is derived from the Greek word proteios, meaning “of first importance.” Proteins are the most versatile molecules in the body, and many discussions in this book will draw on your understanding of protein structure and behavior.
A protein is a polymer of amino acids. An amino acid has a central carbon atom with an amino ( NH2) and a carboxyl ( COOH) group bound to it (fig. 2.23a). The 20 amino acids used to make proteins are identical except for a third functional group called the radical (R group) attached to the central carbon. In the simplest amino acid, glycine, R is merely a hydrogen atom, while in the largest amino acids it includes rings of carbon. Some R groups are hydrophilic and some are hydrophobic. Being composed of many amino acids, proteins as a whole are therefore often amphiphilic. The 20 amino acids involved in proteins are listed in table 2.8 along with their abbreviations. A peptide is any molecule composed of two or more amino acids joined by peptide bonds. A peptide bond, formed by dehydration synthesis, joins the amino group of one amino acid to the carboxyl group of the next (fig. 2.23b). Peptides are named for the number of amino acids they have—for example, dipeptides have two and tripeptides have three. Chains of fewer than 10 or 15 amino acids are called oligopeptides,23 and chains larger than that are called polypeptides. An example of an oligopeptide is the childbirth-inducing hormone oxytocin, composed of 9 amino acids. A representative polypeptide is adrenocorticotropic hormone (ACTH), which is 39 amino acids long. A protein is a polypeptide of 50 amino acids or more. A typical amino acid has a molecular weight of about 80 amu,and the molecular weights of the smallest proteins are around 4,000 to 8,000 amu. The average protein weighs in at about 30,000 amu, and some of them have molecular weights in the hundreds of thousands.
Proteins have complex coiled and folded structures that are critically important to the roles they play. Even slight changes in their conformation (three-dimensional shape) can destroy protein function. Protein molecules have three to four levels of complexity, from primary through quaternary structure (fig. 2.24). Primary structure is the protein’s sequence of amino acids. Their order is encoded in the genes (see chapter 4). Secondary structure is a coiled or folded shape held together by hydrogen bonds between the slightly negative C O group of one peptide bond and the slightly positive N H group of another peptide bond some distance away. The most common secondary structures are a springlike shape called the helix and a pleated, ribbonlike shape, the sheet (or -pleated sheet). Many proteins have multiple -helical and -pleated regions joined by short segments with a less orderly geometry. A single protein molecule may fold back on itself and have two or more -pleated regions linked to each other by hydrogen bonds. Separate, parallel protein molecules also may be hydrogenbonded to each other through their -pleated regions. Tertiary24 (TUR-she-air-ee) structure is formed by the further bending and folding of proteins into various globular and fibrous shapes. It results from hydrophobic R groups associating with each other and avoiding water, while the hydrophilic R groups are attracted to the surrounding water. Globular proteins, somewhat resembling a wadded ball of yarn, have a compact tertiary structure well suited for proteins embedded in cell membranes and proteins that must move around freely in the body fluids, such as enzymes and antibodies. Fibrous proteins such as myosin, keratin, and collagen are slender filaments better suited for such roles as muscle contraction and providing strength to skin, hair, and tendons. The amino acid cysteine (Cys), whose R group is CH2 SH (see fig. 2.23), often stabilizes a protein’s tertiary structure by forming covalent disulfide bridges. When two cysteines align with each other, each can release a hydrogen atom, leaving the sulfur atoms to form a disulfide ( S S ) bridge. Disulfide bridges hold separate polypeptide chains together in such molecules as antibodies and insulin (fig. 2.25). Quaternary25 (QUA-tur-nare-ee) structure is the association of two or more polypeptide chains by noncovalent forces such as ionic bonds and hydrophilic-hydrophobic interactions. It occurs in only some proteins. Hemoglobin for example, consists of four polypeptides—two identical chains and two identical, slightly longer chains (see fig. 2.24). One of the most important properties of proteins is their ability to change conformation, especially tertiary structure. This can be triggered by such influences as voltage changes on a cell membrane during the action of nerve cells, the binding of a hormone to a protein, or the dissociation of a molecule from a protein. Subtle, reversible changes in conformation are important to processes such as enzyme function, muscle contraction, and the opening and closing of pores in cell membranes. Denaturation is a more drastic conformational change in response to conditions such as extreme heat or pH. It is seen, for example, when you cook an egg and the egg white protein (albumen) turns from clear to opaque. Denaturation is sometimes reversible, but often it permanently destroys protein function. Conjugated proteins have a non-amino-acid moiety called a prosthetic26 group covalently bound to them. Hemoglobin, for example, not only has the four polypeptide chains described earlier, but each chain also has a complex iron-containing ring called a heme moiety attached to it (see fig. 2.24). Hemoglobin cannot transport oxygen unless this group is present. In glycoproteins, as described earlier, the carbohydrate moiety is a prosthetic group
Proteins have more diverse functions than other macromolecules. These include:
• Structure. Keratin, a tough structural protein, gives
strength to the nails, hair, and skin surface. Deeper
layers of the skin, as well as bones, cartilage, and teeth,
contain an abundance of the durable protein collagen.
• Communication. Some hormones and other cell-tocell signals are proteins, as are the receptors to which
the signal molecules bind in the receiving cell. A
hormone or other molecule that reversibly binds to a
protein is called a ligand27 (LIG-and).
• Membrane transport. Some proteins form channels in
cell membranes that govern what passes through the
membranes and when. Other proteins act as carriers that
briefly bind to solute particles and transport them to the
other side of the membrane. Among their other roles,
such proteins turn nerve and muscle activity on and off.
• Catalysis. Most metabolic pathways of the body are
controlled by enzymes, which are globular proteins
that function as catalysts.
• Recognition and protection. The role of glycoproteins
in immune recognition was mentioned earlier.
Antibodies and other proteins attack and neutralize
organisms that invade the body. Clotting proteins
protect the body against blood loss.
• Movement. Movement is fundamental to all life, from
the intracellular transport of molecules to the
galloping of a racehorse. Proteins, with their special
ability to change shape repeatedly, are the basis for all
such movement. Some proteins are called molecular
motors for this reason
• Cell adhesion. Proteins bind cells to each other, which
enables sperm to fertilize eggs, enables immune cells
to bind to enemy cancer cells, and keeps tissues from
falling apart
Enzymes are proteins that function as biological catalysts. They permit biochemical reactions to occur rapidly at normal body temperatures. Enzymes were initially given somewhat arbitrary names, still with us, such as pepsin and trypsin. The modern system of naming enzymes, however, is more uniform and informative. It identifies the substance the enzyme acts upon, called its substrate; sometimes refers to the enzyme’s action; and adds the suffix -ase. Thus, amylase digests starch (amyl- starch) and carbonic anhydrase removes water (anhydr-) from carbonic acid. Enzyme names may be further modified to distinguish different forms of the same enzyme found in different tissues
A given enzyme may exist in slightly different forms, called isoenzymes, in different cells. Isoenzymes catalyze the same chemical reactions but have enough structural differences that they can be distinguished by standard laboratory techniques. This is useful in the diagnosis of disease. When organs are diseased, some of their cells break down and release specific isoenzymes that can be detected in the blood. Normally, these isoenzymes would not be present in the blood or would have very low concentrations. If their blood levels are elevated, it can help pinpoint what cells in the body have been damaged. For example, creatine kinase (CK) occurs in different forms in different cells. An elevated serum level of CK-1 indicates a breakdown of skeletal muscle and is one of the signs of muscular dystrophy. An elevated CK-2 level indicates heart disease, because this isoenzyme comes only from cardiac muscle. There are five isoenzymes of lactate dehydrogenase (LDH). High serum levels of LDH-1 may indicate a tumor of the ovaries or testes, while LDH-5 may indicate liver disease or muscular dystrophy. Different isoenzymes of phosphatase in the blood may indicate bone or prostate disease.
To appreciate the effect of an enzyme, think of what happens when paper burns. Paper is composed mainly of glucose (in the form of cellulose). The burning of glucose can be represented by the equation C6H12O6 6 O2 → 6 CO2 6 H2O Paper does not spontaneously burst into flame, because few of its molecules have enough kinetic energy to react. Lighting the paper with a match, however, raises the kinetic energy enough to initiate combustion (rapid oxidation). The energy needed to get the reaction started, supplied by the match, is called activation energy (fig. 2.26a). In the body, we carry out the same reaction and oxidize glucose to water and carbon dioxide to extract its energy. We could not tolerate the heat of combustion in our bodies, however, so we must oxidize glucose in a more controlled way at a biologically feasible and safe temperature. Enzymes make this happen by lowering the activation energy—that is, by reducing the barrier to glucose oxidation (fig. 2.26b)—and by releasing the energy in small steps rather than a single burst of heat.
Substrates bind to pockets called active sites in the enzyme surface and create a temporary enzyme-substrate complex. The enzyme may break covalent bonds and convert the substrate to a reaction product, or it may hold two or more substrates close together, in adjacent active sites, thus enabling the substrates to react with each other (fig. 2.27). The enzyme then releases the reaction products and is free to begin the process again. Since enzymes are not consumed by the reactions they catalyze, one enzyme molecule can convert millions of substrate molecules, and at astonishing speeds. A single molecule of carbonic anhydrase, for example, breaks carbonic acid (H2CO3) down to H2O and CO2 at a rate of 36 million molecules per minute. A substrate fits an enzyme somewhat like a key fits a lock. A given enzyme is very selective—that is, it exhibits enzyme-substrate specificity. An enzyme that oxidizes glucose, for example, will not act on the similar sugar galactose, which does not fit its active site. Factors that change the shape of an enzyme—notably temperature and pH—tend to alter or destroy the ability of the enzyme to bind its substrate. They disrupt the hydrogen bonds and other weak forces that hold the enzyme in its proper conformation, essentially changing the shape of the “lock” (active site) so that the “key” (substrate) no longer fits. Enzymes vary in optimum pH according to where in the body they normally function. Thus salivary amylase, which digests starch in the mouth, functions best at pH 7 and is inactivated when it is exposed to stomach acid; pepsin, which works in the acidic environment of the stomach, functions best around pH 2; and trypsin, a digestive enzyme that works in the alkaline environment of the small intestine, has an optimum pH of 9.5. Our internal body temperature is nearly the same everywhere, however, and all human enzymes have a temperature optimum (that is, they produce their fastest reaction rates) near 37°C.
Many enzymes cannot function without nonprotein partners called cofactors—for example, iron, copper, zinc, magnesium, or calcium ions. By binding to an enzyme, a cofactor may stimulate it to fold into a shape that activates its active site. Coenzymes are organic cofactors usually derived from niacin, riboflavin, and other water-soluble vitamins. They accept electrons from an enzyme in one metabolic pathway and transfer them to an enzyme in another pathway. For example, cells partially oxidize glucose through a pathway called glycolysis. A coenzyme called NAD, 28 derived from niacin, shuttles electrons from this pathway to another one called aerobic respiration, which uses energy from the electrons to make ATP (fig. 2.28). If NAD is unavailable, the glycolysis pathway shuts down
A metabolic pathway is a chain of reactions with each step usually catalyzed by a different enzyme. A simple metabolic pathway can be symbolized A → B → C → D where A is the initial reactant, B and C are intermediates, and D is the end product. The Greek letters above the reaction arrows represent enzymes that catalyze each step of the reaction. A is the substrate for enzyme , B is the substrate for enzyme , and C for enzyme . Such a pathway can be turned on or off by altering the conformation of any of these enzymes, thereby activating or deactivating them. This can be done by such means as the binding or dissociation of a cofactor, or by an end product of the pathway binding to an enzyme at an earlier step (product D binding to enzyme and shutting off the reaction chain at that step, for example). In these and other ways, cells are able to turn on metabolic pathways when their end products are needed and shut them down when the end products are not needed.
Nucleotides are organic compounds with three principal components—a single or double carbon-nitrogen ring called a nitrogenous base, a monosaccharide, and one or more phosphate groups. One of the best-known nucleotides is ATP (fig. 2.29a), in which the nitrogenous base is a double ring called adenine, the sugar is ribose, and there are three phosphate groups
Adenosine triphosphate (ATP) is the body’s most important energy-transfer molecule. It briefly stores energy gained from exergonic reactions such as glucose oxidation and releases it within seconds for physiological work such as polymerization reactions, muscle contraction, and pumping ions through cell membranes. The second and third phosphate groups of ATP are attached to the rest of the molecule by high-energy covalent bonds traditionally indicated by a wavy line in the molecular formula. Since phosphate groups are negatively charged, they repel each other. It requires a high-energy bond to overcome that repulsive force and hold them together—especially to add the third phosphate group to a chain that already has two negatively charged phosphates. Most energy transfers to and from ATP involve adding or removing that third phosphate. Enzymes called adenosine triphosphatases (ATPases) are specialized to hydrolyze the third high-energy phosphate bond, producing adenosine diphosphate (ADP) and an inorganic phosphate group (Pi). This reaction releases 7.3 kilocalories of energy for every mole (505 g) of ATP. Most of this energy escapes as heat, but we live on the portion of it that does useful work. We can summarize this as follows:ATP H2O → ADP Pi Energy Work The free phosphate groups released by ATP hydrolysis are often added to enzymes or other molecules to activate them. This addition of Pi, called phosphorylation, is carried out by enzymes called kinases (phosphokinases). The phosphorylation of an enzyme is sometimes the “switch” that turns a metabolic pathway on or off. ATP is a short-lived molecule, usually consumed within 60 seconds of its formation. The entire amount in the body would support life for less than 1 minute if it were not continually replenished. At a moderate rate of physical activity, a full day’s supply of ATP would weigh twice as much as you do. Even if you never got out of bed, you would need about 45 kg (99 lb) of ATP to stay alive for a day. The reason cyanide is so lethal is that it halts ATP synthesis. ATP synthesis is explained in detail in chapter 26, but you will find it necessary to understand the general idea of it before you reach that chapter—especially in understanding muscle physiology (chapter 11). Much of the energy for ATP synthesis comes from glucose oxidation (fig. 2.30). The first stage in glucose oxidation (fig. 2.31) is the reaction pathway known as glycolysis (gly-COLL-ih-sis). This literally means “sugar splitting,” and indeed its major effect is to split the six-carbon glucose molecule into two three-carbon molecules of pyruvic acid. A little ATP is produced in this stage (a net yield of two ATPs per glucose), but most of the chemical energy of the glucose is still in the pyruvic acid. What happens to pyruvic acid depends on whether or not oxygen is available. If not, pyruvic acid is converted to lactic acid by a pathway called anaerobic29 (AN-errOH-bic) fermentation. This pathway has two noteworthy disadvantages: First, it does not extract any more energy from pyruvic acid; second, the lactic acid it produces is toxic, so most cells can use anaerobic fermentation only as a temporary measure. The only advantage to this pathway is that it enables glycolysis to continue (for reasons explained in chapter 26) and thus enables a cell to continue producing a small amount of ATP. If oxygen is available, a more efficient pathway called aerobic respiration occurs. This breaks pyruvic acid down to carbon dioxide and water and generates up to 36 more molecules of ATP for each of the original glucose molecules. The reactions of aerobic respiration are carried out in the cell’s mitochondria (described in chapter 3), so mitochondria are regarded as a cell’s principal “ATP factories.”
Guanosine (GWAH-no-seen) triphosphate (GTP) is another nucleotide involved in energy transfers. In some reactions, it donates phosphate groups to other molecules. In some pathways, it donates its third phosphate group to ADP to regenerate ATP. Cyclic adenosine monophosphate (cAMP) (see fig. 2.29b) is a nucleotide formed by the removal of both the second and third phosphate groups from ATP. In some cases, when a hormone or other chemical signal (“first messenger”) binds to a cell surface, it triggers an internal reaction that converts ATP to cAMP. The cAMP then acts as a “second messenger” to activate metabolic effects within the cell.
Nucleic (new-CLAY-ic) acids are polymers of nucleotides. The largest of them, deoxyribonucleic acid (DNA), is typically 100 million to 1 billion nucleotides long. It constitutes our genes, gives instructions for synthesizing all of the body’s proteins, and transfers hereditary information from cell to cell when cells divide and from generation to generation when organisms reproduce. Three forms of ribonucleic acid (RNA), which range from 70 to 10,000 nucleotides long, carry out those instructions and synthesize the proteins, assembling amino acids in the right order to produce each protein “described” by the DNA. The detailed structure of DNA and RNA and the mechanisms of protein synthesis and heredity are described in chapter 4
The sex hormone testosterone stimulates muscular growth and aggressive behavior, especially in males. In Nazi Germany, testosterone was given to SS troops in an effort to make them more aggressive, but with no proven success. In the 1950s, pharmaceutical companies developed compounds related to testosterone, called anabolic-androgenic30 steroids, to treat anemia, breast cancer, osteoporosis, and some muscle diseases, and to prevent the shrinkage of muscles in immobilized patients. By the early 1960s, athletes were using anabolic-androgenic steroids to stimulate muscle growth, accelerate the repair of tissues damaged in training or competition, and stimulate the aggressiveness needed to excel in some contact sports such as football and boxing. The doses used by athletes, however, are 10 to 1,000 times higher than the doses prescribed for medical purposes, and they can have devastating effects on one’s health. They raise cholesterol levels, which promotes fatty degeneration of the arteries (atherosclerosis). This can lead to coronary artery disease, heart and kidney failure, and stroke. Deteriorating blood circulation also sometimes results in gangrene, which may require amputation of the extremities. As the liver attempts to dispose of the high concentration of steroids, liver cancer and other liver diseases may ensue. In addition, steroids suppress the immune system, so the user is more subject to infection and cancer. They cause a premature end to bone elongation, so people who use anabolic steroids in adolescence may never attain normal adult height. Anabolic-androgenic steroids have the same effect on the brain as natural testosterone. Thus, when steroid levels are high, the brain and pituitary gland stop producing the hormones that stimulate sperm production and testosterone secretion. In men, this leads to atrophy of the testes, impotence (inability to achieve or maintain an erection), low sperm count, and infertility. Ironically, anabolic-androgenic steroids have feminizing effects on men and masculinizing effects on women. Men may develop enlarged breasts (gynecomastia), while in some female users the breasts and uterus atrophy, the clitoris enlarges, and ovulation and menstruation become irregular. Female users may develop excessive facial and body hair and a deeper voice, and both sexes show an increased tendency toward baldness. Especially in men, steroid abuse can be linked to severe emotional disorders. The steroids themselves stimulate heightened aggressiveness and unpredictable mood swings, so the abuser may vacillate between depression and violence. It surely doesn’t help matters that impotence, shrinkage of the testes, infertility, and enlargement of the breasts are so incongruous with the self-image of a male athlete who abuses steroids. Partly because of the well documented adverse health effects, the use of anabolic-androgenic steroids has been condemned by the American Medical Association and American College of Sports Medicine and banned by the International Olympic Committee, National Football League, and National Collegiate Athletic Association. But in spite of such warnings and bans, many athletes continue to use steroids and related performance-enhancing drugs, which remain available through unscrupulous coaches, physicians, Internet sources, and foreign mailorder suppliers. By some estimates, as many as 80% of weight lifters, 30% of college and professional athletes, and 20% of male high-school athletes now use anabolic-androgenic steroids. The National Institutes of Health finds increasing usage among high school students in recent years, and increasing denial that anabolic-androgenic steroids present a significant health hazard.
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