Audesirk Biology: Life on Earth, 6/e Chapter 2

Audesirk Biology: Life on Earth, 6/e Chapter 2

U N I T
O N E
The Life of a Cell
Single cells can be complex, independent organisms, such as these two
ciliates, of the kingdom Protista. A large Euplotes (about 300 mm in length)
prepares to eat a much smaller Paramecium. Both are covered with cilia,
short, beating structures used to move and ingest prey.
“If there is magic on this planet, it is contained in water.”
Loren Eiseley in The Immense Journey (1957)
Could something this delicious possibly be good for you? Chocolate, made from seeds found
inside cacao pods (inset), contains high levels of antioxidants.
2
Atoms, Molecules,
and Life
AT A G L A N C E
Case Study: Health Food?
1 What Are Atoms?
Hydrogen Bonds Are Weaker Electrical Attractions
Between or Within Molecules with Polar Covalent Bonds
Atoms, the Basic Structural Units of Matter, Are
Composed of Still Smaller Particles
2 How Do Atoms Interact to Form Molecules?
Atoms Will Interact with Other Atoms Only When There
Are Vacancies in Their Outermost Electron Shells
Charged Atoms Called Ions Interact to Form Ionic Bonds
Uncharged Atoms Can Become Stable by Sharing
Electrons, Forming Covalent Bonds
3 Why Is Water So Important to Life?
Water Interacts with Many Other Molecules
Water Molecules Tend to Stick Together
Water Can Form H+ and OH– Ions
Water Moderates the Effects of Temperature Changes
Water Forms an Unusual Solid: Ice
Case Study Revisited: Health Food?
CASESTUDYCASESTUDYCASESTUDYCASESTUDYCASESTUDYC
W
Health Food?
hen 35 million people in the United
States gave their loved ones boxes of
chocolates last Valentine’s Day, they knew
they were giving sweet comfort—but health
food? Sometimes described as “sinfully delicious,” chocolate has often been a source of
guilt for those who indulge (or overindulge) in
it. Chocolate candy is certainly a significant
source of fat and sugar calories, but recent
research suggests that chocolate itself—the
dark, bitter powder made from the seeds within cacao pods (see inset) may also be a significant source of protective molecules. Medical
scientists have known for some time that
many things that go wrong with our bodies
can be traced to destructive molecules called
free radicals. Many free radicals contain oxygen in a form that reacts strongly with, and
damages, various biological molecules and
their cellular structures. This process is called
a What Are Atoms?
Atoms, the Basic Structural Units of Matter,
Are Composed of Still Smaller Particles
If you took a diamond (a form of carbon) and cut it into
pieces, each piece would still be carbon. If you could
make finer and finer divisions, you would eventually
produce a pile of carbon atoms. Atoms are the fundamental structural units of matter. Atoms themselves,
however, are composed of a central atomic nucleus
(often called simply the nucleus; plural, nuclei, but don’t
confuse it with the nucleus of a cell). The nucleus con-
oxidative stress. Oxidative stress tears up cell
membranes, breaks down DNA, and destroys
enzymes, resulting in many aspects of aging,
cancer, heart disease, and nervous system
disorders. Unfortunately, oxidative stress is a
fact of life, because as our cells use energy,
they naturally produce free radicals. So
what’s a person to do? Well, maybe eat
chocolate! n
tains two types of subatomic particles of equal weight:
positively charged protons and uncharged neutrons.
Subatomic particles called electrons orbit the atomic nucleus (Fig. 2-1). Electrons are lighter, negatively charged
particles. An atom by itself has an equal number of electrons and protons and is therefore electrically neutral.
There are 92 types of atoms that occur naturally. Each
type of atom forms the structural unit of a different element. An element is a substance that can neither be broken down nor converted to other substances by ordinary
chemical means. The number of protons in the nucleus,
called the atomic number, is a characteristic of each element. For example, every hydrogen atom has one proton
in its nucleus, every carbon atom has six protons, and
21
2.1 Interactive Atoms
MEDIATUTOR
22
Chapter 2
Atoms, Molecules, and Life
every oxygen atom has eight. Each element has unique
chemical properties based on the number and configuration of its subatomic particles. Some elements, such as
oxygen and hydrogen, are gases at room temperature;
others, such as lead, are extremely dense solids. Most elements are quite rare, and relatively few are essential to
life on Earth. Table 2-1 lists the most common elements
in the universe, on Earth, and in the human body. Notice
how differently these elements are distributed.
Atoms of the same element may have different numbers of neutrons; when this occurs, the atoms are called
isotopes of each other. Some, but not all, isotopes are radioactive; that is, they spontaneously break apart, forming different types of atoms and releasing energy in the
process. Radioactive isotopes are extremely useful as
“labels” in studying biological processes (see “Scientific
Inquiry: Radioactivity in Research”).
Electrons Orbit the Nucleus at Fixed Distances,
Forming Electron Shells That Correspond
to Different Energy Levels
As you may know from experimenting with a magnet,
like poles repel each other and opposite poles attract
each other. In a similar way, electrons repel one another, owing to their negative electrical charge, and they
are drawn to the positively charged protons of the nucleus. However, because of their mutual repulsion, only
limited numbers of electrons can occupy the space closest to the nucleus. Large atoms can accommodate many
electrons because their electrons orbit at increasing distances from their nucleus. The electrons orbit through a
three-dimensional space; the orbits, which correspond
(a)
e-
(b)
electron
shell
e-
p+
p+
Figure 2-1 Atomic models
Structural representations of the two smallest atoms, (a) hydrogen
and (b) helium. In these simplified models, the electrons are represented as miniature planets, circling in specific orbits around a nucleus that contains protons and neutrons.
to different energy levels, are called electron shells (Figs.
2-1 and 2-2, p. 24).
The electron shell closest to the atomic nucleus is the
smallest and can hold only two electrons. The second
shell can hold up to eight electrons. The electrons in an
atom normally fill the shell closest to the nucleus and
then begin to occupy the next shell. Thus, a carbon
atom, with six electrons, has two electrons in the first
shell, closest to the nucleus, and four electrons in its second shell (see Fig. 2-2). Nuclei and electron shells play
complementary roles in atoms. Nuclei (assuming they
are not radioactive) provide stability, while the electron
shells allow interactions, or bonds, with other atoms.
Nuclei resist disturbance by outside forces. Ordinary
Percent in
Universeb
Percent in
Earthb
Symbol
Hydrogen
H
1
91
Helium
He
2
9
Carbon
C
6
0.02
Nitrogen
N
7
0.04
Trace
Oxygen
O
8
0.06
47
0.14
Trace
0.03
Percent in
Human Bodyb
9.5
Trace
18.5
3.3
65
Sodium
Na
11
Trace
2.8
0.2
Magnesium
Mg
12
Trace
2.1
0.1
Phosphorus
P
15
Trace
0.07
Sulfur
S
16
Trace
0.03
0.3
1
Chlorine
Cl
17
Trace
0.01
0.2
Potassium
K
19
Trace
2.6
0.4
Calcium
Ca
20
Trace
3.6
Iron
Fe
26
Trace
5
1.5
Trace
Atomic number = number of protons in the atomic nucleus.
b
n
Helium (He)
Hydrogen (H)
Element
a
n
e-
nucleus
Table 2-1 Common Elements Important in Living Organisms
Atomic
Numbera
p+
Approximate percentage of atoms of this element, by weight, in the universe, in Earth’s crust, and in the human body.
What Are Atoms?
23
Scientific Inquiry
Radioactivity in Research
As you read this text, you will encounter many statements that
may cause you to wonder, How do they know that? How do
biologists know that DNA is the genetic material of cells (Chapter 9)? How do paleontologists measure the ages of fossils
(Chapter 17)? How do botanists know that sugars made in
plant leaves during photosynthesis are transported to other
parts of the plant in the sieve tubes of phloem (Chapter 23)?
These discoveries, and many more, have been possible only
through the use of radioactive isotopes.
Although all atoms of a particular element have the same
number of protons, the number of neutrons may vary. Neutrons don’t affect the chemical reactivity of an atom very much,
but they add to the atom’s mass, which can be detected by sophisticated instruments, such as mass spectrometers. Nuclei
with “too many” neutrons break apart spontaneously, or
decay, often emitting radioactive particles in the process. Those
particles can also be detected—for example, with Geiger counters. The process by which a radioactive isotope spontaneously
breaks apart is called radioactive decay.
A particularly fascinating and medically important use of
radioactive isotopes is positron emission tomography, more
commonly known as PET scans (Fig. E2-1a). In one common
application of PET scans, a subject is given the sugar glucose
that has been labeled with (that is, attached to) a harmless
radioactive isotope of fluorine. When the nucleus of radioactive fluorine decays, it emits two bursts of energy that travel
in opposite directions along the same line. Energy detectors
arranged in a ring around the subject look at a “slice” of the
brain, recording the nearly simultaneous arrival of the two
energy bursts (Fig. E2-1b). A powerful computer then calculates the location within the subject at which the decay must
have occurred and generates a color-coded map of the frequency of decays. Because the radioactive fluorine is attached to glucose molecules, this map reflects the glucose
concentrations within the subject’s brain. Since the brain
uses this sugar for energy, the more active a brain cell is, the
more glucose it uses, and the more radioactivity is concentrated there. For example, tumor cells (Fig. E2-1c) or brain regions in which epileptic seizures originate generally have
excessively high glucose utilization and show up in PET scans
as “hot spots.” Normal brain cells activated by a specific
mental task will also have higher glucose demands, which
can be detected by PET scans.
Biology and medicine have profited immensely from close
interactions with the other sciences, especially chemistry and
physics. The development of PET scans required close cooperation with chemists (developing and synthesizing the radioactive probes), physicists (understanding interactions between
electrons and short-lived, positively charged particles called
positrons, as well as the geometry of the resulting energy
emissions), and engineers (designing and building the electronic apparatus). Continued teamwork among scientists
promises further advances in both the fundamental understanding of biological processes and applications in medicine
and agriculture.
(b)
(a)
(c)
detector ring
Subject's head is placed
within a ring of detectors.
Radioactive decay releases
energetic particles that
activate the detectors.
Figure E2-1 How positron emission tomography works
Red indicates the highest
radioactivity; blue is least.
A malignant brain tumor
shows clearly in red.
24
Chapter 2
Atoms, Molecules, and Life
2e–
5e–
8e–
4e–
6e–
8e–
8e–
2e–
2e–
2e–
2e–
6p+
6n
8p+
8n
15p+
16n
20p+
20n
Carbon (C)
Oxygen (O)
Phosphorus (P)
Calcium (Ca)
C
O
P
Ca
Figure 2-2 Electron shells in atoms
Most biologically important atoms have at least two shells of electrons. The first shell, closest to the nucleus, can hold two
electrons; the next shell can hold a maximum of eight electrons. More distant shells can also hold eight electrons each.
sources of energy, such as heat, electricity, and light,
hardly affect them at all. Because its nucleus is stable, a
carbon atom remains carbon whether it is part of a diamond, carbon dioxide, or sugar. Electron shells, however, are dynamic; as you will soon see, atoms bond with
one another by gaining, losing, or sharing electrons.
How Do Atoms Interact
b to Form Molecules?
Atoms Interact with Other Atoms When There
Are Vacancies in Their Outermost Electron Shells
A molecule consists of two or more atoms, of either the
same or of different elements, held together by interactions among their outermost electron shells. A substance whose molecules are formed of different types of
atoms is called a compound. Atoms interact with one
another according to two basic principles:
1. An atom will not react with other atoms when its outermost electron shell is completely full or empty. Such
an atom is described as being inert.
2. An atom will react with other atoms when its outermost electron shell is only partially full. Such atoms
are described as reactive.
To demonstrate these principles, consider three atoms:
hydrogen, oxygen, and helium (see Fig. 2-1). Hydrogen
(the smallest atom) has one proton in its nucleus and
one electron in its single (and therefore outermost)
electron shell, which can hold up to two electrons. The
oxygen atom has six electrons in its outer shell, which
can hold eight. In contrast, helium has two protons in its
nucleus, and two electrons fill its single electron shell.
Therefore, we predict that hydrogen and oxygen atoms,
with partially empty outer shells, should be reactive,
while helium atoms, with a full shell, should be stable.
We might further predict that hydrogen and oxygen
atoms could gain stability by reacting with each other.
Table 2-2 Chemical Bonds
Type of Bond
Bond Forms:
Weak Bonds: allow interactions between individual atoms or molecules
Ionic bonds
Between positive and negative ions
Hydrogen bonds
Between a hydrogen atom involved in a polar covalent bond and
another atom involved in a polar covalent bond
Hydrophobic interactions
Because interactions between water molecules exclude hydrophobic
molecules
Strong Bonds: hold atoms together within molecules
Covalent bonds
By the sharing of electron pairs; equal sharing produces nonpolar
covalent bonds; unequal sharing produces polar covalent bonds
How Do Atoms Interact to Form Molecules?
The single electrons from each of the two hydrogen
atoms would fill the outer shell of oxygen, forming H2O:
water (see Fig. 2-4c). As predicted, hydrogen can react
readily with oxygen—in fact, the reaction is an explosive one. The space shuttle and many other rockets use
liquid hydrogen as fuel to power liftoff. The hydrogen
fuel reacts with oxygen, releasing water as a by-product.
In contrast, helium, with a full outer shell, is almost
completely inert. Our case study mentioned free radicals. Free radicals are atoms or molecules that lack one
or more electrons in their outer shells, making them
highly reactive.
An atom with an outermost electron shell that is partially full can gain stability by losing electrons (emptying the shell completely), by gaining electrons (filling
the shell), or by sharing electrons with another atom
(with both atoms behaving as though they had full outer
shells). The results of losing, gaining, and sharing electrons are chemical bonds; attractive forces that hold
atoms together in molecules. Each element has chemical bonding properties that arise from the configuration
of electrons in its outer shell (Table 2-2). Chemical reactions, the making and breaking of chemical bonds to
form new substances, are essential for the maintenance
of life and for the working of modern society. Whether
they occur in a plant cell as it captures solar energy, your
brain as it forms new memories, or your car’s engine as
it guzzles gas, chemical reactions consist of making new
chemical bonds and/or breaking existing ones.
Charged Atoms Called Ions Interact
to Form Ionic Bonds
Both atoms that have an almost empty outermost electron shell and atoms that have an almost full outermost
shell can become stable by losing electrons (emptying
their outermost shell) or by gaining electrons (filling
their outermost shell). The formation of table salt (sodium chloride) demonstrates this principle. Sodium (Na)
has only one electron in its outermost electron shell,
and chlorine (Cl) has seven electrons in its outer shell—
one electron short of being full (Fig. 2-3a). Sodium,
therefore, can become stable by losing the electron to
chlorine from its outer shell, leaving that shell empty;
chlorine can fill its outer shell by gaining the electron.
Atoms that have lost or gained electrons, altering the
balance between protons and electrons, are charged.
These charged atoms are called ions. To form sodium
chloride, sodium loses an electron and thereby becomes
a positively charged sodium ion (Na+); chlorine picks
up an electron and becomes a negatively charged chloride ion (Cl–) (Fig. 2-3b,c).
Opposite charges attract; therefore, sodium ions and
chloride ions tend to stay near one another. They form
crystals that contain repeating orderly arrangements of
the two ions (Fig. 2-3c). The electrical attraction between oppositely charged ions that holds them together
Na
(a)
Cl
Chlorine atom (neutral)
_ _
Sodium atom (neutral)
_
_
_
_
_
_
_
_
11p+
11n
25
_
_
_
_
_
_
_
_
_
_
_
_
_
17p+
18n
_
_
_
_
_
Electron transferred
Na+
Cl–
Sodium ion (+)
(b)
_
_
_
_
_
_
Chloride ion (–)
_ _
_
_
_
11p+
11n
_
_
_
_
_
_
_
_
_
_
_
_
17p+
18n
_
_
_
_
_
Attraction between
opposite charges
(c)
Na+
Cl–
An ionic compound: NaCl
Figure 2-3 The formation of ions and ionic bonds
(a) Sodium has only one electron in its outer electron shell; chlorine has seven. (b) Sodium can become stable by losing an electron, and chlorine can become stable by gaining an electron.
Sodium becomes a positively charged ion and chlorine a negatively charged ion. (c) Because oppositely charged particles attract
one another, the resulting sodium ions (Na+) and chloride ions
(Cl–) nestle closely together in a crystal of salt, NaCl.
in crystals is called an ionic bond. Ionic bonds are weak
and easily broken, as occurs when salt is dissolved in
water (see Table 2-2).
Uncharged Atoms Can Become Stable
by Sharing Electrons, Forming Covalent Bonds
An atom with a partially full outermost electron shell
can also become stable by sharing electrons with another atom, forming a covalent bond. Consider the
26
Chapter 2
Atoms, Molecules, and Life
(a)
nonpolar covalent
bonding
_
_
+
+
(c)
polar covalent
bonding
Hydrogen (H H or H2),
a nonpolar molecule
(slightly negative)
_
_
_
_
_
(b)
_
_
_
_
_
_
8p+
8n
_
_
_
_
_
8p+
8n
_
_
_
_
_
_
+
+
8p+
8n
_
_
_
_
Oxygen (O O or O2),
a nonpolar molecule
(slightly positive)
Water (H O H or H2O),
a polar molecule
Figure 2-4 Covalent bonds
Electrons are shared between atoms to form covalent bonds. (a) In hydrogen gas, one electron from each hydrogen
atom is shared, forming a single covalent bond. The resulting molecule of hydrogen gas is represented as H–H or H2.
(b) In oxygen gas, two oxygen atoms share four electrons, forming a double bond (O=O or O2). (c) Oxygen lacks two
electrons to fill its outer shell, so oxygen can make one bond with each of two hydrogen atoms to form water (H–O–H or
H2O). Oxygen exerts a greater pull on the electrons than does hydrogen, so the oxygen end of the molecule has a slight
negative charge and the hydrogen end has a slight positive charge. This is an example of polar covalent bonding. The
water molecule with its slightly charged ends is called a polar molecule.
hydrogen atom, which has one electron in a shell built
for two. A hydrogen atom can become reasonably stable if it shares its single electron with another hydrogen atom, forming a molecule of hydrogen gas, H2 (Fig.
2-4a). Because the two hydrogen atoms are identical,
neither nucleus can exert more attraction and capture
the other’s electron. So the two electrons orbit around
both nuclei for equal amounts of time, forming a single
covalent bond; each hydrogen atom behaves almost as
if it had two electrons in its shell. Two oxygen atoms
also share electrons equally, producing a molecule of
oxygen gas, O2, with a double covalent bond (Fig.
2-4b). In such a bond, each atom contributes two electrons. If atoms share three pairs of electrons, a triple
covalent bond is formed, as in nitrogen gas, N2.
All covalent bonds are strong compared with ionic
bonds, but some are stronger than others, depending on
the atoms involved (see Table 2-2). Some covalent
bonds, such as those in water (H2O; Fig. 2-4c) and carbon
dioxide (CO2), are extremely stable—that is, it takes
a lot of energy to break the bonds. Other bonds, such
as those in oxygen gas (Fig. 2-4b) or gasoline, are less
stable, and come apart more easily. When a chemical
reaction occurs in which less stable bonds are broken
and more stable bonds are formed (such as burning
gasoline with oxygen to form carbon dioxide and water),
energy is released, as we shall describe in Chapter 6.
Most Biological Molecules Utilize Covalent Bonding
Covalent bonds are crucial to life because the atoms in
most biological molecules are joined by covalent
bonds. The molecules in proteins, sugars, bone, and
cellulose are formed of atoms held together by covalent bonds. Hydrogen, carbon, oxygen, nitrogen, phosphorus, and sulfur are the most common atoms found
in biological molecules. Except for hydrogen, each of
these atoms needs at least two electrons to fill its outermost electron shell and can share electrons with two or
Why Is Water So Important to Life?
27
Table 2-3 Bonding Patterns of Atoms Commonly Found in Biological Molecules
Atom
Capacity of
Outer Electron Shell
Electrons in
Outer Shell
Number of Covalent
Bonds Normally Formed
Hydrogen
2
1
1
Carbon
8
4
4
Nitrogen
8
5
3
Oxygen
8
6
2
Phosphorus
8
5
5
P
Sulfur
8
6
2
S
more other atoms. Hydrogen can form a covalent bond
with one other atom; oxygen and sulfur with two other
atoms; nitrogen with three; and phosphorus and carbon
with up to four. (Phosphorus is unusual; although it has
only three spaces in its outer shell, it can form up to five
covalent bonds with up to four other atoms.) This diversity of bonding arrangements permits biological
molecules to be constructed in an almost infinite variety and complexity. Double and triple bonds increase
the variety of shapes and functions of biological molecules. Table 2-3 summarizes bonding patterns in biological molecules.
Polar Covalent Bonds Form When Atoms
Share Electrons Unequally
In hydrogen gas the two nuclei are identical, and the
shared electrons spend equal time near each nucleus.
Therefore, not only is the molecule as a whole electrically neutral, but each end, or pole, of the molecule is
also electrically neutral. Such an electrically symmetrical bond is called a nonpolar covalent bond and the
compound formed with such nonpolar bonds is a nonpolar molecule, such as hydrogen (H2) or oxygen (O2)
(see Fig. 2-4a,b). But electron sharing in covalent bonds
is not always equal. In many molecules, one nucleus
may initially have a larger positive charge, and therefore attract the electrons more strongly, than does the
other nucleus. This situation produces a polar covalent
bond. Although the molecule as a whole is electrically
neutral, it has charged parts: The atom that attracts the
electrons more strongly then picks up a slightly negative charge (the negative pole of the molecule), and the
other atom has a slightly positive charge (the positive
pole). In water, for example, oxygen attracts electrons
more strongly than does hydrogen, so the oxygen end
of a water molecule is negative and each hydrogen is
positive (see Fig. 2-4c). Water with its charged ends is a
polar molecule.
Common Bonding
Patterns
H
C
C
C
N
N
C
N
O
O
Hydrogen Bonds Are Weaker Electrical
Attractions Between or Within Molecules
with Polar Covalent Bonds
Because of the polar nature of their covalent bonds,
nearby water molecules attract one another. The partially negatively charged oxygens of some water molecules attract the partially positively charged hydrogens
of other water molecules. This electrical attraction is
called a hydrogen bond (Fig. 2-5; see Table 2-2). As we
shall see shortly, hydrogen bonds between molecules
give water several unusual properties that are essential
to life on Earth.
Hydrogen bonds are common and important in biological molecules, as well as in water. They may occur
whenever polar covalent bonds produce slightly negative and slightly positive charges that then attract one
another. Both nitrogen and oxygen atoms attract electrons more strongly than do hydrogen atoms. Therefore,
the nitrogen or oxygen pole of a nitrogen–hydrogen or
oxygen–hydrogen bond is slightly negative, and the hydrogen pole is slightly positive. The resulting polar parts
of the molecules can form hydrogen bonds with water,
with other biological molecules, or with polar parts of
the same molecule. Although individual hydrogen bonds
are quite weak, many of them working together are
quite strong. As we shall see in Chapter 3, hydrogen
bonds play crucial roles in shaping the three-dimensional
structures of proteins. In Chapter 9 you’ll discover their
importance in DNA
.
c Why Is Water So Important to Life?
Water is extraordinarily abundant on Earth, has unusual properties, and is so essential to life that it merits special consideration. Life is very likely to have arisen in
28
Chapter 2
Atoms, Molecules, and Life
2.2 Water and Life
MEDIATUTOR
H
(+)
H
(+)
O
(_)
H
(+)
O
(_ )
H
(+)
hydrogen
bonds
Figure 2-5 Hydrogen bonds
The partial charges on different parts of water molecules produce
weak attractive forces called hydrogen bonds (dotted lines) between the hydrogens of one water molecule and the oxygens of
other molecules.
and negative poles. If a salt crystal is dropped into
water, the positively charged hydrogen ends of water
molecules will be attracted to and will surround the
negatively charged chloride ions, and the negatively
charged oxygen poles of water molecules will surround
the positively charged sodium ions. As water molecules
enclose the sodium and chloride ions and shield them
from interacting with each other, the ions separate
from the crystal and drift away in the water—and the
salt dissolves (Fig. 2-6).
Water also dissolves molecules held together by
polar covalent bonds. Its positive and negative poles
are attracted to oppositely charged regions of dissolving molecules. Ions and polar molecules are termed hydrophilic (Greek for “water-loving”) because of their
electrical attraction for water molecules. Many biological molecules, including sugars and amino acids, are
hydrophilic and dissolve readily in water (Fig. 2-7).
Water also dissolves gases such as oxygen and carbon
dioxide. The fish swimming below the ice on a frozen
lake rely on oxygen that dissolved before the ice
formed, and they release CO2 into solution in the
water. By dissolving such a wide variety of molecules,
the watery substance inside a cell provides a suitable
environment for the countless chemical reactions essential to life on Earth.
Molecules that are uncharged and nonpolar, such as
fats and oils, usually do not dissolve in water and hence
are called hydrophobic (“water-fearing”). Nevertheless,
water has an important effect on such molecules. Oils,
for example, form globules when spilled into water. Oil
the waters of the primeval Earth. Living organisms still
contain about 60% to 90% water, and all life depends
intimately on the properties of water. Why is water so
crucial to life?
Na+
Water Interacts with Many Other Molecules
Water enters into many of the chemical reactions that
occur in living cells. The oxygen that green plants release into the air is derived from water during photosynthesis. In manufacturing a protein, fat, nucleic acid,
or sugar, your body produces water in the process; conversely, when you digest proteins, fats, and sugars in the
foods you eat, water is used in the reactions. Why is
water so important in biological chemical reactions?
Water is an extremely good solvent—that is, it is capable of dissolving a wide range of substances, including protein, salts, and sugars. Water or other solvents
containing dissolved substances are called solutions.
Recall that a crystal of table salt is held together by the
electrical attraction between positively charged sodium
ions and negatively charged chloride ions (see Fig.
2-3c). Because water is a polar molecule, it has positive
Cl–
H
H
O–
Cl–
H
Na+
H
O
Na+
Figure 2-6 Water as a solvent
The polarity of water molecules allows water to dissolve polar and
charged substances. When a salt crystal is dropped into water,
the water surrounds the sodium and chloride ions with oppositely
charged poles of the water molecules. Thus insulated from the attractiveness of other molecules of salt, the ions disperse, and the
whole crystal gradually dissolves.
Why Is Water So Important to Life?
water
hydrogen bond
glucose
hydroxyl
group
Figure 2-7 Water dissolves many biological molecules
Many biological molecules dissolve in water because they have
polar parts—for example, OH– (hydroxyl) groups—that can form
hydrogen bonds with water molecules. As shown, hydrogen bonds
can form between the hydroxyl groups on a glucose molecule (a
simple sugar) and surrounding water molecules.
molecules in water disrupt the hydrogen bonding
among adjacent water molecules. When oil molecules
encounter one another in water, their nonpolar surfaces
nestle closely together, surrounded by water molecules
that form hydrogen bonds with one another but not
with the oil. To separate again, the oil molecules would
have to break apart the hydrogen bonds that link sur-
(a)
29
rounding water molecules. Thus, the oil molecules remain together, forming a glistening droplet that floats
on the water’s surface. The tendency of oil molecules to
clump together is described as a hydrophobic interaction
(see Table 2-2). As we shall discuss in Chapter 4, the
membranes of living cells owe much of their structure
to hydrophobic interactions.
Water Molecules Tend to Stick Together
In addition to interacting with other molecules, water
molecules interact with each other. Because hydrogen
bonds interconnect individual water molecules, liquid
water has high cohesion—that is, water molecules have
a tendency to stick together. Cohesion among water
molecules at the water’s surface produces surface tension, the tendency for the water surface to resist being
broken. If you’ve ever experienced the slap and sting of
a belly flop into a swimming pool, you’ve discovered
firsthand the power of surface tension. Surface tension
can support fallen leaves, some spiders and water insects, and even a running lizard (Fig. 2-8a).
A more important role of cohesion in water occurs
in the life of land plants. Since a plant absorbs water
through its roots, how does the water reach the aboveground parts, especially if the plant is a 100-meter-tall
redwood (Fig. 2-8b)? As we shall see in Chapter 23,
water molecules are pulled up by the leaves. Water
fills tiny tubes that connect the leaves, stem, and roots.
Water molecules that evaporate from the leaves pull
(b)
Figure 2-8 Cohesion among water molecules
(a) With webbed feet bearing specialized scales, the basilisk lizard of South America makes use of surface tension,
caused by cohesion, to support its weight as it races across the surface of a pond. (b) In giant redwoods, cohesion
holds water molecules together in continuous strands from the roots to the topmost leaves even 300 feet (about
100 meters) above the ground.
30
Chapter 2
Atoms, Molecules, and Life
water up the tubes, much like a chain being pulled up
from the top. The system works because the hydrogen
bonds interconnecting water molecules are stronger
than the weight of the water in the tubes, even 100
meters’ worth; thus, the water “chain” doesn’t break.
Without the cohesion of water, there would be no
land plants as we know them, and terrestrial life
would undoubtedly have evolved quite differently.
You may have realized by now that the “common
bond” producing the sting of a belly flop, the ability of
a lizard to run on water, and the movement of water
up a tree is actually the hydrogen bond between water
molecules.
Water exhibits another property, adhesion, a word
that describes its tendency to stick to polar surfaces
having slight charges that attract polar water molecules. Adhesion helps water move within small spaces,
such as the thin tubes in plants that carry water from
roots to leaves. If you stick the end of a narrow glass
tube into water, the water will move a short distance up
the tube. Put some water in a narrow glass bud vase or
test tube and you’ll see that the upper surface is curved;
water pulls itself up the sides of the glass by its adhesion to the surface of the glass and by the cohesion
among water molecules.
Water Can Form H+ and OH– Ions
Although water is generally regarded as a stable compound, individual water molecules constantly gain, lose,
and swap hydrogen atoms. As a result, at any given time
about two of every billion water molecules are ionized—that is, broken apart into hydrogen ions (H+) and
hydroxide ions (OH–):
(–)
O
H
water
(H2O)
O
H
(+)
+
H
H
hydroxide ion
(OH–)
hydrogen ion
(H+)
A hydroxide ion has gained an electron from the hydrogen ion, giving it a negative charge, while the hydrogen
ion, which has lost its electron, now has a positive
charge. Pure water contains equal concentrations of hydrogen ions and hydroxide ions.
In many solutions, however, the concentrations of
H+ and OH– are not the same. If the concentration of
H+ exceeds the concentration of OH–, the solution is
acidic. An acid is a substance that releases hydrogen
ions when it is dissolved in water. When hydrochloric
acid (HCl), for example, is added to pure water, almost
all of the HCl molecules separate into H+ and Cl–.
Therefore, the concentration of H+ greatly exceeds the
concentration of OH–, and the resulting solution is
acidic. (Many acidic substances, such as lemon juice
and vinegar, have a sour taste. This is because the sourtaste receptors on your tongue are specialized to respond to the excess of H+.)
If the concentration of OH+ is greater, the solution
is basic. A base is a substance that combines with hydrogen ions, reducing their number. If, for instance,
sodium hydroxide (NaOH) is added to water, the
NaOH molecules separate into Na+ and OH–. The
OH– combine with H+, reducing their number. The solution is then basic.
The degree of acidity is expressed on the pH scale
(Fig. 2-9), in which neutrality (equal numbers of H+ and
OH–) is assigned the number 7. Acids have a pH below
7; bases have a pH above 7. Pure water, with equal concentrations of H+ and OH–, has a pH of 7. Each unit on
the pH scale represents a tenfold change in the concentration of H+. Thus, a cola drink (pH = 3) has a concentration of H+ 10,000 times that of water (pH = 7)—no
wonder it is bad for your teeth!
A Buffer Helps Maintain a Solution
at a Relatively Constant pH
In most mammals, including humans, both the cell interior (cytoplasm) and the fluids that bathe the cells are
nearly neutral (pH about 7.3 to 7.4). Small increases or
decreases in pH may cause drastic changes in both the
structure and function of biological molecules, leading
to the death of cells or entire organisms. Nevertheless,
living cells seethe with chemical reactions that take up
or give off H+. How, then, does the pH remain constant
overall? The answer lies in the many buffers found in
living organisms. A buffer is a compound that tends to
maintain a solution at a constant pH by accepting or releasing H+ in response to small changes in H+ concentration. If the H+ concentration rises, buffers combine
with them; if the H+ concentration falls, buffers release
H+. The result is that the concentration of H+ is restored to its original level. Common buffers in living organisms include bicarbonate (HCO3–) and phosphate
(H2PO4– and HPO42–), both of which can accept or release H+, depending on the circumstances. If the blood
becomes too acidic, for example, bicarbonate accepts
H+ to form carbonic acid:
HCO3 2
(bicarbonate)
1
H1
(hydrogen ion)
→
H2CO3
(carbonic acid)
If the blood becomes too basic, carbonic acid liberates
hydrogen ions, which combine with the excess hydroxide ions, forming water:
H2CO3
(carbonic acid)
1
OH 2
(hydroxide ion)
→
HCO3 2
(bicarbonate)
1 H2O
(water)
In either case, the result is that the blood pH remains
near its normal value.
Why Is Water So Important to Life?
H + concentration
(moles/liter)
pH
value
0 1-molar hydrochloric
acid (HCI)
10 –1
1 stomach acid
lime juice
10 –2
10 –3
10 –4
10 –5
10
increasingly acidic (H+ > OH–)
10 0
–6
10 –7
2 lemon juice
3 "acid rain" (2.5–5.5)
vinegar, cola, orange juice,
tomatoes
4 beer
5 black coffee, tea
6
neutral
_
(H+ = OH )
normal rain (5.6)
urine (5.7)
7 pure water (7.0)
saliva
blood, sweat (7.4)
8 seawater (7.8–8.3)
10 –9
9 baking soda
10 –10
10 –11
10 –12
increasingly basic (H+ < OH–)
10 –8
10 phosphate detergents
chlorine bleach
milk of magnesia
11 household ammonia
some detergents
(without phosphates)
12
washing soda
10 –13
13 oven cleaner
10 –14
14 1-molar sodium
hydroxide (NaOH)
Figure 2-9 The pH scale
The pH scale expresses the concentration of hydrogen ions in a
solution on a scale of 0 (very acidic) to 14 (very basic). Each unit of
change in pH on the pH scale represents a tenfold change in the
concentration of hydrogen ions. Lemon juice, for example, is about
10 times more acidic than orange juice, and the most severe acid
rains in the northeastern United States are almost 1000 times more
acidic than normal rainfall. Except for the inside of your stomach,
nearly all the fluids in your body are finely adjusted to a pH of 7.4.
The color coding corresponds to a common pH indicator dye,
bromthymol blue, widely used by aquarium owners to monitor the
pH of water for their fish.
31
Water Moderates the Effects
of Temperature Changes
Your body and the bodies of other organisms can survive only within a limited temperature range. As we
shall see in Chapter 6, high temperatures may damage
enzymes that guide the chemical reactions essential to
life. Low temperatures are also dangerous, because enzyme action slows as the temperature drops. Subfreezing temperatures within the body are usually lethal,
because spearlike ice crystals can rupture cells.
Water has important properties that moderate the
effects of temperature changes. These properties help
keep the bodies of organisms within tolerable temperature limits. Also, large lakes and the oceans have a
moderating effect on the climate of nearby land, making it warmer in winter and cooler in summer. First,
some background: Temperature reflects the speed of
molecules; the higher the temperature, the greater their
average speed. Generally speaking, if heat energy enters a system, the molecules of that system move more
rapidly, and the temperature of the system rises. Individual water molecules, however, are weakly linked to
one another by hydrogen bonds (see Fig. 2-5). When
heat enters a watery system such as a lake or a living
cell, much of the heat energy goes into breaking hydrogen bonds rather than speeding up individual molecules. Thus, 1 calorie of energy will heat 1 gram of water
1 °C, whereas it takes only 0.6 calorie per gram to heat
alcohol 1 °C, 0.2 calorie for table salt, and 0.02 calorie
for common rocks such as granite or marble. So the energy required to heat a pound (a pint) of water only
1 °C would raise a pound of rock by 50 °C. If a lizard
wants to warm up, it will seek out a rock, rather than a
puddle. Because the human body is mostly water, a
sunbather can absorb a lot of heat energy without
sending his or her body temperature soaring—and
many hot sunbathers can jump into a swimming pool to
cool off without raising the temperature of the water
very much. (The energy required to heat a gram of a
substance by 1 °C is called its specific heat—water has a
high specific heat.)
Second, water moderates the effects of high temperatures because it takes a great deal of heat, 539 calories per gram, to convert liquid water to water vapor.
This, too, is due to the hydrogen bonds that interconnect individual water molecules. For a water molecule
to evaporate, it must move quickly enough to break all
the hydrogen bonds that hold it to the other water
molecules in the solution. Only the fastest-moving
water molecules, carrying the most energy, can break
their hydrogen bonds and escape into the air as water
vapor. The remaining liquid is cooled by the loss of
these high-energy molecules. As children romp
through a sprinkler on a hot summer day, water coats
their bodies. Heat energy is transferred from their skin
to the water and from the water to the vapor as the
32
Chapter 2
Atoms, Molecules, and Life
water evaporates. Evaporating just 1 gram of water
cools 539 grams of a person’s body 1 °C, so water is a
very effective coolant. This also means that evaporating perspiration produces a great loss of heat without
much loss of water. (The heat required to vaporize
water is called its heat of vaporization—water’s heat of
vaporization is one of the highest known!)
Third, water moderates the effects of low temperatures because an unusually large amount of energy must
be removed from molecules of liquid water before they
form the precise crystal arrangement of ice. As a result,
water freezes more slowly than do many other liquids at
a given temperature and loses more heat to the environment in the process. (This property of water is called its
heat of fusion, which is very high.)
Water Forms an Unusual Solid: Ice
Water, of course, will become a solid after prolonged exposure to temperatures below its freezing point. But
even solid water is unusual. Most liquids become denser
when they solidify, and the solid sinks. Ice is rather
unique because it is less dense than liquid water. When
a pond or lake starts to freeze in winter, the ice stays on
top, forming an insulating layer that delays the freezing
of the rest of the water. This insulation allows fish and
other lake residents to survive in the liquid water below.
If ice were to sink, ponds and lakes in much of North
America would freeze solid, from the bottom up, during
the winter, killing fish and underwater plants and making drinking water far less available to animals.
REVISITEDCASESTUDYREVISITED CASESTUDYREVISITEDCAS
Health Food?
You now know that a free radical is an
atom or molecule containing an unfilled
outer electron shell, which reacts vigorously with other molecules to fill its outer
shell. Normal cellular activities produce a
variety of molecules that contain an oxygen atom with an unfilled outer shell.
Such molecules react vigorously with
other molecules, damaging them in the
process. If you’ve seen iron turn to rust
(iron oxide), you’ve witnessed oxygen’s
reactive power. Oxidative stress refers to
a kind of ”biological rusting” in which
free radicals that contain oxygen damage
cells. Substances that react with these
oxygen-containing free radicals and render them harmless are called antioxi-
dants. Although cells contain some of
their own antioxidants, your eating habits
can also make a difference. Fruits and
vegetables, particularly those with yellow,
orange, and red colors, are good sources
of antioxidant compounds; vitamins C
and E are antioxidants also. Now, amazingly, researchers have given us an excuse
to eat chocolate and feel good about it:
cocoa powder contains high concentrations of flavenoids, which are powerful
antioxidants. This research is in its early
stages, and no studies have been done to
determine whether high consumption of
chocolate actually reduces the risk of
cancer or heart disease, but there will certainly be no shortage of volunteers for
this research. Although becoming fat by
eating too much chocolate candy could
counteract any positive effects of the
cocoa powder itself, slim “chocoholics”
have reason to relax and enjoy.
As you will learn in Chapter 8, oxygen is
important for harvesting the maximum
amount of energy from molecules such as
sugar. Ross Hardison, a researcher at
Pennsylvania State University, stated eloquently: “Keeping oxygen under control
while using it in energy production has
been one of the great compromises
struck in the evolution of life on Earth.”
What did he mean by this?
Summary of Key Concepts
1 What Are Atoms?
An element is a substance that can neither be broken
down nor converted to different substances by ordinary
chemical means. The smallest possible particle of an element is the atom, which is itself composed of a central nucleus, containing protons and neutrons, and electrons
outside the nucleus. All atoms of a given element have the
same number of protons, which is different from the number of protons in the atoms of every other element. Electrons orbit the nucleus in electron shells, at specific
distances from the nucleus. Each shell can contain a fixed
maximum number of electrons. The chemical reactivity of
an atom depends on the number of electrons in its outermost electron shell: An atom is most stable, and therefore
least reactive, when its outermost shell is either completely full or empty.
2 How Do Atoms Interact to Form Molecules?
Atoms may combine to form molecules. The forces holding
atoms together in molecules are called chemical bonds.
Atoms that have lost or gained electrons are negatively or
positively charged particles called ions. Ionic bonds are
electrical attractions between charged ions, holding them
together in crystals. When two atoms share electrons, covalent bonds form. In a nonpolar covalent bond, the two
atoms share electrons equally. In a polar covalent bond,
one atom may attract the electron more strongly than the
other atom does; in this case, the strongly attracting atom
bears a slightly negative charge, and the weakly attracting
atom bears a slightly positive charge. Some polar covalent
bonds give rise to hydrogen bonding, the attraction between charged regions of individual polar molecules or distant parts of a large polar molecule.
Thinking Through the Concepts
3 Why Is Water So Important to Life?
Properties of the water molecule that are important to living organisms include its ability to interact with many
other molecules and to dissolve many polar and charged
substances; to force nonpolar substances, such as fat, to as-
33
sume certain types of physical organization; to participate
in chemical reactions; to cohere to itself by using hydrogen
bonds between water molecules; and to maintain a fairly
stable temperature in the face of wide temperature fluctuations in the environment.
Key Terms
acid p. 30
acidic p. 30
atom p. 21
atomic nucleus p. 21
atomic number p. 21
base p. 30
basic p. 30
buffer p. 30
calorie p. 31
chemical bond p. 25
chemical reaction p. 25
cohesion p. 29
compound p. 24
covalent bond p. 25
electron p. 21
electron shell p. 22
element p. 21
hydrogen bond p. 27
hydrophilic p. 28
hydrophobic p. 28
hydrophobic interaction p. 29
ion p. 25
ionic bond p. 25
isotope p. 22
molecule p. 24
neutron p. 21
nonpolar covalent bond p. 27
pH scale p. 30
polar covalent bond p. 27
proton p. 21
radioactive p. 22
solvent p. 28
surface tension p. 29
Thinking Through the Concepts
Multiple Choice
1. What is the purest form of matter that cannot be separated
into different substances by chemical means?
a. compounds
b. molecules
c. atoms
d. elements
e. electrons
2. Which phrase best describes chemical bonds?
a. physical bridges
b. attractive forces
c. shared protons
d. atomic reactions
e. all of these phrases are equally descriptive
3. When an atom ionizes, what happens?
a. It shares one or more electrons with another atom.
b. It emits energy as it loses extra neutrons.
c. It gives up or takes up one or more electrons.
d. It shares a hydrogen atom with another atom.
e. none of the above
4. If electrons in water molecules were equally attracted to hydrogen nuclei and oxygen nuclei, water molecules would be
a. more polar
b. less polar
c. unchanged
d. triple bonded
e. unable to form
5. A covalent bond forms
a. when two ions are attracted to one another
b. between adjacent water molecules, producing surface
tension
c. when one atom gives up its electron to another atom
d. when two atoms share electrons
e. between water molecules and fat globules
6. What is the defining characteristic of an acid?
a. It donates hydrogen ions.
b. It accepts hydrogen ions.
c. It will donate or accept hydrogen ions, depending on
the pH.
d. It has an excess of hydroxide ions.
e. It has a pH greater than 7.
? Review Questions
1. What are the six most abundant elements that occur in living organisms?
2. Distinguish among atoms and molecules; elements and
compounds; and protons, neutrons, and electrons.
3. Compare and contrast covalent bonds and ionic bonds.
4. Why can water absorb a great amount of heat with little
increase in its temperature?
5. Describe how water dissolves a salt. How does this phenomenon compare with the effect of water on a hydrophobic substance such as corn oil?
6. Define acid, base, and buffer. How do buffers reduce
changes in pH when hydrogen ions or hydroxide ions are
added to a solution? Why is this phenomenon important
in organisms?
34
Chapter 2
Atoms, Molecules, and Life
Applying the Concepts
1. Many “over-the-counter” substances are sold to bring relief from “acid stomach” or “heartburn.” What is the
chemical basis for these compounds? Why do they work?
3. What would the effects be for aquatic life if the density of
ice were greater than that of liquid water? What would be
the impact on terrestrial organisms?
2. Fats and oils do not dissolve in water; polar and ionic molecules dissolve easily in water. Detergents and soaps help
clean by dispersing fats and oils in water so that they can
be rinsed away. From your knowledge of the structure of
water and the hydrophobic nature of fats, what general
chemical structures (for example, polar or nonpolar parts)
must a soap or detergent have, and why?
4. How does sweating help you regulate your body temperature? Why do you feel hotter and more uncomfortable on
a hot, humid day than on a hot, dry day?
For More Information
Atkins, P. W. Molecules. New York: Scientific American Library, 1987.
A layperson’s introduction to atoms and molecules, with superb illustrations.
Raloff, J. “Chocolate Hearts.” Science News, March 18, 2000. Describes
recent research indicating that chocolate is high in antioxidants.
Glasheen, J. W., and McMahon, T. A. “Running on Water.” Scientific
American, September 1997. Answers the question, “How does the
basilisk lizard run on water?”
Storey, K. B., and Storey, J. M. “Frozen and Alive.” Scientific American,
December 1990. By triggering ice formation here, suppressing it
there, and loading up their cells with antifreeze molecules, some animals, including certain lizards and frogs, can survive with 60% of
their body water frozen solid.
Morrison, P., and Morrison, P. Powers of Ten. New York: W. H. Freeman, 1982. A fascinating journey from the universe to the nucleus of
an atom.
Answers to Multiple-Choice Questions
1. d
2. b
3. c
4. b
5. d
6. a
MediaTutor
35
MEDIATUTOR
Atoms, Molecules, and Life
CD Activities
Web Investigations
Estimated time: 5 minutes
This tutorial provides an introduction to atomic
structure and chemical bonding. An understanding
of simple chemistry is critical to understanding biology. This tutorial will provide the background to
learn about more complex biological molecules.
The bonds that hold together large, complex, biologically important molecules like DNA are the
same as the bonds we will learn about here.
Estimated time: 10 minutes
Ah, creamy, sweet, delicious chocolate! Most people love it. Some even think it’s addictive. This exercise takes a closer, scientific look at an old favorite.
•
Activity 2.1: Interactive Atoms
Activity 2.2: Water and Life
Estimated time: 5 minutes
Water is essential for life. In fact, water may be the
only absolute essential for living systems. This tutorial explores the properties of water and how they
relate to living systems.
Start the MediaTutor Student CD-ROM and enter
the activity number in the Quick Search box to be
taken directly to that activity.
•
Case Study: Health Food?
Go to http://www.prenhall.com/audesirk6, the Audesirk Companion Web site. Select Chapter 2 and the
Web Investigation to begin.