20
Chapter 2
or more other atoms. Each type of atom forms a characteris-
tic number of covalent bonds, which depends on the number
of electrons in its outermost orbit. The number of chemical
bonds formed by the four most abundant atoms in the body
are hydrogen, one; oxygen, two; nitrogen, three; and car-
bon, four. When the structure of a molecule is diagrammed,
each covalent bond is represented by a line indicating a pair of
shared electrons. The covalent bonds of the four elements just
mentioned can be represented as
A
A
H—
—O—
—N—
—C—
A
A molecule of water, H
2
O, can be diagrammed as
H—O—H
In some cases, two covalent bonds—a double bond—form
between two atoms when they share two electrons from each
atom. Carbon dioxide (CO
2
) contains two double bonds:
O
P
C
P
O
Note that in this molecule the carbon atom still forms four
covalent bonds and each oxygen atom only two.
Molecular Shape
When atoms are linked together, they form molecules with
various shapes. Although we draw diagrammatic structures of
molecules on fl at sheets of paper, molecules are three-dimen-
sional. When more than one covalent bond is formed with a
given atom, the bonds are distributed around the atom in a
pattern that may or may not be symmetrical (
Figure 2–2
).
Molecules are not rigid, infl exible structures. Within
certain limits, the shape of a molecule can be changed with-
out breaking the covalent bonds linking its atoms together. A
covalent bond is like an axle around which the joined atoms
can rotate. As illustrated in
Figure 2–3
, a sequence of six car-
bon atoms can assume a number of shapes by rotating around
various covalent bonds. As we will see, the three-dimensional,
fl exible shape of molecules is one of the major factors govern-
ing molecular interactions.
Ions
A single atom is electrically neutral because it contains equal
numbers of negative electrons and positive protons. If, how-
ever, an atom gains or loses one or more electrons, it acquires
a net electric charge and becomes an
ion.
For example, when
a sodium atom (Na), which has 11 electrons, loses one elec-
tron, it becomes a sodium ion (Na
+
) with a net positive charge;
it still has 11 protons, but it now has only 10 electrons. On
the other hand, a chlorine atom (Cl), which has 17 electrons,
can gain an electron and become a chloride ion (Cl
) with
a net negative charge—it now has 18 electrons but only 17
protons. Some atoms can gain or lose more than one electron
to become ions with two or even three units of net electric
charge (for example, the calcium ion Ca
2+
).
Although the number of neutrons in the nucleus of an
atom is often equal to the number of protons, many chemical
elements can exist in multiple forms, called
isotopes,
which
differ in the number of neutrons they contain. For example,
the most abundant form of the carbon atom,
12
C, contains
6 protons and 6 neutrons, and thus has an atomic number
of 6. Protons and neutrons are approximately equal in mass.
Therefore,
12
C has an atomic weight of 12. The radioactive
carbon isotope
14
C contains 6 protons and 8 neutrons, giving
it an atomic number of 6 but an atomic weight of 14.
One
gram atomic mass
of a chemical element is the
amount of the element, in grams, equal to the numerical value
of its atomic weight. Thus, 12 g of carbon (assuming it is all
12
C) is 1 gram atomic mass of carbon.
One gram atomic mass
of any element contains the same number of atoms.
For example,
1 g of hydrogen contains 6
×
10
23
atoms, and 12 g of carbon,
whose atoms have 12 times the mass of a hydrogen atom, also
has 6
×
10
23
atoms (the so-called Avogadro’s number).
Atomic Composition of the Body
Just four of the body’s essential elements (see Table 2–1)—
hydrogen, oxygen, carbon, and nitrogen—account for over 99
percent of the atoms in the body.
The seven essential mineral elements are the most abun-
dant substances dissolved in the extracellular and intracellu-
lar fl uids. Most of the body’s calcium and phosphorus atoms,
however, make up the solid matrix of bone tissue.
The 13 essential
trace elements
are present in extremely
small quantities, but they are nonetheless essential for normal
growth and function. For example, iron plays a critical role in
the blood’s transport of oxygen.
Many other elements, in addition to the 24 listed in
Table 2–1, may be detected in the body. These elements enter
in the foods we eat and the air we breathe but are not essential
for normal body function and may even interfere with normal
body chemistry. For example, ingested arsenic has poisonous
effects.
Molecules
Two or more atoms bonded together make up a
molecule.
For
example, a molecule of water contains two hydrogen atoms
and one oxygen atom, which can be represented as H
2
O. The
atomic composition of glucose, a sugar, is C
6
H
12
O
6
, indicat-
ing that the molecule contains 6 carbon atoms, 12 hydrogen
atoms, and 6 oxygen atoms. Such formulas, however, do not
indicate how the atoms are linked together in the molecule.
Covalent Chemical Bonds
The atoms in molecules are held together by chemical bonds,
which form when electrons transfer from one atom to another
or when two atoms share electrons. The strongest chemical
bond between two atoms, a
covalent bond,
forms when one
electron in the outer electron orbit of each atom is shared
between the two atoms (
Figure 2–1
). The atoms in most
molecules found in the body are linked by covalent bonds.
The atoms of some elements can form more than one
covalent bond and thus become linked simultaneously to two
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