Essential Cell Biology 5th edition

Chemical Bonds
45
Four electrons can be shared, for example, two coming from each participating
atom; such a bond is called a double bond. Double bonds are
shorter and stronger than single bonds and have a characteristic effect
on the geometry of molecules containing them. A single covalent bond
between two atoms generally allows the rotation of one part of a molecule
relative to the other around the bond axis. A double bond prevents
such rotation, producing a more rigid and less flexible arrangement of
atoms (Figure 2–10). This restriction has a major influence on the threedimensional
shape of many macromolecules.
Some molecules contain atoms that share electrons in a way that produces
bonds that are intermediate in character between single and
double bonds. The highly stable benzene molecule, for example, is
made up of a ring of six carbon atoms in which the bonding electrons
are evenly distributed, although the arrangement is sometimes depicted
as an alternating sequence of single and double bonds. Panel 2–1
(pp. 66–67) reviews the covalent bonds commonly encountered in biological
molecules.
Electrons in Covalent Bonds Are Often Shared
Unequally
When the atoms joined by a single covalent bond belong to different elements,
the two atoms usually attract the shared electrons to different
degrees. Covalent bonds in which the electrons are shared unequally in
this way are known as polar covalent bonds. A polar structure (in the electrical
sense) is one in which the positive charge is concentrated toward
one atom in the molecule (the positive pole) and the negative charge is
concentrated toward another atom (the negative pole). The tendency of
an atom to attract electrons is called its electronegativity, a property
that was first described by the chemist Linus Pauling.
Knowing the electronegativity of atoms allows one to predict the nature
of the bonds that will form between them. For example, when atoms
with different electronegativities are covalently linked, their bonds will
be polarized. Among the atoms typically found in biological molecules,
oxygen and nitrogen (with electronegativities of 3.4 and 3.0, respectively)
attract electrons relatively strongly, whereas an H atom (with an
electronegativity of 2.1) attracts electrons relatively weakly. Thus the
covalent bonds between O and H (O–H) and between N and H (N–H) are
polar (Figure 2–11). An atom of C and an atom of H, by contrast, have
similar electronegativities (carbon is 2.6, hydrogen 2.1) and attract electrons
more equally. Thus the bond between carbon and hydrogen, C–H,
is relatively nonpolar.
Covalent Bonds Are Strong Enough to Survive the
Conditions Inside Cells
We have already seen that the covalent bond between two atoms has
a characteristic length that depends on the atoms involved (see Figure
2–10). A further crucial property of any chemical bond is its strength.
Bond strength is measured by the amount of energy that must be supplied
to break the bond, usually expressed in units of either kilocalories
per mole (kcal/mole) or kilojoules per mole (kJ/mole). A kilocalorie
is the amount of energy needed to raise the temperature of 1 liter of
water by 1°C. Thus, if 1 kilocalorie of energy must be supplied to break
6 × 10 23 bonds of a specific type (that is, 1 mole of these bonds), then the
strength of that bond is 1 kcal/mole. One kilocalorie is equal to about
4.2 kJ, which is the unit of energy universally employed by physical scientists
and, increasingly, by cell biologists as well.
(A) ethane
(B) ethene
Figure 2–10 Carbon–carbon double
bonds are shorter and more rigid than
carbon–carbon single bonds. (A) The
ethane molecule, with a single covalent
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bond between the two carbon atoms, shows
the tetrahedral arrangement of the three
single covalent bonds between each carbon
atom and its three attached H atoms. The
CH 3 groups, joined by a covalent C–C
bond, can rotate relative to one another
around the bond axis. (B) The double
bond between the two carbon atoms in a
molecule of ethene (ethylene) alters the
bond geometry of the carbon atoms and
brings all the atoms into the same plane;
the double bond prevents the rotation of
one CH 2 group relative to the other.
H
δ –
O
δ + δ +
water
H
O
oxygen
O
Figure 2–11 In polar covalent bonds,
the electrons are shared unequally.
Comparison of electron distributions in the
polar covalent bonds in a molecule of water
(H 2 O) and the nonpolar covalent bonds in a
molecule of oxygen (O 2 ). In H 2 O, electrons
are more strongly attracted to the oxygen
nucleus than to the H nucleus, as indicated
by the distributions of the partial negative
(δ – ) and partial positive (δ + ) charges.
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