Linus Pauling. In addition to single covalent bonds, double and triple bonds also exist.
These kinds of bonds are more exactly described by the quantum-mechanical atom
model, in which the electron shells of an atom can be described by one of several dif-
ferently shaped orbitals that represent the areas where the electrons are located with
highest probability (electron clouds). A covalent bond is then described by molecule
orbitals, which are derived from atom orbitals. Furthermore, if single and double
bonds are altered in a single molecule or a double bond is in the direct vicinity of an
atom with a free electron pair, then one electron pair of the double bond and the free
electron pair can delocalize across the participating atoms, e.g., the three electron
pairs in benzol (Fig. 2.2b) or the double bond between C and O and the free electron
pair of N in a peptide bond (Fig. 2.6 a). Such electrons are called de-localized p-elec-
trons. For a more detailed description, please consult books about general and anor-
ganic chemistry or introductory books about biochemistry.
Hydrogen atoms with a positive partial charge that are bound to oxygen or nitro-
gen (as in H
2
OorNH
3
) are able to interact with free electron pairs of atoms with a
negative partial charge. These attractions are called hydrogen bonds and are rela-
tively weak compared to solid-state ionic bonds or covalent bonds. To break a hydro-
gen bond, only about 4 kJ mol
–1
is required. Therefore, hydrogen bonds separate
readily at elevated temperatures, which is often the reason that proteins such as
enzymes lose their function during heating. Likewise, the hydrogen bonds that hold
together the double strands of nucleic acids (see Section 2.2.3) can be separated at
high temperatures. This fact is utilized for several molecular biological methods,
e.g., polymerase chain reaction (PCR) and radioactive labeling of DNA (deoxyribonu-
cleic acid) fragments. Hydrogen bonds also explain why water is liquid at room tem-
perature and boils at 100 8C. Small alcohols, such as methanol or ethanol, are fully
soluble in water due to their hydroxyl group, which interacts with the hydrogen
bonds of water, whereas larger alcohols, such as hexanol or heptanol, are weakly so-
luble or insoluble in water due to their longer unpolar carbohydrate tail. As we have
seen, polarized functional groups can interact with water, which is why they often
are called hydrophilic (or lipophobic), while nonpolar molecules or molecule parts
are called hydrophobic (or lipophilic).
Also critical to structures and interactions of biological molecules are the van der
Waals forces. The electron clouds surrounding atoms that are held together by cova-
lent bonds are responsible for these forces. Momentary inequalities in the distribu-
tion of electrons in any covalent bond, due to chance, can make one end of the cova-
lent bond more negative or positive than the other for a short moment, which results
in rapid fluctuations in the charge of the electron cloud. These fluctuations can set
up opposite fluctuations in nearby covalent bonds, thus establishing a weak attrac-
tive force. This attractive force is stronger the closer the electron clouds are, but if
the outermost electron orbitals begin to overlap, the negatively charged electrons
strongly repel each other. Thus, van der Waals forces can be either attractive or repul-
sive. Their binding affinity is, at 0.4 kJ mol
–1
in water, even lower than that of hydro-
gen bonds. The optimal distance for maximum van der Waals forces of an atom is
called its van der Waals contact radius. The van der Waals repulsions have an impor-
tant influence on the possible conformations of a molecule.
25
2.2 Molecular Biology of the Cell