Download Electronegativity:

APPENDIX
6
Electronegativity:
A Quick Reference for the Biology Student
Electronegativity and the
Periodic Table of the Elements
Electronegativity and
Chemical Bonds
Chemist Linus Pauling, winner of the Nobel prize,
first defined the concept of electronegativity in
the 1930s. The electronegativity of an element
is a measure of its ability to attract electrons in
a bond. This chemical property is relevant only
to bonding between atoms, not to individual
atoms. Pauling created a quantitative scale of
the electronegativity values of the elements
listed in the periodic table. In general,
electronegativity increases from left to right
along each row (period) of the periodic table.
Electronegativity also tends to increase from
the bottom to the top of each column (group).
Fluorine has the highest electronegativity
value (4.0), so fluorine is said to be the most
electronegative element. Francium and cesium
each have the lowest electronegativity value (0.7),
so they are the least electronegative elements.
The following table lists the electronegativity
values of some biologically important elements.
Two atoms of the same element have the same
electronegativity, or the same ability to attract
electrons in a bond. Therefore, if two atoms of
the same element form a bond, as in a molecule
of hydrogen (H2 ), oxygen (O2 ), or nitrogen (N2 ),
the bonding electrons are equally shared
between the atoms. A bond in which electrons
are shared between atoms is called a covalent
bond. If the electrons are equally shared, the
bond is said to be non-polar covalent.
If a covalent bond is formed between atoms
of two different elements, such as hydrogen and
chlorine in hydrogen chloride (HCl), the
electronegativities of the elements determine
how the bonding electrons are distributed.
Because the electronegativities are different, the
bonding electrons are shared unequally. A
covalent bond in which electrons are shared
unequally is called a polar covalent bond. As
shown in Table A6.1, chlorine is more
electronegative than hydrogen, so chlorine has
the stronger attraction for the bonding electrons
in a hydrogen chloride molecule. For this
reason, the chlorine atom becomes slightly
negative, and the hydrogen atom becomes
slightly positive. The slightly charged ends of
the bond are known as poles. Because two
poles are present, the bond is said to have a
dipole. The slightly positive and negative ends
of the polar covalent bond can be represented
using the symbols δ+ and δ−, as shown in
Figure A6.1.
Table A6.1
Electronegativity values of some biologically
important elements
Symbol
Electronegativity
potassium
K
0.8
sodium
Na
0.9
calcium
Ca
1.0
iron
Fe
1.8
hydrogen
H
2.1
phosphorus
P
2.2
carbon
C
2.5
sulfur
S
2.6
chlorine
CI
3.0
nitrogen
N
3.0
oxygen
O
3.5
Element
δ+ δ−
H Cl
Figure A6.1 The bond between hydrogen and chlorine
in a hydrogen chloride molecule is polar covalent.
If two elements form a covalent bond, the
difference between their electronegativities
determines how polar the bond is. Table A6.1
shows that the electronegativity difference
between carbon and oxygen is 1.0, whereas the
Appendix 6 • MHR
563
electronegativity difference between carbon and
hydrogen is 0.4. Therefore, the sharing of
electrons is more unequal between carbon and
oxygen than between carbon and hydrogen. In
other words, a carbon–oxygen bond is more
polar than a carbon–hydrogen bond.
When two elements have very different
electronegativities, the bonds the elements form
can be described in terms of electron transfers,
rather than electron sharing. For example,
sodium and chlorine have an electronegativity
difference of 2.1. When these elements react,
electrons are transferred from sodium atoms
to chlorine atoms to form positively charged
sodium ions (Na+ ) and negatively charged
chloride ions (Cl−). These ions attract each
other. Bonds created by the transfer of electrons
from one type of atom to another are called
ionic bonds.
There is no sharp distinction between polar
covalent bonds and ionic bonds. The greater
the difference in the electronegativities of two
bonded atoms, the more polar is the bond
between them. When sharing is very unequal,
chemists describe the electrons as being
transferred from one atom to the other. For
convenience, chemists have chosen an
approximate value of the electronegativity
difference that distinguishes polar covalent
bonds from ionic bonds. In general, if the
electronegativity difference equals or exceeds
1.7, the elements will tend to form ionic bonds.
If the electronegativity difference is less than 1.7,
the elements will tend to form covalent bonds.
If the electronegativity difference is greater than
0 but less than 1.7, the bonds will be polar
covalent. Strictly speaking, bonds are non-polar
covalent only if the electronegativity difference
is exactly 0, that is, if two atoms of the same
element form a bond. However, when the
electronegativity difference is less than 0.5
(for example, 0.4 for C and H), the electrons
are shared almost equally, and the bond dipole
is weak.
The electronegativity of carbon is 2.5, which
is very close to the middle of the range of
electronegativity values, from 0.7 to 4.0. Carbon
can, therefore, form covalent bonds with
elements that are more electronegative, such
as oxygen, and with elements that are less
electronegative, such as hydrogen. Carbon
atoms are also able to form covalent bonds with
564
MHR • Appendix 6
other carbon atoms to produce a wide range of
compounds that contain chains or rings of
carbon atoms.
Molecular Polarity
You learned earlier that the bond in a hydrogen
chloride molecule (HCl) is polar covalent.
Because the bond is polar, and there is only one
bond in the molecule, the whole molecule is
also polar. However, most covalent molecules
contain more than two atoms and more than one
bond. In such cases, the polarity of a molecule
depends on the polarity of the individual
bonds and on the overall shape of the molecule.
Two examples that show the importance of
molecular shapes are carbon dioxide (CO2 ) and
water (H2O).
In a carbon dioxide molecule, the two highly
electronegative oxygen atoms form double
bonds to a less electronegative carbon atom,
as shown in Figure A6.2. The carbon–oxygen
bonds are polar covalent. The shape of the
molecule is straight or linear. As a result, the
two dipoles counterbalance each other, and the
molecule has no net dipole. Thus, even though
a carbon dioxide molecule contains polar bonds,
the molecule is non-polar.
δ−δ+δ+ δ−
O C O
Figure A6.2 A linear carbon dioxide molecule contains
polar covalent bonds, but the molecule is non-polar.
In a water molecule, one highly
electronegative oxygen atom forms single
bonds to two less electronegative hydrogen
atoms, as shown in Figure A6.3. The hydrogen–
oxygen bonds are polar covalent. In contrast
to the linear carbon dioxide molecule, a water
molecule is V-shaped or bent. As a result, the
two dipoles do not counterbalance each other,
and the molecule has a net dipole. Thus, a
water molecule contains polar bonds, and the
molecule is polar.
δ−δ−
O
δ+
δ+
H
H
Figure A6.3 Because a water molecule is bent and
contains polar covalent bonds, the molecule is polar.
Hydrocarbons, which are composed of
carbon and hydrogen, are an important class of
compounds in organic chemistry and biology.
Hydrocarbon molecules, some of which contain
large numbers of atoms, include carbon–hydrogen
bonds that are slightly polar. However, because
of the shapes of hydrocarbon molecules, these
molecules are either non-polar or very low
in polarity.
Some Biological Applications
The presence or absence of dipoles can have
a great effect on the physical and chemical
properties of covalent molecules. Therefore,
an understanding of dipoles is very important
in explaining many biological processes. Some
examples of how electronegativity is important
in biological processes are given below.
Functional Groups
Some important functional groups are listed in
Figure 1.24 in section 1.4. All of the functional
groups listed contain covalent bonds between
elements with different electronegativities, so
the functional groups contain polar covalent
bonds. The presence of these polar bonds in
biological molecules influences the physical
interactions of the molecules and their
chemical properties.
A hydroxyl group (−OH) contains a single
covalent bond between a highly electronegative
oxygen atom and a less electronegative hydrogen
atom. As in a water molecule, the oxygen atom
is at the slightly negative end of the dipole, and
hydrogen is at the slightly positive end. Hydroxyl
groups can attract some ions and polar molecules
in a type of interaction called a hydrogen bond,
which will be described below.
In a carbonyl group, highly electronegative
oxygen forms a double bond with less
electronegative carbon. Because the carbon
atom is at the more positive end of the dipole,
the carbon atom tends to react with substances
that are electron donors. The oxygen atom at the
more negative end of the dipole tends to react
with substances that are electron acceptors.
A carboxyl group contains both a carbonyl
group and a hydroxyl group connected together.
Because a carboxyl group contains two highly
electronegative oxygen atoms, the positive
charge on the hydrogen atom is larger than in a
hydroxyl group alone. In fact, carboxyl groups
partially dissociate to form hydrogen ions (H+ ).
Therefore, carboxyl groups are present in
organic acids, such as amino acids. An amino
acid molecule also contains an amino group
(−NH2). In this functional group, the nitrogen
atom is at the slightly negative end of each
nitrogen–hydrogen dipole. The nitrogen atom
can act as an electron donor, for example to a
hydrogen ion. Thus, an amino acid contains a
functional group (a carboxyl group) that can
release hydrogen ions and another functional
group (an amino group) that can react with them.
Hydrogen Bonds
When hydrogen atoms are bonded to small,
highly electronegative atoms, such as oxygen,
a weak force of attraction can exist between
molecules. This force is known as a hydrogen
bond. A hydrogen bond can be formed when a
negative ion (such as Cl−), or a slightly negative
atom in a dipole (such as the oxygen in an −OH
group) attracts a slightly positive hydrogen atom
in another dipole. Therefore, water molecules
form hydrogen bonds with each other, as shown
in Figure A6.4. Hydrogen bonds are only about
one-twentieth as strong as covalent bonds, but
hydrogen bonds have important effects on the
physical and chemical behaviour of many
molecules in biological systems.
δ−
hydrogen
bond
δ+
H
O
H
Figure A6.4 Hydrogen bonds exist between water
molecules and account for many of the physical and
chemical properties of water.
The hydrogen bonds between water molecules
make water an ideal medium for life processes
on Earth. Scientists believe that if there were no
hydrogen bonds present, then water would be a
gas, not a liquid, under the conditions in which
we live. The polarity of water molecules, and
their ability to form hydrogen bonds, explains
Appendix 6 • MHR
565
the fact that water can act as a solvent for so
many substances in living and non-living
systems. When ionic and polar substances
dissolve in water, they may form hydrogen
bonds to water molecules. Hydrogen bonds can
also bind the components of complex biological
substances, such as the chains of nucleic acids
that make up a DNA molecule. These two long
intertwined chains are held to each other by
thousands of hydrogen bonds. Proteins, such as
enzymes, also require hydrogen bonding to
maintain their shapes and their functions.
+
Energy for
synthesis of
2e−
ATP
ort
nsp
tra
in
cha
MHR • Appendix 6
2H+
on
ctr
566
2H
(from food)
Ele
Aerobic Cellular Respiration
Aerobic cellular respiration involves the
breakdown of glucose molecules (food) to form
energy-rich molecules called ATP, adenosine
triphosphate. Molecules of ATP are used to fuel
many other chemical reactions in cells. ATP
molecules are formed in several systems in
aerobic cellular respiration. The system that
produces the most ATP is called the electron
transport chain (see Figure A6.5). In the
electron transport chain, electrons (e−) flow
along the chain from one protein molecule to
another. As electrons move along the chain,
they release energy. This energy is used to
make ATP molecules. Once electrons reach the
end of the chain, however, they need to be
removed from the system. Removal of the
electrons is necessary in order to allow new
electrons to enter at the start of the chain. The
element that accepts the electrons at the end
of the electron transport chain is oxygen. As
you have learned, oxygen is one of the most
electronegative elements in the periodic table.
Oxygen accepts the electrons and combines
with hydrogen ions (H+ ), obtained from the
breakdown of glucose, to form water molecules.
Without oxygen to accept the electrons, the
electron transport chain could not function
and, as a result, could not produce ATP
molecules. This is why oxygen is so important
in cellular processes.
2e−
1
O
2 2
+
2H
H2O
Figure A6.5 The electron transport chain