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Periodic Trends: Electronegativity
Learning Objectives
• Define electronegativity.
• Describe how the electronegativity changes within a period.
• Describe how the electronegativity changes within a group.
• Analyze the importance of electronegativity in determining bond polarity
Periodic trends are specific patterns that are present in the periodic table that illustrate different
aspects of a certain element, including its size and its electronic properties. Major periodic trends
include: electronegativity, ionization energy, electron affinity, atomic radius, melting point, and
metallic character. Periodic trends, arising from the arrangement of the periodic table, provide
chemists with an invaluable tool to quickly predict an element's properties. These trends exist
because of the similar atomic structure of the elements within their respective group families or
periods, and because of the periodic nature of the elements. We have already studied metallic
character, atomic radius, and ionization energy. We are now going to have a closer look at
electronegativity.
Electronegativity Trends
Electronegativity can be understood as a chemical property describing an atom's ability to attract
and bind with electrons. Because electronegativity is a qualitative property, there is no
standardized method for calculating electronegativity. However, the most common scale for
quantifying electronegativity is the Pauling scale (Table A2), named after the chemist Linus
Pauling. The numbers assigned by the Pauling scale are dimensionless due to the qualitative nature
of electronegativity. Electronegativity values for each element can be found on certain periodic
tables. An example is provided below.
Electronegativity measures an atom's tendency to attract and form bonds with electrons. This
property exists due to the electronic configuration of atoms. Most atoms follow the octet rule
(having the valence, or outer, shell comprise of 8 electrons). Because elements on the left side of
the periodic table have less than a half-full valence shell, the energy required to gain electrons is
significantly higher compared with the energy required to lose electrons. As a result, the elements
on the left side of the periodic table generally lose electrons when forming bonds. Conversely,
elements on the right side of the periodic table are more energy-efficient in gaining electrons to
create a complete valence shell of 8 electrons. The nature of electronegativity is effectively
described thus: the more inclined an atom is to gain electrons, the more likely that atom will pull
electrons toward itself.
• From left to right across a period of elements, electronegativity increases. If the
valence shell of an atom is less than half full, it requires less energy to lose an electron than
to gain one. Conversely, if the valence shell is more than half full, it is easier to pull an
electron into the valence shell than to donate one.
• From top to bottom down a group, electronegativity decreases. This is because atomic
number increases down a group, and thus there is an increased distance between the valence
electrons and nucleus, or a greater atomic radius.
• Important exceptions of the above rules include the noble gases, lanthanides, and
actinides. The noble gases possess a complete valence shell and do not usually attract
electrons. The lanthanides and actinides possess more complicated chemistry that does not
generally follow any trends. Therefore, noble gases, lanthanides, and actinides do not have
electronegativity values.
• As for the transition metals, although they have electronegativity values, there is little
variance among them across the period and up and down a group. This is because their
metallic properties affect their ability to attract electrons as easily as the other elements.
Patterns of electronegativity in the Periodic Table
The distance of the electrons from the nucleus remains relatively constant in a periodic table row,
but not in a periodic table column. The force between two charges is given by Coulomb’s law.
F=kQ1Q2
r2
In this expression, Q represents a charge, k represents a constant and r is the distance between the
2 2 2
charges. When r = 2, then r = 4. When r = 3, then r = 9. When r = 4, then r = 16. It is readily seen
from these numbers that, as the distance between the charges increases, the force decreases very
rapidly. This is called a quadratic change.
The result of this change is that electronegativity increases from bottom to top in a column in the
periodic table even though there are more protons in the elements at the bottom of the column.
Elements at the top of a column have greater electronegativities than elements at the bottom of a
given column.
The overall trend for electronegativity in the periodic table is diagonal from the lower left corner
to the upper right corner. Since the electronegativity of some of the important elements cannot be
determined by these trends (they lie in the wrong diagonal), we have to memorize the following
order of electronegativity for some of these common elements.
F > O > Cl > N > Br > I > S > C > H > metals
The most electronegative element is fluorine. If you remember that fact, everything becomes easy,
because electronegativity must always increase towards fluorine in the Periodic Table.
According to these two general trends, the most electronegative element is fluorine, with 3.98
Pauling units. The least is cesium, at 0.79. Francium is rated as lower, but as it is a radioactive
element, its reactivity is not typically a consideration.
SUMMARY
Trends
• Electronegativity refers to the ability of a nucleus to attract electrons or to retain electrons
during chemical bonding.
• The electronegativity of the elements within a period generally increases from left to
right. This is because the nuclear charge is increasing faster than the electron shielding,
so the attraction that the atoms have for the valence electrons increases.
• The electronegativity of the elements within a group generally decreases from top to
bottom. This is because as you go from top to bottom down a group, the atoms of each
element have an increasing number of energy levels. The electrons in a bond are thus
farther away from the nucleus and are held less tightly.
• Atoms with low ionization energies have low electronegativities because their nuclei do
not have a strong attraction for electrons. Atoms with high ionization energies have high
electronegativities because the nucleus has a strong attraction for electrons.
• Although the noble gases possess very high ionization energies, He, Ne and Ar do not
have listed electronegativity values, as they do not bond with other elements. Kr and Xe
do form
QUESTIONS:
1) What is electronegativity?
2) Considering also the periodic trends in atomic radius and ionization energy, explain why
fluorine has the highest electronegativity.
3) Why are there no values of EN for He, Ne and Ar?
4) a. What is the trend in EN across a period (row) from left to right?
b. What causes this trend?
5) a. What is the trend in EN down a group (column) from top to bottom?
b. What causes this trend?
6) In each pair, select the element which has the higher electronegativity:
a. N and As
b. Mg and Sr
c. Na and S
d. K and Br
7) Which group would generally have the lowest electronegativity?
a. Transition Metals (Groups 3-12)
b. Alkali Metals (Group 1)
c. Noble Gases (Group 18)
d. Alkaline Earth Metals (Group 2)
e. Halogens (Group 17)
Justify your response.
9) Low electronegativty is considered a property of
a. Metals
b. Nonmetals
Justify your response
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