Chemistry - A Molecular Science

Chapter 3 Atomic Structure and Properties

Figure 3.6 shows the periodic trend in ionization energies of the main group elements.

Note that the rise in the ionization energies fro

m left to right within in a period follows

that of the effective nuclear charge, while th

e steady decrease within a group is due to an

increase in the n quantum number. In general,

Metals have low ionization energies because they have low effective nuclear charges. Consequently, they readily lose valence electrons to form cations. Nonmetals have high ionization energies, so they do not form cations in chemical processes

.

Electron configurations in which sublevels are filled (Groups 2A and 8A) or half-filled

(Group 5A) are unusually stable

, so removing an electron from these elements is more

difficult and results in deviations from the exp

ected periodicity. For example, the effective

nuclear charge of B is greater than that of

Be, but the ionization energy of Be is greater

than that of B because the electron must be removed from a filled 2s sublevel in Be. Similarly, the ionization energy of N is greater than that of O because the 2p sublevel of N is half-filled.

3.6

ELECTRONEGATIVITY

Electronegativity

(χ

) is a measure of an atom's ability to attract bonding electrons, so

electron density in a bond accumulates near those atoms with higher electronegativities. Bonding electrons reside in orbitals involving

the valence orbitals of the atoms, especially

those that are unfilled, and electrons seek to minimize their energy, so an atom that is highly electronegative is simply one whose va

lence orbitals, especially those that are

unfilled, are low in energy. E

(^) n
∝^
-(Z
)eff
2 /n
2 , so atoms with large
Z
(nonmetals) and eff
valence orbitals with low n quantum numbers have high electronegativities.
Electronegativity depends upon the energy of
the valence orbitals. It increases as the
energy of the valence orbitals decreases.
Consider the orbital energies of Li, C, and F shown in Figure 3.7. Li is a metal with a low Zeff
, so its orbital energy is high. The electron in
it is readily lost (low ionization energy),
but bonding electrons are not drawn to the high-energy orbital, so Li has a very low electronegativity. F is a nonmetal with a high Z
, so its orbital energy is low. Thus, it is eff
very difficult to remove a 2p electron from F (high ionization energy), but bonding electrons are drawn to the low energy unfilled orbital, so F is highly electronegative. The 2p orbital energy of C is about half way between the valence orbital energies of Li and F, and its electronegativity is also about half
way between these two extremes in the period.
0
2500200015001000500
Atomic Number
Ionization Energy
(kJ/mol)
H
He Be Li
N C B
F O
Ne Na
Mg Al
P SSi
Ar Cl K
Ca Ga
Ge
Br Se
Kr
As
Sr Rb
Sb SnIn
XeI Te
Ba Cs
Pb Tl
At PoBi
Rn
Figure 3.6 Ionization energies of the main group elements Circles of the same color are used for elements in the same group.
Li 2s
C2p
F2p
0
Energy (kJ/mol)
-500 -1000-1500
c
= 1.0 c
= 2.5
c
= 4.0
Figure 3.7 Electronegativity and orbital energy Electronegativities of Li, C, and F. Atoms with lower energy valence orbitals have higher electronegativities.
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