Chemistry Reference
In-Depth Information
5
Ne
4.5
F
4
O
3.5
Ar
N
Kr
Cl
S
P
Si
Al
Mg
Na
3
Br
Se
A
Ge
Ga
C
Xe
I
Te
S
Sn
In
2.5
B
2
Be
1.5
Ca
K
Li
1
Sr
Rb
0.5
2
3
4
5
Period
FIGURE 3.4
Spectroscopic electronegativities for groups I-VIII of representative elements.
(Adapted from L.C. Allen, “Electronegativity Is the Average One-Electron Energy of
the Valence-Shell Electrons in Ground-State Free Atoms.
J. Am. Chem. Soc
. 111, no. 25
(1989):9003-9014.)
numerically similar to Pauling electronegativities. The one-electron energies can be
determined directly from spectroscopic data that are available for almost all ele-
ments. This method allows the estimation of electronegativities for elements that
cannot be generated otherwise. However, it is unclear what should be considered the
valence electrons for the d- and f-block elements. This leads to an ambiguity for their
electronegativities calculated by the Allen method.
For the Allen electronegativity, several general trends can be noted:
•
Noble gases are the most electronegative elements in their period
(Figure 3.4).
•
Neon has the highest electronegativity of all elements, followed by fluorine,
helium, and oxygen.
•
Elements with
n
= 2 valence orbitals are significantly more electronegative
than the other elements in their respective groups.
•
Electronegativity decreases down a group, except: B > Al < Ga and C > Si < Ge.
This last trend is called the
alternation effect
and is caused by the increased
nuclear charge that accompanies the filling of the 3d orbitals. The 3d electrons shield
the 4p electrons poorly, making Ga and Ge more electronegative.
Electronegativity was used as descriptor in some QSAR studies related to the
metal ion toxicity (Somers 1959; Biesinger and Christensen 1972; Khangarot and
Ray 1989; Enache et al. 2003).
Modern definitions of electronegativity were proposed on the basis of the density
functional theory (DFT).
