James Lauer, Matt Rose,
Tom Greenwood, Eric Gu
The ability to pull electrons in a covalent bond
towards the nucleus of an atom
Affected by atomic number and the distance valence
electrons are from the nucleus
Cannot be directly measured and must be calculated from other sources,
mostly through measuring how the energy of a bond changes when an
atom is added.
Xa - Xb = {[Ed(AB) - 1/2[Ed(AA) + Ed(BB)]}^1/2
Most commonly measured in Pauling units
Pauling scale runs from 0.7 to
3.98, with smaller quantities
experiencing less attraction and
larger quantities experiencing
greater attraction
Electronegativity generally increases from left groups
to right groups and it decreases when going down the
periods of the periodic table.
Increases from left groups to right groups because of
the increase in the nucleus charge.
This means the atom will have an increased attraction
for its outermost electrons.
Decreases from the top period to the bottom period
because of they have more valence shells and
electrons are further away from the nucleus.
Noble gases, Lanthanides, and Actinides have a
complete valence shell so they tend to not attract
Transition metals’ properties affect their ability to
attract electrons as easily
as the other elements
thus is little variance
among them.
Gallium and Germanium have higher electronegativity
than Aluminum and Silicon which are above them. This
is because of the D-Block Contraction.
D-Block: The D electrons are not good at shielding the
nuclear charge, so the atomic radius
does not change much as electrons
are added. Almost like disregarding
the D electrons being added.