The electronic configurations of the first row d-block
elements are given in the table below, along with those for the M2+
and M3+
ions. Because the 3d orbitals are all of the same energy, they are each filled
with a single electron first. Only after
each of the five 3d orbitals is singly filled (3d5), do the electrons
start pairing up, from 3d6 onwards.
For a free atom, the 4s orbital is normally filled
before the 3d orbitals. This means that most first row d-block elements have
the electronic configuration 3dn 4s2.
Chromium (3d5 4s1)
and copper (3d10 4s1)
are exceptions to this.
The 4s orbital is more diffuse than the 3d orbitals and
is affected more by the presence of other atoms or by the charge on the metal.
As a result, in an ion or a compound, the 4s orbital is higher in energy than
the 3d orbitals. This means that, when electrons are lost to form ions, it is
the 4s electrons that are lost first. This is reflected in the electronic
configurations of the d-block elements in ions and compounds , as these never
contain 4s electrons unless the d
orbitals are full. For example, the electronic configuration of V2+
is 3d3
not 3d1
4s2,
and the electronic configuration of Cr(0) in a compound is 3d6,
not 3d5
4s1
as it is in atomic chromium.
IMPORTANT: In a compound of a first row d-block element,
the 4s orbital is higher in energy than the 3d orbitals. For oxidation states
of +2 and higher, the electronic configuration can be found by removing the 4s
electrons, plus the appropriate number of 3d electrons, but for lower oxidation
states the 4s electrons must be transferred into 3d orbitals before removing
electrons.
The video below gives examples of how to obtain electronic
configurations for:
•Fe3+,
•Fe(0), which is
different from atomic iron,
•Ni2+,
•Ni(0),
•Co, and,
•Co+.
