Electrons in atoms
It is the electrons in the atoms which participate in bonding and are exchanged during chemical
reactions. Having a model for how these electrons are arranged in the atom helps us
understand the energy of chemical reactions as well as the shapes of molecules. We write
electron configurations to describe the electron arrangement in an atom or ion.
First: Dispense with the idea of an electron in an atom being a particle. In atoms, it is more
useful to think of the electron as a wave. In particular, a standing wave whose general shape is
fixed in space.
These standing waves are the modern idea of orbitals. Bohr’s old model of point-like electrons
orbiting the nucleus is no longer accepted as correct, although the arrangement of electrons at
different distances from the nucleus survives in the size of the orbitals. This is the principal
quantum number n, which corresponds to the period or row number on the periodic table.
There are several different orbital shapes which show
up.
The lowest energy orbital at any distance is the s
orbital, spherical in shape.
Beginning with the second period, there are p orbitals
with 2 lobes and a node between where the electron
is never found. Beginning with the third level, there
are four-lobed d orbitals as shown at right.
The last type of orbital which has electrons in the
ground state of any atom is the f orbital with 8 lobes
shown below.
As we consider atoms with higher atomic numbers,
the nucleus has a greater positive charge and pulls
the electrons in tighter, shrinking the size of the
orbitals. This results in the outermost level of orbitals
only increasing very slightly through the periodic table
and the inner orbitals shrinking inside the electron
cloud.
This shrinking of the orbitals is accompanied by a drop
in energy of the electrons in them.
As electrons are placed in the atom, they fill from the
lowest energy up. A diagram showing the energies of
the orbitals in a heavier atom is given at right.
As an example of an electron configuration, consider a neutral oxygen atom, O, which has 8
electrons.
Put the electrons in the orbitals starting at the lowest energy orbital.
Each orbital can hold 2 electrons with opposite spins.
The 1s orbital is filled first, with 2 electrons with opposite spins (opposite magnetic poles).
The next 2 electrons fill the 2s orbital leaving 4 electrons in the 2p orbitals.
This is written O: 1s2 2s2 2p4 or, as a box diagram, _↑↓_ _↑↓_ _↑↓ _↑_ _↑_
1s
2s
2p
The notation 2p4 indicates that the 2p orbitals has 4 electrons in them. In the box diagram, the
arrows represent electrons with spin up ↑ or spin down ↓.
When placing electrons in the p,d or f orbitals, the electrons are placed in different orbitals in the
set to keep the negative charges separated, keeping the energy low. This means that a d5
configuration will have one electron in each d orbital.
More examples of electron configurations:
N : 1s2 2s2 2p3 _↑↓_ _↑↓_ _↑_ _↑_ _↑_
1s
2s
2p
Mn: 1s2 2s2 2p6 3s2 3p6 4s2 3d5 _↑↓ ↑↓_ ↑↓ ↑↓_ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑_ ↑_ ↑_ ↑_ ↑_
1s 2s
2p
3s 3p
4s
3d
2
5
short form: [Ar] 4s 3d
There are a couple of ways to remember the filling order. One is a list of the orbital types by
primary level which is then traversed in a diagonal direction to give the filling order:
Another way is to associate the orbitals with the periodic table regions:
The f-block at the bottom should really be inserted between the 6s,7s and 5d,4d rows. Notice
that the structure of the periodic table, based on chemical similarities of the elements,
corresponds exactly to the highest energy orbitals of those elements as above.
There is one further wrinkle encountered in the middle of the d block. In this region, the d
orbitals and the previous s orbitals are nearly the same energy (remember I said the energy of
the inner orbitals drops as the atomic number increases). This leads to a stability of half filled
shells (one electron per orbital). A diagram of the energies of the orbitals as a function of
atomic number shows the
situation:
Look in the region marked A
in the diagram. Just prior to
this rectangle, the 4s orbital
(blue) is lower in energy
than the 3d (red) and so it
fills first. As we start into the
first row of the d block, the
energy of the 3d orbitals
drops below that of the 4s,
that is why the 4s electrons
are removed before the 3d
electrons. Also, from atomic
numbers 22 to 29, the
energies of the 3d and 4s
orbitals are so close that it is
cheaper (lower energy) to
spread electrons out among
the 3d and 4s orbitals.
Notice this also occurs for
the 4d and 5d blocks and for
the f block, which you are
likely not to be asked to deal
with.
This means, for example, that the electron configuration for chromium would be
Cr: [Ar] 4s1 3d5
instead of
[Ar] 4s2 3d4 .
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