The Parts of the Periodic Table

 

The electron configuration of an element is a list of the
atomic orbitals which are occupied by electrons, and how many electrons
are in each of those orbitals.  The rules for writing electron
configurations are known as the aufbau principle (German for
“building up”):

• Each electron that is added to an atom is placed in the
  lowest-energy orbital that is available.  The orbitals are
  filled in the order:

1s, 2s, 2p, 3s, 3p,
4s, 3d, 4p, 5s, 4d, 5p, 6s,
4f, 5d, 6p, 7s, 5f

• Each orbital can hold no more than two electrons. 
  Two electrons in the same orbital must have opposite spins (the
  Pauli exclusion principle).

• If two or more orbitals are available at the same
  energy level (degenerate orbitals), one electron is placed in
  each orbital until the available orbitals are occupied by one
  electron; any additional electrons are placed in the half-filled
  orbitals.

Electron configurations are written as a list of orbitals which are
occupied, followed by a superscript to indicate how many electrons are
in those orbitals.

H
1s1

He
1s2

Li
1s2
2s1

Be
1s2
2s2

B
1s2
2s2 2p1

C
1s2
2s2 2p2

N
1s2
2s2 2p3

O
1s2
2s2 2p3

F
1s2
2s2 2p4

Ne
1s2
2s2 2p5

Na
1s2
2s2 2p6 3s1

 

Electron configurations in which all of the electrons are in their
lowest-energy configurations are known as ground state configurations. 
If an electron absorbs energy, it can move into a higher-energy orbital,
producing an excited state configuration.

For atoms with a large number of electrons, the complete electron can
be very cumbersome, and not very informative.  For instance, the
complete configuration of the element radium is

Ra:  1s2 2s2 2p6
3s2 3p6 3d10 4s2 4p6
4d10 4f14 5s2 5p6 5d10
6s2 6p6 7s2

(With a description like that, you’d be radioactive too!)  Since
everything up to the 6p6 is the same electron configuration
as the noble gas radon, the configuration can be abbreviated as

Ra:  [Rn] 7s2

Abbreviated electron configurations are always based on the nearest
preceding noble gas.

Electron configurations can be written directly from the periodic
table, without having to memorize the aufbau scheme, using the following
patterns:

 

Half-filled and filled subshells are especially stable, leading to
some anomalous electron configurations:

 

Predicted configuration
Actual configuration

Cr
[Ar] 3d4
4s2
[Ar] 3d5
4s1

Cu
[Ar] 3d9 4s2
[Ar] 3d10 4s1

Ag
[Kr] 4d9 5s2
[Kr] 4d10 5s1

Au
[Xe] 4f14 5d9 6s2
[Xe] 4f14 5d10 6s1

In the case of chromium, an electron from the 4s orbital moves
into a 3d orbital, allowing each of the five 3d orbitals
to have one electron, making a half-filled set of orbitals.  In the
case of copper, silver and gold, an electron from the highest-occupied
s orbital moves into the d orbitals, thus filling the d
subshell.  Many anomalous electron configurations occur in the
heavier transition metals and inner transition metals, where the energy
differences between the s, d, and f subshells is
very small.