The AUFBAU principle (from the German for building up) is a set of basic rules for working out the electron arrangments in atoms. The rules are easy to apply:
↑ | ↑ | ↑ |
The arrangement below is not allowed, the p-orbitals are all occupied singularly but the electrons do not have parallel spins.
↓ | ↑ | ↑ |
Also the arrangement below is not allowed since this time the electrons are paired up in one of the p-orbitals when an empty degenerate orbital is available.
↑ ↓ | ↑ |
The electron energy levels or shells will fill up according to the rules set out above in the aufbau principle.
The diagram opposite shows the energy levels for each of the
sub-shells and orbitals in a multi-electron atom.
If we start at the bottom, that is the sub-level which is lowest in energy
we can clearly
work out the order in which to place the electrons in the sub-levels and orbitals,
they will fill in the
following order:
1s → 2s → 3s → 3p → 4s → 3d → → 4p → 5s → 4d → 5p → 6s → 4f
→ 5d → 6p → 7s → 5f → 6d
You may also have seen a diagram similar to the one opposite left, it simply shows an easier way to remember the order in which the sub-shells or sub-levels fill. Simply start at the top and follow the arrows downwards to get the same order as shown above.
The easiest way to get the hang of working out electron arrangements is simply practice writing them out. If you do this you will quickly notice some rather obvious patterns across the periodic table which should make sure you get the electron arrangements correct every time. The table below gives the electron arrangements for the first 10 elements. Why not work them out yourself first and then check your answers with the ones below?
element | atomic number | 1s orbital | 2s orbital | 2p orbital | electron arrangement | ||
---|---|---|---|---|---|---|---|
H | 1 | ↑ | 1s^{1} | ||||
He | 2 | ↑ ↓ | 1s^{2} | ||||
Li | 3 | ↑ ↓ | ↑ | 1s^{2}2s^{1} | |||
Be | 4 | ↑ ↓ | ↑ ↓ | 1s^{2}2s^{2} | |||
B | 5 | ↑ ↓ | ↑ ↓ | ↑ | 1s^{2}2s^{2}2p^{1} | ||
C | 6 | ↑ ↓ | ↑ ↓ | ↑ | ↑ | 1s^{2}2s^{2}2p^{2} | |
N | 7 | ↑ ↓ | ↑ ↓ | ↑ | ↑ | ↑ | 1s^{2}2s^{2}2p^{3} |
O | 8 | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ | ↑ | 1s^{2}2s^{2}2p^{4} |
F | 9 | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ | 1s^{2}2s^{2}2p^{5} |
Ne | 10 | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑↓ | 1s^{2}2s^{2}2p^{6} |
A good way to tell if you are getting the electron arrangements correct is that the Noble gases always have filled p-orbitals, that is they are always np^{6} (except helium). In fact the periodic table can help a lot in checking you have the correct electron arrangements. The periodic table can be divided up into blocks. These blocks are called the s-block, the d-block and the p-block simply because the outer valence electrons in each block are in s, p or d sub-shells. These blocks are shown in the image of the periodic table below and again at the foot of the page:
Element number 11 is sodium, it is in
period 3 in the periodic table, this means its outer valence electrons
are in the third electron shell and since it is in the s-block these
electrons will be in the 3s sub-shell. Writing out
electron arrangements can become a bit tedious after a while, so we can
use a shortened version of the electron arrangement. We can shorten
the electronic configuration by simply writing out the inner
electron arrangement from the preceding noble
gas e.g. The noble gas before sodium
is neon, its electronic
configuration is 1s^{2}2s^{2}2p^{6}, sodium the
next element will be
1s^{2}2s^{2}2p^{6}3s^{1} or [Ne]3s^{1}. The noble gas Argon has
an atomic number of 18,
its electronic configuration will be 1s^{2}2s^{2}2^{6}3s^{2}3p^{6},
so the next element potassium will be 1s^{2}2s^{2}2^{6}3s^{2}3p^{6}4s^{1}
or [Ar]4s^{1}. Remember the 4s sub-shell is lower in energy than
the 3d sub-shell so it fills first.
The table below gives the electronic configuration of the elements sodium to calcium.
element | atomic number | 1s orbital | 2s orbital | 2p orbital | 3s orbital | 3p orbital | 4s orbital | electron arrangement | ||||
---|---|---|---|---|---|---|---|---|---|---|---|---|
Na | 11 | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ | [Ne]3s^{1} | ||||
Mg | 12 | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | [Ne]3s^{2} | ||||
Al | 13 | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ | [Ne]3s^{2}3p^{1} | |||
Si | 14 | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ | ↑ | [Ne]3s^{2}3p^{2} | ||
P | 15 | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ | ↑ | ↑ | [Ne]3s^{2}3p^{3} | |
S | 16 | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ | ↑ | [Ne]3s^{2}3p^{4} | |
Cl | 17 | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ | [Ne]3s^{2}3p^{5} | |
Ar | 18 | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | [Ne]3s^{2}3p^{6} | |
K | 19 | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ | [Ar]4s^{1} |
Ca | 20 | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | [Ar]4s^{2} |
After element 20, calcium the 4s sub-shell is full and we enter the d-block of the periodic table. This block contains the transition metals. Many of the characteristic properties of the transition metals are due to the d-electrons. There are 5 d-orbitals, each holding 2 electrons, so there are 10 transition metals. The electronic configuration of the first row of the transition metals is shown below:
element | atomic number | 1s orbital | 2s orbital | 2p orbital | 3s orbital | 3p orbital | 4s orbital | 3d | electron arrangement | ||||||||
---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|
Sc | 21 | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ | [Ar]4s^{2}3d^{1} | ||||
Ti | 22 | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ | ↑ | [Ar]4s^{2}3d^{2} | |||
V | 23 | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ | ↑ | ↑ | [Ar]4s^{2}3d^{3} | ||
Cr | 24 | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ | ↑ | ↑ | ↑ | ↑ | ↑ | [Ar]4s^{1}3d^{5} |
Mn | 25 | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ | ↑ | ↑ | ↑ | ↑ | [Ar]4s^{2}3d^{5} |
Fe | 26 | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ | ↑ | ↑ | ↑ | [Ar]4s^{2}3d^{6} |
Co | 27 | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑↓ | ↑ ↓ | ↑ | ↑ | ↑ | [Ar]4s^{2}3d^{7} |
Ni | 28 | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ | ↑ | [Ar]4s^{2}3d^{8} |
Cu | 29 | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | [Ar]4s^{1}3d^{10} |
Zn | 30 | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | ↑ ↓ | [Ar]4s^{2}3d^{10} |
You may have noticed something odd with the electron arrangements of two of the d-block metals. Chromium for example has an electron configuration of [Ar]4s^{1}3d^{5} whereas you might have expected it to be [Ar]4s^{2}3d^{4}, similarly copper has an electron configuration of [Ar] 4s^{1}3d^{10} whereas you might have expected it to be [Ar] 4s^{2}3d^{9}. In both cases an electron from the 4s sub-shell has been promoted into the 3d sub-shell, the reason for this is due to the unusual stability of half-filled and full d-sub-shells. By promoting an electron from the 4s sub-shell in each case we end up with either half-filled 3d sub-shell in the case of chromium and a full 3d sub-shell in the case of copper. This transfer of an electron from the 4s sub-shell lowers the overall energy of the atom.
While working out the electron configurations for the elements you will have no doubt noticed that the periodic table is divided into blocks based on the outer shell electrons. An outline of this is shown below: