Electron Structure

Atoms consist of electrons circling the nucleus of the atom in paths called shells or orbitals.

Electron shells increase in energy, as they get further away from the nucleus.
There are four types of sub shells; s, p, d and f. Each sub-shell contains orbitals:

• The s sub-shell contains 1 x s orbital
• The p sub-shell contains 3 x p orbitals
•  The d sub-shell contains 5 x d orbitals
•  The f sub-shell contains 7 x f orbitals

An orbital is a region of space where electrons with “opposite spin” are found.

Each orbital holds a maximum of 2 electrons so:

• The s sub-shell can hold 2 electrons
• The p sub-shell can hold 6 electrons
• The d sub-shell can hold 10 electrons
• The f sub-shell can hold 14 electrons

The table below gives a summary of the information above but also relates it to how the periodic table is built up. For example, electron shell 1 is relating to the elements along period 1 of the periodic table, electron shell 2 is relating to the elements along period 2, and so on. Please note that it only goes up to the element Krypton, atomic number = 36.

Electron shell	Maximum number of electrons	Sub-shells
		1 x s	3 x p	5 x d	7 x f
1	2	2	-	-	-
2	8	2	6	-	-
3	18	2	6	10	-
4	32	2	6	10	14

Electrons are fed into the shells of an atom with the lowest energy sub-levels filled up first. The superscripts show the maximum number of electrons that sub-level can hold. Note that the 4s sub-level is filled before the 3d:

            1s²          2s²          2p⁶          3s²          3p⁶          4s²          3d¹⁰..           

            Low energy                                                           High energy

Lets take the magnesium atom, which has 12 electrons and an electron configuration of 2.8.2. The electrons are then fed into the sub-levels, giving an electronic configuration of:

            1s²          2s²             2p⁶            3s²

Another example would be krypton, which has 36 electrons and its electron configuration is 2.8.18.8. The electronic configuration is:

            1s²             2s²          2p⁶         3s²          3p⁶         4s²         3d¹⁰         4p⁶

Electrons fill sub-shells singly before pairing up due to lower repulsion when unpaired (more stable). This can be shown by using the ‘electrons in boxes’ notation to help make this clearer:

            e.g. The oxygen atom has an electronic configuration of 1s² 2s² 2p⁴, the 8 electrons occupy the sub-shells in the following way:

                                        1s           2s                    2p

Note! The electrons fill up the 2p sub-shells singly before pairing up.

Below shows the ‘electrons in boxes’ notation from the lowest energy sub-level, 1s, up to higher energy sub-level, 4d: