Atomic Structure - Electron Arrangement (A-Level Chemistry)

Electron Arrangement

Electron Notations

Sub-Shell Notations

You need to understand sub-shell notation, as shown in the diagram below:

Electron-in-Boxes Notation

Within each sub-shell, we can draw out the electrons in orbitals. Each box represents one orbital. The electrons are represented by the arrows located inside the boxes. The arrows always point in opposite directions to indicate that the electrons are spinning in opposite directions.

Electron Configurations

Rules to Work Out Electron Configurations

In exams you’ll often have to write electron configurations for different elements. You can answer these questions by following the following rules:

1. Electrons start filling up with the lowest energy sub-shells

Electrons fill up the 1s sub shell first, then the 2s, then the 2p and so on. When an electron is at its lowest possible energy level we say it is in its ground state. If it were found at a higher energy level, we would say it is in its excited state.

Exception: The 4s sub shell fills up before the 3d sub shell and therefore in electron configuration 4s is written before the 3d. Despite the principal number quantum number of 4d being larger, this is overridden by the lower energy level of the 4d sub shell, meaning it is filled first.

2. Electrons fill each orbital in each shell individually before they start pairing and sharing an orbital with another electron.

When we fill orbitals in a particular sub-shell, we try to avoid having pairs of orbitals. This is because having two electrons in a single orbital requires more energy, as energy is needed to overcome the repulsion between the electrons.

Consider Nitrogen, which has 7 electrons. First, we will up 1s2 and 2s2 with 2 electrons each, leaving 3 electrons for the 2p sub-shell. p subshells contain 3 orbitals. We fill 1 electron in each orbital, rather than putting 2 electrons in a single orbital, as shown below.

Atoms, molecules or ions with unpaired electrons in outer shell configuration are known as free radicals.

Electron Configurations of Ions

To find the electron configuration of ions from either the s and p block of the periodic table, electrons are eliminated from or added to the subshell with the highest energy. For example if we look at the electron configuration of Calcium (forms a positive ion), we would remove 2 electrons from the 4s sub-shell:

Ca = 1s2 2s2 2p6 3s2 3p6 4s2.                         Ca2+ = 1s2 2s2 2p6 3s2 3p6

Another example is fluorine (forms a negative ion), for which we add an electron to the highest energy sub-shell.:

Fl = 1s2 2s2 2p5                                                  Fl- = 1s2 2s2 2p6

Transition Metals

Transition metals have two unusual rules:

1) A half-full or full 3d sub-shell is especially stable.

Chromium and Copper have unusual properties as their sub shells fill differently to usual. Chromium and Copper donate one of their 4s electrons to the 3d sub-shell, making a half-full (chromium) and full (copper) d sub-shell, which is particularly stable.

2) When forming ions, they lose 4s electrons before 3d electrons.

Usually when atoms form ions they lose electrons from the 3d sub-shell first and then the 4s sub-shell as the 3d sub shell is at a higher energy level. However, transition metals lose electrons from the 4s first.