Electrons in an atom aren't scattered randomly around the nucleus — they occupy specific energy levels, often called electron shells. Understanding how electrons are arranged in shells helps explain why elements behave the way they do, and why the periodic table has the shape it does.
What is an electron shell?
An electron shell (main energy level) is a region around the nucleus where electrons of similar energy are found. Shells are numbered n = 1, 2, 3, 4... starting closest to the nucleus. Electrons fill shells according to two simple rules:
- Lower shells fill first — electrons occupy the lowest available energy level before moving to higher ones.
- Each shell has a maximum capacity, given by the formula:
maximum electrons in shell n = 2n²
So shell 1 holds up to 2 electrons, shell 2 holds up to 8, shell 3 holds up to 18 (though in practice, for the elements you'll meet early on, shell 3 often "settles" for 8 before the next shell starts filling — you'll refine this once you learn about subshells).
Writing shell (electron) arrangements
A quick way to describe an atom's electrons is to write the number in each shell, separated by commas — e.g. sodium (11 electrons) is written 2,8,1.
Worked Example 1
Deduce the electron shell arrangement of a magnesium atom (12 electrons).
- Shell 1: 2 electrons
- Shell 2: 8 electrons
- Shell 3: 2 electrons remaining
Arrangement: 2,8,2 (Talbot, p. 132)
Shells and the Periodic Table
The number of shells an atom uses tells you its period, and the number of electrons in the outermost (valence) shell tells you its group (for main-group elements). This is why elements in the same group behave similarly — they have the same number of valence electrons.
Worked Example 2
Element X has the shell arrangement 2,8,7. Identify its period and group, and name the element.
- Three shells are used → Period 3
- 7 electrons in the outer shell → Group 17
- This matches chlorine (Talbot, p. 132)
Shells and ion formation
Atoms often gain or lose electrons to achieve a full outer shell, matching the stable arrangement of the nearest noble gas.
- Metals (groups 1, 2, 13) tend to lose electrons, forming cations.
- Non-metals (groups 15, 16, 17) tend to gain electrons, forming anions.
Worked Example 3
Deduce the ion formed by a phosphorus atom (2,8,5) and its electron arrangement.
- Phosphorus needs 3 more electrons to reach a full outer shell of 8.
- It forms P³⁻, with arrangement 2,8,8 — the same as argon (Talbot, p. 132).
This pattern explains why:
- Group 1 metals form +1 ions (e.g. Na⁺, arrangement 2,8)
- Group 2 metals form +2 ions (e.g. Mg²⁺, arrangement 2,8)
- Group 16 non-metals form −2 ions (e.g. S²⁻, arrangement 2,8,8)
Carbon and silicon (group 14) are exceptions — with 4 valence electrons, it's energetically unfavourable to lose or gain 4 electrons, so they mostly form covalent bonds instead of simple ions (Talbot, p. 132).
Key takeaways
- Electrons occupy shells at set energy levels, filling from the lowest shell outward, with a maximum of 2n² electrons per shell.
- An atom's shell arrangement (e.g. 2,8,1) directly links to its position in the periodic table: number of shells = period, outer shell electrons = group (for main groups).
- Atoms gain or lose electrons to achieve a full outer shell like the nearest noble gas, forming ions.
- Metals lose electrons to form cations; non-metals gain electrons to form anions.
- Group 14 elements (like carbon and silicon) usually don't form simple ions — they favour covalent bonding instead.