Electron Configuration
Now that you know how the periodic table is organised into blocks, groups, and periods, it's time to connect that structure to something more precise: electron configuration — the exact arrangement of electrons in an atom's orbitals. This is the key that links an element's position on the periodic table to its chemical behaviour.
Why electron configuration matters
The chemical properties of an element depend on its electron configuration, and elements in the same group have similar configurations of their valence electrons (the outermost electrons). This is exactly why elements in the same group behave similarly — for example, all Group 1 elements have the configuration ending in ns¹ (Talbot, p.240).
Reading configurations from the periodic table
You can predict an atom's electron configuration just from where it sits on the table:
- s-block: number of valence electrons = group number (e.g. Na, group 1, is [Ne] 3s¹)
- p-block: number of valence electrons = group number − 10 (e.g. Se, group 16, is [Ar] 3d¹⁰ 4s² 4p⁴)
- The nearest noble gas before the element gives the "core" configuration, written in square brackets, and you build the rest onto it (Talbot, p.240–241)
Electrons fill orbitals in order of increasing energy, not strictly by shell number — this is why the 4s sublevel fills before 3d.
Worked Examples
Example 1: Calcium (Ca) Calcium is in Period 4, Group 2 (s-block). Its configuration should end in 4s².
- Nearest noble gas before Ca: Argon (Ar)
- Condensed: [Ar] 4s²
- Full: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² (Talbot, p.241)
Example 2: Tin (Sn) Tin is in Period 5, Group 14 (p-block, second column of the p-block), so its configuration ends in 5p².
- Nearest noble gas before Sn: Krypton (Kr)
- Condensed: [Kr] 5s² 4d¹⁰ 5p²
- Full (in order of filling): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d¹⁰ 5p² (Talbot, p.241)
Example 3: From atom to ion Sodium's full configuration is 1s² 2s² 2p⁶ 3s¹. When sodium loses its one valence electron to form Na⁺, it becomes 1s² 2s² 2p⁶ — identical to neon. This is a general rule: all Group 1 cations (+1) have the same configuration as the previous noble gas (Talbot, p.132).
Similarly, phosphorus (1s² 2s² 2p⁶ 3s² 3p³) gains three electrons to form P³⁻, giving 1s² 2s² 2p⁶ 3s² 3p⁶ — the same as argon. Group 15 anions (−3) match the configuration of the next noble gas (Talbot, p.132).
A useful exception
Carbon and silicon (Group 14) have 4 valence electrons and don't easily lose or gain electrons to form simple ions — instead they form covalent bonds. Keep this in mind: not every element's story ends neatly with ion formation (Talbot, p.132).
Key Takeaways
- Electron configuration shows how electrons are distributed among orbitals, filling in order of increasing energy.
- You can predict configurations directly from an element's block and position: s-block → group number = valence electrons; p-block → group number − 10 = valence electrons.
- Configurations are often written in condensed form using the previous noble gas as a shorthand core.
- Group 1 cations and Group 15 (and other) anions often adopt noble gas configurations when they form ions.
- Carbon and silicon are exceptions — they prefer covalent bonding over simple ion formation.