Electronic structure of the first 20 elements
How an atom reacts is decided almost entirely by its electron arrangement — especially the electrons in the outer shell. Knowing how to write electron arrangements is essential for the rest of the spec.
Shells (energy levels)
Electrons orbit the nucleus in shells (energy levels). Each shell can hold a fixed maximum number of electrons:
- Shell 1 (closest to nucleus): maximum 2 electrons.
- Shell 2: maximum 8 electrons.
- Shell 3: maximum 8 electrons (for the first 20 elements; transition metals later complicate this).
Electrons fill from the innermost shell outwards. Once shell 1 is full, electrons go in shell 2; once shell 2 is full, they go in shell 3, and so on.
Writing electronic structures
The convention is to write the number of electrons in each shell, separated by commas:
- Hydrogen (1 electron): 1
- Helium (2): 2
- Carbon (6): 2,4
- Sodium (11): 2,8,1
- Chlorine (17): 2,8,7
- Argon (18): 2,8,8
- Calcium (20): 2,8,8,2
Tip: the number of shells = the period; the outer shell electrons (for groups 1–7) = the group number.
Why electronic structure matters
The number of outer-shell electrons controls:
- Group placement — Group 1 = 1 outer electron, Group 7 = 7 outer electrons, Group 0 = 8 outer (or 2 for He).
- Reactivity patterns within a group.
- Bonding tendencies (lose, gain, or share electrons to reach a full outer shell).
A full outer shell is the most stable arrangement — the same as a noble gas. Atoms react to achieve a full outer shell:
- Group 1 atoms (1 outer electron) lose 1 electron to form +1 ions.
- Group 7 atoms (7 outer electrons) gain 1 electron to form −1 ions.
- Group 0 atoms (full outer shell) are unreactive.
Predicting reactivity from electronic structure
Group 1 (alkali metals)
Each has 1 outer electron. Loses it easily to form a +1 ion. Reactivity increases down the group because:
- The outer electron is further from the nucleus.
- More inner shells shield the outer electron.
- Less attraction → easier to lose → more reactive.
Group 7 (halogens)
Each has 7 outer electrons. Gains 1 to form a −1 ion. Reactivity decreases down the group because:
- The shell receiving the electron is further from the nucleus.
- More shielding by inner electrons.
- Weaker pull on incoming electrons → less reactive.
Group 0 (noble gases)
Full outer shell → no need to gain, lose or share electrons → unreactive.
✦Worked example— Worked example — predict bonding
Sodium (2,8,1) loses 1 electron → Na⁺ (2,8) — same as neon. Chlorine (2,8,7) gains 1 electron → Cl⁻ (2,8,8) — same as argon.
Result: Na⁺ and Cl⁻ attract each other → NaCl ionic compound. (Detail in C2.)
Drawing diagrams
A typical electron-shell diagram has concentric circles for shells with crosses or dots for electrons. Pair electrons up in each shell as you draw — e.g. 8 electrons in shell 2 are usually drawn as four pairs.
⚠Common mistakes
- Putting more than 2 in shell 1, or more than 8 in shell 2. Capacities are fixed.
- Not closing brackets / using full stops instead of commas. Use commas: 2,8,1.
- Saying "carbon has 4 electrons". Carbon has 6 electrons total; 4 are in the outer shell.
- Forgetting that an ion has a different electron count from the neutral atom.
Links
Foundation for C1.5–1.8 (groups), C2 (bonding) and C4 (chemical changes / reactivity series).
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