The atom is not the solid sphere we sometimes picture. It has internal structure, and understanding it is the key to explaining all chemical properties.
The Bohr model (starting point)
In 1913 Niels Bohr proposed a planetary model: the nucleus sits at the center, electrons orbit on discrete shells. Each shell has a maximum capacity:
- K shell (n = 1): 2 electrons
- L shell (n = 2): 8 electrons
- M shell (n = 3): 18 electrons
- N shell (n = 4): 32 electrons
Electrons fill inner shells first. For sodium (Na, Z = 11), with 11 electrons: (K)² (L)⁸ (M)¹.
This model is enough to understand why elements in the same group share properties: they have the same number of electrons on their outermost shell (valence shell).
From planetary to quantum
The Bohr model works for hydrogen but breaks down for heavier atoms. Experimental data (emission spectra, ionization energies) demands a finer model: quantum mechanics.
In the quantum model, the electron has no definite path. We use the notion of atomic orbital: a region of space where the probability of finding the electron is high. Orbitals are classified by:
- principal quantum number n (size, energy): 1, 2, 3...
- secondary quantum number ℓ (shape): s, p, d, f
- magnetic quantum number mₗ (orientation)
- spin mₛ: +1/2 or −1/2
Electron configuration
The Pauli principle forbids two electrons from sharing all four quantum numbers. The Aufbau rule gives the filling order:
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p
For carbon (Z = 6): 1s² 2s² 2p². For iron (Z = 26): [Ar] 4s² 3d⁶.
One atom, many faces
Depending on the precision you need:
- The planetary model is enough to discuss shells and valence.
- The quantum model is needed to explain covalent bonding, molecular geometry, the color of compounds, magnetism.
In high school, you mostly use the simplified quantum model (s, p, d). At the undergraduate level, you go deep into orbitals and their geometry. Up to you.