Electrons in atoms
Electrons occupy energy levels and orbitals following quantum rules. Their arrangement explains the shape of the periodic table and the colours of light atoms emit.
Light and energy levels
When atoms are heated, they emit light only at specific wavelengths, unique to each element, which is why fireworks and neon signs have characteristic colours. Bohr explained this by proposing that electrons occupy fixed energy levels; light is emitted when an electron drops to a lower level, with photon energy \(E = hf\).
E = hfQuantum numbers and orbitals
In the modern quantum model, electrons occupy orbitals, regions where an electron is likely to be found. Orbitals are described by quantum numbers: the principal level \(n\) (size and energy), the shape (s, p, d, f), the orientation, and the electron's spin. Each orbital holds at most two electrons with opposite spins (the Pauli exclusion principle).
Subshells
Level \(n\) contains \(n\) subshells: \(n = 1\) has 1s; \(n = 2\) has 2s and 2p; \(n = 3\) has 3s, 3p, 3d. An s subshell has 1 orbital (2 electrons), p has 3 (6 electrons), d has 5 (10 electrons), and f has 7 (14 electrons). These capacities explain the widths of the periodic table's blocks.
Electron configurations
Electrons fill orbitals from lowest energy upward (the aufbau principle): 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, and so on. Oxygen (8 electrons) is \(1s^2 2s^2 2p^4\). Within a subshell, electrons occupy separate orbitals before pairing up (Hund's rule).
Valence electrons
Valence electrons are those in the outermost shell, and they govern chemical behaviour. Main-group elements in the same group have the same number of valence electrons: all halogens have 7, which is why they each gain one electron to form \(-1\) ions.
The noble gas core shorthand
Configurations can be shortened using the preceding noble gas: sodium is \([\text{Ne}]3s^1\), and calcium is \([\text{Ar}]4s^2\). This highlights the valence electrons that matter for reactions. Some transition metals, like chromium \([\text{Ar}]3d^54s^1\), are exceptions because half-filled subshells are especially stable.
Worked example
Write the configuration of phosphorus (15 electrons): \(1s^2 2s^2 2p^6 3s^2 3p^3\), or \([\text{Ne}]3s^23p^3\). It has 5 valence electrons, matching its group 15 position, and the three 3p electrons occupy separate orbitals by Hund's rule.
Why the table looks the way it does
The s block (groups 1 and 2) is two columns wide, the p block six, the d block ten, and the f block fourteen, exactly matching subshell capacities. The periodic table is a map of electron configurations.
Key takeaways
Electrons occupy quantised energy levels; transitions emit or absorb photons with \(E = hf\). Orbitals hold two electrons each; s, p, d, f subshells hold 2, 6, 10, 14. Configurations follow the aufbau principle, Pauli exclusion, and Hund's rule. Valence electrons determine chemistry and explain the table's layout.
Flashcard
Photon energy
Check your understanding
How many electrons can a p subshell hold?