Simple Electronic Structure

Electrons are arranged in energy levels or quantum shells, numbered by principal quantum number $n$. The maximum electron capacity for each shell is:

• Shell 1: up to 2 electrons
• Shell 2: up to 8 electrons
• Shell 3: up to 18 electrons
• Shell 4: up to 32 electrons
  The electronic configuration represents the distribution of electrons in an atom. For example, lithium has 3 electrons (2 in n=1, 1 in n=2), neon has 10 (2 in n=1, 8 in n=2), and chlorine has 17 (2 in n=1, 8 in n=2, 7 in n=3).

Evidence for Electronic Structure

Ionisation energy (IE) is the energy required to remove an electron from an atom. The first ionisation energy (IE1) is defined for one mole of atoms in gas phase forming one mole of gaseous 1+ ions. Measured in kJ mol−1, for example, $IE1 = 590$ for calcium. Successive ionisation energies (IE2, IE3, etc.) increase due to increased nuclear charge acting on remaining electrons.

Factors Influencing Ionisation Energy

1. Nuclear Charge: Higher proton number increases positive charge, enhancing attraction to electrons.
2. Distance from Nucleus: Electrons in shells farther from nucleus experience weaker attraction.
3. Shielding Effect: Inner shell electrons repel outer shell electrons, reducing effective nuclear charge felt by outer electrons.
4. Spin-Pair Repulsion: Electrons paired in the same orbital repel each other, making removal easier for the paired electron.

Successive Ionisation Energies

A graph of ionisation energies can reveal the number of outer shell electrons and confirm the group of the element in the Periodic Table.

Quantum Sub-shells and Atomic Orbitals

Principal quantum shells are divided into sub-shells (s, p, d, f), with energy levels increasing as follows: s < p < d. The number of orbitals is specific to each sub-shell:

• s: 1 orbital (max 2 electrons)
• p: 3 orbitals (max 6 electrons)
• d: 5 orbitals (max 10 electrons)
  Shapes of orbitals affect electron configuration; s orbitals are spherical, while p orbitals are hourglass-shaped.

Electronic Configuration

Elements follow a consistent order in electronic configuration, reflecting sub-shell filling: $1s^2$, $2s^2$, $2p^6$, etc. Some exceptions exist like chromium and copper due to stability considerations in their d sub-shells.

Periodic Patterns in Radii and Ionisation Energies

Atomic radius increases down a group and decreases across a period, influenced by nuclear charge and electron shielding. Ionisation energy generally increases across a period and decreases down a group due to similar factors. Significant jumps in ionisation energy indicate removal from a fuller, closer shell.

Summary

Electron shells can be categorized into sub-shells (s, p, d) with particular electron capacities. The factors influencing ionisation energy include distance from the nucleus, nuclear charge, shielding effects, and spin-pair repulsion. Successive ionisation energies allow reasoning of electronic configurations and element grouping in the Periodic Table.