Atomic Structure

Protons define what element you have. Six protons is carbon, and nothing else can be carbon. Neutrons change the mass but not the chemistry — that is what isotopes are. Electrons do all the chemistry.

Every reaction you will ever study is electrons rearranging. The nucleus sits there unchanged unless you are doing nuclear physics.

Electrons cannot sit anywhere they like. They occupy discrete energy levels — shells — filling from the inside out. Shell capacities follow 2n²: 2, 8, 18, 32.

Only the outermost shell — the valence shell — takes part in bonding. This single fact explains the entire shape of the periodic table: elements in the same group have the same number of valence electrons, so they behave the same way.

Slide from Li (3) to Na (11) to K (19). All three have exactly one valence electron — that is why they are all violently reactive alkali metals, and why they sit in the same column.

Atomic radius shrinks across a period. Counter-intuitive, but correct: you are adding protons and electrons to the same shell, so the growing nuclear charge pulls that shell in tighter. Going down a group you add a whole new shell, so atoms get bigger.

Ionisation energy — the energy to remove the outermost electron — rises across a period (held more tightly) and falls down a group (further from the nucleus and shielded by inner shells).

Electronegativity — the pull on shared electrons — follows the same pattern, peaking at fluorine.

Switch between the three properties and watch the heat map flip direction. Radius brightens toward the bottom-left; electronegativity and ionisation energy brighten toward the top-right.

Not every atom of an element is identical. Isotopes are atoms of the same element — the same number of protons — with different numbers of neutrons, and so different mass numbers. Chlorine, for example, exists as chlorine-35 and chlorine-37. Because they have the same number of electrons arranged in the same way, isotopes are chemically identical; only their masses (and physical properties that depend on mass, like density) differ.

Since a real sample is a mixture of isotopes, the mass we quote for an element is a weighted average called the relative atomic mass (A_r) — the average mass of its atoms compared with 1/12 of the mass of a carbon-12 atom. The weighting is by how abundant each isotope is.

The simple shells (2, 8, 8 …) are really made of sub-shells labelled s, p, d and f, which hold up to 2, 6, 10 and 14 electrons respectively. Each sub-shell is divided into orbitals, regions that each hold at most two electrons spinning in opposite directions.

Sub-shells fill in order of increasing energy, which is why 4s fills before 3d: 1s 2s 2p 3s 3p 4s 3d 4p … We write the arrangement as an electron configuration, for example sodium is 1s² 2s² 2p⁶ 3s¹. The outer electrons this notation reveals are exactly the ones that decide an element's chemistry.

The first ionisation energy is the energy needed to remove one electron from each atom in a mole of gaseous atoms. Removing successive electrons needs steadily more energy, because each electron is pulled from an increasingly positive ion. The big jumps in successive ionisation energies are strong evidence for shells: a sudden leap means you have started removing electrons from a shell closer to the nucleus.

Across a period, first ionisation energy generally rises (more nuclear charge, similar shielding, so electrons are held tighter); down a group it falls (the outer electron is further out and more shielded, so easier to remove). These are the same forces that drive reactivity.