Atoms and the periodic table

An atom has a dense nucleus of protons and neutrons, surrounded by electrons. Protons carry a positive charge, electrons an equal negative charge, and neutrons none. Almost all the mass sits in the nucleus while almost all the volume is the electron region, which is why Rutherford's experiment showed most particles passing straight through gold foil.

The number of protons, the atomic number, defines the element. Change it and you have a different element. Change the neutron count and you have an isotope of the same element, with the same chemistry but a different mass. Change the electron count and you have an ion, with the same element but a net charge.

Electrons occupy regions called orbitals, grouped into shells and subshells. They fill from lowest energy upward, no orbital holds more than two electrons, and electrons spread across equal-energy orbitals before pairing up.

Chemistry is mostly about the outermost shell. Those valence electrons determine how an atom bonds, because filled outer shells are stable and atoms behave in ways that move them toward that arrangement, by losing, gaining or sharing electrons.

Periods, the horizontal rows, correspond to electron shells. Groups, the vertical columns, gather elements with the same number of valence electrons, which is exactly why they behave similarly. Group 1 metals all have one outer electron to lose, and all react vigorously with water. Group 18 already have full outer shells and barely react at all.

The blocks reflect which subshell is filling: the two left columns and helium fill s orbitals, the right-hand block fills p, the transition metals fill d, and the two detached rows fill f. The table's odd shape is a direct picture of electron structure.

• Atomic radius decreases across a period, because rising nuclear charge pulls the same shell in tighter, and increases down a group as shells are added.
• Ionisation energy, the energy to remove an electron, increases across a period and decreases down a group.
• Electronegativity, the pull an atom exerts on shared electrons, increases toward fluorine at the top right.
• Metallic character increases down and to the left, non-metallic character up and to the right.

Almost every periodic trend reduces to two competing factors: how strongly the nucleus pulls, and how far away and how shielded the outer electrons are. Across a period the pull increases while the distance barely changes. Down a group the distance and shielding increase and win.

If you can state those two factors you can reconstruct any trend under exam pressure, including the exceptions, which usually occur where a subshell has just started filling or has just become half full.