Periodic Properties & Electron Configuration

n (principal quantum number)

energy level (shell)

l (angular momentum)

subshell (s, p, d, f)

mₗ (magnetic)

orbital orientation

mₛ (spin)

+½ or −½

Pauli Exclusion Principle

No two electrons have the same 4 quantum numbers; max 2 electrons per orbital (opposite spins)

Aufbau Principle

Fill lowest energy orbitals first

Hund’s Rule

Fill orbitals singularly before pairing

s =, p=, d=, f=

2, 6, 10, 14

Filling Order

1s 2s 2p 3s 3p 4s 3d 4p 5s …

Noble Gas Notation

Replace core electrons with nearest noble gas (e.g. Silicon — [Ne] 3s² 3p²)

Cations…

lose electrons from highest n level first (e.g. Fe²⁺ loses 4s before 3d)

Paramagnetic

unpaired electrons (attracted to magnet)

Diamagnetic

all paired (not attracted)

Effective Nuclear Charge (Zₑff)

Zeff = Z − S (Z = protons, S = shielding electrons)

Atomic Radius Trend

➡ Decreases across a period, ⬇ Increases down a group, because more protons pull electrons closer (higher Zeff)

Ion Size

Cations: smaller (lost electrons), Anions: larger (gained electrons). More protons = smaller size. Isoelectronic: Sr²⁺ < Br⁻ < Se²⁻

Ionization Energy (IE)

Energy required to move an electron. Smaller atoms: higher IE; removing inner electrons: very high IE. Same trend as Atomic Radius. Sb < As < Br

Electron Affinity (EA)

Energy change when gaining an electron. ➡ Increases across a period, ⬇ Slight change down group. Exceptions: Group 2A → full s (stable), Group 5A → half-filled p, Group 8A → full (very stable)

Metals

Conduct electricity, malleable & shiny, form cations, low ionization energy

Nonmetals

Poor conductors, brittle (if solid), form anions, high electronegativity

Metalloids

Mixed properties, semiconductors

Akali Metals (Group 1A)

Very reactive, low ionization energy, react with water

Alkaline Earth Metals (2A)

Less reactive than Group 1A

Halogens (7A)

Very reactive nonmetals, form negative ions

Noble Gases (8A)

Very stable, do not react easily,

Atomic Size

Br < Se < Te