EXAM #3 HAS BEEN MOVED TO MONDAY, NOVEMBER 9 TH Bring a Periodic Table to class this week November 2, 2009

Four Quantum Numbers n (1,2, …)  size/energy of the orbital l (0,1,2,…)  shape of the orbital- s,p,d,f… m l (-l to l)  orientation of the orbital m s (- ½, ½)  spin up/down (magnetic moment) How do we know these things? Absorption and emission spectra- electron energies Zeeman effect- spectrum splits when magnetic field applied; separates orbitals at the same energy level and led to discovery of electron spin

Electron Spin is the Source of Magnetism in Materials Diamagnetic Paramagnetic Ferromagnetic (“real magnets”)

Pauli Exclusion Principle No two electrons in an atom can have the same 4 quantum numbers n, ℓ, m ℓ define an orbital  Therefore: an orbital can hold two electrons, with opposite spins because m s can only be +1/2 or -1/2

Orbital Energies Only depends on distance from the nucleus Electron-electron repulsion affects energy Different for different orbital shapes

1s ___ 1 2p ___ ___ ___ 2s ___ 2 3d ___ ___ ___ ___ ___ 3p ___ ___ ___ 3s ___ 3 ENERGY For most atoms: Energy increases as n increases: 1 < 2 < 3 < 4 … Energy increases as subshells go from s < p < d < f At the same main shell level, a p orbital will be at a higher energy than an s orbital 4f ___ ___ ___ ___ ___ ___ ___ 4d ___ ___ ___ ___ ___ 4p ___ ___ ___ 4s ___ 4

Rules for filling orbitals 1. Pauli Exclusion Principle No two electrons can have the same 4 quantum numbers An orbital has a maximum of 2 electrons of opposite spin 2. Aufbau/Build-up Principle Lower energy levels fill before higher energy levels 3. Hund’s Rule Electrons only pair after all orbitals at an energy level have 1 electron 4. Madelung’s Rule Orbitals fill in the order of the value of n + l

Orbital Filling Order

Electron Configurations General Rule: electrons fill lowest energy orbitals first Sodium, Na as an example Na has 11 electrons. Fill 2 electrons per orbital till you run out A box represents an orbital. A arrow represents an electron.