• EXAM #3 HAS BEEN MOVED TO MONDAY, NOVEMBER 9TH

•Bring a Periodic Table to class this week

November 2, 2009

Four Quantum Numbers

n (1,2, …) size/energy of the orbitall (0,1,2,…) shape of the orbital- s,p,d,f…ml (-l to l) orientation of the orbital

ms (- ½, ½) spin up/down (magnetic moment)

How do we know these things? Absorption and emission spectra- electron energiesZeeman 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 ParamagneticFerromagnetic (“real magnets”)

Pauli Exclusion Principle

No two electrons in an atom can have the same 4 quantum numbers

n, ℓ, mℓ define an orbital Therefore: an orbital can hold two electrons,

with opposite spins because ms 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 ___11

2p ___ ___ ___2s ___22

3d ___ ___ ___ ___ ___

3p ___ ___ ___3s ___

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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 ___

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Rules for filling orbitals

1. Pauli Exclusion PrincipleNo two electrons can have the same 4 quantum numbersAn orbital has a maximum of 2 electrons of opposite spin

2. Aufbau/Build-up PrincipleLower energy levels fill before higher energy levels

3. Hund’s RuleElectrons only pair after all orbitals at an energy level have

1 electron

4. Madelung’s RuleOrbitals 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.