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EXAM #3 HAS BEEN MOVED TO MONDAY, NOVEMBER 9 TH Bring a Periodic Table to class this week November 2, 2009
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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
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Electron Spin is the Source of Magnetism in Materials Diamagnetic Paramagnetic Ferromagnetic (“real magnets”)
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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
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Orbital Energies Only depends on distance from the nucleus Electron-electron repulsion affects energy Different for different orbital shapes
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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
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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
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Orbital Filling Order
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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.
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