1 / 9

Understanding Quantum Numbers in Chemistry

Learn about the significance of quantum numbers in chemistry, electron spin, magnetic properties, orbital energies, and filling rules for orbitals. Understand how to apply these concepts in electron configurations and orbital filling order.

Télécharger la présentation

Understanding Quantum Numbers in Chemistry

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript


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

  2. Four Quantum Numbers • n (1,2, …) size/energy of the orbital • l (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 energies Zeeman effect- spectrum splits when magnetic field applied; separates orbitals at the same energy level and led to discovery of electron spin

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

  4. 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

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

  6. 3d ___ ___ ___ ___ ___ 3p ___ ___ ___ 3s ___ 2p ___ ___ ___ 2s ___ 3 2 1s ___ 1 ENERGY 4f ___ ___ ___ ___ ___ ___ ___ 4d ___ ___ ___ ___ ___ 4p ___ ___ ___ 4s ___ 4 • 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

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

  8. Orbital Filling Order

  9. 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.

More Related