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Sections 7.1 – 7.3 Electron Spin, Orbital Energies and Electron Configurations

Sections 7.1 – 7.3 Electron Spin, Orbital Energies and Electron Configurations. In these sections… Electron Spin and Magnetism Energies of Orbitals Electron Configurations of Atoms. Atomic Electronic Structure.

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Sections 7.1 – 7.3 Electron Spin, Orbital Energies and Electron Configurations

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  1. Sections 7.1 – 7.3 Electron Spin, Orbital Energies and Electron Configurations

  2. In these sections… • Electron Spin and Magnetism • Energies of Orbitals • Electron Configurations of Atoms Atomic Electronic Structure

  3. Electron Spin: Electrons exhibit a magnetic fieldWe think of them as spinning.They can spin only two ways: think of it as left or rightSpin quantum number: ms can be +1/2 or -1/2

  4. Magnetic Properties come from additive effects ofelectron spins. Diamagnetic: all electrons are pairedParamagnetic: 1 or more unpaired electronsFerromagnetic (real magnets): unpaired electrons all lined up in the same direction

  5. 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 only two electrons, with opposite spins because ms can only be +1/2 or -1/2

  6. Pauli Exclusion Principle What’s allowed?

  7. Orbital Energies Single Electron Atoms Multi-electron Atoms Why? With a single electron, energy depends only on how far from the nucleus. With multiple electrons, e-e- repulsions also play a role and differ depending on orbital shape.

  8. Single Electron Atoms Multi-electron Atoms For most atoms: Energy increases as n increases: 1 < 2 < 3 < 4 … Energy increases as subshells progress: s < p < d < f

  9. Atomic Electron Configurations An atom has lots of electrons and lots of orbitals. Which orbitals do the electrons occupy?

  10. Atomic Electron Configurations An atom has lots of electrons and lots of orbitals. Which orbitals do the electrons occupy? Electrons fill the lowest energy orbitals first. Electron Configuration: a listing of how many electrons occupy each orbital.

  11. 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. An arrow represents an electron.

  12. Electron Configurations:Three Notation Types 2. spdf (or spectroscopic) notation: List subshells and how many electrons they contain: 1s22s22p63s1 3. Noble gas notation: short [Ne]3s1 Where [Ne] = 1s22s22p6 1.

  13. Electron Configurations and thePeriodic Table Examples using Electron Configuration Simulation • Periodic Blocks • Hund’s Rule (using the p block) • n value increases as you move down table • Anomalies: Cr and Cu

  14. Electron Configurations and thePeriodic Table: Periodic Blocks

  15. Electron Configurations and the Periodic Table II

  16. Periodic Table and the Order of Filling In what order are subshells filled?

  17. Hund’s Rule: Subshells are filled to give the maximum number of unpaired electrons

  18. Using Periodic Blocks: C

  19. Using Periodic Blocks: Cl

  20. Noble Gas Notation: Mg

  21. Noble Gas Notation: Mg

  22. Diamagnetic vs. Paramagnetic Elements

  23. d-Block Elements: Fe

  24. Two Anomalies: Cr and Cu

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