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Electron Configuration

Electron Configuration. An orbital is a region within an atom where there is a probability of finding an electron. This is a probability diagram for the s orbital in the first energy level…. Orbital shapes are defined as the surface that contains 90% of the total electron probability.

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Electron Configuration

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  1. Electron Configuration

  2. An orbital is a region within an atom where thereis a probability of finding an electron. This is a probability diagram for the s orbital in the first energy level… Orbital shapes are defined as the surface that contains 90% of the total electron probability.

  3. Sizes of s orbitals Orbitals of the same shape (s, for instance) grow larger as n increases… Nodes are regions of low probability within an orbital.

  4. Electron Probability and Shape of Orbitals: s orbitals

  5. Electron Probability and Shape of Orbitals: p orbitals

  6. d orbital shapes Things get a bit more complicated with the five d orbitals that are found in the d sublevels beginning with n = 3. To remember the shapes, think of “double dumbells” …and a “dumbell with a donut”!

  7. Shape of the f orbital

  8. Orbital Diagrams • diagram used to show how the electrons are distributed among the orbitals of a subshell • orbital is represented by a circle or square • electrons represented by an arrow • Example: Hydrogen

  9. Pauli Exclusion Principle No two electrons in an atom can have the same four quantum numbers. And, no two fans in the Superdome should have the same 4 seat numbers! Wolfgang Pauli

  10. Pauli Exclusion Principle • An orbital can hold at most two electrons, and then only if the electrons have opposite spin • each electron is an atom has a unique set of quantum numbers • Example: Helium

  11. Sample Problem • Based on the Pauli exclusion principle, which of the following orbital diagrams are possible?

  12. Hund’s Rule The most stable arrangement of electrons is that with the most unpaired electrons all with the same spin

  13. Hund’s Rule • the lowest energy arrangement is obtained by putting e- in separate orbital of a subshell with parallel spin before pairing e- • Example: Carbon • Z = 6: 1s2 2s2 2p2 • You try Oxygen • Z = 8

  14. A. General Rules Pauli Exclusion Principle Each orbital can hold TWO electrons with opposite spins.

  15. A. General Rules Hund’s Rule Within a sublevel, place one e- per orbital before pairing them. “Empty Bus Seat Rule” RIGHT WRONG

  16. Aufbau Principle Electrons in an atom will occupy the lowest-energy orbitals available (aufbau = “building up”)

  17. Diagonal Rule (AufbauPrinciple) Sometimes the arrows are hard to follow. Here is a list of the sublevels in order. All you need to do is count up the number of electrons in order until you get to the number that matches or is slightly above the number you are looking for. 1s2 - 2s2 - 2p6 - 3s2 - 3p6 - 4s2 - 3d10 - 4p6 - 5s2 - 4d10 - 5p6 - 6s2 - 4f14 - 5d10 - 6p6 - 7s2 - 5f14 - 6d10 - 7p6

  18. Orbital filling table

  19. Orbital Filling Order

  20. A. General Rules Aufbau Principle Electrons fill the lowest energy orbitals first. “Lazy Tenant Rule”

  21. Orbital Diagrams Orbital Diagrams are models of electron arrangements showing configuration, subshell, aufbau, hunds, and pauli H:[ ] 1S

  22. 2s 2p 1s B. Notation • Orbital Diagram O 8e- • Electron Configuration 1s2 2s22p4

  23. 1st column of s-block 1st Period s-block C. Periodic Patterns Example - Hydrogen 1s1

  24. Core Electrons Valence Electrons B. Notation • Longhand Configuration S 16e- 2p6 2s2 1s2 3s2 3p4 • Shorthand Configuration S 16e- [Ne]3s2 3p4

  25. C. Periodic Patterns Example - Germanium [Ar] 4s2 3d10 4p2

  26. Orbital occupancy for the first 10 elements, H through Ne He and Ne have filled outer shells: confers chemical inertness

  27. Sample Problem • Electron Configuration of Vanadium • Quantum Numbers • n = • l = • ml = • ms = 1s 2s 2p 3s 3p 4s 3d

  28. Exceptions • Some transition metals do not follow the diagonal rule! • Example: chromium • Z = 24 • diagonal rule • 1s2 2s2 2p6 3s2 3p6 4s2 3d4 • -filled and half-filled orbitals are more stable than unevenly filled orbitals. • true electron configuration • 1s2 2s2 2p6 3s2 3p6 4s1 3d5

  29. D. Stability - Exceptions • Copper • EXPECT: [Ar] 4s2 3d9 • ACTUALLY: [Ar] 4s1 3d10 • Copper gains stability with a full d-sublevel. • Chromium • EXPECT: [Ar] 4s2 3d4 • ACTUALLY: [Ar] 4s13d5 • Chromium gains stability with a half-full d-sublevel.

  30. Electronic Structure of Atoms Shells and Orbitals • Shells of an atom contain a number of stacked orbitals

  31. Electronic Structure of Atoms p s f d Relative Energies for Shells and Orbitals • Some orbital subshells overlap others in different energy levels. Relative Energies of the orbitals

  32. IONS • Fe: [Ar] 4s2 3d6 Fe2+: [Ar] 3d6 • Notice for iron the 4s electrons are lost - typical • for the transition metals; the 4s (or 5s) electrons • are lost first

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