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Periodic properties of the elements

Periodic properties of the elements. Chapter 8. Quantum numbers. Principal quantum number, n Determines size and overall energy of orbital Positive integer 1, 2, 3 . . . Corresponds to Bohr energy levels. Quantum numbers. Angular momentum quantum number, l Determines shape of orbital

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Periodic properties of the elements

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  1. Periodic properties of the elements Chapter 8

  2. Quantum numbers • Principal quantum number, n • Determines size and overall energy of orbital • Positive integer 1, 2, 3 . . . • Corresponds to Bohr energy levels

  3. Quantum numbers • Angular momentum quantum number, l • Determines shape of orbital • Positive integer 0, 1, 2 . . . (n–1) • Corresponds to sublevels

  4. Quantum numbers • Magnetic quantum number, ml • Determines number of orbitals in a sublevel and orientation of each orbital in xyz space • integers –l . . . 0 . . . +l

  5. What type of orbital is designated by each set of quantum numbers? • n = 5, l = 1, ml = 0 5p • n = 4, l = 2, ml = –2 4d • n = 2, l = 0, ml = 0 2s • Write a set of quantum numbers for each orbital • 4s n = 4, l = 0, ml = 0 • 3d n = 3, l = 2, ml = –2, –1, 0, +1, or +2 • 5p n = 5, l = 1, ml = –1, 0, or +1

  6. Electron configurations • Electrons exist within orbitals, given by three quantum numbers n, l, and ml

  7. Electron configurations Configuration shows which orbitals are occupied • Aufbau principle: e– takes lowest available energy • Hund’s rule: if there are 2 or more orbitals of equal energy (degenerate orbitals), e– will occupy all orbitals singly before pairing

  8. Electron configurations • Electron has spin, either “up” or “down” • Electron spin given by 4th quantum number, ms • Pauli exclusion principle: no two e– in an atom can have the same set of 4 quantum numbers ⇒ 2 e– per orbital, one up ↑ and one down ↓

  9. Magnetic properties • Atom or ion with unpaired e– is attracted to a magnetic field = paramagnetic • Atom or ion with all e– paired is slightly repelled by a magnetic field = diamagnetic

  10. Effective nuclear charge: Zeff • Electron experiences attraction of nucleus and repulsion of other e– in the atom • Outer e– is partially shielded from full charge of nucleus • Zeff = actual nuclear charge – charge shielded by other e–

  11. Effective nuclear charge: Zeff • Core e– effectively shield outer e– from nuclear charge • Outer e– do not shield other outer e– very efficiently • Thus, • Li outer e– experiences Zeff ≈ 3–2 = +1 • Be outer e– experience Zeff ≈ 4–2 = +2

  12. Trends in atomic radius • Atomic radius increases down a group • Same Zeff • Outer e– in higher principal energy level = larger orbital • Atomic radius decreases across a period • Same principal energy level • Increasing Zeff pulls in outer e–

  13. Trends in atomic radius • Transition metal radii stay roughly constant across a period • Outer e– stay same • Adding protons to nucleus and electrons to n–1 (core) orbital, so Zeff stays about constant

  14. Ions and ionic radii • Elements may lose or gain outer e– to form ions • Metals lose e–→ cations • Nonmetals gain e– → anions • The ions in these examples are isoelectronic (same e– configuration)

  15. Transition metal ions • Transition metals lose their ns e– before losing their (n–1)d e–

  16. Ions and ionic radii • A cation is much smaller than its parent atom

  17. Ions and ionic radii • An anion is much larger than its parent atom

  18. Ions and ionic radii • For isoelectronic species, the one with the highest nuclear charge will have the smallest radius

  19. Ionization energy • Ionization energy (IE) = energy needed to remove e– from atom/ion in gaseous state • IE always positive (endothermic) • Successive IE values always increase • IE increases dramatically when begin to remove core e–

  20. Trends in 1st ionization energy • IE1 decreases down a group • Same Zeff • Outer e– in higher principal energy level = farther from nucleus, easier to remove • IE1 generally increases across a period • Increasing Zeff pulls outer e– closer to nucleus, harder to remove

  21. Two important exceptions • 1st e– in p sublevel • 3s orbital penetrates closer to nucleus • 3p e– somewhat more shielded from nuclear charge, easier to remove • e– begin pairing in p sublevel P S

  22. Electron affinity (EA) • EA = energy change when gaseous atom/ion gains an e– • Trends less regular but generally becomes more exothermic across a period

  23. Summary of periodic trends • Across a period • atomic radius decreases • IE1 (always endothermic) increases (with two important exceptions) • EA generally becomes more exothermic • Down a group • atomic radius increases • IE1 decreases

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