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Chapter 8 Periodic Properties of the Elements

Chapter 8 Periodic Properties of the Elements. For an atom, electrons are in atomic orbitals. Energy of atomic orbitals. Orbital Energy Levels for the Hydrogen Atom. H atom: E only depends on n. degenerate. E depends on n and l same n, l ↑ ↔ E↑. A Picture of the Spinning Electron.

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Chapter 8 Periodic Properties of the Elements

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  1. Chapter 8 Periodic Properties of the Elements

  2. For an atom, electrons are in atomic orbitals. Energy of atomic orbitals

  3. Orbital Energy Levels for the Hydrogen Atom H atom: E only depends on n degenerate

  4. E depends on n andl same n, l↑ ↔ E↑

  5. A Picture of the Spinning Electron

  6. Spin quantum number ms ms = +1/2 or −1/2 4 quantum numbers are used to specify an electron. How do electrons fill up atomic orbitals?

  7. Pauli Exclusion Principle In a given atom, no two electrons can have the same set of four quantum numbers. An orbital can hold only two electrons, they must have opposite spins.

  8. H atom electron configuration 1s1 Lowest energy: ground state ↑ 1s 2s1 ↑ 2s Excited states 2p1 ↑ 2p orbital diagram

  9. Now we can write the ground state electron configurations and draw orbital diagrams according to Pauli principle.

  10. Hund’s rule For degenerate orbitals, the lowest energy is attained when the number of electrons with the same spin is maximized. Valence electrons: electrons in the outermost shell. involved in bonding Core electrons: inner electrons

  11. Elements in the same group have similar valence electron configuration — similar chemical properties. Number of valence electrons = main group number Number of filled shells = period number Noble gases have 8 (He 2) valence electrons. Stable structure. Metals: tend to lose valence electrons to reach 8(2) valence electron. Nonmetals: tend to gain electrons to reach 8(2) valence electrons.

  12. Periodic trends in atomic properties • Atomic radius

  13. Atomic Radii (in Picometers) for Selected Atoms

  14. Atomic radius In a period: decreases from left to right In a group: increases from top to bottom

  15. EXAMPLE 8.5 Atomic Size On the basis of periodic trends, choose the larger atom in each pair (if possible). Explain your choices. (a) N or F (b) C or Ge (c) N or Al (d) Al or Ge

  16. EXAMPLE 8.7 Ion Size Choose the larger atom or ion from each pair. (a) S or S2–(b) Ca or Ca2+(c) Br– or Kr

  17. Periodic trends in atomic properties • Atomic radius • Ionization energy

  18. Ionization energy Energy required to remove an electron from a gaseous atom or ion. X(g)  X+(g) + e− first ionization energy X+(g)  X2+(g) + e− second ionization energy

  19. Ionization energy In a period: increases from left to right In a group: decreases from top to bottom (general trend)

  20. Periodic trends in atomic properties • Atomic radius • Ionization energy • Electron affinity

  21. Electron affinity Energy change associated with the addition of an electron to a gaseous atom. X(g) + e−  X−(g) X(g) + e− E Ei X−(g) Ef ∆E = Ef − Ei = EA < 0

  22. Electron affinity In a period: increases from left to right In a group: no clear trend (very rough trend)

  23. Periodic trends in atomic properties • Atomic radius • Ionization energy • Electron affinity Remember the trends

  24. 1. Arrange the following groups of atoms in order • of increasing size. • Te, S, Se; b) K, Br, Ni; c) Ba, Si, F 2. Arrange the atoms in Ex. 1 in order of increasing first ionization energy.

  25. Chapter 8 Problems 4, 7, 43, 45, 51, 53, 55, 61, 63, 65, 67, 69, 73, 75

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