Plan for Wed, 5 Nov 08

1. Plan for Wed, 5 Nov 08 • Lab Stuff • Volumetric Analysis reports will be returned by Monday • Today’s lab (Calorimetry) will be written up as a formal report… • Your Calorimetry formal report will be due in class on Friday, Nov 14th • No lab next week! • Lecture: • Electron spin and the Pauli principle (7.8) • Orbital energies for polyelectronic atoms (7.9) • Aufbau (“filling-up”) principle (7.11) • Periodic trends (7.12) • Exam 2, Quiz 4 regrades, and Quiz 5 returned

2. Which orbital is expected to have the greatest energy? 1s 2s 2p 3p

3. Electron Spin • Experiments demonstrated the need for one more quantum number. • Specifically, some particles (electrons in particular) demonstrated inherent angular momentum… • Basically, this means that electrons have two ways of interacting with an applied magnetic field. Interpretation: the electron is a bundle of “spinning” charge “spin up” “spin down”

4. Electron Spin (cont.) • The new quantum number is ms (analagous to ml). • For the electron, ms has two values: +1/2 and -1/2 ms = 1/2 ms = -1/2

5. Pauli Exclusion Principle Defn: No two electrons may occupy the same quantum state simultaneously. In other words: electrons are very territorial. They don’t like other electrons horning in. In practice, this means that only two electrons may occupy a given orbital, and they must have opposite spin.

6. Quantum Number Summary • n: principal quantum number • index of size and energy of electron orbital • can have any integral value: 1, 2, 3, 4, … • l: angular momentum quantum number • related to the shape of the orbitals • can have integral values 0 … n - 1 • ml: magnetic quantum number • related to orbital orientation (relative to the other l-level orbitals) • can have integral values –l … 0 … +l • ms: electron spin quantum number • related to the “magnetic moment” of the electron • can have half-integral values –1/2 or +1/2

7. Polyelectronic Atoms • For polyelectronic atoms, a direct solution of the Schrodinger Eq. is not possible. • When we construct polyelectronic atoms, we use the hydrogen-atom orbital nomenclature to discuss in which orbitals the electrons reside. • This is an approximation (and it is surprising how well it actually works). No solution for polyelectronic atoms!!

8. + Polyelectronic Atoms e- “Screening”: The presence of other electrons means a given electron does not feel the attraction of the nucleus as strongly as it would in hydrogen. “Penetration”: Orbitals that have some probability density close to the nucleus will be energetically favored over orbitals that do not.

9. The Aufbau Principle • When placing electrons into orbitals in the construction of polyelectronic atoms, we use the Aufbau Principle. • This principle states that in addition to adding protons and neutrons to the nucleus, one simply adds electrons to the hydrogen-like atomic orbitals • Pauli exclusion principle: No two electrons may have the same quantum numbers. Therefore, only two electrons can reside in an orbital (differentiated by ms).

10. 4s 3s 3p 3d 2s 2p 1s Orbital Energies H has only one electron, so all of the sublevels in a given principal level have the same energy...they are degenerate. Energy In many-electron atoms, a given electron is simultaneously attracted to the nucleus and repelled by other electrons, causing the energies of the sublevels to change relative to H. When we put electrons in orbitals, we fill them in order of increasing energy, not n.

11. Let’s fill some orbitals RULES • Orbitals are filled starting from the lowest energy. • The two electrons in an orbital must have opposite spin. • Example: Hydrogen (Z = 1) 1s1 1s 2s 2p • Example: Helium (Z = 2) 1s2 1s 2s 2p

12. Let’s fill some more orbitals • Lithium (Z = 3) 1s22s1 1s 2s 2p • Berillium (Z = 4) 1s22s2 1s 2s 2p • Boron (Z = 5) 1s22s22p1 1s 2s 2p

13. Filling Orbitals (cont.) • Carbon (Z = 6) 1s22s22p2 REVISED RULES • Orbitals are filled starting from the lowest energy. • The two electrons in an orbital must have opposite spin. • Hund’s Rule: the orbitals in degenerate series (such as 2p in the example above) must each have an electron before any of them can have two. 1s 2s 2p Hund’s Rule: Lowest energy configuration is the one in which the maximum number of unpaired electrons are distributed amongst a set of degenerate orbitals.

14. Filling Orbitals (cont.) • Carbon (Z = 6) 1s22s22p2 1s 2s 2p • Nitrogen (Z = 7) 1s22s22p3 1s 2s 2p

15. Filling Orbitals (cont.) • Oxygen (Z = 8) 1s22s22p4 1s 2s 2p • Fluorine (Z = 9) 1s22s22p5 1s 2s 2p • Neon (Z = 10) 1s22s22p6 full 1s 2s 2p

16. Ne 3s • Compare to Neon (Ne) (Z = 10) 1s22s22p6 full 1s 2s 2p Filling Orbitals (cont.) • Sodium (Z = 11) 3p 1s 2s 2p 3s 1s22s22p63s1 [Ne]3s1

17. Filling Orbitals (cont.) • Sodium (Z = 11) [Ne]3s1 Ne 1s22s22p63s1 3s • Phosphorus (P) (Z = 17) [Ne]3s23p3 Ne 3s 3p • Argon (Z = 18) [Ne]3s23p6 Ne 3s 3p

18. Filling Orbitals (cont.) • We now have the orbital configurations for the first 18 elements. • Elements in same column have the same # of valence electrons! Valence Electrons: The total number of s and p electrons in the highest occupied energy level.

19. Suppose there are four possible spin values instead of two. In this case, what would be the electron configuration of Cl? 1s42s42p9 B. 1s22s22p63s23p5 C. 1s32s32p93s2 D. 1s62s62p5

20. The Aufbau Principal (cont.) • Similar to Sodium, we begin the next row of the periodic table by adding electrons to the 4s orbital. • Why not 3d before 4s? • 3d is closer to the nucleus • 4s allows for closer approach; therefore, is energetically preferred.

21. Back to Filling Orbitals • Elements Z=19 and Z= 20: Z= 19, Potassium: 1s22s22p63s23p64s1 = [Ar]4s1 Ar 4p 4s 1s22s22p63s23p64s2 = [Ar]4s2 Z= 20, Calcium: Ar 4p 4s

22. 3d 3d 4p 4p Filling Orbitals (cont.) • Elements Z=21to Z=30 have occupied d orbitals: Z= 21, Scandium: 1s22s22p63s23p64s23d1 = [Ar] 4s23d1 Ar 4s 1s22s22p63s23p64s23d10 = [Ar] 4s23d10 Z= 30, Zinc: Ar 4s

23. The Aufbau Principal (cont.) • Elements Z=19 and Z= 20: Z= 19, Potassium: 1s22s22p63s23p64s1 = [Ar]4s1 Z= 20, Calcium: 1s22s22p63s23p64s2 = [Ar]4s2 • Elements Z = 21 to Z = 30 have occupied d orbitals: Z= 21, Scandium: 1s22s22p63s23p64s23d1 = [Ar] 4s23d1 Z = 24, Chromium: [Ar] 4s13d5 exception Z= 30, Zinc: 1s22s22p63s23p64s23d10 = [Ar] 4s23d10

24. Write down the orbitals for each n on separate lines. Arrows drawn as shown will give you the order in which the orbitals should be filled. Note that this scheme fills 4s before 3d, as expected. What if you forget the orbital-filling order?

25. This orbital filling scheme gives rise to the modern periodic table. Periodic Table

26. After Lanthanum ([Xe]6s25d1), we start filling 4f. Periodic Table

27. After Actinium ([Rn]7s26d1), we start filling 5f. Periodic Table

28. Row headings correspond to the highest occupied energy level for any element in that period. Periodic Table

29. Column headings give total number of valence electrons for any element in that group. Periodic Table “Valence” only refers to s and p electrons in the highest occupied energy level.

30. What is the electron configuration for the indicated element? A. 1s22s22p63s23p64s23d3 B. 1s22s22p63s23p64s24d3 C. 1s22s22p63s23p64s23d2 D. 1s22s22p73s23p64s23d2

31. Valence Electrons • The total number of s and p electrons in the highest occupied energy level. • As we’ll see, all the “action” happens at the valence electrons. • Elements in the same group (column) in the periodic table have the same number of valence electrons. • This means elements in the same group tend to have similar chemical properties.

32. C Valence Electrons (cont.) Chemists use Lewis dot symbols to indicate the number of valence electrons in an atom. The valence electrons are drawn as dots around the atomic symbol, with orbital occupancy indicated...that is, electrons that occupy the same orbital appear as paired dots. HOWEVER, we will encounter situations where it is more convenient to spread the dots out around the element symbol. C