Chap 3 Electron Configurations & Quantum Numbers

1. Chap 3 Electron Configurations & Quantum Numbers

2. Quantum Numbers Help us locate all the electrons.

3. The number and relative energies of all hydrogen electron orbitals through n=3 • At ordinary temperatures essentially all hydrogen atoms are in their ground states • The electron may be promoted to an excited state by the absorbtion of a photon with the appropriate quantum of energy

4. Emission lines As excited electrons relax to their ground state they give off light waves at very specific wavelengths called emission lines

5. Quantum Mechanical Model • Proposed by Schrodinger to account for matters’ wave-like behavior. • Estimates Probability of finding an electron in an area 90% of the time. • Replaces Bohr’s planetary orbits with orbitals, shown as fuzzy clouds

6. Quantum Numbers: n, l, m, s • n: the primary energy level (quanta) • average distance from nucleus • l: the subLevel • s, p, d, f • m: the number of orbitals within a sublevel • 1,3,5,7 • s: the electron Spin • up  & down 

7. Sublevels • Number of sublevels increase as radius increases (as n increases) energy # sublevels name of level n = n sublevels n=1 1 sublevel s n=2 2 sublevel s, p n=3 3 sublevel s, p, d n=4 4 sublevel s, p, d, f

8. Orbitals The different sublevels can hold different # of Orbitals Sublevel # of Orbitals s 1 p 3 d 5 f 7

9. OrbitalsHave specific shapes and quantities

10. Orbital Shapes

11. The Electron Configuration Notation

12. Electron Filling Rules • The Aufbau Principle • Electrons are added one at a time to the lowest orbital available until all of the electrons are used. • The Pauli Exclusion Principle • An orbital can have a maximum of two electrons. • To occupy the same orbital, two electrons must spin in opposite directions • Hund’s Rule • Electrons occupy equal energy orbitals so that the maximum number of unpairedelectrons result.

13. Element # of Electrons in Element Electron Configuration He 2 1s2 Li 3 1s22s1 Be 4 1s22s2 O 8 1s22s22p4 Cl 17 1s22s22p63s23p5 K 19 1s22s22p63s23p64s1 Electron Configurations

14. Chlorine Electron Configuration • The electron configuration for chlorine is • 1s2 2s2 2p6 3s2 3p5 • The large numbers represent the energy level. • The letters represent the sublevel. • The superscripts indicate the number of electrons in the sublevel.

15. Filling using an Aufbau Diagram

16. H 1 Li 3 Na 11 K 19 Rb 37 Cs 55 Fr 87 1s1 1s22s1 1s22s22p63s1 1s22s22p63s23p64s1 1s22s22p63s23p64s23d104p65s1 1s22s22p63s23p64s23d104p65s24d10 5p66s1 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s1

17. He 2 1s2 1s22s22p6 1s22s22p63s23p6 1s22s22p63s23p64s23d104p6 1s22s22p63s23p64s23d104p65s24d105p6 1s22s22p63s23p64s23d104p65s24d10 5p66s24f145d106p6 Ne 10 Ar 18 Kr 36 Xe 54 Rn 86

18. Electron Configurations repeat • The shape of the periodic table is a representation of this repetition. • When we get to the end of the column the outermost energy level is full. • This is the basis for our shorthand.

19. Yes there is a shorthand!

20. The Shorthand • Write the symbol of the noble gas before the element. • Then the rest of the electrons. • Aluminum - full configuration. • 1s22s22p63s23p1 • Ne is 1s22s22p6 • so Al is [Ne] 3s23p1

21. More examples • Ge = 1s22s22p63s23p64s23d104p2 • Ge = [Ar] 4s23d104p2 • Hf=1s22s22p63s23p64s23d104p65s2 4d105p66s24f145d2 • Hf=[Xe]6s24f145d2

22. Writing Electron configurations Using the periodic table

23. S- block s1 • Alkali metals all end in s1 • Alkaline earth metals all end in s2 • really have to include He but it fits better later. • He has the properties of the noble gases. s2

24. The P-block p1 p2 p6 p3 p4 p5

25. Transition Metals -d block d4 d9 d1 d2 d3 d5 d6 d7 d8 d10

26. f6 f13 f1 f2 f3 f4 f5 f7 f8 f10 f12 f14 f11 f9 F - block • inner transition elements

27. 1 2 3 4 5 6 7 • Each row (or period) is the energy level for s and p orbitals.

28. D orbitals fill up after previous energy level so first d is 3d even though it’s in row 4. 1 2 3 4 5 6 7 3d

29. 1 2 3 4 5 6 7 • f orbitals start filling at 4f 4f 5f

30. The Shorthand Again Sn- 50 electrons The noble gas before it is Kr Takes care of 36 Next 5s2 Then 4d10 Finally 5p2 [ Kr ] 5s2 4d10 5p2

31. Configuration of Ions • Ions always have noble gas configuration. • Na is 1s12s22p63s1 • Forms a +1 ion - 1s12s22p6 • Same configuration as neon. • Metals form ions with the configuration of the noble gas before them - they lose electrons.

32. Configuration of Ions • Non-metals form ions by gaining electrons to achieve noble gas configuration. • They end up with the configuration of the noble gas after them.

33. IA Li 1s22s1 Li+ 1s2 Na 1s22s22p63s1 Na+ 1s22s22p6 K 1s22s22p63s23p64s1 K+ 1s22s22p63s23p6 The electrons in the outermost shell (the ones with the highest value of n) are the most energetic, and are the ones which are exposed to other atoms. This shell is known as the valence shell. The inner, core electrons (inner shell) do not usually play a role in chemical bonding. Elements with similar properties generally have similar outer shell configurations. For instance, we already know that the alkali metals (Group I) always form ions with a +1 charge; the "extra" s1 electron is the one that's lost:

34. IVA IIA C Be 1s22s2 1s22s22p2 Be2+ C4- 1s22s22p6 1s2 VA Mg N 1s22s22p63s2 1s22s22p3 N3- Mg2+ 1s22s22p6 1s22s22p6 VIA IIIA Al O 1s22s22p4 1s22s22p63s23p1 O2- Al3+ 1s22s22p6 1s22s22p6 VIIA F 1s22s22p5 F- 1s22s22p6 • The Group IIA and IIIA metals also tend to lose all of their valence electrons to form cations. • The Group IV - VII non-metals gain electrons until their valence shells are full (8 electrons).

35. VIIIA Ne 1s22s22p6 Ar 1s22s22p63s23p6 • The Group VIII noble gases already possess a full outer shell, so they have no tendency to form ions.

36. n l ml Number oforbitals OrbitalName Number ofelectrons 1 0 0 1 1s 2 2 0 0 1 2s 2 1 -1, 0, +1 3 2p 6 3 0 0 1 3s 2 1 -1, 0, +1 3 3p 6 2 -2, -1, 0, +1, +2 5 3d 10 4 0 0 1 4s 2 1 -1, 0, +1 3 4p 6 2 -2, -1, 0, +1, +2 5 4d 10 3 -3, -2, -1, 0, +1, +2, +3 7 4f 14 Table of Allowed Quantum Numbers