Electrons exhibit a magnetic field We think of them as spinning They can spin only two ways: think of it as left or ri

1. Electrons exhibit a magnetic fieldWe think of them as spinningThey can spin only two ways: think of it as left or rightSpin quantum number: ms can be +1/2 or -1/2

2. Magnetic Properties come from additive effects ofelectron spins. Diamagnetic: all electrons are pairedParamagnetic: 1 or more unpaired electronsFerromagnetic (real magnets): unpaired electrons all lined up in the same direction

3. Pauli Exclusion Principle No two electrons in an atom can have the same 4 quantum numbers n, l, ml define an orbital Therefore: an orbital can hold two electrons, with opposite spins because ms can only be +1/2 or -1/2

4. Orbital Energies

5. Orbital Energies

6. Electron Configurations General Rule: electrons fill lowest energy orbitals first Sodium, Na as an example

7. Electron Configurations:Three Notation Types

8. Electron Configurations and thePeriodic Table Examples using Electron Configuration Simulation Periodic Blocks Hund�s Rule (using the p block) n value increases as you move down table Anomalies: Cr and Cu

9. Electron Configurations and thePeriodic Table I

10. Electron Configurations and the Periodic Table II

11. Electron Configurations and the Periodic Table III

12. Notes: There�s no known reason electrons have spin, or have only two of them. The other stuff about orbitals is theoretically derived from Schrod. Equn., but the whole spin thing is just something we see. Can�t explain it- just know it�s true. Like gravity or Coulombic attractions. The reason why different subshells have different energies: for example: The energy of the 2s subshell has to do with how well the 2s electrons are attacted to the nucleus minus how much they are repelled by the 1s electrons. Same thing for the 2p electrons. Difference is, the 1s electrons repel the 2p electrons more than the 2s electrons, so the 2p electrons are less stable, and higher energy. Same reasoning happens when you go to higher subshells (e.g. d > p) Why does the 4s subshell come before the 3d subshell? The reason above about d being higher in energy plus the fact that as you go up in n value, the orbital energies all get closer together. So, 2 is much higher than 1; 3 is less higher than 2; 4 is not much higher than 3, etc. This comes from the En = -constant/n2 The reason for Hund�s Rule: there is less e-e repulsion if electrons are in different orbitals because they are in different Places. That�s why they go to different orbitals in a subshell first. I don�t know why they go with the same spin. Have them fill in the blanks for a set of elements as you use the simulation. Be sure to include a Hund�s rule one: B, C, N, or something like it. Answer to hard question: the pt looks the same, but is half as wide for each block because each orbital can only hold a single electron.

13. Predicting Electron Configurations

14. Predicting Electron Configurations

15. Predicting Electron Configurations

16. Predicting Electron Configurations

21. Electron Configurations of Cations

22. Electron Configurations of Anions

23. Transition Metal Cations: Lose s electrons first

24. Diamagnetic vs. Paramagnetic Elements

25. Periodic Properties of the Elements All Depend on energies of outermost orbitals Atomic Size Ionization Energy Electron Affinity Ion Size

31. Which atom is the smallest of all? H He Cs Rn

32. Which of these atoms is largest? K Ca Rb Sr

34. Which atom has the largest ionization energy? K Ca Rb Sr

35. Which ionization energy for Mg will see the largest jump? 1st 2nd 3rd 4th

36. Why the breaks in the line?

40. Which atom has the smallest common ion? H Na F Cl

42. Which atom has the largest common ion? Na K F Cl

44. Which of the following isoelectronic species is smallest? Mg2+ Na+ Ne F- O2-

46. Which of the following isoelectronic species has the lowest ionization energy? Mg2+ Na+ Ne F- O2-