1. Electron Configurations 4.3

2. Pauli Exclusion Principle  No more than two e- can occupy a single orbital at a time  e- spin in opposite directions  Spin quantum number can be +½ or -½ 2

3. Electron Configuration  A description of the electron orbitals in an atom 3

4. e- occupy lower energy levels first  Aufbau principle – electrons in an atom will occupy the lowest energy orbitals available. 4

5. Electron Configurations  Remember, the smaller the principal quantum number, the lower the energy, and the smaller the l quantum number, the lower the energy 5

6. Electron Configurations  The order in which energy levels fill is … 1s<2s<2p<3s<3p 6

7. Electron Configurations  After this, the energy levels are less straightforward  The E levels of the 3d orbitals are slightly higher than those of the 4s orbitals. 1s<2s<2p<3s<3p<4s≈3d 7

8. Electron Configurations  The next irregularity is 5s and 4d are close in energy 1s<2s<2p<3s<3p<4s≈3d<4p<5s≈4d  Still more irregularities exist with higher energy orbitals. 8

9. Rules for Writing e- Configs 1. Determine the # of e- the atom has (atomic #) Fluorine as an example has 9 electrons 9

10. Rules for Writing e- Configs 2. Fill orbitals in order of increasing energy (follow the diagram on the next slide) S orbital – 2 e- P orbitals – 6 e- D orbitals – 10 e- F orbitals – 14 e- 10

11. Rules for Writing e- Configs 11

12. Rules for Writing e- Configs  The configuration for F is 1s 2 2s 2 2p 5 3. Make sure the total number of e- in the config match the atomic number 12

13. Electron Configurations  Tells us how the 16 e- of S are configured  The electron configuration for S is  S=1s 2 2s 2 2p 6 3s 2 3p 4  Each s orbital has 2 e-, each p orbital can have 6 e- (2 per orbital) 13

14. Electron Configurations  Do the electron configurations for the following elements…  O, Ar, Ca, V, Sr, and Sn  Remember the d orbital is off by 1 E level 14

15. The f orbital  If you end in the f orbital – make sure you indicate d 1 first.  If you go through the f orbital, count the s orbital, then all of the f then go to d and beyond. 15

16. Electron Configurations  There are still some irregularities with higher energy orbitals.  Chromium is an example of this  Cr =1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5  There is one unfilled d orbital and a filled s orbital 16

17. Electron Configurations  There are e- configs listed in the p.t. in the back of your book.  These are the ground state configs of the isolated atoms in the gas phase.  Under other conditions, the configs could be different 17

18. Orbital Diagrams  Hund’s Rule says the maximum stability for e- is when you have the maximum number of unpaired electrons when they have the same quantum number. 18

19. Orbital Diagrams  Use boxes to show e- location  Boron = 1s 2 2s 2 2p 1  The orbital diagram looks like… 19 1s 2s { 2p   

20. Orbital Diagrams  What is the Carbon orbital diagram.  The config is 1s 2 2s 2 2p 6 3s 2 20

21. Electron Configurations  To save space, some configurations are written like  S = [Ne]3s 2 3p 4  This means take neon’s configuration and add 3s 2 3p 4 to the end of it. 21