2. Electron Configurations Niels Bohr “Any one who is not shocked by Quantum theory does not understand it.”

3. Electron Configurations The quantum mechanical model of the atom predicts energy levels for electrons; it is concerned with probability, or likelihood, of finding electrons in a certain position.

4. Electron Configurations Regions where electrons are likely to be found are called orbitals. EACH ORBITAL CAN HOLD UP TO 2 ELECTRONS!

5. Electron Configurations In quantum theory, each electron is assigned a set of quantum numbers analogy: like the mailing address of an electron

6. Describes the energy level that the electron occupies n=1, 2, 3, 4 The larger the value of n, the farther away from the nucleus and the higher the energy of the electron. **Electrons always start filling in the lowest possible energy level available (AUFBAU PRINCIPLE) Principal Quantum Number (n) n = 1 n = 2 n = 3 n = 4

7. Sublevels ( l ) The number of sublevels in each energy level is EQUAL to the quantum number, n, for that energy level. Sublevels are labeled with a number that is the principal quantum #, and a letter: s, p, d, f (ex: 2 p is the p sublevel in the 2 nd energy level) ***REMEMBER THIS… let’s start filling in the chart

8. Sublevels ( l ) Principal Energy LevelSublevelsOrbitals n = 1 n = 2 n = 3 n = 4 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f one (1s) one (2s) three (2p) one (3s) three (3p) five (3d) one (4s) three (4p) five (4d) seven (4f)

9. Sublevels ( l ) Sublevel# of orbitalsMax # of electrons s p d f 1 3 5 7 2 6 10 14 ONLY 2 electrons per 1 orbital!!!!!

11. Electron Configurations NOTICE!!!!!! 3d is on the 4 th energy level (d block is n-1)

13. Complete electron configurations helium boron neon aluminum Uranium (follow your periodic table) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 6 7s 2 5f 4 1s 2 2s 2 2p 6 3s 2 3p 1 1s 2 2s 2 2p 6 1s 2 2s 2 2p 1 1s 2

14. Abbreviated electron diagrams (AKA: noble gas config.) helium boron aluminum cobalt uranium N 3- Se 2- Mg 2+ 1s 2 [He]2s 2 2p 1 [Ne]3s 2 3p 1 [Ar]4s 2 3d 7 [Rn]7s 2 5f 4 [He]2s 2 2p 6 = [Ne] = same # of electrons as a Kr atom: [Ar] 4s 2 3d 10 4p 6 = same # of electrons as a Ne atom: [He] 2s 2 2p 6

15. Spin quantum number (m s ) Labels the orientation of the electron Electrons in an orbital spin in opposite directions; these directions are designated as +½ and -½

16. Pauli Exclusion Principle States that no 2 electrons have an identical set of four quantum #’s to ensure that no more than 2 electrons can be found within a particular orbital.

17. Hund’s Rule Orbitals of equal energy are each occupied by one electron before any pairing occurs. Repulsion between electrons in a single orbital is minimized (ex: you want your own bedroom before you would have to share a bedroom with a sibling) All electrons in singly occupied orbitals must have the same spin (such as in the p sublevel). When 2 electrons occupy the same orbital they must have opposite spins (want to be a little different. You wouldn’t want to wear the exact same clothes as your brother or sister, right?!)

18. Orbital Diagrams Each orbital is represented by a box. Each electron is represented by an arrow.

19. Orbital Diagrams hydrogen helium carbon 1s 2s2p