1. Electron Configuration Mapping the electrons

2. Electron Configuration The way electrons are arranged around the nucleus.

3. Quantum Mechanical Model 1920’s Werner Heisenberg (Uncertainty Principle) Louis de Broglie (electron has wave properties) Erwin Schrodinger (mathematical equations using probability, quantum numbers)

4. Heisenberg uncertainty principle it is impossible to determine simultaneously both the position and velocity of an electron or any other particle with any great degree of accuracy or certainty.

5. Erwin Schrodinger  Formulated equation that describes behavior and energies of subatomic particles.  Leads to Quantum Mechanics: we cannot pinpoint an electron in an atom but we can define the region where electrons can be in a particular time……… called a Probability map….a 3-dimensional area in space called an ORBITAL

6. Orbital The space where there is a high probability that it is occupied by a pair of electrons. Orbitals are solutions of Schrodinger’s equations.

7. Principal Quantum Number, n Each orbital can hold 2 electrons! Principle Quantum Number Number of sublevels Type of sublevel n=111s (1 orbital) n=222s (1 orbital), 2p (3 orbitals) n=333s(1 orbital), 3p (3 orbitals), 3d (5 orbitals) n=444s (1 orbital), 4p (3 orbitals), 4d (5 orbitals), 4f (7orbitals)

8. Orbitals in Sublevels Sublevel # Orbitals # electrons s12 p36 d510 f714 g918

9. Orbitals-different shapes

10. Three rules are used to build the electron configuration: Aufbau principle Pauli Exclusion Principle Hund’s Rule

11. Aufbau Principle Electrons occupy orbitals of lower energy first.

12. Filling Order diagram

13. -Pauli Exclusion Principle -Electron Spin Quantum Number An orbital can hold only two electrons and they must have opposite spin.

14. Hund’s Rule Electrons will fill the orbitals in a way that would give the maximum number of parallel spins (maximum number of unpaired electrons). Analogy: Students could fill each seat of a school bus, one person at a time, before doubling up.

15. Orbital Diagram for Hydrogen

16. Orbital Diagram for Helium

17. Orbital Diagram for Lithium

18. Orbital Diagram for Beryllium

19. Orbital Diagram for Boron

20. Orbital Diagram for Carbon

21. Orbital Diagram for Nitrogen

22. Standard Notation of Fluorine Main Energy Level Numbers 1, 2, 2 Sublevels Number of electrons in the sub level 2,2,5 1s 2 2s 2 2p 5

23. Blocks in the Periodic Table

24. Electron Dot Diagrams Represent only Valence Electrons (electrons in the highest energy level, reactive electrons).

25. Examples: Carbon: Fluorine:

26. Shorthand Notation Use the last noble gas that is located in the periodic table right before the element. Write the symbol of the noble gas in brackets. Write the remaining configuration after the brackets. Ex: Fluorine: [He] 2s 2 2p 5

27. Oxidation Numbers + or – number assigned to an atom to indicate its degree of oxidation (loss of electrons) or reduction (gain electrons). Metals lose electrons. Non-metals gain electrons. Octet Rule: rule of 8. Electrons want to get to 8 in outer shell to become stable. Noble gases generally don’t react. Have 8!

28. Oxidation Numbers

29. The 4 Quantum Numbers The Pauli Exclusion Principle: No two electrons in an atom can have the same set of four 2 quantum numbers.