1. Atom video http://www.youtube.com/watch?v=x qNSQ3OQMGI&feature=share

2. Basic Principle: electrons occupy lowest energy levels available

4. Aufbau Principle -- “Bottom Up Rule”

5. Electron spin How could an orbital hold two electrons without electrostatic repulsion?  Stern-Gerlach Experiment

6. 1 1 s value of energy level sublevel no. of electrons spdf NOTATION for H, atomic number = 1 spdf Notation Orbital Box Notation Arrows show electron spin (+½ or -½) ORBITAL BOX NOTATION for He, atomic number = 2 1s1s 2 1s1s  2 ways to write electron configurations

7. Example: Determine the electron configuration and orbital notation for the ground state neon atom. An orbital can contain a maximum of 2 electrons, and they must have the opposite “spin.” Pauli exclusion principle

8. Hund’s Rule - Write the ground state configuration and the orbital diagram for oxygen in its ground state

9. Outer electron configuration for the elements

10. Using the periodic table to know configurations Period 1 2 3 4 5 6 7 Ne Ar Kr Xe

11. Valence e ’ s for “main group” elements

12. Rules for Filling Orbitals Bottom-up (Aufbau’s principle) Fill orbitals singly before doubling up (Hund’s Rule) Paired electrons have opposite spin (Pauli exclusion principle) Basic Principle: electrons occupy lowest energy levels available

13. Identify examples of the following principles: 1) Aufbau 2) Hund’s rule 3) Pauli exclusion

14. Shorthand notation practice Examples ●Aluminum: 1s 2 2s 2 2p 6 3s 2 3p 1 [Ne]3s 2 3p 1 ●Calcium: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 [Ar]4s 2 ●Nickel: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 8 [Ar]4s 2 3d 8 {or [Ar]3d 8 4s 2 } ●Iodine: [Kr]5s 2 4d 10 5p 5 {or [Kr]4d 10 5s 2 5p 5 } ●Astatine (At): [Xe]6s 2 4f 14 5d 10 6p 5 {or [Xe]4f 14 5d 10 6s 2 6p 5 } [ Noble Gas Core ] + higher energy electrons

15. Electron configuration for As

16. Note: Not written according to Aufbau, but grouping according to n

17. Orbital energy ladder s p n = 2 s d p n = 3 f s d p n = 4 s n = 1 Energy

18. Phosphorus Symbol: P Atomic Number: 15 Full Configuration: 1s 2 2s 2 2p 6 3s 2 3p 3 Valence Configuration: 3s 2 3p 3 Shorthand Configuration: [Ne]3s 2 3p 3    1s 2s 2p 3s 3p Box Notation

19. Quantum numbers and orbital energies Each electron in an atom has a unique set of quantum numbers to define it { n, l, m l, m s } n = principal quantum number –electron’s energy depends principally on this l = azimuthal quantum number –for orbitals of same n, l distinguishes different shapes (angular momentum) m l = magnetic quantum number –for orbitals of same n & l, m l distinguishes different orientations in space m s = spin quantum number –for orbitals of same n, l & m l, m s identifies the two possible spin orientations

20. Energy levelSublevel# of orbitals/sublevel n = 11s (l = 0)1 (m l has one value) n = 2 2s (l = 0) 1 (m l has one value) 2p (l = 1) 3 (m l has three values) n = 3 3s (l = 0) 1 (m l has one value) 3p (l = 1) 3 (m l has three values) 3d (l = 2) 5 (m l has five values) n = principal quantum number (energy) l = azimuthal quantum number (shape) m l = magnetic quantum number (orientation) Quantum numbers and orbital energies Each atom’s electron has a unique set of quantum numbers to define it { n, l, m l, m s }

21. 21 Concept: Each electron in an atom has a unique set of quantum numbers to define it { n, l, m l, m s }

22. Quantum numbers: unique set for each e - s orbitals p orbitals d orbitals f orbitals l = 0 l = 1 l = 2 l = 3 m l = 0 m l = -1, 0, 1 m l = -2, -1, 0, 1, 2 m l =-3,-2,-1,0,1,2,3 An s subshell A p subshell A d subshell An f subshell One s orbital Three p orbitals Five d orbitals Seven f orbitals For n=1 l=0 an s subshell (with 1 orbital) For n=2 l=0,1 an s subshell and a p subshell (with 3 orbitals) For n=3 l=0,1,2 an s subshell, a p subshell, a d subshell (with 5 orbitals) For n=4 l=0,1,2,3 an s subshell, a p subshell, a d subshell, an f subshell (with 7 orbitals)

23. Electronic configuration of Br 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 5 [Ar] 3d 10 4s 2 4p 5 [Ar] = “noble gas core” [Ar]3d 10 = “pseudo noble gas core” (electrons that tend not to react) Atom’s reactivity is determined by valence electrons valence e’s in Br: 4s 2 4p 5 highest n electrons

24. Valence e - shells for transition metalsmain group elements transition metals v. main group elements d orbitals sometimes included in valence shell d orbitals not included in valence shell (pseudo noble gas cores)

25. Rule-of-thumb for valence electrons Examples ●Sulfur: 1s 2 2s 2 2p 6 3s 2 3p 4 or [Ne]3s 2 3p 4 valence electrons: 3s 2 3p 4 ●Strontium: [Kr]5s 2 valence electrons: 5s 2 ●Gallium: [Ar]4s 2 3d 10 4p 1 valence electrons: 4s 2 4p 1 ●Vanadium: [Ar]4s 2 3d 3 valence electrons: 4s 2 or 3d 3 4s 2 Identify all electrons at the highest principal quantum number (n) Use on exams, but recognize limitations Use Table 8.9 for online HW

26. Selenium’s valence electrons Pseudo noble gas core includes:  noble gas electron core  d electrons (not very reactive) Written for increasing energy:

27. Core and valence electrons in Germanium Pseudo noble gas core includes:  noble gas core  d electrons Written for increasing energy:

28. d-block: some exceptions to the Aufbau principle Fig. 8.9: Use this table for online homework

29. Paramagnetic : atoms with unpaired electrons that are weakly attracted to a magnet. Diamagnetic : atoms with paired electrons that are not attracted to a magnet. Paramagnetic : atoms with unpaired electrons that are weakly attracted to a magnet. Diamagnetic : atoms with paired electrons that are not attracted to a magnet. Electron spin & magnetism For the ground state oxygen atom: spdf configuration: orbital box notation:

30. Apparatus for measuring magnetic properties