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FAIR USE STATEMENT :

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FAIR USE STATEMENT :

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  1. FAIR USE STATEMENT: Please feel free to edit and use this presentation in your classroom. Please do not remove the credit line on the title page or republish the file in whole or in part as your own. Please do not distribute the file to individuals or at conferences or workshops. I am more than willing to share the presentation with anyone that contacts me at rhondaa@cox-internet.com. The images used in the presentation are not original and the presentation is distributed freely but only for classroom instruction. Rhonda Alexander

  2. Quantum Mechanics Principle quantum number Depends on energy level n=(1,2,3, …)

  3. Azimuthal quantum number Shape of orbital-depends on location on periodic chart If s block then l=0 If p block then l=1 If d block then l=2 If f block then l=3

  4. Atomic Orbitals 1 – s 3 - p 5 – d 7 - f

  5. Atomic Orbitals d - orbitals l ranges from 0 to n-1. If n=1, then l=0 If n=2 then l=0,1 If n=3 then l= 0, 1, 2

  6. AngularQuantum Number • ml represents the individual orbitals of a given type. • ml ranges from –l to +l • ml tells you which p, d, or f orbital the electron is in

  7. Magnetic Quantum Number (ms): Spin

  8. QuantumNumbers

  9. Quantum Numbers • n = integer from 1 to 7 • l = 0 to n-1 • ml = -l to +l • ms =

  10. HOW DO WE DESCRIBE THE LOCATION OF EACH ELECTRON IN AN ATOM? WITH A SYMBOLIC NOTATION CALLED AN ELECTRON CONFIGURATION.

  11. An Electron Configuration is a shorthand method of listing the location of the electrons in an atom. The system locates each electron by energy level and sublevel. The number of electrons in each sublevel is indicated with a superscript. For instance, the electronic configuration of Sodium is1s22s22p63s1.This indicates that there are two electrons located in the 1s orbital, two electrons in the 2s orbital, six electrons in the 2p orbital and a single electron in the 3s orbital.

  12. Remember:

  13. Filling Orbitals Follow rules of modern atomic model: Aufbau Principle -electrons fill from lowest energy level first Hund’s Rule -have maximum number of unpaired electrons Pauli Exclusion Theory -no electron has same set of quantum numbers because of electron spin

  14. RULES FOR PLACING ELECTRONS IN ORBITALS 1.   Electrons occupy lowest energy orbitals first. 2.  An orbital can hold a maximum of 2 electrons. The Pauli Exclusion Principle must be obeyed. 3.   Hund's Rule must be obeyed; when placing electrons into degenerate orbitals, there must be one electron in each orbital before any pairing of electrons can take place. {Degenerate orbitals are orbitals of the same energy level and sublevel.}

  15. Aufbau Diagram

  16. The Order Electrons Fill Orbitals

  17. Orbital Diagrams __ ____ __ __ __ __ __ __ 1s2s2p3s3p __ __ __ __ __ __ __ __ __ 4s 3d4p __ __ __ __ __ __ __ __ __ __ 5s 4d5p 6s

  18. Electron Configuration of Zr 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d2

  19. Practice • Write the electron configuration of the following on a separate page. • Oxygen • Lithium • Iron • Bromine • Tin

  20. Core Notation of Sn • Locate Sn on the periodic table

  21. Electron Configuration of Sn • Sn [Kr] • The noble gas core is Kr

  22. Sn [Kr]5s2 • The noble gas core is Kr • From Kr, go 2 spaces across the s-block in the 5th row  5s2

  23. Sn [Kr]5s24d105p2 • The noble gas core is Kr • From Kr, go 2 spaces across the s-block in the 5th row  5s2 • Then go 10 spaces across the d-block on the 5th row  4d10 • Finally go 2 spaces into the p-block on the 5th row 5p2

  24. Steps • Locate the element • Go to the end of the row and up 1 noble gas • Write the Noble gas core in brackets • Continue electron configuration with the next period

  25. Nobel Gas (Core) & Orbital Diagram • Refer to a periodic table and write the electron configurations of these atoms. Use the noble gas core. • Zn • I [Ar]4s23d10 [Kr]5s24d105p5

  26. Zn [Ar]4s23d10 • I [Kr]5s24d105p5       4s 3d          5s 4d 5p

  27. Quiz - Quantum Numbers .n = 3 .l = 2 .ml = +2 .ms = -1/2 • Zn • Sb • Cs • .n = 4 • .l = 3 • .ml = -3 • .ms = -1/2 .n = 5 .l = 1 .ml = +1 .ms = + 1/2 .n = 6 .l = 0 .ml = 0 .ms = +1/2 Tb

  28. The Periodic Table • Is a table that arranges the elements according to similarities in their properties.

  29. Dmitri Mendeleev Father of the Periodic Table

  30. Mendeleev’s Periodic Table (63 known elements) Developed the Periodic Law that said: - columns arranged by increasing atomic mass (not correct) - rows arranged by chem. & physical properties

  31. Mendeleev’s Table (cont.) • Concluded gaps in table were elements yet to be discovered. • - led to the search for missing elements • - predicted existence of aluminum, boron, silicon, germanium

  32. Predicted PropertiesObserved Properties Atomic weight 72 72.61 Density 5.5 g/cm3 5.32 g/cm3 Melting point 825 C 938 C Oxide formula RO2 GeO2 Density of oxide 4.7 g/cm3 4.70 g/cm3 Chloride formula RCl4 GeCl4

  33. Mendeleev’s Original Table

  34. Modern Periodic Table - Developed by Henry Moseley • Solved problems in Mendeleev’s • table. • Periodic law is based on increasing • atomic #, NOT atomic mass.

  35. Structure of the Periodic Table periods: rows going across, numbered 1-7 groups: columns going down, numbered 1-18; aka families • Elements w/in groups have similar • physical and chemical properties.

  36. The Metals • located left of zig-zag or stair-step Properties - good conductors - lusterous - malleable - ductile

  37. The Nonmetals • located right of zig-zag or stair-step • Properties: dull; brittle; poor conductors,

  38. The Metalloids • Border zig-zag / stair-step except aluminum and polonium • Properties: - some metallic - some nonmetallic - semi-conductors

  39. Representative Elements • 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 • IA IIA IIIA IVA VA VIA VIIA VIIIA

  40. metals nonmetals Elements: MetalsNonmetals MetalloidsTransition Metals Inner Transition Groups/Families:

  41. Key Terms • Atomic radius – ½ distance btwn nuclei of two atoms in a molecule. • Ionization energy – amt. of energy needed to remove an e- from an atom. • Electronegativity – tendency for an atom to attract e- • Valence e- - e- on outermost energy level.

  42. Valence Electrons 6 valence electrons 5 valence electrons 7 valence electrons 1 valence electron 4 valence electrons 8 valence electrons 2 valenceelectrons 3 valence electrons Electrons in the outermost level are called valence electrons.

  43. Element Families Group 1)Alkali Metals (I A) - highly reactive, especially with water - 1 valence e- - loses valence e- - become 1+ ions Li, Na, K, Rb, Cs, Fr

  44. Group 2)Alkaline Earth Metals (II A) • - very reactive but less than group 1 • 2 valence e- • lose valence e-’s • become 2+ ions Be, Mg, Ca, Sr, Ba, Ra

  45. Groups 3-12)Transition Metals (B) - most have 2 valence e-, some with 1 or more - all lose valence e- - several become 2+ ions 3 4 5 6 7 8 9 10 11 12

  46. Group 17)Halogens (VII A) - highly reactive - form cmpds calledhalides - fluorine most reactive element - 7 valence e- - Gain 1 e- - become 1- ions F, Cl, Br, I, At

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