1. Do Now: 1) Take out and turn in the homework (reading summary) 2) Put these terms in order of most general to most specific  Orbital, Sublevel, PEL, Spin 3) Write the quantum atom electron configuration AND box diagram for calcium

2. Do now Review (pre quiz!)

3. Quiz time! Adele says: “Put everything away and no talking. Cheating isn’t cool.”

4. Kernel vs. Valence and Periodic Trends

5. The Kernel, Valence Electrons & Lewis Dot Diagrams  A very powerful tool for showing how different elements bond together  Lewis (Electron) Dot Diagrams  developed by a chemist named G.N. Lewis.

6. The Kernel, Valence Electrons & Lewis Dot Diagrams  Many times in chemical bonding:  only the electrons on the outside or in the highest energy level, called the VALENCE SHELL, actually become involved in bonding to form compounds.  We are going to look at the Lewis (Electron) Dot Diagrams and elements in this next section.

7. The Kernel, Valence Electrons & Lewis Dot Diagrams  Valence Shell: Outermost energy level of an element (highest number)  Oxygen: O 2-6 the 2 nd principle energy level is the valence shell

8. The Kernel, Valence Electrons & Lewis Dot Diagrams  Valence Electrons (VE): electrons located in the outermost energy level  These are the electrons which are the furthest to the right in the electron configuration found on the PT  How many VE does Oxygen have? 66

9. The Kernel, Valence Electrons & Lewis Dot Diagrams  Valence Electrons (VE): electrons located in the outermost energy level  the 6 electrons are the valence electrons  In a dot diagram, the VE are represented by dots:   Kernel: This is the nucleus and all the other electrons located in the inner energy levels  In a dot diagram, the chemical symbol represents the kernel: O

10. Sidebar: Visualizing Kernel Valence

11. The Kernel, Valence Electrons & Lewis Dot Diagrams  The valence shell always contains only 4 orbitals.  These orbitals are represented by the four sides around the symbol.  The top side represents s orbital which has less energy than the other three p orbitals.  The other three p orbitals have the same amount of energy. Represented by the other 3 sides.  First 2 electrons are always placed first in this top orbital.  Going around clockwise, one electron is placed on each remaining side until you need to pair the electrons up.

12. 4 sides of the “Kernel” Lower Energy than p – fill w/ 2e - 1st O (Ovals are just drawn to show the orbitals) O S P P P Place 1 e - in each P 1 st before pairing b/c same Energy

13. Period 2 Dot Diagrams SymbolLiBeBC Electron Configur ation Dot Diagram

14. Period 2 Dot Diagrams SymbolNOFNe Electron Configur ation Dot Diagram

15. Guided Practice: look for a pattern! SymbolElectron Configuration (Circle the Valence Electrons (VE) Numbe r of VE Dot Diagra m S Se Na K

16. Guided Practice: look for a pattern! SymbolElectron Configuration (Circle the Valence Electrons (VE) Numbe r of VE Dot Diagra m Si Ge Cl Br

17. Guided Practice: look for a pattern! SymbolElectron Configuration (Circle the Valence Electrons (VE) Numbe r of VE Dot Diagra m Ar Kr

18. What is the pattern in Dot diagrams?

19. PERIODIC TABLE: PROPERTIES OF ELEMENTS and TRENDS  One of the most useful tools to a chemist is the Periodic Table.  There are several important trends or patterns.  Electronegativity is an important tool for predicting the type of bonding that will be present.  We will look at these trends and then use them to help explain and predict bonding later in the course.

20. Electronegativity  Attraction for electrons in a bond  found on Reference Table S  based on Fluorine (F), 4.0 For Metals (left side of PT) it tends to be LOW because inner electrons shield the attraction of the nucleus for other electrons in the valence shell For Nonmetals (right side of PT) it tends to be high because nuclear charge is large compared to small size and attracts the electrons closer

21. Electonegativity Trend  Across a period (L->Rt) Increases  Down a Group Decreases

22. Ionization Energy:  Energy needed to remove the outermost electron  found on Reference Table S  For Metals (left side of PT) it tends to be LOW  For Nonmetals (right side of PT) it tends to be HIGH

23. Ionization Energy:  Across a period (L->Rt) Increases  Down a Group Decreases  Why?  Metals are larger and electrons are attracted less so more easily removed than smaller nonmetal atoms  Down a group, more principle energy levels so outer electrons are further from the attraction of nucleus, more easily removed

24. How do the trends for electronegativity and ionization energy compare?

25. Atomic Radii:  Size of the neutral atom  found on Reference Table S radii Increases as you go down a group (add) on the PT, because there are more energy levels present radii Decreases as you go across the Periodic Table Left  Right because increased nuclear charge attracts energy levels closer

26. Ionic Radii:  Size of the atom after it gains or loses electrons  Metals form __ (+) __ ions (cations) by losing  electrons & get smaller because losing the valence electrons leaves the energy level empty

27. Ionic Radii of Metals:  Bohr Model of Neutral Atom  Bohr model of Li ion of Li (Li +1 )

28. Ionic Radii of Nonmetals  Nonmetals form __ (-) __ ions (anions) by _ Gaining ____  electrons & get Bigger because the added electron(s) expand the atom because there are more electrons than protons

29. Ionic Radii of Nonmetals  Bohr Model of Neutral Bohr Model an ion of Atom of FF (F -1 )

30. Do now:  FINISH YOUR ATOM BOOKS (should be able to do the dot diagrams and what not now)  Finish the homework