1. Chemistry Review You need to remember some basic things

2. The Atom Smallest possible unit that maintains properties of the element Made of: ▫Protons – positively charged particles ▫Neutrons- neutral particles  Together form the atomic nucleus ▫Electrons- negatively charged particles  Fly around the nucleus Each element has a unique number of protons (atomic number)

4. Electron Orbitals/Shells Electrons are found in characteristic areas around the nucleus, called an orbital ▫Each one represents a different energy level Simplifying things, orbitals are grouped into “shells”

5. Electron Shells The innermost shell is filled first Electrons are distributed to each orbital in a shell before filling each orbital The outermost shell is called the valence shell

6. Electron Shells Con. The first shell has only 1 orbital, so it can hold only 2 electrons

7. The 2 nd /3 rd Shell Consists of 4 orbitals, so each shell can hold 8 electrons

9. Draw on your Whiteboard A neutral boron atom (for the nucleus you can just write B) A neutral fluorine atom

10. Using the Periodic Table Ignore the D block (the metals) The row tells you the # of shells the atom should have The column tells you the # of valence electrons a neutral atom should have in its valence shell

11. Draw A neutral magnesium atom A neutral phosphorus atom

12. Ions Aka charged atoms + ions occur when there are more protons than electrons - ions occur when there are more electrons than protons Atoms can gain and lose electrons

13. Draw the ions on your Whiteboard Na+ P 3- Si 2- H +

14. Filling Valence Shells Generally chemical reactions occur that fill valence electron shells Either by gaining/losing electrons OR By sharing electrons with other atoms

15. 6a. Covalent Bond Sharing of electrons between two atoms A single bond consists of 2 shared electrons, which occupy the valence shell of both atoms ▫Double bond = 4 electrons ▫Triple bond = 6 electrons

16. Guidelines of Bonding Atoms almost always will end up with 8 electrons in their valence shell (may be lone pairs or shared electrons) So an atom that normally has 6 valence electrons needs to get 2 more from bonding (only showing the valence electrons)

17. The column can be used to figure out how many bonds an atom will normally form 4 3 2 1 0

18. Lewis Structures A line represents 2 electrons, usually shared in a covalent bond Dots represent electrons that are held by only one atom (lone pairs) Only valence electrons are shown Each atom should have a total of 8 electrons (except H and He which hold 2)

19. Guidelines for Drawing Lewis Structures Carbons make up core Add Hs last (they can’t connect anything) Remember how many bonds each atom will make (using the periodic table

20. On your Whiteboard Draw or make: C 3 H 8 CH 3 OH CF 2 O 2 H 2

21. Draw : H 3 CCH 2 OH H 2 NCH 2 OH

22. Double Bonds Use double bonds or triple bonds when there aren’t enough atoms to form the proper number of bonds i.e. oxygen O=O, if it was a single bond O would not have the correct number of bonds

23. Draw and or make: N 2 HCOOHHCN

24. Drawing ionic molecules (think about total valence electrons present) C should bring 4 valence electrons, N should bring 5. If neutral there would be 9. BUT if it’s – charged there should be 10 total Notice atoms won’t form the correct number of bonds NH 4 + H 3 O + OH -

25. 6b.Polar vs. Non-Polar Covalent Bonds NonpolarPolar Electrons shared equally Both atoms have similar electronegativity (affinity for electrons) Neither atom ends up with any charge Electrons are not shared equally 1 atom is more electronegative (O, F, N, Cl are the usual culprits) Electronegative atom ends up with a partial – charge since they have the electron more often Other atom ends up with a partial + charge as they are deprived of the electron

26. Non-PolarPolar

27. 10. Ion Formation Some atoms more easily give up electrons (1 st and 2 nd columns) to end up with a full valence shell These electrons can be stripped by atoms in the 7 th column (need 1 e) Forms ions

28. 6c. Ionic Bonding Opposites attract! Significantly weaker than a covalent bond Can also occur between ionic molecules

30. Hydrogen Bonds Weak attraction between the partial charges of polar covalently bonded molecules In water, between O and H Means partial