1. Chemical Bonding Bonding within a molecule is called intramolecular attraction –Ionic bonds –Covalent bonds –Polar covalent bonds

2. Electronegativity Electronegativity is the tendency of an atom to attract electrons. Group 1 and 2 lose electron easily, these elements have a low electonegativity Group 17 and 16 can attain electrons easily, these elements have high electronegativities Delta represents a partial charge or fractional charge

3. Ionic Bonds Steals an e-, this causes one atom to have a positive charge and the other a negative charge. This is the reason for the attraction Opposites attract!! Group 1 and 2 lose electrons to group 16 and 17 (usually) When the difference in electonegativities between atoms is greater than 1.7 the molecules is ionic

4. Example NaCl Na =.9 Cl = 3.0 3.0 -.9 = 2.1 2.1 > 1.7 = ionic bond Cl would be the more negative atom because it has a higher electronegativity (it steals the e- from Na)

5. Covalent Bonds Atoms share electrons because they have similar electronegativities Atoms share electron to fill octet rule Hydrogen form stable molecules were it shares two electrons, this is the duet rule Single bond formed when atoms share one pair of e- Double bond formed when atoms share two pair of e- Triple bond formed when atoms share three pair of e- The electonegativites differences are 0 -.3

6. Example O 2 O = 3.5 3.5 – 3.5 = 0 0 falls between 0 and.3 therefore O 2 is a covalent bond

7. More examples of covalent bonds Ø

8. Polar covalent bonds Electrons not always shared equally. The atom with the higher electronegativity attract the shared electron pair more strongly pulling it away from the other atom. The shared pair is shifted from the center between the two participating atoms making one end of the molecule positive and the other end negative. The bond is polarized. (Dipole – one side of molecule is slightly negative and one part slightly positive) The difference in electronegativity among the atoms is.4 to 1.7

9. Now practice with worksheet #55 Example PO 3 P = 2.1 O = 3.5 3.5 – 2.1 = 1.4 1.4 falls between.4 and 1.7 therefore PO3 is a polar covalent bond Oxygen would be the more negative atom because of the greaterelectonegativity

11. Lewis Structures Represents individual valence electrons Represents a nonbonding pair or lone pair of electrons Represents a pair of e-

12. B. Lewis Structures Electron Dot Diagrams –show valence e - as dots –distribute dots like arrows in an orbital diagram –Show a single line for a single bond, double line for a double bond, triple line for a triple bond –4 sides = 1 s-orbital, 3 p-orbitals –EX: oxygen 2s2p O X

13. Lewis Structures 1) Count all valence electrons in all the atoms in the molecule; it doesn’t matter which atoms they come from. 2) If the compound has more than 2 atoms, the least electronegative atom is the central atom, often a single atom. If carbon is present, it is almost always the central atom. Hydrogen is never a central atom since it can only form one bond.) 3) Make a bond between the central atom and the other atoms using a dash (-) to show a pair of shared electrons. 4) Place the remaining valence electrons around each of the atoms so that they have an octet.

14. Lewis structure When more than one Lewis structure can be drawn for a particular molecule this molecule exhibits resonance

15. Example of an electron dot diagram for CCl 4 First step determine the number of valence electrons –C has 4 (remember 2 from 2s and 2 from 2p) –Cl has 7 (4) (remember the 7 comes from 2e- from the 3s and 5ee from 3p) the 4 comes from Cl 4 –4 + 7 (4) = 32 valence electrons CL OO O Draw the e- dot diagram Now count to see if you have 32 valence e-

16. Your turn, draw the e- dot diagram for CH 3 I Step 1 – count valence e- Step 2 – draw e- dot diagram remember (If the compound has more than 2 atoms, the least electronegative atom is the central atom, often a single atom. If carbon is present, it is almost always the central atom. Hydrogen is never a central atom since it can only form one bond.) Step 3 count to see if you have the correct number of Valence e-

17. Did you remember to: Check to see if all of the valence electrons got used? Check to make sure every atom has an octet of electrons (or 2 for Hydrogen)

18. There are 14 Valence e- (4 +3 + 7) H only needs two e- to satisfy the duet rule Iodine and Carbon needs 8 e- to satisfy the octet rule C H I H H How did you do? Now practice with worksheet #56 and #57

19. Metallic Bonding Metallic bonding is the attraction between metal atoms and a sea of surrounding valence electrons A result of this bonding is mobile electons which gives rise to the excellent electrical conductivity of most metals.

21. Intermolecular Forces Intermolecular forces – bond that holds molecules together. –Effect boiling and freezing points –http://www.bcpl.net/~k drews/interactions/inter actions.htmlhttp://www.bcpl.net/~k drews/interactions/inter actions.html

22. Intermolecular Forces Dipole – Dipole –Molecules are attracted to each other as a result of partial charges of dipole molecules this is called a dipole-dipole intermolecular force –In general, intermolecular forces are about 1% as strong as intramolecular forces.

23. Intermolecular Forces Hydrogen Bonding –Especially strong dipole- dipole are called hydrogen bonding, example = water –When F,O and N are attached to hydrogen they will form hydrogen bonds with other molecules –Common molecules that form hydrogen bonds are HF, H2O and NH3. –Molecules with an O- Hbond, like alcohols also exhibit hydrogen bonding

24. Hydrogen Bonding – type of attraction that holds two water molecules together. Cohesion – attractive force between particles of the same kind Adhesion – attractive force between unlike substances (meniscus)

25. Intermolecular Forces London Dispersion Forces (LDF) –The force that holds noble gas atoms and nonpolar molecules together –No permanent dipoles however, at any given time, the e- can be mostly on one side of the molecule creating a slightly negative charge on one side compared to the other –Weak and easily broken force call LDF –Larger the molecule the more e- it has therefore the stronger LDF

26. Bonding forces and boiling and freezing points Strongest bonding force is a hydrogen bond, followed by a dipole-dipole force, and then LDF. H-Bonds have highest boiling, melting and freezing points Dipole-dipole force have intermediate boiling, freezing and melting points LDF have lowest boiling, freezing and melting points

27. Now practice with worksheet #58

28. VSEPR model Valance Shell Electron Pair Repulsion (VSEPR) is a 3D model of a molecule 1.Draw the Lewis structure for the molecule 2.Count the electron pairs and arrange them in the way that minimizes repulsion(that is, put the pairs as far apart as possible) 3.Determine the positions of the atoms from the way the electron pairs are shared 4.Determine the name of the molecular structure from the positions of the atoms. 5.http://wunmr.wustl.edu/EduDev/ Vsepr/table1.htmlhttp://wunmr.wustl.edu/EduDev/ Vsepr/table1.html

29. Linear Number of bonding pairs around central atom = 2 Number of lone pairs around central atom = 0 180 bond Example- BeCl2, CO2

30. Triangle (Trigonal planar structure) Number of bonding pairs around central atom = 3 Number of lone pairs around central atom = 0 Examples- BF3, SO3

31. tetrahedral Number of bonding pairs around central atom = 4 Number of lone pairs around central atom = 0 109.5 bond angle Example- CH4

32. Trigonal pyramid Number of bonding pairs around central atom = 3 Number of lone pairs around central atom = 1 120 bond angle Example-NH3

33. Bent or v-shaped Number of bonding pairs around central atom = 2 Number of lone pairs around central atom = 1 or 2 106 bond angle Example H2O, SO2