1. Chemical Bonding Warm-up What determines the reactivity of a metal? What determines the reactivity of a non-metal?

2. Chemical Bonding Warm-up What determines the reactivity of a metal? Ionization energy What determines the reactivity of a non-metal? Electronegativity or electron affinity

3. Chemical Bonding Unit 4 Learning Targets 1.) Compare and contrast properties and characteristics of ionic, covalent, and metallic bonds 2.) Draw Lewis dot structures and explain how they predict the types of bonds that will stabilize an atom 3.) Use Valence Shell Electron Pair Repulsion (VSEPR) principles to predict and model the geometries and polarity of molecules 4.) Describe intermolecular forces of attraction

4. Chemical Bonding Unit 4 Vocabulary Ionic bondLewis Structure Covalent bondIonic compound PolarPolyatomic ion MoleculeMetallic bonding Diatomic moleculeIntermolecular forces Activity Make a Frayer model for each new vocabulary word in your notebook or on notecards. The terms are located in Chapter 6 pgs. 160-194 in your textbook.

5. Chemical Bonding Learning Target #1 – Compare and contrast properties and characteristics of ionic, covalent, and metallic bonds.

6. Chemical Bonding Ionic bonds – chemical bond between a positive and negative ions due to a transfer of electrons that forms ionic compounds Covalent bonds – chemical bond between atoms that have shared one or more pairs of electrons to form a molecule Metallic bonds – chemical bond that results from the attraction between metal atoms and the surrounding delocalized electrons

7. Chemical Bonding

8. Ionic compound Molecule Metal

9. Chemical Bonding Polar – a polar bond is a covalent bond in which there is a difference of 0.3-1.7 in the atoms’ electronegativity When the difference is less than 0.3 the bond is nonpolar When the difference is greater than 1.7 the atoms will typically transfer electrons to form an ionic compound

10. Chemical Bonding Warm up Using the shell models and electron configurations below notate on your main block periodic table how many valence shell electrons each of the first 18 elements has. What trend do you see? Use the trend to fill in the # of valence electrons for the rest of the main block.

11. Chemical Bonding Warm up The members of a column or group all have the same # of valence electrons. When each element forms an ion it will gain or lose valence electrons until it has a full outer shell. Each electron gained adds a -1 charge while each electron lost adds a +1 charge.

12. Chemical Bonding Atoms will gain or lose electrons to form ions according to their electron configuration. An ion with a full outer shell of eight electrons like a noble gas is the most stable. This is called an octet.

13. Chemical Bonding Warm up 1.Why are metals such good conductors of electricity? 2.How does an IONIC BOND stabilize the atoms that are bonded together? 3.How does a COVALENT BOND stabilize the atoms that are bonded together?

14. Chemical Bonding Warm up 1.Why are metals such good conductors of electricity? The overlapping mostly empty electron shells allow electrons to flow freely throughout the metal. 2.How does an IONIC BOND stabilize the atoms that are bonded together? When the electrons are transferred the resulting ions have a full valence shell of electrons. 3.How does a COVALENT BOND stabilize the atoms that are bonded together? Electrons are shared giving each atom a full valence shell of electrons.

15. Chemical Bonding Learning Target #2– Draw Lewis dot structures and explain how they predict the types of bonds that will stabilize an atom

16. Chemical Bonding Electron dot notation– Elemental symbol plus “dots” for every valence shell electron

17. Chemical Bonding Lewis dot structure – Diagram that shows how a molecule is bonded together; elemental symbols stand for atomic nuclei and inner-shell electrons, dashes stand for shared pairs or bonding pairs of electrons, and dots stand for unshared or lone pairs of electrons

18. Chemical Bonding Guidelines for drawing Lewis structures 1.Arrange atoms as they will be bonded together, central atom is usually the least electronegative but cannot be H 2.Calculate the total # of valence electrons in the molecule 3.If the molecule is charged, add or subtract electrons accordingly 4.Draw single covalent bonds connecting outer atoms with the central atom 5.Using the remaining electrons fill in the octets of the outer atoms 6.If you have extra electrons place additional lone pairs around the central atom (only if in period 3, 4, 5, 6, or 7) 7.If atoms do not have an octet share additional pairs of electrons in double or triple bonds 8.EXCEPTIONS! H can only have 2 valence electrons B usually only has 6 valence electrons

19. Chemical Bonding Activity As a class we will draw Lewis Structures for F 2, O 2, N 2, CH 4, BF 3, NH 4 +, SF 6, C 2 H 4

20. Chemical Bonding Warm-up 1.) Draw the Lewis Structure for O 3 and CF 2 Cl 2 2.)Draw a 2 nd VALID Lewis Structure for ozone

21. Chemical Bonding Warm-up 1.) Draw the Lewis Structure for O 3 and CF 2 Cl 2 2.)Draw a 2 nd VALID Lewis Structure for ozone Resonance

22. Chemical Bonding Warm up Watch the following video closely... http://www.youtube.com/watch?v=s8akB9NGhYs&safe=active From our Lewis structures can you name a molecule that is... Linear Trigonal Planar Tetrahedral

23. Chemical Bonding Learning Target #3 Use Valence Shell Electron Pair Repulsion (VSEPR) principles to predict and model the geometries and polarity of molecules

24. Chemical Bonding Valence Shell Electron Pair Repulsion Key principles: 1.Electrons repel each other 2.Electrons are found in bonding pairs and lone pairs around bonding atoms 3.These electron pairs will orient as far apart from each other as possible 4.This determines the 3D shape of a molecule around a central atom 5.The electron domain is the # of shared pairs and unshared pairs around a central atom 6.Lone pairs repel MORE than bonding pairs

25. Chemical Bonding Valence Shell Electron Pair Repulsion CO 2 has two electron domains The furthest spread is 180⁰ and results in a LINEAR shape

26. Chemical Bonding Valence Shell Electron Pair Repulsion BH 3 has three electron domains The furthest spread is 120⁰ and results in a TRIGONAL PLANAR shape

27. Chemical Bonding Valence Shell Electron Pair Repulsion CH 4 has four electron domains The furthest spread is 109.5⁰ and results in a TETRAHEDRAL shape

28. Chemical Bonding Valence Shell Electron Pair Repulsion PF 5 has five electron domains The furthest spread is a mix of 120⁰and 90⁰and results in a TRIGONAL BIPYRAMIDAL shape

29. Chemical Bonding Valence Shell Electron Pair Repulsion SF 6 has six electron domains The furthest spread is a 90⁰and results in a OCTAHEDRAL shape

30. Chemical Bonding Valence Shell Electron Pair Repulsion Other molecular shapes are possible when bonding pairs are replaced by lone pairs Bent Tetrahedral Trigonal pyramidal

31. Chemical Bonding

32. This is a great review video for VSEPR and includes 3D models! http://ocw.mit.edu/resources/res-tll- 004-stem-concept-videos-fall- 2013/videos/representations/vsepr/ Valence Shell Electron Pair Repulsion Shorter video with just the 3D models http://www.youtube.com/watch?v=i3F CHVlSZc4&safe=active

33. Chemical Bonding Warm-up What is the molecular geometry around... Carbon #2 Carbon #9 Nitrogen #8 All of the other Nitrogen atoms? Remember to count the lone pairs and bonds around each atom! My good friend, caffeine!

34. Chemical Bonding Warm-up What is the molecular geometry around... Carbon #2 tetrahedral Carbon #9 Trigonal planar Nitrogen #8 bent All of the other Nitrogen atoms? Trigonal pyramidal My good friend, caffeine!

35. Chemical Bonding Warm-up 1.) Rank the following bonds from least polar to most polar using the differences in electronegativity. H-O, H-H, H-F, H-N 2.) The bonds of CO 2 are polar but the molecule itself is not. Using your knowledge of Lewis Structures and VSEPR make a guess why that would be the case.

36. Chemical Bonding Warm-up 1.) Rank the following bonds from least polar to most polar using the differences in electronegativity. H-H, H-N, H-O, H-F 2.) The bonds of CO 2 are polar but the molecule itself is not. Using your knowledge of Lewis Structures and VSEPR make a guess why that would be the case. Since the bonds are symmetrical the polar bonds cancel each other out.

37. Chemical Bonding Learning Target #4 Describe intermolecular forces of attraction

38. Chemical Bonding Dipole-dipole forces The forces of attraction between polar molecules. Requires a bond or bonds involving atoms with a difference of greater than 0.4 in electronegativity The dipoles must be arranged in an additive manner (not be equal and opposite)

39. Chemical Bonding Hydrogen bonding The forces of attraction between a hydrogen bonded to a highly electronegative atom and a lone pair of electrons. Requires a bond or bonds between H-F, H-O, or H-N Requires a molecule with lone pairs of electrons Strongest IMF

40. Chemical Bonding London Dispersion Forces The forces of attraction caused by temporary dipoles resulting from the constant movement of electrons Present in ALL molecules Only IMF present in non-polar molecules Weakest IMF

41. Chemical Bonding Effects of Intermolecular Forces Stronger IMFs hold molecules more closely together Molecules with stronger IMFs have higher melting points and boiling points Molecules with strong IMFs like water display properties like surface tension and higher viscosity

42. Chemical Bonding Warm-up What IMFs are present in the following molecules? Which molecule will have the highest boiling point? Which will have the lowest boiling point? Butane Propanol Propanal