1. To get back in the groove.. What did you do for spring break? What is an electron configuration? How do electron configurations relate to reactivity?

2. How many drops of water do you believe I can stack on this penny? How many pennies can I fit in this beaker? How do you suppose water will react when interacting with a charge?

3. Why is water so unique? Why can water bugs run across a pond? Why does water have such a high boiling point compared to other liquids? Why can we live on earth?

4. In this unit we will be able to understand how the Chemical Bonds in a substance determine physical properties Why water is so unique How the bonds that compose a substance determine the properties within How shampoo works How household cleaners work effectively

5. Lets set up your Lab Book Title: Properties Lab Purpose: To study the physical properties of common solids and to investigate the relationship between the type of bonding in a substance and its properties. –Volatility –Melting Point –Solubility –Brittleness –Conductivity

6. Procedure: See Handout Volatility-Waft the substance Solubility (Hexane and Water) in well plate Conductivity (RED &GREEN LIGHT MEANS CONDUCTIVE) Melting Point Heat Watch Glass on beaker of water & test tube in bunsen burner Brittleness (MORTAR AND PESTLE STATION)

7. Data Lauric Acid Aluminum Glucose NaCl SiO 2 Record Observations in Table can make own or paste in lab hand out table Draw a watch glass and record which substances melted (LABEL TO KEEP TRACK OF SUBSTANCES)

8. Disposal Rescue Aluminum if possible Rinse out Sand in Garbage Everything else can go down Sink

9. What did you discover in the periodic properties lab? Which substance was the most volatile? Which substances had the lowest melting point? Which substances conducted electricity? Which substances dissolved in water? Hexane? Which substances do you believe had the strongest bonds? Why? Which substances do you believe have the weakest bonds? Why?

10. How are substances held together? Why are we able to live on the earth? Why is water so “unique” Why can bugs run across the water? Why do metals conduct electricity?

11. It’s all about Chemical Bonds Definition: The force that holds two atoms together. Why does a bond form? So that an atom: 1. becomes more stable 2. takes on a noble gas configuration

12. To determine the type of bond Electronegativity: measure of how strongly an atom attracts the electrons that are shared in a bond

13. NaCl FeNO 3 KCl CsSO 4 All these substances contain Ionic Bonds What rules could you determine about ionic bonds from examining these compounds?

14. Types of Bonds 1. Ionic The attraction between oppositely charged ions Atoms become ions by adding or losing electrons They form these charges to reach a noble gas configuration in their outer energy level Usually a Metal and a Non Metal –Can include Polyatomics Electronegativity difference is larger than

15. These compounds have covalent bonds. What rules could you determine about covalent bonds? CO 2 H 2 O CH 4 SiO 2

16. 2. Covalent Bonds A sharing of a pair of electrons between two atoms Mostly non metals Individual atoms attain a noble gas configuration with the shared electrons in their outer energy level Electronegativity difference smaller than

17. Lewis Dot Drawings Show the sharing or transfer of electrons Also called “electron dot” drawings. Involve only valence electrons (those in the outermost energy level) –Think about Electron Configurations Show the type of bond formed (either ionic or covalent) All atoms will satisfy the “octet” rule (except for hydrogen (duet rule) and metals)

18. Elementvalence electronsLewis dot N O F C

19. Lewis dot drawings for 1.Ionic bonds Show electrons being transferred Include brackets and charges on ions examples: H and F Na and Cl Na and OH -

20. Lewis Structure for Covalent bonds Technique: Place the atom with the largest number of unpaired electrons in the middle. (Never put H in the middle of a molecule!!) Determine how the electrons will be shared so that all atoms are stable (Octet Rule) H 2 O CH 4 SCl 2

21. Try these.. Mg MgCl 2 Mg(OH) 2

22. Double and Triple Bonds Example: HCN Make a table: atomhaveneed H 1 2 C 4 8 N 5 8 total 10 18 Difference: 18-10=8 divide by 2 = 4 You need 4 bonds in this structure Sharing 4 or 6 electrons (Double or Triple bonds allow this to happen)

23. Try These Examples C 3 H 6 SO 2

24. Electron dot drawings for polyatomic ions Always include brackets and charges, but have covalent bonds inside the ion Count the number of valence electrons for each and the add or subtract and electron to make the correct charge NH 4 + OH - SO 4 2- Draw NH 4 OH

25. Exceptions to the octet rule 1.Metals MgH 2 BH 3 2. Molecules with an odd number of electrons NO NO 2

26. 3. Some Nonmetal atoms because of their size, they can have more than an octet of electrons (due to the presence of empty “d” orbitals which can be used for bonding). SF 6 PCl 5 DON’T FOCUS ON THESE BUT KNOW THEY OCCUR!

27. Try these…. Mg(OH) 2 C 3 H 6 O 2

28. Note: Not all covalent bonds have equal sharing of electrons… There are electron hogs!!! Elements that hold on to the electrons more tightly than others You can determine if a bond is ionic,covalent and if there is an electron hogs, through looking at a characteristic property.

29. To determine the type of bond Lets remind ourselves… Electronegativity: measure of how strongly an atom attracts the electrons that are shared in a bond The difference of electronegativity will determine the type of bond

30. What would you predict are the trends in electronegativity? in families? in periods? What family has the highest electronegativity? What family has the lowest electronegativity? What period has the highest electronegativity? What family has the highest electronegativity

31. Electronegativity Allows you to predict the nature of the bond between two atoms To determine where the electrons tend to spend the most time in the molecule

32. To determine the type of bond When the difference in electronegativity (ΔE.N.) is 2.0 or greater, the bond is ionic Examples: NaCl KF Where are these atoms on the periodic table in relation to one another?

33. When the ΔE.N. is less than 2.0, the bond is covalent Examples: H 2 O NO 2 This means the electrons spend more time around one of the elements giving it a partial charge Draw a picture of how you think the electrons would be distributed for each of these molecules. When the electrons are shared equally ex: H 2 NCl 3 the bond is pure covalent and has no partial charge Why do you think there would not be a partial charge on these bonds? Which covalent bond do you think is stronger? H 2 or N Cl

34. These bonds are called intramolecular forces Have various strengths –Ionic (STRONGEST) –Polar Covalent (NEXT STRONGEST) –Covalent (STRENGHTH DEPENDS ON ELECTRONEGATIVITY DIFFERENCE)

35. Draw Partial Charge distribution for the following bonds C-FSi-HP-H

36. Classify the typeof bonds in each CH 4 H 2 O Na 3 PO 4 F 2

37. Why are molecules a particular shape? Get out your lab…Discuss Questions with seat mate I will roll the dice for questions 1-3

38. Shapes of Molecules/Compounds In Molecules Shape is Determined by # of bonds Lone pair Electrons Ionic substances Ions stack together, anions alternating with cations

39. The structure of Ionic solids

40. Possible Shapes Linear Trigonal Planar Tetrahedral Triganol bipyramidal octahedral

41. 2. Covalent compounds The shape of the molecule is determined by the repulsion between the electrons that the atoms share Understanding the placement of electrons can help us determine the shape of a molecule

42. VSEPR Theory Valence Shell Electron Pair Repulsion The shape of a molecule is determined by the repulsion between the electron of the bonded atoms

43. Molecules containing a central atom Electrons want to be as far apart as possible # of bonds(from central atom) influences shape Lone pairs of electrons on a central atom influence shape –They are more repulsive than bonded electrons because they flare (take up more space)

44. Different shapes 1.Molecules with only two atoms will always be linear Ex: COHCl 2.Molecules with two bonds can have two different shapes Example: BeCl 2 H2OH2O

45. This molecule is linear. The H 2 O molecule is “bent” or “angular”

46. Effect of the lone pairs on shape or H 2 O

47. Molecules with 3 bonds Can have two shapes Ex: NH 3 BF 3

48. This one is trigonal pyramidal because of the lone pair of electrons This one is trigonal planar due to the absence of lone pairs on the central atom

49. Molecules with 4 bonds If there are no lone pairs on the central atom: Ex: CH 4 This is called tetrahedral This makes all bonds equidistant from each other

50. Molecules with 5 bonds (and no lone pairs on the middle atom) Ex.: PCl 5 This is called trigonal bipyramidal

51. Molecules with 6 bonds and no lone pairs on the central atom Ex.: SF 6 This is called octahedral

52. Predict the shapes of these molecules PH 3 SO 2 CO 2

53. Practice Make electron dot drawings of the following substances and predict the shape F 2 SiO 2 PF 5 BF 3

54. Back to water Water is so unusual because of the forces “within” the molecule” These are called “intra molecular forces”