Write the electron configurations for the following

1. Write the electron configurations for the following • S2- • Ca • Br • O2-

2. Draw Box diagrams for the following • Co • Al • B

3. What did you discover in the periodic properties lab? • What data can you use to support your claim?

4. Alkali metals are more reactive than Alkaline earth metals • As you go down the period things become more reactive.

5. What patterns did you discover in the graphing activity? • Atomic Size • Increases as you go down the periodic table • Decreases as you move from left to right across the periodic table • Ionization Energy • Decreases as you move down the periodic table • Increases from left to right • Increases when an orbital is full

6. How does this relate to reactivity? • Elements are more reactive as you move down the periodic table • Relates to the amount of Electrons in an orbital. • Atoms that only “need a few” electrons or can give away “a few” electrons to have a full orbital are more reactive.

7. 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?

8. Lets go back to the penny challenge… • Why does water hold so many drops compared to hexane? • Why does it bend toward charge? • How does this relate to it being so unique?

9. 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?

10. 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

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

12. NaCl • FeNO3 • KCl • CsSO4 All these substances contain Ionic Bonds What rules could you determine about ionic bonds from examining these compounds?

13. 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

14. These compounds have covalent bonds. What rules could you determine about covalent bonds? • CO2 • H2O • CH4 • SiO2

15. 2. Covalent Bonds • A sharing of a pair of electrons between two atoms • Individual atoms attain a noble gas configuration with the shared electrons in their outer energy level

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

17. Element valence electrons Lewis dot N O F C

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

19. 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) H2O CH4 SCl2

20. Try these..

21. Double and Triple Bonds Example: HCN Make a table: atom have need H 1 2 C 4 8 N 58 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)

22. Try These Examples C3H6 SO2

23. 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 NH4+ OH- SO42- Draw NH4OH

24. Exceptions to the octet rule • Metals MgH2 BH3 2. Molecules with an odd number of electrons NO NO2

25. 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). SF6 PCl5 DON’T FOCUS ON THESE BUT KNOW THEY OCCUR!

26. Try these…. • Mg(OH)2 C3H6 O2

27. 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.

28. 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

29. 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

30. 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

31. 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?

32. When the ΔE.N. is less than 2.0, the bond is covalent Examples: H2O NO2 • 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: H2 NCl3 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? H2 or N Cl

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

34. Draw Partial Charge distribution for the following bonds • C-F Si-H P-H

35. Classify the typeof bonds in each • CH4 • H2O • Na3PO4 • F2

36. Why are molecules a particular shape?

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

38. The structure of Ionic solids

39. Possible Shapes • Linear Ex: • Trigonal Planar • Tetrahedral • Triganol bipyramidal • octahedral

40. 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

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

42. 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)

43. Different shapes • Molecules with only two atoms will always be linear Ex: CO HCl • Molecules with two bonds can have two different shapes Example: BeCl2 H2O

44. This molecule is linear. The H2O molecule is “bent” or “angular”

45. Effect of the lone pairs on shape or H2O

46. Molecules with 3 bonds Can have two shapes Ex: NH3 BF3

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

48. Molecules with 4 bonds If there are no lone pairs on the central atom: Ex: CH4 This is called tetrahedral This makes all bonds equidistant from each other

49. Molecules with 5 bonds (and no lone pairs on the middle atom) Ex.: PCl5 This is called trigonal bipyramidal

50. Molecules with 6 bonds and no lone pairs on the central atom Ex.: SF6 This is called octahedral