1. UNIT 7: BONDING What previous knowledge will help us understand bonding? How can we describe energy involved in a chemical bond? How can we explain and draw ionic bonds? How can we explain and draw covalent bonds? What are metallic bonds and why are they good conductors? What is the difference between bond polarity and molecule polarity? How are molecules geometrically arranged? How does the VESPR theory influence the geometry of molecules? What are the different forces that hold molecules together? What are sigma and pi bonds?

2.   ENDOTHERMIC:   EXOTHERMIC:   POTENTIAL ENERGY:   Why do atoms become ions?   How do atoms become ions?   How do metals form ions?   How do nonmetals form ions? chemical reaction that absorbs heat, producing products with more PE than the reactants chemical reaction that produces heat with less PE than the reactants stored energy based on position or composition To become stable Gain or loss of electrons Lose electrons – become positive Gain electrons – become negative

3. BONDING  Chemical bonds provide the glue that holds all compounds together  The electron structure of atoms helps explain many aspects of chemical bonding Ionic Bonding

4. ENERGY AND CHEMICAL BONDS  Chemical bonds are the forces that holds atoms together.  Energy is required to overcome these attractive forces and separate the atoms in a compound  Breaking a chemical bond is an endothermic process  ENDOTHERMIC: chemical reaction that absorbs energy producing products with more potential energy than the reactants Ex) N 2 + ENERGY  N + N

5. ENERGY AND CHEMICAL BONDS  Formation of a bond is an exothermic process. EXOTHERMIC: chemical reaction that releases energy producing products with less potential energy than the reactants Ex) N + N  N 2 + ENERGY

6. Bonding and stability  When bonds are formed their products are more stable  The compounds have smaller amount of potential energy POTENTIAL ENERGY: stored energy based on position or composition  The bonded elements of a compound are more stable than the individual atom or ions because the atoms have filled their valence electron shell

7. Types of bonds  There are three types of bonds: 1.Ionic 2.Covalent 3.Metallic  They differ in the types of elements involved.  Also how the valence electrons are handled

8. IONS  Atoms become ions so that they can become stable. Atoms become ions by either gaining or losing electrons

9. IONS Metals form ions by losing electrons and becoming a positive ion with a smaller atomic radius. Positive ions are called cations. Nonmetals form ions by gaining electrons to become a negative ion with a larger atomic radius. Negative ions are also called anions

10. IONIC BONDING The bond that involves the transfer of one or more electrons from a metal atom to a nonmetal atom to form ions. The positive ion and negative ion attract each other and create a bond.

11. IONIC BONDING There is a large electronegativity difference (E.D.) between a nonmetal and a metal. The nonmetal rips away the valence electron from the metal atom. Nonmetal becomes a negative ion or anion and the metal becomes a positive ion or cation

12. IONIC BONDING Examples: 1 - EX) when Na and Cl atoms come together  Na loses electrons and becomes Na +1 Cl gains electrons and becomes Cl -1 They attract to form ionic compound NaCl (aka: salts or electrolytes ) Ionic bonds have the highest polarity (most unequal type of bonding) and the most ionic character

13. CLUE FOR RECOGNIZING IONIC BONDS 1.Metal and nonmetal 2. Electronegativity difference is greater than 1.7 (approx)

14. Do Now January 5, 2016  Take out homework from last night   What is an ionic bond?   What are the clues for recognizing an ionic bond?   What happened to electrons in an ionic bond?

15. Properties of Ionic Bonds 1. High melting and boiling point. 2. 2. Good electrical conductor as a liquid or when dissolved in water -ionic breaks into ions, can cause current to flow 3. Not good electrical conductor as solid 4. Hard substances -attract each other strongly because of opposite charges

16. Electron Dot Structure for Elements  1. Write the symbol for the element in question  2. Determine the number of valence electrons for that atom (group #)  3. Use dots (. ) to represent the number of valence electrons  Place the dots on the top/bottom/left/right sides of the symbol Ex: N N : Electron Dot Diagram...

17. Lewis Dot Structure  Examples:  Sodium  Calcium  Fluorine  Oxygen  Neon

18. In-Class Practice  Drawing Lewis Dot structures- packet pg. 6-7 HOMEWORK -finish packet p.6-7 -quiz Fri. on packet p.1-11 -Periodic Table Activity Lab due Friday

19. Wednesday, January 6, 2016 DO NOW  Take out your HW please!  On looseleaf:  Draw the Lewis dot structure for Na. How many electrons is Na likely to gain/lose?  Draw the Lewis dot structure for Cl. How many electrons is Cl likely to gain/lose? Homework: Page 11 of your packet! -Quiz on Friday- pkt pgs. 1-11 -Periodic Table Activity lab due Friday

20. Check HW  Display answers to note packet pg. 6 & 7 Display answers to note packet pg. 6 & 7 Display answers to note packet pg. 6 & 7

21. DRAWING LEWIS DOT STRUCTURES FOR IONIC BONDS   Write the metal symbol with no dots in brackets   Place the charge at the top right of the bracket   Write the nonmetal symbol with 8 dots around it (except H!)   Draw brackets around the symbol and place the charge of the ion at the top right of the bracket Ex: Sodium Chloride [Na] +1 [ :Cl: ] -1 :

22. DRAWING LEWIS DOT STRUCTURES FOR IONIC BONDS   Example: Draw the Lewis dot structure of the following compound NaF

23. DRAWING LEWIS DOT STRUCTURES FOR IONIC BONDS   Example: Draw the Lewis dot structure of the following compound MgCl2

24. DRAWING LEWIS DOT STRUCTURES FOR IONIC BONDS   Example: Draw the Lewis dot structure of the following compound Na2O

25. PRACTICE KH MgI 2 AlBr 3 KF Na 2 O MgO BeS

26. Conclusion  On the index card provided write down the following.  1. Lewis Dot Diagram for Barium chloride  2. List anything you do not understand Homework: Page 11 of your packet! -Quiz on Friday- pkt pgs. 1-11 -Periodic Table Activity lab due Friday

27. Do Now January 7, 2016  Answer the questions on page 8 & 9 of your packet. (There are 5 of them!)

28. IN-CLASS ASSIGNMENT January 7, 2016  Ionic Bonding Practice Ionic Bonding Practice Ionic Bonding Practice  KEY KEY

29. DO NOW January 8, 2016  Hand in lab- Periodic Table Activity  Make sure name and # are on it  Get ready for quiz! Set up a barrier  HW- Midterm Review packet pgs.1- 2

30. Bonding COVALENT/METALLIC

31. ionic Covalent molecular metallic Bonding m,nmall nm all m +,- transfershare M-SOME Polar covalent Non polar covalent

32. AIM - AIM - What is it about the structure of noble gases that leads to their stability? *NOT IN PACKET!! WRITE IN A BLANK AREA!!   Noble gases (group 18) are stable and undergo few chemical reactions – lack reactivity   ( Argon and Xenon combine with fluorine – rare)   Why? – All have 8 valence electrons except He with 2   Octet – configuration of 8 valence electrons represents max # of valence electrons an atom can have (except H and He – max 2)

33. Aim- How can we understand covalent (molecular) bonds?  The sharing of electrons in order to obtain a stable octet   Octet Rule – states atoms generally react by gaining, losing, or sharing electrons in order to achieve a complete octet of 8 valence electrons – noble gas configuration  Covalent bonds between binary compounds result in one atom having a greater attraction for electrons in a bond. (greater electronegativity)

34. Covalent Bonding Covalent Bonding Video

35. DO NOW January 11, 2016  Take out your Periodic Table and your midterm review packet HW  How is a covalent bond different than an ionic bond?  What is the octet rule?  HW- Packet part 2, pg. 2 (says “Covalent Bonding” at top

36.  All Diatomic Molecules are covalent. They are N 2 - Is the only diatomic molecule that has a triple bond at room temperature. (Very stable but unusable) COVALENT BONDING Br 2, I 2, N 2, Cl 2 H 2, O 2, F 2,

37. PROPERTIES   Exist in gas, liquid, or solid state   Good insulators   Poor conductors   Low melting points   Many are soft substances

38. CONSTRUCTING LEWIS DOT STRUCTURE FOR SINGLE BINARY COVALENT MOLECULAR COMPOUNDS 1. 1. Determine valence electrons in total (add them up for each element in the compound 2. 2. Divide by 2 to determine the number of pairs of electrons in total for the compound 3. 3. Place first pair between the two elements (use a dash – to represent the shared pair) 4. 4. Place remaining pairs around each elements making sure not to violate the octet rule (Remember H can have a max of 2 electrons)

39. CONSTRUCTING LEWIS DOT STRUCTURE FOR SINGLE BINARY COVALENT MOLECULAR COMPOUNDS 1. 1. Determine valence electrons for each element in the compound 2. 2. Distribute dots around the symbols 3. 3. Use dashed line to indicate bond (pair of shared electrons)

40. H2H2 Cl 2 Br 2

41. HClHBr

42.  CONSTRUCTING LEWIS DOT STRUCTURES FOR SINGLE TERNANRY COVALENT MOLECULAR COMPOUNDS (more than two elements involved) 1. 1. Determine valence electrons in total (add them up for each element in the compound 2. 2. Divide by 2 to determine the number of pairs of electrons in total for the compound 3. 3. Determine the most electronegative element and place it in the middle 4. 4. Place the other elements around it 5. 5. Start placing pairs (as dash lines) between the central atom and the terminal atoms 6. 6. Place remaining around each elements making sure not to violate the octet rule (Remember H can have a max of 2 electrons)

43.  CONSTRUCTING LEWIS DOT STRUCTURES FOR MULTIPLE COVALENT BONDS (more than two elements involved) 1. 1. Identify the single atom and use it as the central atom 2. 2. Determine the valence electrons of the central atom 3. 3. Distribute around the symbol 4. 4. Determine the valence electrons of all other atoms in the bond 5. 5. Distribute around the symbol 6. 6. Use dashes to indicate the bond

44. NH 3 CH 4 H2OH2O

45. CONSTRUCTING LEWIS DOT STRUCTURES FOR MULTIPLE COVALENT MOLECULAR COMPOUNDS (more than two elements involved) 1. 1. Determine valence electrons in total (add them up for each element in the compound 2. 2. Divide by 2 to determine the number of pairs of electrons in total for the compound 3. 3. Determine the most electronegative element and place it in the middle 4. 4. Place the other elements around it 5. 5. Start placing pairs (as dash lines) between the central atom and the terminal atoms 6. 6. Place remaining around each elements making sure not to violate the octet rule (Remember H can have a max of 2 electrons) 7. 7. If octet rule is not yet reached you can make additional pairs of electrons into double or triple bonds until octet rule is obeyed by all elements *can only be done with CNOPS

46. CO 2 O2O2 N2N2 More practice: Multiple covalent bonds

47. Do Now January 12 th, 2016  Take out homework from last night (part 2 of packet, pg. 2) and a Periodic Table  Answer questions 1-8 on pg. 7 HOMEWORK:  RB pg.105 #13-17, p.107 #18-23  Test on Tuesday, 1/19  Midterm on Monday, 1/25!!

48. How can we understand and recognize metallic bonding?  Contains positively charged metals  Metals are arranged in a crystalline lattice structure immersed in a: sea of mobile electrons  Electrons are delocalized. This means no one atom owns any electrons they belong to the whole crystal. MSOME Metals are good conductors because of mobile ions

50. Sea of Mobile Electrons

51. Classwork  Read page 102-103 answer questions 10-12 on page 103

52. Do atoms in covalent compounds always share the electrons equally?  https://www.youtube.com/watch?v=15Nvv wuTr5E https://www.youtube.com/watch?v=15Nvv wuTr5E https://www.youtube.com/watch?v=15Nvv wuTr5E

53. How can we use the 1.7 rule to predict bond polarity?  1.7 rule is applied primarily to binary compounds.  Determine the electronegativities of all atoms in the bond.  Take the difference between the bonded atoms  If the difference is: >1.7 then ionic < 1.7 but not “0”, then polar covalent < 1.7 but not “0”, then polar covalent =0 non polar covalent

54. Nonpolar covalent Polar covalent

55. Nonpolar covalent bond Polar covalent bond Ionic bond Covalent bonds

56. MOLECULE POLARITY: Nonpolar/polar shapes SNAP PAD  SNAP : S ymmetrical N onpolar A symmetrical P olar  PAD: P olar A symmetrical D ipole  OPEN: O dd P olar E ven N onpolar

57. How can a molecule be both polar and non polar?  CX 4 tetrahedral shape. They are non polar.  The individual or bonds are polar  Conclusion – nonpolar/polar.

58. Classwork  Page 11 of your packet

59. Do Now January 13 th, 2016  Take out homework (RBp.105 and 107)  Take out Reference Tables  Answer the questions on the handout

60. What Determines the Shape of a Molecule?  Simply put, electron pairs, whether they be bonding or nonbonding, repel each other.  By assuming the electron pairs are placed as far as possible from each other, we can predict the shape of the molecule. © 2009, Prentice-Hall, Inc.

61. MOLECULAR SHAPES -Linear -Bent (angular) -Pyramidal -Tetrahedral

62. Intermolecular Forces The attractions between molecules are not nearly as strong as the intramolecular attractions that hold compounds together. © 2009, Prentice-Hall, Inc.

63. Intermolecular Forces They are, however, strong enough to control physical properties such as boiling and melting points, vapor pressures, and viscosities. © 2009, Prentice-Hall, Inc.

64. Intermolecular Forces These intermolecular forces as a group are referred to as van der Waals forces. © 2009, Prentice-Hall, Inc.

65. How Do We Explain This?  The nonpolar series (SnH 4 to CH 4 ) follow the expected trend.  The polar series follows the trend from H 2 Te through H 2 S, but water is quite an anomaly. © 2009, Prentice-Hall, Inc.

66. Hydrogen Bonding  The dipole-dipole interactions experienced when H is bonded to N, O, or F are unusually strong.  We call these interactions hydrogen bonds. © 2009, Prentice-Hall, Inc.

67. Hydrogen Bonding  Hydrogen bonding arises in part from the high electronegativity of nitrogen, oxygen, and fluorine. © 2009, Prentice-Hall, Inc. Also, when hydrogen is bonded to one of those very electronegative elements, the hydrogen nucleus is exposed.

68. Intermolecular Forces Affect Many Physical Properties The strength of the attractions between particles can greatly affect the properties of a substance or solution. © 2009, Prentice-Hall, Inc.

69. ionic Covalent molecular metallic Bonding m,nmall nmall m +,- Polar Covalent Non Polar Covalent