1. CHE1101, Chapter 8 Learn, 1 The Basics of Chemical Bonding Ionic Bonding Octet Rule Lewis Symbols Covalent Bonds Polarity & Electronegativity Lewis Structures Resonance

2. CHE1101, Chapter 8 Learn, 2 Chemical Bonds Three basic types of bonds –Ionic Ions held together via unlike charges –Covalent Sharing of electrons. –Metallic Metal atoms bonded to several other atoms.

3. CHE1101, Chapter 8 Learn, 3

4. CHE1101, Chapter 8 Learn, 4 Lewis Symbols G. N. Lewis developed a method to denote potential bonding electrons by using one dot for every valence electron around the element symbol. When forming compounds, atoms tend to gain, lose, or share electrons until they are surrounded by eight valence electrons (the octet rule).

5. CHE1101, Chapter 8 Learn, 5 Octet rule – atoms gain or lose or share electrons in an effort to obtain 8 valence electrons What is so special about 8 valence electrons ? 8 valence electrons = noble gas configuration noble gas configuration is energetically stable !

6. CHE1101, Chapter 8 Learn, 6 Determine the empirical formula expected for a compound containing Ca and F.

7. CHE1101, Chapter 8 Learn, 7 Covalent Bonding In a covalent “single” bond, 2 electrons are “shared” between 2 atoms Covalently bound species are different than ionic –exist as individual, discrete species (vs. 3-D crystal lattice structure for ionic) –tend to exhibit much lower melting and boiling points (vs. ionic)

8. CHE1101, Chapter 8 Learn, 8 Covalent Bonding

9. CHE1101, Chapter 8 Learn, 9 Lewis Structures Sharing electrons to make covalent bonds can be demonstrated using Lewis structures. We start by trying to give each atom the same number of electrons as the nearest noble gas by sharing electrons. The simplest examples are for hydrogen, H 2, and chlorine, Cl 2, shown below.

10. CHE1101, Chapter 8 Learn, 10 Electrons on Lewis Structures Lone pairs: electrons located on only one atom in a Lewis structure Bonding pairs: shared electrons in a Lewis structure; they can be represented by two dots or one line

11. CHE1101, Chapter 8 Learn, 11 Multiple Bonds single bond – 2 electrons (1 pair of electrons) are shared between 2 atoms double bond – 4 electrons (2 pairs of electrons) are shared between 2 atoms triple bond – 6 electrons (3 pairs of electrons) are shared between 2 atoms

12. CHE1101, Chapter 8 Learn, 12

13. CHE1101, Chapter 8 Learn, 13 a double bond is shorter and stronger than a single bond a triple bond is shorter and stronger than a double bond Multiple Bonds

14. CHE1101, Chapter 8 Learn, 14 Polar Covalent Bonds The electrons in a covalent bond are not always shared equally. Fluorine pulls harder on the electrons it shares with hydrogen than hydrogen does. Therefore, the fluorine end of the molecule has more electron density than the hydrogen end.

15. CHE1101, Chapter 8 Learn, 15 Cl 2 nonpolar bond – electrons are shared equally in the bond polar bond – electrons are NOT shared equally HCl

16. CHE1101, Chapter 8 Learn, 16 electronegativity increases from left right electronegativity decreases from top bottom Electronegativity – the ability of an element to attract electron density to itself in a molecule

17. CHE1101, Chapter 8 Learn, 17

18. CHE1101, Chapter 8 Learn, 18 Arrange the following in order of increasing electronegativity: Na, F, O, K, Al, Si, Mg.

19. CHE1101, Chapter 8 Learn, 19 Electronegativity and Polar Covalent Bonds When two atoms share electrons unequally, a polar covalent bond results. Electrons tend to spend more time around the more electronegative atom. The result is a partial negative charge (not a complete transfer of charge). It is represented by δ–. The other atom is “more positive,” or δ+.

20. CHE1101, Chapter 8 Learn, 20 Rank the following in order of increasing bond polarity: H─FH─BrF─FNa─Cl a)H─F < H─Br < F─F < Na─Cl b)F─F < H─F < H─Br < Na─Cl c)H─Br < H─F < F─F < Na─Cl d)F─F < H─Br < H─F < Na─Cl e)Na─Cl < H─F < H─Br < F─F

21. CHE1101, Chapter 8 Learn, 21 Metallic Bonds The relatively low ionization energy of metals allows them to lose electrons easily. The simplest theory of metallic bonding involves the metal atoms releasing their valence electrons to be shared as a pool by all the atoms/ions in the metal. –An organization of metal cation islands in a sea of electrons –Electrons delocalized throughout the metal structure Bonding results from attraction of cation for the delocalized electrons.

22. CHE1101, Chapter 8 Learn, 22 Metallic Bonding

23. CHE1101, Chapter 8 Learn, 23 a) H 2 O b) C 6 H 12 O 6 c) NO 2 d) Al e) CaCO 3 Which of the following compounds has ionic bonding?

24. CHE1101, Chapter 8 Learn, 24 Write the Lewis Structure for Arsenic.

25. CHE1101, Chapter 8 Learn, 25 Drawing Lewis Structures 1.count the total number of valence electrons 2.make an intelligent guess as to the central element and connectivity a) heavier element is often the central element b) many molecules are symmetric 3.Add electron pairs to satisfy octet rule 4.start making multiple bonds (first double, then triple if single bonds not getting the job done.) 5.Do NOT (under any circumstance…..ever) form a multiple bond to a halogen or hydrogen

26. CHE1101, Chapter 8 Learn, 26 Writing Lewis Structures (Covalent Molecules) 1.Sum the valence electrons from all atoms, taking into account overall charge. –If it is an anion, add one electron for each negative charge. –If it is a cation, subtract one electron for each positive charge. PCl 3 Keep track of the electrons: 5 + 3(7) = 26

27. CHE1101, Chapter 8 Learn, 27 Writing Lewis Structures 2.Write the symbols for the atoms, show which atoms are attached to which, and connect them with a single bond (a line representing two electrons). Keep track of the electrons: 26 − 6 = 20 Heavier element is often the central element

28. CHE1101, Chapter 8 Learn, 28 Writing Lewis Structures 3.Complete the octets around all atoms bonded to the central atom. Keep track of the electrons: 26 − 6 = 20; 20 − 18 = 2

29. CHE1101, Chapter 8 Learn, 29 Writing Lewis Structures 4.Place any leftover electrons on the central atom. Keep track of the electrons: 26 − 6 = 20; 20 − 18 = 2; 2 − 2 = 0

30. CHE1101, Chapter 8 Learn, 30 Writing Lewis Structures 5. If there are not enough electrons to give the central atom an octet, try multiple bonds. Do NOT (under any circumstance…..ever) form a multiple bond to a halogen or hydrogen

31. CHE1101, Chapter 8 Learn, 31 Writing Lewis Structures Then assign formal charges. –For each atom, count the electrons in lone pairs and half the electrons it shares with other atoms. –Subtract that from the number of valence electrons for that atom: the difference is its formal charge.

32. CHE1101, Chapter 8 Learn, 32 Writing Lewis Structures The best Lewis structure… –…is the one with the fewest charges. –…puts a negative charge on the most electronegative atom.

33. CHE1101, Chapter 8 Learn, 33 Draw the Lewis structure for F 2. bonding pair of e - – e - that hold two atoms (bonding pair) together nonbonding pair of e - – e - that are NOT holding (lone pair) 2 atoms together

34. CHE1101, Chapter 8 Learn, 34 Draw the Lewis structure for H 2 O

35. CHE1101, Chapter 8 Learn, 35 Draw the Lewis structure for ethene, C 2 H 4

36. CHE1101, Chapter 8 Learn, 36 Draw the Lewis structure for NO +

37. CHE1101, Chapter 8 Learn, 37 The Best Lewis Structure? Following our rules, this is the Lewis structure we would draw for ozone, O 3. However, it doesn’t agree with what is observed in nature: Both O to O connections are the same.

38. CHE1101, Chapter 8 Learn, 38 Resonance One Lewis structure cannot accurately depict a molecule like ozone. We use multiple structures, resonance structures, to describe the molecule.

39. CHE1101, Chapter 8 Learn, 39 Resonance The electrons that form the second C—O bond in the double bonds below do not always sit between that C and that O, but rather can move among the two oxygens and the carbon. They are not localized; they are delocalized.

40. CHE1101, Chapter 8 Learn, 40 Resonance The organic compound benzene, C 6 H 6, has two resonance structures. It is commonly depicted as a hexagon with a circle inside to signify the delocalized electrons in the ring. Localized electrons are specifically on one atom or shared between two atoms; Delocalized electrons are shared by multiple atoms.

41. CHE1101, Chapter 8 Learn, 41 Determine the Lewis structure for NO 2 - What are your bond expectations for nitrite ?

42. CHE1101, Chapter 8 Learn, 42 A single Lewis structure can NOT be drawn to describe the “real” nitrite species go to lab and measure the actual bond lengths in a real nitrite anion The N-O bonds in nitrite are identical (in every sense; same length; same strength)

43. CHE1101, Chapter 8 Learn, 43 The Real molecule is somewhere in between these two extremes

44. CHE1101, Chapter 8 Learn, 44 Draw the Lewis structure for HNO 3

45. CHE1101, Chapter 8 Learn, 45 Exceptions to the Octet Rule There are three types of ions or molecules that do not follow the octet rule: –ions or molecules with an odd number of electrons, –ions or molecules with less than an octet, –ions or molecules with more than eight valence electrons (an expanded octet).

46. CHE1101, Chapter 8 Learn, 46 Odd Number of Electrons Though relatively rare and usually quite unstable and reactive, there are ions and molecules with an odd number of electrons.

47. CHE1101, Chapter 8 Learn, 47 Draw the Lewis structure for BF 3

48. CHE1101, Chapter 8 Learn, 48 Fewer Than Eight Electrons Consider BF 3 : –Giving boron a filled octet places a negative charge on the boron and a positive charge on fluorine. –This would not be an accurate picture of the distribution of electrons in BF 3.

49. CHE1101, Chapter 8 Learn, 49 Fewer Than Eight Electrons Therefore, structures that put a double bond between boron and fluorine are much less important than the one that leaves boron with only 6 valence electrons.

50. CHE1101, Chapter 8 Learn, 50 Fewer Than Eight Electrons The lesson is: If filling the octet of the central atom results in a negative charge on the central atom and a positive charge on the more electronegative outer atom, don’t fill the octet of the central atom.

51. CHE1101, Chapter 8 Learn, 51 Draw the Lewis structure for PF 5 octet expansion – some atoms can exceed 8 valence electrons (usually P & S)

52. CHE1101, Chapter 8 Learn, 52 More Than Eight Electrons The only way PCl 5 can exist is if phosphorus has 10 electrons around it. It is allowed to expand the octet of atoms on the third row or below. –Presumably d orbitals in these atoms participate in bonding.

53. CHE1101, Chapter 8 Learn, 53 Draw the Lewis structure for PO 4 3-

54. CHE1101, Chapter 8 Learn, 54 More Than Eight Electrons Even though we can draw a Lewis structure for the phosphate ion that has only 8 electrons around the central phosphorus, the better structure puts a double bond between the phosphorus and one of the oxygens.

55. CHE1101, Chapter 8 Learn, 55 More Than Eight Electrons This eliminates the charge on the phosphorus and the charge on one of the oxygens. The lesson is: When the central atom is on the third row or below and expanding its octet eliminates some formal charges, do so.