1. 1 Chapter 8 “Covalent Bonding”

2. 2 Bonds Forces that hold groups of atoms together and make them function as a unit: 1) Ionic bonds – transfer of electrons (gained or lost) 2) Covalent bonds – sharing of electrons. The resulting particle is called a “molecule”.

3. 3 Molecules  Many elements found in nature are in the form of molecules:  a neutral group of atoms joined together by covalent bonds.  For example, air contains oxygen molecules, consisting of two oxygen atoms joined covalently  Called a “diatomic molecule”

4. 4 How does H 2 form? l The nuclei repel each other, since they both have a positive charge, and like charges repel. ++

5. 5 How does H 2 form? ++ l But, the nuclei are attracted to the electrons l They share the electrons, and this is called a “covalent bond”, and involves only NONMETALS!

6. 6 Covalent bonds l Nonmetals hold on to their valence electrons. l They can’t give away electrons to bond. l Still want noble gas configuration. l Get it by sharing valence electrons with each other = covalent bonding l By sharing, both atoms get to count the electrons toward a noble gas configuration.

7. 7 Covalent bonding l Fluorine has seven valence electrons F

8. 8 Covalent bonding l Fluorine has seven valence electrons l A second atom also has seven FF

9. 9 Covalent bonding l Fluorine has seven valence electrons l A second atom also has seven l By sharing electrons… FF

10. 10 Covalent bonding l Fluorine has seven valence electrons l A second atom also has seven l By sharing electrons… l …both end with full orbitals FF

11. 11 Covalent bonding l Fluorine has seven valence electrons l A second atom also has seven l By sharing electrons… l …both end with full orbitals FF 8 Valence electrons

12. 12 Covalent bonding l Fluorine has seven valence electrons l A second atom also has seven l By sharing electrons… l …both end with full orbitals FF 8 Valence electrons

13. 13 Molecular Compounds l Compounds that are bonded covalently (like water and carbon dioxide) are called molecular compounds l Molecular compounds tend to have relatively lower melting and boiling points than ionic compounds.

14. 14 Molecular Compounds l Thus, molecular compounds tend to be gases or liquids at room temperature –Ionic compounds were solids l A molecular compound consists of a molecular formula: –Shows how many atoms of each element a molecule contains

15. 15 Molecular Compounds l The formula for water is written as H 2 O –The subscript 2 behind hydrogen means there are 2 atoms of hydrogen; if there is only one atom, the subscript 1 is omitted l Molecular formulas do not tell any information about the structure (the arrangement of the various atoms).

16. 16 A Single Covalent Bond is... l A sharing of two valence electrons. l Only nonmetals and hydrogen. l Different from an ionic bond because they actually form molecules. l Two specific atoms are joined. l In an ionic solid, you can’t tell which atom the electrons moved from or to

17. 17 How to show the formation… l It’s like a jigsaw puzzle. l You put the pieces together to end up with the right formula. l Another example: lets show how water is formed with covalent bonds, by using an electron dot diagram

18. 18 Water H O Each hydrogen has 1 valence electron Each hydrogen wants 1 more The oxygen has 6 valence electrons The oxygen wants 2 more They share to make each other complete

19. 19 Water l Put the pieces together l The first hydrogen is happy l The oxygen still wants one more H O

20. 20 Water l A second hydrogen attaches l Every atom has full energy levels H O H Note the two “unshared” pairs of electrons

21. 21 Multiple Bonds l Sometimes atoms share more than one pair of valence electrons. l A double bond is when atoms share two pairs of electrons (4 total) l A triple bond is when atoms share three pairs of electrons (6 total) l Table 8.1, p.222 - Know these 7 elements as diatomic: Br 2 I 2 N 2 Cl 2 H 2 O 2 F 2 What’s the deal with the oxygen dot diagram?

22. 22 Dot diagram for Carbon dioxide l CO 2 - Carbon is central atom ( more metallic ) l Carbon has 4 valence electrons l Wants 4 more l Oxygen has 6 valence electrons l Wants 2 more O C

23. 23 Carbon dioxide l Attaching 1 oxygen leaves the oxygen 1 short, and the carbon 3 short O C

24. 24 Carbon dioxide l Attaching the second oxygen leaves both oxygen 1 short, and the carbon 2 short O C O

25. 25 Carbon dioxide l The only solution is to share more O C O

26. 26 Carbon dioxide l The only solution is to share more O C O

27. 27 Carbon dioxide l The only solution is to share more O CO

28. 28 Carbon dioxide l The only solution is to share more O CO

29. 29 Carbon dioxide l The only solution is to share more O CO

30. 30 Carbon dioxide l The only solution is to share more l Requires two double bonds l Each atom can count all the electrons in the bond O CO

31. 31 Carbon dioxide l The only solution is to share more l Requires two double bonds l Each atom can count all the electrons in the bond O CO 8 valence electrons

32. 32 Carbon dioxide l The only solution is to share more l Requires two double bonds l Each atom can count all the electrons in the bond O CO 8 valence electrons

33. 33 Carbon dioxide l The only solution is to share more l Requires two double bonds l Each atom can count all the electrons in the bond O CO 8 valence electrons

34. 34 How to draw them? l Add up all the valence electrons. l Count up the total number of electrons to make all atoms happy. l Subtract; then Divide by 2 l Tells you how many bonds to draw l Fill in the rest of the valence electrons to fill atoms up.

35. 35 Example l NH 3, which is ammonia l N – central atom; has 5 valence electrons, wants 8 l H - has 1 (x3) valence electrons, wants 2 (x3) l NH 3 has 5+3 = 8 l NH 3 wants 8+6 = 14 l (14-8)/2= 3 bonds l 4 atoms with 3 bonds N H

36. 36 NHH H Examples l Draw in the bonds; start with singles l All 8 electrons are accounted for l Everything is full – done with this one.

37. 37 Example: HCN l HCN: C is central atom l N - has 5 valence electrons, wants 8 l C - has 4 valence electrons, wants 8 l H - has 1 valence electron, wants 2 l HCN has 5+4+1 = 10 l HCN wants 8+8+2 = 18 l (18-10)/2= 4 bonds l 3 atoms with 4 bonds – this will require multiple bonds - not to H however

38. 38 HCN l Put single bond between each atom l Need to add 2 more bonds l Must go between C and N (Hydrogen is full) NHC

39. 39 HCN l Put in single bonds l Needs 2 more bonds l Must go between C and N, not the H l Uses 8 electrons – need 2 more to equal the 10 it has NHC

40. 40 HCN l Put in single bonds l Need 2 more bonds l Must go between C and N l Uses 8 electrons - 2 more to add l Must go on the N to fill its octet NHC

41. 41 Another way of indicating bonds l Often use a line to indicate a bond l Called a structural formula l Each line is 2 valence electrons HHO = HHO

42. 42 Other Structural Examples H CN C O H H

43. 43 Bond Dissociation Energies... l The total energy required to break the bond between 2 covalently bonded atoms l High dissociation energy usually means the chemical is relatively unreactive, because it takes a lot of energy to break it down.

44. 44 Resonance is... l When more than one valid dot diagram is possible. l Consider the two ways to draw ozone (O 3 ) l Which one is it? Does it go back and forth? l It is a hybrid of both, like a mule; and shown by a double-headed arrow l found in double bond structures

45. 45 Resonance in Ozone Neither structure is correct, it is actually a hybrid of the two. To show it, draw all varieties possible, and join them with a double-headed arrow. Note the different location of the double bond

46. 46 Resonance Occurs when more than one valid Lewis structure can be written for a particular molecule (due to position of double bond) These are resonance structures of benzene. The actual structure is an average (or hybrid) of these structures.

47. 47 Resonance in a carbonate ion: Resonance in an acetate ion: Polyatomic ions – note the different positions of the double bond.

48. 48 Exceptions to Octet rule l For some molecules, it is impossible to satisfy the octet rule –usually when there is an odd number of valence electrons –NO 2 has 17 valence electrons, because the N has 5, and each O contributes 6 l It is impossible to satisfy octet rule, yet the stable molecule does exist

49. 49 Exceptions to Octet rule Another exception: Boron Page 228 shows boron trifluoride, and note that one of the fluorides might be able to make a coordinate covalent bond to fulfill the boron But fluorine has the highest electronegativity of any element, so this coordinate bond does not form Top page 229 examples exist because they are in period 3 or beyond

50. 50 Bond Polarity l Covalent bonding means shared electrons –but, do they share equally? l Electrons are pulled, as in a tug-of- war, between the atoms nuclei –In equal sharing (such as diatomic molecules), the bond that results is called a nonpolar covalent bond

51. 51 Bond Polarity l When two different atoms bond covalently, there is an unequal sharing –the more electronegative atom will have a stronger attraction, and will acquire a slightly negative charge –called a polar covalent bond, or simply polar bond.

52. 52 Electronegativity The ability of an atom in a molecule to attract shared electrons to itself. The ability of an atom in a molecule to attract shared electrons to itself. Linus Pauling 1901 - 1994

53. 53 Table of Electronegativities

54. 54 Bond Polarity l Refer to Table 6.2, page 177 l Consider HCl H = electronegativity of 2.1 Cl = electronegativity of 3.0 –the bond is polar –the chlorine acquires a slight negative charge, and the hydrogen a slight positive charge

55. 55 Bond Polarity l Only partial charges, much less than a true 1+ or 1- as in ionic bond l Written as: HCl l the positive and minus signs (with the lower case delta: ) denote partial charges.    and  

56. 56 Bond Polarity l Can also be shown: –the arrow points to the more electronegative atom. Table 8.3, p.238 shows how the electronegativity can also indicate the type of bond that tends to form HCl

57. 57 Polar molecules l Sample Problem 8.3, p.239 l A polar bond tends to make the entire molecule “polar” –areas of “difference” l HCl has polar bonds, thus is a polar molecule. –A molecule that has two poles is called dipole, like HCl

58. 58 Polar molecules l The effect of polar bonds on the polarity of the entire molecule depends on the molecule shape –carbon dioxide has two polar bonds, and is linear = nonpolar molecule!

59. 59 Polar molecules l The effect of polar bonds on the polarity of the entire molecule depends on the molecule shape –water has two polar bonds and a bent shape; the highly electronegative oxygen pulls the e - away from H = very polar!

60. 60 Polar molecules l When polar molecules are placed between oppositely charged plates, they tend to become oriented with respect to the positive and negative plates. l Figure 8.24, page 239

61. 61 Attractions between molecules l They are what make solid and liquid molecular compounds possible. l The weakest are called van der Waal’s forces - there are two kinds: 1. Dispersion forces weakest of all, caused by motion of e - increases as # e - increases halogens start as gases; bromine is liquid; iodine is solid – all in Group 7A

62. 62 2. Dipole interactions l Occurs when polar molecules are attracted to each other. l 2. Dipole interaction happens in water –Figure 8.25, page 240 –positive region of one molecule attracts the negative region of another molecule.

63. 63 2. Dipole interactions l Occur when polar molecules are attracted to each other. l Slightly stronger than dispersion forces. l Opposites attract, but not completely hooked like in ionic solids. HFHF  HFHF 

64. 64 2. Dipole Interactions    

65. 65 3. Hydrogen bonding? l …is the attractive force caused by hydrogen bonded to N, O, F, or Cl l N, O, F, and Cl are very electronegative, so this is a very strong dipole. l The hydrogen partially share with the lone pair in the molecule next to it. l This is the strongest of the intermolecular forces.

66. 66 3. Hydrogen bonding defined: l When a hydrogen atom is: a) covalently bonded to a highly electronegative atom, AND b) is also weakly bonded to an unshared electron pair of a nearby highly electronegative atom. –The hydrogen is left very electron deficient, thus it shares with something nearby –Hydrogen is also the ONLY element with no shielding for its nucleus when involved in a covalent bond!

67. 67 Hydrogen Bonding H H O ++ -- ++ H H O ++ -- ++

68. 68 Hydrogen bonding H H O H H O H H O H H O H H O H H O H H O

69. 69 Attractions and properties l Why are some chemicals gases, some liquids, some solids? –Depends on the type of bonding! –Table 8.4, page 244 l Network solids – solids in which all the atoms are covalently bonded to each other

70. 70 Attractions and properties l Figure 8.28, page 243 l melts at very high temperatures, or not at all –Diamond does not really melt, but vaporizes to a gas at 3500 o C and beyond –SiC, used in grinding, has a melting point of about 2700 o C

71. 71 Covalent Network Compounds Some covalently bonded substances DO NOT form discrete molecules. Diamond, a network of covalently bonded carbon atoms Graphite, a network of covalently bonded carbon atoms