Bonding: Part Two. Three types of bonds: Ionic Bond. transfer valence e - Metallic bond. (NaCl) (Fe) mobile valence e - Covalent bond

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1 Bonding: Part Two Three types of bonds: Ionic Bond transfer valence e - Metallic bond mobile valence e - Covalent bond (NaCl) (Fe) shared valence e - (H 2 O) 1

2 Single Covalent Bond H + H H H H-atoms H 2 molecule Electrons are shared by the two H atoms 2

3 Single Covalent Bond Structural Formula H H Single covalent bond (2 shared e - ) 3

4 Ionic vs. Covalent Compounds Covalent: discrete molecules H O, CH, CO Ionic: cations and anions (occupy crystal lattice points) formula unit NaCl, MgI 2 4

5 Covalent Bond Covalent bonds usually form between nonmetal atoms (Groups 14, 15, 16, and 17). Each atom tries to attain the e - configuration of a noble gas by sharing electrons. ( octet rule ) 5

6 Back to H 2 H + H H H H-atoms Each H has e - config. of He 6

7 Fluorine: F 2 or F F F + F F F Each F has e - config of [Ne] 7

8 Lone Pairs Paired valence e - not in the bond F F Each F has an octet lone pairs or nonbonding pairs of e - 8

9 Water (H 2 O) O H H O H H Each atom has Noble Gas configuration 9

10 You Try It! Draw dot structures for: 1.Ammonia (NH 3 ) 2.Chlorine gas (Cl 2 ) 3.Methane (CH 4 ) 10

11 Multiple Bonds Sometimes more than one pair of bonding electrons are needed in a bond to attain a noble gas configuration. Double bond: two pairs of e - Triple bond: three pairs of e - 11

12 Double Bond: O 2 O + O O O Still no octet, so form a double bond O O or O O 12

13 Triple Bond Try nitrogen (N 2 ) N N N N 13

14 Coordinate Covalent Bond Sometimes both e - in the bond come from just one of the atoms. 14

15 Coordinate Covalent Bond e.g. CO no octet O + C O C e - pair came from oxygen O C 15

16 Coordinate Covalent Bond The ammonium ion, NH 4 + H H + H + + N H H N H H H no e - 16

17 Dot Structure Rules e.g. NF 3 1.Arrange the atoms with least electronegative element in center. The central atom is never hydrogen. F N F F 17

18 Dot Structure Rules 2.Count total valence electrons. Account for charges in polyatomic ions. F N F 5 + 3(7) = 26 F 18

19 Dot Structure Rules 3.Connect atoms with single covalent bonds. Then complete octets (H has only 2 not 8). F N F F 19

20 Dot Structure Rules 4.If octet rule is not satisfied for the central atom, try double or triple bonds, using lone pairs from surrounding atoms. F N F OK F 20

21 Dot Structure HNO 3 Step 1. Skeletal structure O N O H O 21

22 HNO 3 Step 2. Number of valence e - O N O H O N, O, H 5 + 3(6) + 1 = 24 22

23 HNO 3 Step 3. Add bonds and complete octets O N O H O Out of e -, but no octet for N 23

24 HNO 3 Step 4. Add multiple bonds O N O H O 24

25 You Try It!!! Draw the e- dot structures for: 1.Hydrogen chloride HCl 2.Hydrogen peroxide H 2 O 2 3.Hydronium ion H 3 O + 4.Ozone O 3 25

26 Lewis Structures 26

27 Resonance: Ozone O 3 O O O O O O Each is called a resonance structure. The bonds are equal (~1.5 bond) 27

28 Resonance: Try It!!! Draw resonance structures for SO 2 28

29 Exception to the Octet Rule: Odd # e - Try to write the dot structure of nitrogen monoxide. With an odd number of valence electrons (11), it is impossible to have octets around both atoms. NO is paramagnetic. 29

30 Predicting Molecular Shapes linear triatomic trigonal planar bent triatomic trigonal pyramid tetrahedral others 30

31 VSEPR Theory Valence-shell e - pair repulsion All valence electron pairs (bonding & nonbonding pairs) repel each other. Predicts geometry. 31

32 VSEPR Theory Methane (CH 4 ) is drawn as: H H H C H or H C H H H Actually CH 4 is 3-D 32

33 VSEPR Theory Maximum repulsion of e - pairs C o tetrahedron 33

34 VSEPR Theory Ammonia (NH 3 ) 1 nonbonding e- pair H N H H 3 bonding e- pairs 34

35 VSEPR Theory Ammonia tetrahedral electrons trigonal pyramidal atoms 35

36 Water (H 2 O) VSEPR Theory 2 nonbonding e- pairs H O H 2 bonding e- pairs 36

37 VSEPR Theory Water e - : tetrahedral atoms: bent linear 37

38 VSEPR: Rules 1.Draw the Lewis dot structure. 2.Move e - pairs (bonding and lone pairs) as far apart as possible. 3.Treat double and triple bonds as if they were single bonds. 4.Distinguish between shape of e - pairs and molecular shape 38

39 Possible Molecular Shapes linear bent tetrahedron trigonal pyramid trigonal planar See problem set. 39

40 Predict the Molecular Shape Hydrogen sulfide: H 2 S Carbon tetrachloride: CCl 4 Sulfur dioxide: SO 2 Sulfur trioxide: SO 3 Nitrogen tribromide: NBr 3 40

41 Is Breaking a Bond Endo- or Exothermic? Endothermic! It takes energy to break a bond. atoms Energy is given off when bonds form (exothermic). molecule 41

42 Bond Energy The energy needed to break a bond is called bond dissociation energy or bond energy. H H H + H DH = +435 kj (per mol) 42

43 Bond Dissociation Energy Bond Bond Energy (kj) Bond Length (pm) H H C H C C C C C C

44 Bond Strength vs. Length Multiple bonds are stronger than single bonds. Multiple bonds have shorter bond lengths! Why? C C C C C C weakest, longest strongest, shortest 44

45 Bond Polarity Covalent bonds involve sharing e -, however the two bonded atoms don t always share equally. 45

46 Bond Polarity How do you know which atom wins the tug-of war for the bonding electrons? The more electronegative one. Regents Table S 46

47 Bond Polarity In some cases neither atom wins. Both atoms have same electronegativity. Hydrogen Nitrogen Oxygen Nonpolar Chlorine 47

48 Polar Molecules EN = 2.2 EN = 4.0 H F d+ d- Bonding electrons shift toward F. Thus HF is polar. 48

49 Water EN = 3.4 O H H EN = 2.2 Each bond in water is polar, and the overall molecule is polar because of its shape (nonsymmetrical). Distinguish between bond polarity and molecular polarity. 49

50 Nonpolar Molecules CO 2 O C O Where is the average center of positive charge, and where is the center of negative charge? Even though CO 2 has polar bonds, it is a nonpolar molecule. 50

51 Summary: Polar Molecules If only 2 atoms in the molecule: the molecule is polar if the atoms have different electronegativities. H-Cl polar Br-Br nonpolar 51

52 Summary: Polar Molecules If more than 2 atoms in molecule: Draw dot structure If the central atom has a lone pair or if the outside atoms are different, the molecule is polar (not symmetrical). 52

53 Summary: Polar Molecules CH 4 is nonpolar (symmetrical) CH 2 F 2 is polar (nonsymmetrical) NH 3 is polar (nonsymmetrical) CH 4 CH 2 F 2 NH 3 53

54 Polar Molecule: Yes or No? Symmetrical nonpolar Nonsymmetrical polar CHCl 3 HI NI 3 Br 2 SO 3 CI 4 54

55 Bond Polarity & Bond Type If the difference in electronegativity is greater than ~2.0, one atom takes all the bonding e - and the compound is ionic. What types of elements would have a big difference in EN? 55

56 Difference in Electronegativity covalent ionic more ionic more covalent H 2 HCl NaCl CsF nonpolar polar 56

57 Identify Bond Types nonpolar covalent polar covalent ionic for these pairs of atoms: Cs & F P & O H & Br 57

58 Ionic vs. Covalent Compounds Ionic Covalent Unit formula unit molecule Bond e - transfer e - sharing Elements metal + nonmetals nonmetal D EN > 2.0 < 2.0 solid S, L, G M. P. high low 58

59 59

60 Warm-up Write the e - dot symbol of calcium chloride. Define the metallic bond. How does this bonding explain why metals conduct electricity? Draw dot structure for methane, CH 4. 60

61 Warm-up What is the Lewis dot structure of PH 3? 61

62 Warm-up: Write dot structure: ammonium ion nitrate ion 62

63 Warm-up Draw the e - dot structure for NO 2 and draw its resonance structures. 63

64 Warm-up Predict the shape of: SeO 3 CS 2 PCl 3 NO 2 64

65 Warm-up What is the molecular shape of: NBr 3 SCl 2 CO

66 Warm-up Write the dot formula for the bromine atom and the bromide ion. What is the shape of the nitrate ion? 66

67 Warm-up Draw the correct dot structure and state whether the molecule is polar or nonpolar? SO 2 PBr 3 CH 2 Br 2 67

68 Warm-up Draw a diagram showing at least 12 atoms of metallic potassium in the solid state. Write the dot formula for sodium oxide. 68