Lewis Structures of Simple Molecules

1. H H C H H .. .. F .. .. .. F C F .. .. .. .. F Lewis Structures of Simple Molecules H H .. CH4 .. H C C O H Methane H H Ethyl Alcohol (Ethanol) .. .. .. .. O K+ .. Cl .. .. .. .. .. .. O O .. CF4 .. KClO3 Potassium Chlorate Carbon Tetrafluoride

2. Resonance: Delocalized Electron-Pair Bonding - I Ozone : O3 .. .. .. .. .. O O .. .. .. .. .. .. O O O O .. II I Resonance Hybrid Structure .. O .. .. .. .. O O One pair of electron’s resonances between the two locations!!

3. Resonance: Delocalized Electron-Pair Bonding - II H H C C H H C C C C H H C C H H C C H H H C C C H H C H H C H C C H C Benzene Resonance Structure H

4. .. .. Cl .. .. B .. .. .. .. Cl Cl Lewis Structures for Octet Rule Exceptions .. .. .. .. .. .. F .. .. .. .. F Cl .. .. F Each chlorine atom has 8 electrons associated. Boron has only 6! Each fluorine atom has 8 electrons associated. Chlorine has 10 electrons! . .. .. .. .. .. .. N .. .. .. .. Cl Be Cl .. O O Each chlorine atom has 8 electrons associated. The beryllium has only 4 electrons. NO2 is an odd electron atom. The nitrogen has 7 electrons.

5. Resonance Structures - Expanded Valence Shells .. .. O O .. .. .. .. H O S O H H O S O H .. .. O O .. .. .. .. .. .. .. .. .. .. .. .. .. .. F F F F .. .. .. .. .. .. .. F S F P .. .. F .. F .. .. .. .. .. .. .. .. F F .. F p = 10e- S = 12e- Sulfur hexafluoride Phosphorous pentafluoride .. .. .. .. .. Resonance Structures .. .. .. .. .. Sulfuric acid S = 12e-

6. Lewis Structures of Simple Molecules . . Sulfate . . -2 O . . . . Resonance Structures-V . . . . O S O . . . . -2 . . . . O o * o o o o O Plus 4 others for a total of 6 x o o o o o x O S O o * x o o o . . x o o -2 o o . . . . O x x . . . . o o O o o O S O o o . . . . . . . . O . . x = Sulfur electrons o = Oxygen electrons

7. Fig. 9.14

8. Figure 9.15: Electronegatives of the elements.

9. The Periodic Table of the Elements 2.1 He 0.9 1.5 2.0 2.5 3.0 3.5 4.0 Ne Electronegativity 0.9 1.2 1.5 1.8 2.1 2.5 3.0 Ar 0.8 1.0 1.3 1.5 1.6 1.6 1.5 1.8 1.8 1.8 1.9 1.6 1.6 1.8 2.0 2.4 2.8 Kr 0.8 1.0 1.2 1.4 1.6 1.8 1.9 2.2 2.2 2.2 1.9 1.7 1.7 1.8 1.9 2.1 2.5 Xe 0.7 0.9 1.1 1.3 1.5 1.7 1.9 2.2 2.2 2.2 2.4 1.9 1.8 1.8 1.9 2.0 2.2 Rn 0.7 0.9 1.1 Ce Pr Nd Pm Yb Lu 1.1 1.1 1.1 1.2 1.2 1.1 1.2 1.2 1.2 1.2 1.2 1.2 1.2 1.3 1.3 1.5 1.7 1.3 1.3 1.3 1.3 1.3 1.3 1.3 1.3 1.3 1.5 Th Pa U Np No Lr

10. Fig. 9.16

11. Fig. 9.17

12. Determining Bond Polarity from Electronegativity Values Problem: (a)Indicate the polarity of the following bonds with a polarity arrow: O - H, O - Cl, C - N, P - N, N - S, C - Br, As - S (b) rank those bonds in order of increasing polarity. Plan: (a) We use Fig. 9.16 to find the EN values, and point the arrow toward the negative end. (b) Use the EN values. Solution: a) the EN of O = 3.5 and of H = 2.1: O - H the EN of O = 3.5 and of Cl = 3.0: O - Cl the EN of C = 2.5 and of P = 2.1: C - P the EN of P = 2.1 and of N = 3.0: P - N the EN of N = 3.0 and of S = 2.1: N - S the EN of C = 2.5 and of Br = 2.8: C - Br the EN of As = 2.0 and of O = 3.5: As - O b) C - Br < C - P < O - Cl < P - N < N - S < O - H < As - O 0.3 < 0.4 < 0.5 < 0.9 < 0.9 < 1.4 < 1.5

13. Fig. 9.18

14. Percent Ionic Character as a Function ofElectronegativity Difference (En) Fig. 9.19

15. The Charge Density of LiF Fig. 9.20

16. Figure 9.15: Electronegatives of the elements.

17. The Periodic Table of the Elements 2.1 He 0.9 1.5 2.0 2.5 3.0 3.5 4.0 Ne Electronegativity 0.9 1.2 1.5 1.8 2.1 2.5 3.0 Ar 0.8 1.0 1.3 1.5 1.6 1.6 1.5 1.8 1.8 1.8 1.9 1.6 1.6 1.8 2.0 2.4 2.8 Kr 0.8 1.0 1.2 1.4 1.6 1.8 1.9 2.2 2.2 2.2 1.9 1.7 1.7 1.8 1.9 2.1 2.5 Xe 0.7 0.9 1.1 1.3 1.5 1.7 1.9 2.2 2.2 2.2 2.4 1.9 1.8 1.8 1.9 2.0 2.2 Rn 0.7 0.9 1.1 Ce Pr Nd Pm Yb Lu 1.1 1.1 1.1 1.2 1.2 1.1 1.2 1.2 1.2 1.2 1.2 1.2 1.2 1.3 1.3 1.5 1.7 1.3 1.3 1.3 1.3 1.3 1.3 1.3 1.3 1.3 1.5 Th Pa U Np No Lr

18. Fig. 9.16

19. Fig. 9.17

20. Determining Bond Polarity from Electronegativity Values Problem: (a)Indicate the polarity of the following bonds with a polarity arrow: O - H, O - Cl, C - N, P - N, N - S, C - Br, As - S (b) rank those bonds in order of increasing polarity. Plan: (a) We use Fig. 9.16 to find the EN values, and point the arrow toward the negative end. (b) Use the EN values. Solution: a) the EN of O = 3.5 and of H = 2.1: O - H the EN of O = 3.5 and of Cl = 3.0: O - Cl the EN of C = 2.5 and of P = 2.1: C - P the EN of P = 2.1 and of N = 3.0: P - N the EN of N = 3.0 and of S = 2.1: N - S the EN of C = 2.5 and of Br = 2.8: C - Br the EN of As = 2.0 and of O = 3.5: As - O b) C - Br < C - P < O - Cl < P - N < N - S < O - H < As - O 0.3 < 0.4 < 0.5 < 0.9 < 0.9 < 1.4 < 1.5

21. Fig. 9.18

22. Percent Ionic Character as a Function ofElectronegativity Difference (En) Fig. 9.19

27. The Charge Density of LiF Fig. 9.20

28. VSEPR: Valence Shell Electron Pair Repulsion: A way to predict the shapes of molecules Pairs of valence electrons want to get as far away from each other as possible in 3-dimensional space.

29. Balloon Analogy for the MutualRepulsion of Electron Groups Two Three Four Five Six Number of Electron Groups

32. AX2 Geometry - Linear .. .. .. .. Molecular Geometry = Linear Arrangement .. .. Cl Be Cl BeCl2 1800 Gaseous beryllium chloride is an example of a molecule in which the central atom - Be does not have an octet of electrons, and is electron deficient. Other alkaline earth elements also have the same valence electron configuration, and the same geometry for molecules of this type. Therefore this geometry is common to group II elements. .. .. .. .. O C O CO2 1800 Carbon dioxide also has the same geometry, and is a linear molecule, but in this case, the bonds between the carbon and oxygens are double bonds.

33. The Two Molecular Shapes of the Trigonal Planar Electron-Group Arrangement

34. S O O AX3 Geometry - Trigonal Planar .. .. .. .. All of the boron Family(IIIA) elements have the same geometry. Trigonal Planar ! .. .. F F BF3 B Boron Trifluoride 1200 .. .. .. F AX2E SO2 .. - .. .. .. .. O .. .. NO3- 1200 .. .. N .. .. .. .. O O The AX2E molecules have a pair of Electrons where the third atom would appear in the space around the central atom, in the trigonal planar geometry. 1200 Nitrate Anion

35. The Three Molecular Shapes of the Tetrahedral Electron-Group Arrangement

36. 107.30 .. N N H H H AX4 Geometry - Tetrahedral H Methane H 109.50 CH4 H C H C H H H H All molecules or ions with four electron groups around a central atom adopt the tetrahedral arrangement H H 109.50 109.50 + H+ H Ammonia is in a tetrahedral shape, but it has only an electron pair in one location, so the smaller angle! all angles are the same! Ammonium Ion H

37. The Four Molecular Shapes of the Trigonal Bipyramidal Electron- Group Arrangement

38. AX5 Geometry - Trigonal Bipyramidal .. .. .. .. .. .. F I .. .. .. 86.20 .. .. .. I 1800 Br .. F .. .. .. .. .. I .. F .. AX3E2 - BrF3 .. AX2E3 - I3- .. .. .. Cl .. .. Cl .. .. P Cl .. AX5 - PCl5 .. .. .. .. .. Cl .. Cl

39. The Three Molecular Shapes of the Octahedral Electron-Group Arrangement

40. AX6 Geometry - Octahedral .. .. .. .. .. .. .. .. .. .. F .. F .. .. .. F .. F .. .. .. F F .. .. .. .. S I .. .. .. .. .. .. F F .. .. F F .. .. .. .. F AX5E Iodine Pentafluoride AX6 Sulfur Hexafluoride .. .. .. .. .. .. F .. F .. Xe .. .. .. .. .. F .. F Xenon Tetrafluoride Square planar shape

41. Using VSEPR Theory to Determine Molecular Shape 1) Write the Lewis structure from the molecular formula to see the relative placement of atoms and the number of electron groups. 2) Assign an electron-group arrangement by counting all electron groups around the central atom, bonding plus nonbonding. 3) Predict the ideal bond angle from the electron-group arrangement and the direction of any deviation caused by the lone pairs or double bonds. 4) Draw and name the molecular shape by counting bonding groups and non-bonding groups separately.

43. Hybrid Orbital Model

45. The sp Hybrid Orbitals in Gaseous BeCl2

49. The sp3 Hybrid Orbitals in NH3 and H2O

50. The sp3d Hybrid Orbitals in PCl5