Exam T Chapters 7 and 8 Review M 5 SL 110

1. Exam T Chapters 7 and 8 Review M 5 SL 110

2. Periodicity of First Ionization Energy (IE1) Like Figure 8-18

3. Fig. 8.15

5. Identifying Elements by Its Successive Ionization Energies Problem: Given the following series of ionization energies (in kJ/mol) for an element in period 3, name the element and write its electron configuration: IE1 IE2 IE3 IE4 580 1,815 2,740 11,600 Plan: Examine the values to find the largest jump in ionization energy, which occurs after all valence electrons have been removed. Use the periodic table! Solution:

6. Identifying Elements by Its Successive Ionization Energies Problem: Given the following series of ionization energies (in kJ/mol) for an element in period 3, name the element and write its electron configuration: IE1 IE2 IE3 IE4 580 1,815 2,740 11,600 Plan: Examine the values to find the largest jump in ionization energy, which occurs after all valence electrons have been removed. Use the periodic table! Solution: The largest jump in IE occurs after IE3 so the element has 3 valence electrons thus it is Aluminum ( Al, Z=13), its electron configuration is : 1s2 2s2 2p6 3s2 3p1

7. Fig. 8.16

8. Ranking Elements by First Ionization Energy Problem: Using the Periodic table only, rank the following elements in each of the following sets in order of increasingIE! a) Ar, Ne, Rn b) At, Bi, Po c) Be, Na, Mg d) Cl, K, Ar Plan: Find their relative positions in the periodic table and apply trends! Solution:

9. Ranking Elements by First Ionization Energy Problem: Using the Periodic table only, rank the following elements in each of the following sets in order of increasingIE! a) Ar, Ne, Rn b) At, Bi, Po c) Be, Na, Mg d) Cl, K, Ar Plan: Find their relative positions in the periodic table and apply trends! Solution: a) Rn, Ar,Ne These elements are all noble gases and their IE decreases as you go down the group.

10. Suggested problems for Chapter 9 19, 20, 21, 22, 23, 25, 27, 29, 35, 37, 39, 43, 47, 49, 51, 53, 55, 56, 60, 62, 65, 79, 85, 93, 95, 96, 97 103, 105, 107

11. Chapter #9 - Models of Chemical Bonding 9.1) Atomic Properties and Chemical Bonds 9.2) The Ionic Bonding Model 9.3) The Covalent Bonding Model 9.4) Between the Extremes: Electronegativity and Bond Polarity 9.5) An Introduction to Metallic Bonding

13. Sodium Chloride

14. Depicting Ion Formation with Orbital Diagrams and Electron Dot Symbols - I Problem: Use orbital diagrams and Lewis structures to show the formation of magnesium and chloride ions from the atoms, and determine the formula of the compound. Plan: Draw the orbital diagrams for Mg and Cl. To reach filled outer levels Mg loses 2 electrons, and Cl will gain 1 electron. Therefore we need two Cl atoms for every Mg atom. Solution: Mg + Mg+2 + 2 Cl- .. 2 Cl . .. .. .. . Cl Cl . .. .. .. .. Mg + Mg+2 + 2 Cl . .. ..

15. Depicting Ion Formation from Orbital Diagrams and Electron Dot Symbols - II Problem: Use Lewis structures and orbital diagrams to show the formation of potassium and sulfide ions from the atoms, and determine the formula of the compound. Plan: Draw orbital diagrams for K and S. To reach filled outer orbitals, sulfur must gain two electrons, and potassium must lose one electron. Solution: 2 K + 2 K+ + S - 2 S . .. . 2 - . .. .. .. .. .. K . + S 2 K+ + S K

16. Three Ways of Showing the Formation ofLi+ and F - through Electron Transfer

17. Lewis Electron-Dot Symbols for Elements in Periods 2 & 3

18. The Reaction between Na and Br to Form NaBr The Elements The Reaction!

19. Vaporizing an Ionic Compound

20. Melting and Boiling Points of Some Ionic Compounds Compound mp( oC) bp( oC) CsBr 636 1300 NaI 661 1304 MgCl2 714 1412 KBr 734 1435 CaCl2 782 >1600 NaCl 801 1413 LiF 845 1676 KF 858 1505 MgO 2852 3600

22. Figure 9.11: Potential-energy curve for H2.

23. Covalent Bonding in Hydrogen, H2

25. Figure 9.10: The electron probability distribution for the H2 molecule.

26. Covalent bonds http://wine1.sb.fsu.edu/chm1045/notes/Bonding/Covalent/Bond04.htm animation http://www.chem.ox.ac.uk/vrchemistry/electronsandbonds/intro1.htm

27. For elements larger than Boron, atoms usually react to develop octets by sharing electrons. H, Li and Be strive to “look” like He. B is an exception to the noble gas paradigm. It’s happy surrounded by 6 electrons so the compound BH3 is stable. Try drawing a Lewis structure for methane.

28. Draw Lewis dot structures for the halogens. Notice that these all follow the octet rule! Try oxygen and nitrogen. These also follow the octet rule!

29. Bond Lengths and Covalent Radius

30. Figure 9.14: The HCl molecule.

31. The Charge Density of LiF Fig. 9.20

32. Figure 9.12: Molecular model of nitro-glycerin. What is the formula for this compound?

33. Rules for drawing Lewis structures 1. Count up all the valence electrons 2. Arrange the atoms in a skeleton 3. Have all atoms develop octets (except those around He)

34. Make some Lewis Dot Structures with other elements: SiH4 H2O NH3 CH2O C2H6 C2H6O

35. CH3I Figure 9.9: Model of CHI3Courtesy of Frank Cox.

36. Make some Lewis Dot Structures with other elements: CH4 H2O NH3 CH2O C2H6 C2H6O

37. Look at all these structures and make some bonding rules:

39. Rules for drawing Lewis structures 1. Count up all the valence electrons 2. Arrange the atoms in a skeleton 3. Have all atoms develop octets (except those around He) 4. Satisfy bonding preferences!

40. A model of ethylene.

41. A model of acetylene.

42. Figure 9.14: The HCl molecule.

43. A model of COCl2.