Introduction:

Unit 7 will discuss the covalent, electron pair bonding theory presented in the 1920's by G.N. Lewis. The three dimensional character of molecules will be discussed from the perspective of the Valence Shell Electron Pair Repulsion Theory, (VSEPR ). The combination of these two ideas, "Lewis Structures" and the prediction of geometry around a "central atom" provides the sufficient background to understand most of structural, organic chemistry involving the atoms H, C, N and O. Because of its numerous applications, Chapter Nine is one of the most important chapters in the text and workbook.

Before you begin this unit, you should learn how to use the "Molecular Models" part of the CD-ROM. Viewing molecules in this software program is necessary for this unit and the next one. Instructions for using the modeling applications is given in the Workbook, Appendix A, "CAChe Visualizer for Education" pages, A1 to A6 for "Mac" users and pages, A7 to A10 is the " Windows Version". Models for most of the molecules in Unit 7 are found in the Model Folder under the file name, INORGANIC and are listed by formula. This modeling portion of the ROM is really quite good, unfortunately some of the applications are beyond the level of an introductory course. You should practice "rotating" and moving the molecular model. Learn how to change the view from "stick" to "space-filling" and how to measure bond distances.

Unit Objectives: After completing Unit 7, you should be able to:

1. discuss the difference between ionic interactions and covalent bonding; and

2. give the Lewis Dot Structures for hydrogen and the "non-metals" group on the periodic table; and

3. explain the "octet rule" and use it to construct Lewis structures for the elements H, C, N, O and F which always obey the rule; and

4. define the term, "resonance" and draw appropriate resonance structures; and

5. predict trends in bond order, bond length and bond energy; and

6. use the electronegativity table to discuss bond and molecular polarity; and

7. draw structures using elements that don't obey the octet rule; and

8. predict geometry around a "central atom" using VSEPR theory.

Reading and Study Assignment:

Text: Chapter 9, "Bonding and Molecular Structure: Fundamental Concepts", pages 398 to 427 and pages 430 to 449. Omit the sub-section on "formal charge".

CD-ROM: Chapter 9. start with Screen 9-5 and study through Screen 9-17. Omit Screens 9.8, Free Radicals, Screen 9.10, Bond Energies and "Delta" H Reaction, and Screen 9.13, Formal Charge.

Workbook Notes:

Screen 9-5 "Drawing Lewis Electron Dot Structures" is the central CD screen for Unit 7. The basic scheme for constructing Lewis structures is presented. The "rules" for octet systems are developed here on the ROM and on page 405 of the text. Practice constructing Lewis structures until you are completely aware how to do this activity. For example, try drawing the Lewis structures of CH4, PF3, CO2 and actetate ion, CH3CO2 (-) then look at the instructions for these structures shown on this screen. "Drawing" is a visual activity and somewhat hard to do in this "computer" instruction mode. I will make sketches and written comments about some of the workbook activities and "post" them in the "answer book" for this unit. See "answer book" for information on problems 1 and 2 for Screen 9.5

Screen 9-6 "Resonance Structures" On page 408 of the text, ozone, O(3) is used as an example of a compound with resonance structures. The concept of resonance arises because Lewis Bonding Theory is not complete and can not explain certain aspects of covalent compounds. In cases where more that one equivalent Lewis structure can be constructed, that compound such as O(3) is said to show "resonance": forms. Another "textbook" example of s structure with resonance forms is carbonate anion, CO3 (2-) shown at the top of page 409. Check the answer book for my drawings of the ions listed in the workbook for Screen 9-6. Do Exercise 9.1 and check in the back of the text or workbook. Carbonate is shown on page 409 in the text as indicated, previously.

Screen 9-7 In some sense, the only interesting example of a "electron deficient" compound in boron trifluoride shown on this screen, You should be able to draw and predict the geometry of BF3 No problems assigned here.

Omit Screen 9-8

Screen 9.9 "Bond Properties" A general rule presented here is a triple bond is shorter and "stronger" than a double bond and the "longest" and "weakest" bond between two atoms is a single bond. Answers for problem 1 and comments about 2 , 3 and 4 are in the answer book. You need to practice using the "Molecular Models" folder for these problems. I hope it works on your computer.

Omit Screen 9-10

Screen 9-11 "Bond Polarity and Electronegativity" See Figure 9.7, page 423, text. You should memorize the values for H, C, N, O, F, Cl, Br and S These are used all the time in organic examples. If you need values for something else, just look them up on a table. Do Exercise 9.3 which is the same as (.9 in the text.

Screen 9-12 "Oxidation Numbers" This is a "review" of oxidation numbers from Unit 2 but gives you a sense of where the concept of an oxidation number originated. See drawing in answer book for problems 1 and 2.

Omit Screen 9-13

Screen 9-14, 9-15 and 9-16 These three screens and related activities show how molecular geometry can be determined using VSEPR.The Lewis Structures indicated how electron pairs are distributed in a given compound as bonding or lone pairs but suggests nothing about the actual "shape' of a molecule. VSEPR is the first of two methods for predicting the shape of molecules. Unit 8. introduces "Valence Bond Theory" which can predict the shape of molecules, also. Check the geometry of various shapes for central atom cases shown in the text on pages 436 and 438. Answers to problem 3 a) and b) are in the answer book. You need to run the "Molecular Models" program for this part. Do Exercise 9.4 and check your answer against the drawing in the answer book or the Appendix of the "workbook".

Screen 9-17 Once you know the "shape" of a particular molecule, you can think about the question of whether or not the molecule is "polar"? Molecular polarity is one of the more difficult aspects of molecular structure for beginning students. I think the simple idea to keep in mind is "symmetry" If the overall molecular structure has "symmetry" it probably is not polar. It is more important for you to know how to draw correct Lewis Structures and make predictions about their geometry than polarity. Decisions about the polar character of molecules comes with experience.

e-mail Task (15)

a) I was able to run the "molecular models" application with this unit. Y or N ? If "Y" was it useful ? Brief comment.

b) What is the relationship between bond order, bond length and bond energy for a series of related bonds such as a C-N, (a single CN bond), C=N (a double CN bond) and a triple CN bond ?

Study Questions:

Workbook (Text), Chapter 9 - 1 (26), 2 (30), 3 (32), 4 (34), 5 (38), 6 (40), 7 (42), 8 (44), 11 (54), 13 (64), 14 (66), 16 (78), 17 (80), 18 (84) and 19 (86).

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