Chapter 9 Reading Guide Bonding I Dr. Baxley Tro 4 th edition 1 See end of this document for suggested (strongly urged) problems. Section 9.1, Why is chemical bonding important? What disease has been tamed to some degree through the help of chemical bonding theory? Section 9.2, What types of bonding exist? What are the three main types of chemical bonds? How are electrons distributed in these types of bonds? Study Figure 9.1. How are the arrangements of atoms/ions in the ionic compound different from the molecular compound with covalent bonding? Section 9.3, What is a group of eight called? Drawing electron diagrams with valence electrons for atoms is an important first step. Can you draw the dot diagram for Li-Ne? What is an octet? NOTE: the octet rule is not a rule, it s more of a guideline (like the pirate s code). In fact, I prefer to call it the octet guideline, not the octet rule. Section 9.4, How does lattice energy explain the energetics of formation of NaCl? This section really gets at the energetics of bonding for ionic compounds. Review how to write electron dot diagrams for ionic compounds, including the square brackets, from chapter 8 if necessary. Note how the diagram for CaCl 2 in Example 9.1 shows 2 chloride ions. Does this help show that CaCl 2 does not consist of Cl 2 2? What is the definition for lattice energy? Here is a video of the reaction discussed in the text: https://www.youtube.com/watch?v=ozdqji-uwys the good stuff starts at 3:45, but make sure you look at the difference in the color of the gas in the flask between t = 0 and 3:45. Figure 9.4 shows a Born-Haber cycle involving 5 different steps from Na(s) and Cl 2(g) to NaCl(s). The important part of examining this figure is the energetics of each step. Ask yourself, which processes absorb energy and which processes release energy? Which step releases the most energy? Note that the author of the text does discuss sodium wanting to lose electrons. What does he have to say about it? Does his note fit my rant in the chapter 8 reading guide?
Rant Rewind: Atoms DON T want anything! 2 Want to watch a Born Haber cycle conceptual video, with comments from me? https://edpuzzle.com/media/57b5ff046b6e2cb616d07ed9 Trends in Lattice Energy: ion size. How and why does the amount of lattice energy change from LiCl to CsCl? Trends in Lattice Energy: ion charge. How and why does the amount of lattice energy change from NaF to CaO? Example 9.2 and the related For Practice should wrap up these concepts and trends in lattice energy. There are some nice examples of medically important ionic compounds in the Chemistry and Medicine box. Section 9.5, What are Lewis structures? The first few pages of the section are a great introduction to the nature of covalent bonds, and how sharing of electrons leads to single, double, and triple bonds. Figure 9.5 displays a critical concept. Conceptual Connection 9.2 is a great question. Section 9.6, Are electrons shared equally in covalent bonds? So, are electrons shared equally in covalent bonds? What symbol is used to show the partial positive and partial negative charges? What is a polar covalent bond? Electronegativity is a critical concept in chapter 9. What does electronegativity mean? Examine Figure 9.8. Which elements have the highest electronegativity values? What are the trends in electronegativity from left to right, and from top to bottom, on the periodic table? Bond Polarity: How do the electronegativities of elements impact the polarity of bonds between two atoms. Table 9.1 gives a rough estimate of bond types and the differences in electronegativity, and Figure 9.9 shows the continuum of bond types. IGNORE the dipole moment calculations and percent ionic character.
3 However, knowing that dipole moments are a way to measure bond polarity is important (see Table 9.2). Example 9.3 and For Practice 9.3 is helpful for determining bond characteristics. Section 9.7, How do I draw Lewis structures? This is the critical skill in chapter 9, and will be used in chapters 10, 11, and 12 also. I posted an interactive video tutorial for Lewis structures, and there is also a video in MC in the study area for Chapter 9. The interactive video: https://edpuzzle.com/media/57b5ff046b6e2cb616d07ed8 You have known me for 7 chapters of gen chem over about 10 weeks of class. I don t make students follow many rules, especially for solving problems. BUT, for Lewis structures, especially at the beginning, FOLLOW THE RULES! You can read these rules in the text, but MY RULES are better. My rules are GUARANTEED to give you better results, 100% of the time. (almost). 1. Add up all valence e from the atoms involved, including the impact of the charge, if an ion 2. Arrange atoms as bonded a. Look at formula, (many around 1) b. What s bonded to what (H O in oxo acids) c. typical # of bonds 3. Place 2 e between each bonded atom 4. Add in e pairs to make octets for outer atoms (not H) 5. Place extra e in pairs around central atom 6. If no octet on central atom, share e to make multiple bonds 7. If all else fails, LOOK UP STRUCTURE 8. Large atoms more likely to violate octet rule (3d orbitals for S, P, Cl) Additional tips: The typical number of bonds for an atom: 8 group number = # of bonds (usually, N can make 4, O can make 3) Carbon atoms are usually bonded together in a chain (C C C, not C H C) Carbon atoms ALWAYS have a share of 8 (eight) electrons in normal, stable structures. Don t make too many double bonds (never for H, B, or F) You will get loads of practice with Lewis structures in Experiment 11 and 12. I do not assign Lewis structures in Mastering Chemistry because the structure drawing is a real pain. Thus, it is imperative that you practice on your own. Here are the suggested problems from chapter 9, Tro 3rd edition: 51, 53, 55, 59 (a is carbon tetraiodide), 61 (a and b), 73b, 75, (treat these like the lab activity, in that you should draw the Lewis structure, identify the molecular shapes at central atoms, and determine if the molecule is polar or non-polar). Section 9.8, What do the color green, mules, and ligers have in common? Resonance is one key concept in this section.
4 Resonance describes how two otherwise equivalent Lewis structures exist as a blend of the two structures rather than either structure on their own. This is very much like how mules (horse+donkey), green (blue+yellow paint), and ligers (lion+tiger) exist as blends rather than sometimes one thing and sometimes the other. It is hard to draw a single Lewis structure that shows a blend of structures, so resonance structures in chemistry are written with double-headed arrows as you see in this section. For Example and Practice 9.7 illustrate this concept. Formal charge is the second key concept in this section. What is the definition of formal charge? What is the formula for calculating formal charge? There is a nice four item summary of formal charge rules in this section, with an example for H C N vs H N C. Formal charge is really only useful for determining which one structure is more likely to be the predominant structure from a group of similar structures. Formal charge is a bit you when you went to a formal dance in high school. There are rules for formal dress (like the formal charge rules) but the rules don t really give an accurate depiction of you (or any kind of charge or oxidation state). You are sweatshirt and jeans person and the formal you is possible for an evening, but it is not really you. Examples 9.8 and 9.9 should help illustrate the concept of resonance and formal charge in determining the most likely Lewis structures. Section 9.9, Are there exceptions to the Octet Rule (Guideline)? Yes! The rule is more of a guideline. There are 118 elements and 41,000,000 known compounds. Only 3 elements must follow the octet guideline in stable compounds. Those elements are C, O, and F. No exceptions for C, O, and F! There are exceptions for some other elements when they form compounds. What is the first example of exceptions in the text? It deals with the missing element in my list of 3 NO Exceptions. Conceptual Connection 9.5 is great practice. What is the second example of exceptions? This type of exception is fairly simple when you think about it, like for H H. Does this apply to very many elements? What is the third type of exception in the text? These are the most common and interesting. The text reveals that the non-metallic elements in row 3, 4, 5, or 6 can use the next empty d orbital set to add extra electrons to the central atom. Can 2 nd row elements have expanded octets?
The text shows sulfuric acid as an example here, but the oxo-anions are notoriously difficult to deal with in terms of formal charges and expanded octets. We won t worry about those. 5 SKIM Section 9.10 This section really emphasizes that breaking bonds absorbs energy, and forming bonds releases energy. Are there any exceptions to this in Table 9.3, where positive values mean absorbing energy for bond breaking? What does the text day about reactions being exothermic or endothermic? The author even cites the biology ATP issue. I ve been citing that example since 2001! I swear, sometimes I think Tro was in one of my classes. I promise I don t just steal stuff from his book, which we have only been using for a few years. The two of us just have some uncanny similarities in how we present material. We will NOT do any calculations like in Example 9.11, but Conceptual Connection 9.7 is a great question. Read the page about bond lengths and bond strengths. As bonds get shorter, what happens to bond strength? Section 9.11, Is this the 8 th Sea? I love the sea of electrons model for metals. It sounds serene and peaceful. Read the whole section! Chapter 9 suggested problems: Make sure you practice drawing lots of Lewis Structures, starting with counting the valence electrons. Simply looking at someone else s structures will not prepare you for the quiz or exam. There won't be any assigned Mastering Chemistry HW on drawing Lewis structures, so PRACTICE on your own. Skipping the suggested problems will be bad. Don't even ask what will happen if you don't practice. The strongly suggested problems are for students who need help with perfecting Lewis structures, or those who want to prove they are already perfect. That should include all of you! Strongly urged problems 39 (a, b, c), 41 (b, c), 43, 45, 49, 51, 53, 55, 59 (a is carbon tetraiodide), 61 (a and b), 63, 73 (b, c), 75, (treat these like the lab activity, in that you should draw the Lewis structure, identify the molecular shapes at central atoms, and determine if the molecule is polar or non-polar), 65 (formal charges) Plus, the following: A: I 3 is a known compound, with a linear geometry. No evidence exists for the formation of F 3 (expected linear geometry). Explain why F 3 does not seem to exist, after drawing Lewis structures. B: Draw Lewis structures for two ions, hypochlorite and perchlorate, so that each chlorine atom follows the octet guideline, and assign formal charges (chapter 8) and oxidation numbers (chapter 4) to the chlorine atoms in each structure. Are formal charges and oxidation numbers necessarily equal? Why are these substances both oxidizing agents?