Guide to Chapter 7. Covalent Bonds and Molecular Structure
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1 Guide to Chapter 7. Covalent Bonds and Molecular Structure We will spend three lecture days on this chapter. In this chapter we will learn more about the structure of the covalent molecules. We will start with a discussion of electronegativity, bond polarity, dipoles, ion electron configuration and relative size, lattice energy, covalent bond energies, and Lewis dot structures. Once we are able to sketch Lewis dot structures for molecules, we will take up some exceptions to the octet rule, the concept of resonance, formal charge, and the shape of molecules. Finally, we will cover expanded octets, hybridization, s and p bonds. This is a chapter study guide, given section-by-section. Work problems on separate sheets of paper and keep them with this guide. When working problems, use plenty of space and when appropriate, show all work. Read the introductory paragraph to Chapter 7. Covalent Bonds and Molecular Structure and Section 7.1 The Covalent Bond Read Section 7.2 Strengths of Covalent Bonds. Learning Objective 1: Know how to use a bond energy table to estimate the strength of covalent bonds. Read Section 7.3 Comparison of Ionic and Covalent Compounds. Learning Objective 2: Know the differences in the physical properties between ionic compounds and molecular compounds (covalent). Know the forces that hold ionic compounds together. Know the forces that hold molecular compounds together. Read Section 7.4 Polar-Covalent Bonds: Electronegativity. Learning Objective 3: Know the meaning of the term electronegativity. Learning Objective 4: Know the periodic trend in electronegativity of the elements. Given a group of elements, use the periodic table to predict the most/least electronegative element or rank them in order of increasing electronegativity. Learning Objective 5: Be able to predict the polarity of a bond (d+ and d-) and the "direction" of the dipole moment for a given bond between two elements. Learning Objective 6: Given a species that has covalent bonds, assign a d+ and a d- to all the atoms in the molecule. Learning Objective 7: Given the formula of a simple compound determine if it is mostly ionic or mostly covalent. Do Problems 1-3 at the end of the section. Do the following end-of-chapter problems: , 42, 44 Problem Club Question A. Using only a periodic table, circle the most electronegative member from each group: Answers in blue (a) P As Sb (d) F Cl Br (b) Ge As Se (e) S Cl Ar (c) Ga Si N (f) C Si Ge Problem Club Question B. Using only a periodic table whenever possible, circle the most polar bond from each group. (In some cases, you may need to use a table from the text.): Answers in blue (a) C-N or C-O (d) C-H or Si-H (almost the same! I had to look up exact EN values) (b) N-F or P-F (e) H-H or F-F (neither) (c) F-Cl or F-Br (f) C-O or C-S Problem Club Question C. When the electronegativity difference between the two atoms is very small or zero, the resulting bond is called a [non-polar covalent] bond, since the bonding electrons are essentially equally shared. (a) ionic (b) polar covalent (c) non-polar covalent (d) metallic Dr. Mattson, General Chemistry, Chm 203, Guide to Chapter 7. Covalent Bonds & Molecular Structure 1
2 Problem Club Question D. Which of these bonds is the most polar? C-N C-O O-F N-F C-F Read Section 7.5 Electron Dot Structures. Learning Objective 8: Determine the number of valence electrons in an atom or ion. Learning Objective 9: Given the formula of a simple covalent species (a molecule or polyatomic ion), draw its Lewis structure. The Lewis structure MUST clearly show all bonding pairs and lone or nonbonding pairs of electrons. Use the guidelines from class notes and text. Learning Objective 10: Use the Lewis structure to determine the ABE formula for a given molecule. Learning Objective 11: Use the Lewis structure to determine the bond order (single, double or triple). Learning Objective 12: Given examples of covalent species, rank the length and strength of the bonds. The structure of ammonia surprised early scientists. Learning Objective 13: Determine by inspection if a structure is paramagnetic or diamagnetic. Do Problems 4 and 5 at the end of the section. Do the following end-of-chapter problems: 48, 56 Problem Club Question E. Draw the Lewis dot structures for the following molecules and polyatomic ions. In some cases, you will have to predict the formula of the compound. (a) NF 3 ; (b) OCl 2 ; (c) TeH 2 ; (d) NH 4 + ; (e) NH 2 - ; (f) SiH?; (g) SeF?; (h) OH 3 + ; (i) CO 3 2- Answer: Using the notation that the central atom is designated by A, the bonding groups by B and the electron groups by E, the answers are: (a) AB 3 E, (b) AB 2 E 2, (c) AB 2 E 2, (d) AB 4, (e) AB 2 E 2, (f) AB 2 E, (g) AB 2 E 2, (h) AB 3 E, and (i) AB 3 Problem Club Question F. Carefully draw the Lewis dot structures of the following molecules and ions: (a) NO 2 - ; (f) SiO4 4- ; (g) PCl4 + ; (h) BF 3 ; (i) SCN - (C is central); (j) NO2; (k) CO 3 Answer: Using the notation that the central atom is designated by A, the bonding groups by B and the electron groups by E, the answers are: (a) AB 2 E, (b) AB 2, (c) AB 2 E 2, (d) AB 3 E, (e) AB 2 E 2, (f) AB 4, (g) AB 4, (h) AB 3, (i) AB 2, (j) AB 2 E, (k) AB 3, and (l) AB 4 (i) Which disobey(s) the octet rule? Answer: h and j (ii) Which has/have one or more lone pairs on the central atom? Answer: a, c, d, e, j (iii) Which has/have resonance? Answer: a, I, j, k (iv) Which is/are paramagnetic? Answer: j (v) Which has/have at least one double or triple bond? Answer: a, b, I, j, k Problem Club Question G. In which of the following is the oxygen-oxygen bond predicted to be the: (a) shortest: O 2 O 2 (b) longest: O 2 O 2 (c) weakest: O 2 O 2 (d) strongest: O 2 O 2 Read Section 7.6 Electron Dot Structures of Polyatomic Molecules. Learning Objective 14: Be able to extend the concepts of Section 7.5 to more complex systems with more than one central element. Do Problems 6-11 throughout this section. Do the following end-of-chapter problems: 52, 58 Dr. Mattson, General Chemistry, Chm 203, Guide to Chapter 7. Covalent Bonds & Molecular Structure 2
3 Read Section 7.7 Electron Dot Structures and Resonance. Learning Objective 15: After drawing your Lewis structure, determine if resonance forms are possible and draw the Lewis structure of each resonance form. Do Problems at the end of the section. Do the following end-of-chapter problems: 50, 54 Problem Club Question H. In which ONE of the following is resonance predicted? XeO 2 TeO 3 SO 3-2 ClO 4 - Problem Club Question I. Which of the structures that you drew in Probem Club Question F has/have resonance? (a) NO 2 - ; (f) SiO4 4- ; (g) PCl4 + ; (h) BF 3 ; (i) SCN - (C is central); (j) NO2; (k) CO 3 Read Section 7.8 Formal Charges. Learning Objective 16: After you have drawn the Lewis structure or you are given a Lewis structure of a molecule or polyatomic ion, determine the formal charge of each element in the Lewis structure. Learning Objective 17: Chemists use formal charges to determine the best Lewis structure of a covalent species that can have more than one possible Lewis structure. Given an example such as N2O or OCN -, use formal charges to determine the best Lewis structure. Do Problems 15 and 16 at the end of the section. Do the following end-of-chapter problems: 60, 62, 64 Problem Club Question J. Nitrous oxide is shown here. The bonding is nitrogen-nitrogen-oxygen. (a) Sketch the Lewis dot structure and assign formal charges to your structure. (b) How many s and p bonds are all together present in nitrous oxide? Answer: Two resonance forms are important: One features N double-bonded to N double bonded to O and has formal charges of , respectively. Another has N triple bonded to N single bonded to O with FCs of 0, +1 and 1, respectively. (b) Bothe forms have two s and two p bonds. Problem Club Question K. Assign formal charges to each atom in the structures you drew in Probem Club Question F: (a) NO 2 - ; (f) SiO4 4- ; (g) PCl4 + ; (h) BF 3 ; (i) SCN - (C is central); (j) NO2; (k) CO 3 Answer: (a) NO 2 - has two resonance forms the single bonded O will have FC = -1 and the other two atoms will have FC = 0; (b) CO 2 all FC = 0; (c) H 2 Te all FC = 0; (d) NF 3 all FC = 0; (e) NH 2 - The formal charges will be 0 for the two Hs and 1 for N; (f) SiO 4 4- FC = 0 for Si and 1 for each O; (g) PCl4 + Cls have FC = 0 and P has FC = +1; (h) BF 3 All FCs = 0 (i) SCN - has two resonance forms one of them features S single bonded to C triple bonded to N with FCs of 1, 0, and 0 for S, C, and N, respectively. The other, better structure has S double-bonded to C which is double-bonded to N with FC = 0 for S and C and 1 for N; (j) NO2 Two resonance forms are important: Both feature N double-bonded to one O and single-bonded to the other O the single bonded O has FC = -1, the double-bonded O has FC = 0 and N has formal charges of +1; (k) CO 3 2- Three resonance forms are important: Each features C double-bonded to one O and single-bonded to the other two Os the single bonded O has FC = -1, the double-bonded O has FC = 0 and C has formal charges of 0; (l) SO 4-2 If one does not expand S s octet, there will be only one structure which features all Os singly-bonded to S with FC = -1 for each O and for S. Read Section 7.9 Molecular Shapes: The VSEPR Model. Dr. Mattson, General Chemistry, Chm 203, Guide to Chapter 7. Covalent Bonds & Molecular Structure 3
4 Learning Objective 18: For a covalent species, draw its Lewis structure first, then use VSEPR theory to predict its molecular geometry (tetrahedral, bent, etc.), draw the shape, determine the bond angles, count the number of bonding pairs and lone pairs of electrons. Do Problems at the end of the section. Do the following end-of-chapter problems: 32, 34, 36, 66, 68, 70, 72 (a-c), 74, 76, 78 Problem Club Question L. Give its ABE VSEPR designation and name each geometric shape for each structure in Problem Club Question F, above: (a) NO 2 - ; (f) SiO4 4- ; (g) PCl 4 + ; (h) BF 3 ; (i) SCN - (C is central); (j) NO2; (k) CO 3 Answer: (a) AB 2 E; (b) AB 2 ; (c) AB 2 E 2 ; (d) AB 3 E; (e) AB 2 E 2 ; (f) AB 4 ; (g) AB 4 ; (h) AB 3 ; (i) AB 2 ; (j) AB2E; (k) AB 3 ; (l) AB 4 Problem Club Question M. What structure (bent, tetrahedral, etc.) is shown here? Give the name and predicted angles. Answer: see-saw with angles of approximately 120, 90 and 180 degrees (some distortion towards smaller angles occurs because of the E group in the equatorial position. Problem Club Question N. Match each compound or ion with the correct ABE formula and structural name. J NO 2 A. AB 3 E structure K SF 6 B. tetrahedral I SCl 2 C. square pyramid A NH 3 D. bond order 3 F SO 3 E. sp 2 and resonance E SO 2 F. trigonal plane H ClF 3 G. AB 2 C BeF 5 H. T-shaped B SO -2 4 I. bent, <109 o angle D N 2 J. paramagnetic G CS 2 K. octahedral Problem Club Question O. (ACS-Style) Answer: D Problem Club Question P. (ACS-Style) Answer: B Problem Club Question Q. (ACS-Style) Answer: C Problem Club Question R. (ACS-Style) Answer: C Problem Club Question S. (ACS-Style) Answer: D These three sections all go together: Read Section 7.10 Valence Bond Theory. Dr. Mattson, General Chemistry, Chm 203, Guide to Chapter 7. Covalent Bonds & Molecular Structure 4
5 Read Section 7.11 Hybridization and sp 3 Hybrid Orbitals. Read Section 7.12 Other Kinds of Hybrid Orbitals. Learning Objective 19: From the Lewis structure, determine the hybridization for each atom in the molecule. Learning Objective 20: Be able to sketch the orbital diagram and energy diagram for the valence electrons in an hybridized molecule/ion. Learning Objective 21: Show (using a sketch) how the hybridized orbitals of two elements overlap to form s and p bonds. Clearly show (using arrows) the shared electrons in the bond. Learning Objective 22: For a given Lewis structure of a molecule, count the number of s and p bonds. Do Problems throughout these sections. Do the following end-of-chapter problems: 84, 86, 88, 90 Problem Club Question T. Draw the Lewis dot structures of the following molecules and ions: (a) NO 2 - ; (b) SF 4 ; (c) XeO 4 ; (d) ClF 3 ; (e) ICl 2 - ; (f) SF5 - ; (g) PCl6 - ; (h) AsO3 - ; (i) OCN - (C is central); (j) NO 2 ; (k) CO 3 2- ; (l) PF5 Answer: (a) AB 2 E; (b) AB 4 E; (c) AB 4 ; (d) AB 3 E 2 ; (e) AB 2 E 3 ; (f) AB 5 E; (g) AB 6 ; (h) AB 3 ; (i) AB 2 ; (j) AB 2 E; (k) AB 3 ; (l) AB 5 (i) Which are non-polar? Answer: c and l (ii) Which have five structural groups (lone pairs, bonding pairs, and/or unpaired electrons) on the central atom? Give their structural names. Answer: b, d, e, l (iii) Which have six structural groups (lone pairs, bonding pairs, and/or unpaired electrons) on the central atom? Give their structural names. Answer: f, g (iv) Which have p bonds? Give the number of s and p bonds for these. Answer: a (2 s and 1 p); h (3 s and 1 p); i (2 s and 2 p); j (2 s and 1 p); k (3 s and 1 p) (v) Which have sp 2 hybridization? Answer: a, h, j, k (vi) Which feature 120 o and 90 o angles? Answer: b, d, e, l (vii) Which have d 2 sp 3 hybridization? Answer: c, f, g (viii) Which have formal charges other than zero on one or more atoms within the molecule/ion? Answer: a, e, f, g, h, i, j, k Problem Club Question U. (ACS-Style) Answer: D Problem Club Question V. (ACS-Style) Answer: A Skip Sections General Questions: Problem Club Question X. In the molecule shown here, ethane, the dark spheres are carbon and the light spheres are hydrogen. (a) What is the formula for ethane? (b) What is the hybridization for each carbon? (c) What bond angles are present? (d) What ABE formula describes each carbon? (e) What is the C-C bond order? (f and g) Sketch the Lewis dot structure using dots and lines for ethane (two drawings). Answer: (a) C 2 H 6 ; (b) sp 3 ; (c) 109 o ; (d) AB 4 ; (e) 1 Dr. Mattson, General Chemistry, Chm 203, Guide to Chapter 7. Covalent Bonds & Molecular Structure 5
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