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Honors Chemistry Name ____________________ Unit #5B Plan – Bonding Concepts and Carbon Chemistry I) II) Assignments / Handouts Lewis structures and naming H-carbons Electronegativity chart Molecular shape guide Chemical bonding and Lewis structures Bonding concepts ws Text problems / resonance structure asst Naming / Drawing H-carbons ws Polymers info and ws Polymers (uses, recycling codes, refineries) Text – chapter 7-8 study guide Polymer lab Building isomers activity More TBA Text Support Chap 7 Chap 8 Chap 22 Chap 23 pg 192-213 pg 220-255 pg 760-788 pg 796-829 (info – pg 4) (info – pg 3) (notes / info – pg 2) (notes – pg 5-6) (required – pg 7-8) (required – pg 8 --- attach to back of packet) (required – pg 9-10) (required – pg 11-14) (info – pg 15-16) (required – pg 17-18) ionic and metallic bonding covalent bonding hydrocarbon compounds functional groups III) Learning Targets (I can…) A) Use an electronegativity chart to predict whether individual bonds in molecules are ionic, polar covalent or nonpolar covalent. B) Predict shapes of molecules based upon their location in the periodic table. C) Use molecular shape and electronegativity differences between atoms to determine whether an overall molecule is polar or nonpolar. D) Draw electron dot diagrams (Lewis structures) to show the bonding of atoms. E) Identify molecules in which hydrogen bonding occurs. F) Build, draw, and name hydrocarbons. G) Build, draw, and identify isomers of hydrocarbons. H) Use my knowledge of carbon chemistry to solve polymerization and biochemistry problems I) Identify and analyze current issues involving carbon chemistry. IV) Tentative Schedule Mon 11/23 Polymer lab Tue 11/24 Polymer lab wrap-up Ionic, polar covalent, and nonpolar covalent bonding Wed 11/25 Lewis (electron dot) structures Mon 11/30 Molecular shapes Distinguishing between polar and nonpolar Tue 12/1 Intro to H-carbons Wed 12/2 Building isomers and naming H-carbons Thu 12/3 Functional groups Mon 12/7 Review Tue 12/8 Exam – bonding concepts and H-carbons Hints for Drawing Lewis (Electron Dot) Structures 1. Add the number of valence electrons in each atom to determine the total number of valence electrons. (For polyatomic anions, add one electron for each unit of negative charge. For polyatomic cations, subtract one electron for each unit of positive charge.) 2. Put electrons around each atom. Start with the 4 base electrons before pairing electrons up. (It is usually best to start with the atom that has the fewest valence electrons --excluding hydrogen.) 3. Atoms bond when electrons need to be shared to complete an octet around each atom. 4. The total number of electrons in Lewis structure should equal the total number of valence electrons in the molecule (or ion). Hydrogen is “full” with 2 electrons. All other atoms need an octet of electrons to be stable. A “shortage” of electrons usually means double or triple bonds are needed. Oh yes, there are exceptions. Rules for Naming Hydrocarbons 1. 2. 3. 4. Count carbons in longest chain to get “root” name. Check for single, double and triple bonds (or OH group). Use appropriate suffix. Number carbons starting at the end nearest the double/triple bond or OH group. Check for “branches”. (If there are no double/triple bonds or OH groups, number carbons starting at end nearest the outermost “branch”.) Put in alphabetical order. 5. Put commas between consecutive numbers; dashes between numbers and words. Alkane Roots (single bonds) Methane CH4 Ethane C2H6 Propane C3H8 Butane C4H10 Pentane C5H12 Hexane C6H14 Heptane C7H16 Octane C8H18 Nonane C9H20 Decane C10H22 Branches methyl ethyl propyl butyl (etc.) fluoro chloro bromo iodo Functional Groups Alkene --- double bond Alkyne --- triple bond Alcohol --- OH group # of Branches 2 = di 3 = tri 4 = tetra More Functional Groups (R or R’refers to carbon chain) Ether R-O-R’ Aldehyde R-(C=O)-H Ketone R-(C=O)-R’ Carboxylic acid R-(C=O)-OH Ester R-(C=O)-O-R’ Amine R-N(H2) or R-N(R’2) or R-N(R’H) Aromatics R-(C6H6) -CH3 -C2H5 -C3H7 -C4H9 -F -Cl -Br -I Chemical Bonding Notes Conditions for Chemical Bonding Atoms must be physically close to each other Atoms must attract each other o Attractive forces > repulsive forces Energy removal o Energy is released (given off) as bonds formed o Energy needs to be absorbed (added) to break bonds Bond Length –distance between nuclei of 2 atoms when their attractive forces exceed their repulsive forces by a maximum amount Chemical Bond – electrons of 2 atoms are simultaneously attracted to the nuclei of both atoms Ionic Bond – unequal sharing of electron pairs between 2 nuclei Nonpolar Covalent Bond – equal sharing of electron pairs between 2 nuclei Polar Covalent Bond – somewhat equal (or unequal) sharing of electron pairs between 2 nuclei Lewis (Electron Dot) Structures Intro Draw Lewis structures for each of the following covalent compounds. Your Hints for Drawing Lewis Structures will be helpful. H2O CF4 CH2F2 Cl2 HF NH3 CH4 NH4+1 OH-1 Unit #5B – Bonding Concepts Name_____________________________ 1. Use your electronegativity chart to predict if the following molecules have ionic, polar covalent or nonpolar covalent bonds. (Look at the individual bonds --- not the overall symmetry.) A. NaCl G. CF4 B. BaCl2 H. CO2 C. OCl2 I. O2 D. H2O J. NH3 E. CH4 K. H2S F. MgO L. Cl2 2. Draw Lewis (electron dot) structures for the following molecules or ions. A. NH3 H. H2S B. Br2 I. NH4+ C. ClF J. BH4- D. PH3 K. CH3OH E. SiH4 L. O2 F. C2H6 M. N2 G. C2H4 N. C2H2 3. Predict the shapes of the following molecules. A. NH3 D. OF2 B. Br2 E. N2 C. BCl3 F. CH2Cl2 4. Are the following molecules polar or nonpolar? (Consider overall symmetry.) A. NH3 D. OF2 B. Br2 E. N2 C. BCl3 F. CH2Cl2 5. For each of the following molecules: Draw a Lewis (electron dot) structure Determine the shape of the molecule Determine whether the molecule is polar or nonpolar A. H2O D. NF3 B. CH2F2 E. F2 C. CF4 F. CH3Cl Text pg 256-258 --- Do problems 51, 52, 53, 68, 79, 82. Text pg 234 --- Identify 3 chemical formulas (on pg 234) that are best represented by resonance structures. Draw all the resonance structures for each of these formulas. Attach to back of this homework packet. Unit 5B Book Assignment (Chapters 7-8) 1. What is the connection between valence electrons and the octet rule? 2. Cations are formed when metals ___________ electrons. 3. Anions are formed when nonmetals ___________electrons. 4. What is a formula unit? 5. The coordination number for both Na+ and Cl- in NaCl is 6. What does this mean? 6. Metals are good conductors of electricity. Why? 7. What is a metallic bond? 8. Define alloy. 9. Give 3 examples of a molecular (covalent) compound. 10. The “representative unit” of an ionic compound is called a formula unit. Why? 11. The “representative unit” of a molecular (covalent) compound is called a molecule. Why? 12. Draw the electron dot structure (Lewis structure) for each of the following molecular compounds. (Include unshared pairs of electrons.) CH4 NH3 H 2O 13. Which molecule contains: (choose from C2H6, C2H4, C2H2) a. Only single bonds b. One double bond c. One triple bond 14. What is a coordinate covalent bond? Give an example. 15. List 3 exceptions to the octet rule. Give an example of each. 16. How is bond dissociation energy and covalent bond strength related? 17. There are 2 resonance structures for ozone? Does ozone oscillate back and forth between these 2 structures? Explain. 18. A single bond has ________ sigma bonds and _________ pi bonds. 19. A double bond has ________ sigma bonds and _________ pi bonds. 20. A triple bond has ________ sigma bonds and _________ pi bonds. 21. What does VSEPR stand for? 22. What is orbital hybridization? 23. What is the difference between a polar covalent and nonpolar covalent bond? 24. List 3 types of intermolecular forces. 25. What is a network solid. Give an example.