Unformatted text preview: AP Chemistry Practice Problems Module 4: Bonding The headings on these problems correspond to the headings on your content pages. You should work on these throughout the unit. Be sure to show all your work, otherwise you won’t know where you went wrong if you made a mistake. Many times it is easier to write out your work using pen and paper, so you may choose to scan your work if you like. Just make sure that it is clear enough for me to see! Scanned work should be saved as one pdf file. Multiple files may be too large for the dropbox. If you type your answers, use a different color font, highlight, textbox, etc. to make the work easier to see. Place the completed document into the Module 4: Bonding Practice Problems dropbox according to your schedule. Types of Bonds, Properties of Ionic Bonds, Properties of Covalent Bonds, Properties of Metallic Bonds 1. Describe what kind of event must occur (involving electrons) if the atoms of two different elements are to react to form (a) an ionic compound or (b) a molecular compound. 2. What holds an ionic compound together? Can we identify individual molecules in an ionic compound? Formation of Ionic Bonds, Lattice Energy, Formation of Covalent Bonds 3. What must be true about the change in the total potential energy of a collection of atoms for a stable compound to be formed from the elements? 4. How is the tendency to form ionic bonds related to the IE and EA of the atoms involved? 5. Magnesium forms the ion Mg2+, but not the ion Mg3+. Why? Polar and Nonpolar Covalent Bonds, Electronegativity, Predicting Bond Type 6. Which elements are assigned electronegativities of zero? Why? 7. If an element has a low electronegativity, is it likely to be a metal or a nonmetal? Explain your answer. 8. Nitrogen and arsenic are in the same group in the periodic table. Arsenic forms both AsCl3 and AsCl5, but with chlorine, nitrogen only forms NCl3. On the basis of the electronic structures of N and As, explain why this is so. 9. Use the table of electronegativities below to choose the atom in each of the following bonds that carries the partial negative charge: (a) Hg–I, (b) P–I, (c) Si–F, (d) Mg–N. 10. Which of the bonds in the previous problem is the least polar? Lewis Symbols, Lewis Structures, Bond Order, Bond Length, and Bond Energy, Bond Length and Bond Order, Bond Strength and Bond Energy, Limitations to Lewis Structure Model 11. Write Lewis symbols for the following atoms: (a) K, (b) Ge, (c) As, (d) Br, (e) Se. 12. How many dots must appear in the Lewis structures of (a) HIO3, (b) H2CO3, (c) HCO3−, and (d) PCl4+? 13. Draw Lewis structures for (a) AsCl4+, (b) ClO2−, (c) HNO2, and (d) XeF2. 14. Draw Lewis structures for (a) HIO3, (b) H2CO3, (c) HCO3−, and (d) PCl4+. 15. Draw Lewis structures for (a) SeO3 and (b) SeO2. 16. Draw Lewis structures for (a) NO+, (b) NO2−, (c) SbCl6−, and (d) IO3−. 17. Draw Lewis structures for (a) GeCl4, (b) CO32−, (c) PO43−, and (d) O22−. 18. Arrange the following in order of increasing C–O bond length: CO, CO32−, CO2, HCO2− (formate ion). Formal Charge 19. Draw the Lewis structure for HClO4. Assign formal charges to each atom in the formula. Determine the preferred Lewis structure for this compound. Resonance 20. Why is the concept of resonance needed? 21. Draw the resonance structures of the benzene molecule. Why is it more stable than one would expect if the ring contained three carbon–carbon double bonds. Determining Molecular Shape Using VSEPR 22. Predict the shapes of (a) SF3+, (b) NO3−, (c) SO42−, (d) O3, and (e) N2O. 23. Predict the shapes of (a) FCl2+, (b) AsF5, (c) AsF3, (d) SbH3, and (e) SeO2. 24. Predict the shapes of (a) CS2, (b) BrF4−, (c) ICl3, (d) ClO3−, and (e) SeO3. 25. Ethylene, a gas used to ripen tomatoes artificially, has the Lewis structure What would you expect the H–C–H and H–C–C bond angles to be in this molecule? (Caution: Don't be fooled by the way the structure is drawn here.) 26. Predict the bond angle for each of the following molecules: (a) HOCl, (b) PH2−, (c) OCN−, (d) O3, (e) SnF2. Polar Molecules 27. Which of the following molecules have a permanent dipole moment? a. H2O b. CO2 c. CH4 d. N2 e. CO f. NH3 28. Which of the following molecules would be expected to be polar: (a) HBr, (b) POCl3, (c) CH2O, (d) SnCl4, (e) SbCl5? Valence Bond Theory, Hybrid Orbitals 29. Use orbital diagrams to describe the bonding in (a) SnCl4 and (b) SbCl5. Be sure to indicate hybrid orbital formation. 30. Draw Lewis structures for the following and use the geometry predicted by the VSEPR model to determine what kind of hybrid orbitals the central atom uses in bond formation: a) SbCl6−, (b) BrCl3, (c) XeF4. Sigma and Pi Bonds, Double and Triple Bonds 31. What kinds of bonds (σ or π) are found in the numbered bonds in the following molecule? 2 32. A nitrogen atom can undergo sp hybridization when it becomes part of a carbon–nitrogen double bond, as in H2C=NH. (a) Using a sketch, show the electron conﬁguration of sp2 hybridized nitrogen just before the overlapping occurs to make this double bond. (b) Using sketches (and the analogy to the double bond in C2H4), describe the two bonds of the carbon–
nitrogen double bond. (c) Describe the geometry of H2C"NH (using a sketch that shows all expected bond angles). ...
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- Fall '19