Chemistry 1A Spring 1998 Exam #4 KEY Chapters 9 & 10

Transcription

1 Chemistry 1A Spring 1998 Exam #4 KEY Chapters 9 & 10 For each of the following, write the word, words, or number in each blank that best completes each sentence. (2 points each) 1. A(n) molecular orbital can be defined as the volume that contains a high percentage of the electron charge for an electron in a molecule or polyatomic ion. It can also be defined as the volume within which an electron in a molecule or polyatomic ion has a high probability of being found. 2. The most common bonding pattern for sulfur atoms is two covalent bond(s) and two lone pair(s). 3. Isomers are molecules that have the same molecular formula but a different structural formula. They have the same atoms but a different order of linking the atoms. 1

2 4. Compare the Valence Bond Model for covalent bonding with the Molecular Orbital Theory for covalent bonding by answering the following questions. (10 points) a. Draw the best Lewis structure for carbon monoxide, CO. b. Evaluate the stability of this Lewis structure in terms of our criteria for stable structures. Three bonds and one lone pair is a rare bonding pattern for both carbon and oxygen. This suggests relative instability. The plus formal charge on the more electronegative oxygen and the minus formal charge on the less electronegative carbon suggest instability. c. Draw a molecular orbital diagram for CO. σ * 2p π * 2p π 2p σ 2p σ * 2s σ 2s σ * 1s σ 1s d. Evaluate the stability of CO in terms of our Molecular Orbital Theory criteria for stable molecules. What is the bond order for CO? The bond order is three, suggesting a very stable bond. e. Carbon monoxide has very strong bonds and forms anytime there is incomplete combustion of carbon containing compounds, like those burned in a car engine. Which model, valence bond or molecular orbital, best explains this stability? 2

3 5. Draw a reasonable Lewis structure for each of the following formulas. (4 points each) a. O 2 F 2 b. HC 2 F c. HNO 2 d. CH 2 CHCH 3 e. CH 3 CH 2 CO 2 CH 3 6. Draw a reasonable Lewis structure for the ozone molecule, O 3. Use the skeleton below. The structure is best described in terms of resonance, so draw all of its reasonable resonance structures and the resonance hybrid that summarizes these structures. (6 points) 3

4 7. Draw two reasonable Lewis structures for sulfuric acid, SO 3. One of the Lewis structures should emphasize the octet rule, and the other structure should minimize formal charges. If either Lewis structure is best described with resonance, draw all of its reasonable resonance structures and the resonance hybrid. (8 points) 8. For the Lewis Structures on the next three pages, do each of the following. (7 points each) a. a. Predict the hybridization for each atom in the structure. b. Describe each bond by stating whether it is sigma, localized pi, or part of a delocalized pi system and by stating which atomic orbitals overlap to form the bonds. c. Describe the electron group geometry about the central atom. e. Sketch the molecule and indicate the approximate bond angles in your sketch. f. Describe the molecular geometry about the central atom. hybridization for the O sp 2 hybridization for the right N sp 2 hybridization for the F sp 3 description of bonding one σ bond between O and N due to sp 2 -sp 2 overlap one π bond between O and N due to p-p overlap one σ bond between N and F due to sp 2 -sp 3 overlap electron group geometry around the N trigonal planar sketch and bond angles molecular group geometry around the N bent 4

5 b. hybridization for each C sp 2 hybridization for each O sp 2 description of bonding one σ bond between C and C due to sp 2 -sp 2 overlap four σ bonds between O and C due to sp 2 -sp 2 overlap one delocalized pi system due to the overlap of p orbitals on each atom electron group geometry around each C sketch and bond angles trigonal planar molecular geometry around each C trigonal planar 9. For each of the following organic compounds, you are given the Lewis structure, a line drawing, or a condensed formula. (5 points each) Identify each as representing an alkane, alkene, alkyne, arene (aromatic), alcohol, carboxylic acid, aldehyde, ketone, ether, ester, amine, or amide. Draw two other ways to show each structure. For example, if you are given a Lewis structure, draw the line drawing and the condensed formula. Ketone a. CH 3 (CH 2 ) 4 COCH 3 5

6 Alkane b. (CH 3 ) 2 CHCH 2 CH(CH 3 )CH 2 CH 3 Carboxylic acid c. CH 3 (CH 2 ) 12 CO 2 H Ether d. (CH 3 ) 2 CHCH 2 OCH 3 Amide e. CH 3 CH 2 CH 2 CONH 2 6

7 Answer the following in short answer form. (6 points each) 10. With reference to the molecular orbital theory, write an explanation for why bonding molecular orbitals are more stable than the separate atomic orbitals that form them. For example, explain why the σ 1s molecular orbital that is formed from two 1s atomic orbitals is more stable than the separate 1s atomic orbitals. Also write an explanation for why antibonding molecular orbitals are less stable than the separate atomic orbitals that form them. For example, explain why the σ* 1s molecular orbital that is formed from two 1s atomic orbitals is less stable than the separate 1s atomic orbitals. According to the linear combination of atomic orbitals model, when two atomic orbitals interact, they form one bonding molecular orbital from an in-phase interaction of the electron waves, and they form an antibonding molecular orbital from the out-of-phase interaction of the electron waves. In-phase interaction leads to an increase in the intensity of the negative charge between two nuclei. This leads to an increase in +/- interaction and an increase in stability and a decreased potential for the electrons. Thus, the bonding molecular orbital (e.g. σ 1s ) is more stable than the separate atomic orbitals (e.g. 1s). Out-of-phase interaction leads to a decrease in the intensity of the negative charge between two nuclei. This leads to a decrease in +/- interaction and a decrease in stability and an increased potential for the electrons. Thus, the antibonding molecular orbital (e.g. σ* 1s ) is less stable than the separate atomic orbitals (e.g. 1s). 11. The Lewis structure for sulfur tetrafluoride, SF 4, is below. With reference to the assumptions of the valence bond model, write an explanation for how sulfur atoms are able to form four equivalent covalent bonds and one lone pair in molecules like SF 4. (4 points) Only the highest energy electrons participate in bonding. The sulfur atom is surrounded by five electron groups (four bonds and a lone pair), so we predict that it is sp 3 d hybridized. To form these bonds, it is as if an electron is promoted from a 3p orbital to an empty 3d orbital, and then the 3s, the three 3p, and the one 3d orbital blend to form sp 3 d hybrid orbitals. Covalent bonds form to pair unpaired electrons, so we get four equivalent bonds and one lone pair. 7