Use sketches of orbitals to show how VB theory explains the bonding in the Cl2 molecule. Illustrate with appropriate orbital diagrams as well. Use these to determine the number of orbitals that overlap in the Cl2 molecule. How many orbitals overlap in Cl2? Choose answer that aplies. a. Bonding will be similar to the bonding in F2, because fluorine and chlorine are in the same group of the periodic table. b. One 3p orbital of each chlorine atom will participate in the bonding in the Cl2 molecule. c. One 2p orbital of each chlorine atom will participate in the bonding in the Cl2 molecule. d. Two sp7 molecular orbitals will participate in the bonding in Cl2 molecule. d. Bonding will be similar to the bonding in F2, because fluorine and chlorine are in the same period of the periodic table.

11. What is the electron-domain geometry of water molecule (H2O)? a. bent b. tetrahedral c. trigonal planar d. linear 12. Write the correct Lewis dot structure for O2. Which statement correctly describes the structure of the whole molecule? a. There is a double bond and four lone pairs. b. There is a double bond and six lone pairs. c. There is a single bond and six lone pairs. d. There is a single bond, a double bond, and six lone pairs. 13. If there are four (4) electron pairs around the central atom of a molecule, these electron pairs are in a/n a. trigonal planar b. trigonal pyramidal arrangement. C. octahedral d. tetrahedral 14 The molecular geometry of the molecule NF3 is a. linear b. trigonal pyramidal c. tetrahedral d. trigonal planar 15 Which of the following species will exhibit tetrahedral molecular geometry? d. PCI5 a. CCI4 b. CO2 C. O3

The PH3 molecule has a tetrahedral arrangement of the electron pairs. Which of the following is TRUE regarding the H-P-H angles in the molecule? Select one: A. The angles are slightly greater than those in a perfect tetrahedron. B. The angles are equal to those in a perfect tetrahedron. C. The angles are slightly less than those in a perfect tetrahedron. D. None of the above.

Draw the Lewis structures for the following chemicals, predicts the angle between the bonds and say what the hybridization is expected for the central atom.   a. KrF2 b. XeO2F4

1. Consider the diatomic molecules ( molecules with 2 atoms ) in this exercise : O2, N2, Cl2, and HCl . Do you think a diatomic molecule can be any shape other than linear ? Explain . because They are only two alems 2. Look at the Lewis structure you drew for SCl2 and CS2 Both of these molecules are made up of three atoms but have different molecular geometries . a . How is the arrangement of electrons around the central atoms different ? Give the number of lone pairs and bonding groups around each central atom Based on your previous answer , explain why SCl2 can't be linear . c . Explain why one molecule is polar and the other is nonpolar

5. Draw the # and molecular orbitals for the discrete and localized T bonds in the following structures. Indicate any polar- ization these bonds may have in your diagram through the shapes of the discrete bonding and antibonding orbitals. e 6. Show any plausible resonance structures for the molecules given in Exercises 1 and 5. If there are no plausible resonance structures, indicate this (note that the molecule corresponding to letter A is the same in each problem). 7. Discuss the hybridization in the C-C and C-H bonds of cyclopropane, given that the H-C-H bond angle is 118°. withara --- totrachlorida b dingl A. B. C. D. 3

Use your knowledge of Lewis structure and valence bond theory to predict the characteristics of acetonitrile (NCCH3) and its isomer methyl isocyanide (CNCH3). a. Determine the best Lewis structure for acetonitrile. Include the formal charge on each carbon and nitrogen atom. b. Using valence bond theory, determine the hybridization on the nitrogen atom in acetonitrile and state the number of pi and sigma bonds in the nitrogen-carbon bond. c. Determine the best Lewis structure for methyl isocyanide. Include the formal charge on each carbon and nitrogen atom. d. Using valence bond theory, determine the hybridization on each carbon atom in methyl isocyanide and state the hybrid orbitals in which you would expect to find nonbonding electrons. e. Based on the Lewis structures which molecule would you expect to be more reactive: acetonitrile or methyl isocyanide? Briefly explain your answer.

6. Consider the molecule sulfur dioxide. Draw the Lewis structure such that it follows the octet rule. Include any formal charges. What is the electron geometry? What is the molecular shape? What are the expected bond angles? What kind of intermolecular forces does it have? How many sigma bonds are there? How many pi bonds? What is the hybridization of the sulfur atom? In the space below, sketch the hybridized orbitals.

Which of the following statements is incorrect? About molecular orbitals:a. Molecular orbitals (MOs) are obtained by linear combination of atomic orbitals (AOs).b. The combination of two atomic orbitals always yields two molecular orbitals.c. A molecular orbital describes the region of space where an electron shared by two atomsis most likely to be found.d. The number of molecular orbitals is always superior to the number of atomic orbitals theywere created from.e. None of the above.

Study the following sketch of a molecular orbital (MO) in a homonuclear diatomic molecule. This MO was formed by combining one 3p atomic orbital from each atom. The dark dots in this sketch are the nuclei. Now use the sketch to complete the table below. Write the symbol for this MO. Is this a bonding or antibonding MO? What is the energy of this MO, compared to the energy of a 3p orbital on one of the separate. atoms? 0 O bonding O antibonding O higher O lower O the same not enough information to decide O 2 00 X 2 뮤 0 Ś