Use the following information to determine the Lewis structure, find the electron and molecular geometry of the molecule, determine the angle of the molecule, and determine the polarity. If the atoms are not the same, you may assume that the difference in their electronegativities are between 0.4 and 2.0. Atom information: A: 6 valence electrons, CANNOT exceed the octet. Further from fluorine on the periodic table than X. X: 5 valence electrons, CANNOT exceed the octet. Closer to fluorine on the periodic table than A. Molecule: AX2 Electron Geometry: Select ] Molecular Geometry: I Select] Bond Angle: [ Select ] Polarity: [ Select]

Use the following information to determine the Lewis structure, find the electron and molecular geometry of the molecule, determine the angle of the molecule, and determine the polarity. If the atoms are not the same, you may assume that the difference in their electronegativities are between 0.4 and 2.0. Atom information: A: 6 valence electrons, CANNOT exceed the octet. Further from fluorine on the periodic table than X. X: 5 valence electrons, CANNOT exceed the octet. Closer to fluorine on the periodic table than A. Molecule: AX2 Electron Geometry: Select ] Molecular Geometry: I Select] Bond Angle: [ Select ] Polarity: [ Select]

Chemistry by OpenStax (2015-05-04)

1st Edition

ISBN:9781938168390

Author:Klaus Theopold, Richard H Langley, Paul Flowers, William R. Robinson, Mark Blaser

Publisher:Klaus Theopold, Richard H Langley, Paul Flowers, William R. Robinson, Mark Blaser

Chapter7: Chemical Bonding And Molecular Geometry

Section: Chapter Questions

Problem 113E: Use the simulation (http://openstaxcollege.org/l/16MolecPolarity) to perform the following exercises...

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Aspartame is an artificial sweetener marketed under the name Nutra-Sweet. A partial Lewis structure for aspartame is shown below. Aspartame can be classified as an organic compound (a compound based on carbon atoms). The majority of Lewis structures for simple organic compounds have all atoms with a formal charge of zero. Therefore, add lone pairs and multiple bonds to the structure above to give each atom a formal charge of zero when drawing the Lewis structure. Also note that the six-sided ring is shorthand notation for a benzene ring (C6H5). Benzene is discussed in Section 4-7. Complete the Lewis structure for aspartame. How many C and N atoms exhibit sp1 hybridization? How many C and O atoms exhibit sp3 hybridization? How many and bonds are in aspartame?
Use the simulation (http://openstaxcollege.org/l/16MolecPolarity) to perform the following exercises for a real molecule. You may need to rotate the molecules in three dimensions to see certain dipoles. (a) Sketch the bond dipoles and molecular dipole (if any) for O3 Explain your observations. (b) Look at the bond dipoles for NH3. Use these dipoles to predict whether N or H is more electronegative. (c) Predict whether there should be a molecular dipole for NH3 and, if so, in which direction it will point. Check the molecular dipole box to test your hypothesis.
Describe the shapes and compare the polarities of HF and HBr molecules. In each case, identify the end of the molecule that is more positive.
CH3COCH3 (acetone) is a common laboratory solvent that is often used in nail polish remover. Its Lewis structure and space-filling molecular image are shown here: Acetone is a polar molecule; the oxygen end has a slightly negative charge (oxygen is more electronegative), whereas the carbon and hydrogen end has a slightly positive charge. In liquid acetone, the molecules are attracted to each other via these polar ends-the positive end of one molecule is attracted to the negative end of its neighbor-and therefore align as shown here: Draw a Lewis structure and space-filling molecular image for CH2Cl2 (dichloromethane), another common laboratory solvent. Is the molecule polar? Which end of the molecule has a slightly negative charge? Which end has a slightly positive charge? Draw several space-filling molecular images of CH2Cl2 showing how they align in liquid dichloromethane.
Vitamin B6 is an organic compound whose deficiency in the human body can cause apathy, irritability, and an increased susceptibility to infections. Below is an incomplete Lewis structure for vitamin B6. Complete the Lewis structure and answer the following questions. Hint: Vitamin B6 can be classified as an organic compound (a compound based on carbon atoms). The majority of Lewis structures for simple organic compounds have all atoms with a formal charge of zero. Therefore, add lone pairs and multiple bonds to the structure below to give each atom a formal charge of zero. a. How many bonds and bonds exist in vitamin B6? b. Give approximate values for the bond angles marked a through g in the structure. c. How many carbon atoms are sp2 hybridized? d. How many carbon, oxygen, and nitrogen atoms are sp3 hybridized? e. Does vitamin B6 exhibit delocalized bonding? Explain.
a model of each of the following molecules: a. Based on your model, draw a bond-line representation with as many atoms as possible in the plane of the paper. Use wedge and dash bonds to represent any atoms that do not lie in the plane of the paper. b. Indicate each unique bond angle and the shape of each unique central atom.
Use VSEPR to explain why the HBH bond angle of BH3 is 120°. (Hint: What is one-third of360°?)
Chlorofluorocarbons have been banned because they cause ozone depletion. The main replacements are hydrofluorocarbons, such as CF3CFH2. Draw a Lewis structure for CF3CFH2, determine its geometry, make a sketch of the molecule, and determine whether it is polar. (The two carbons are bonded to each other in the middle of the molecule, and the other atoms are bound to the two central carbon atoms.)
7.30 The bond in HF is said to be polar, with the hydrogen carrying a partial positive charge. For this to be true, the hydrogen atom must have less than one electron around it. Yet the Lewis dot structure of HF attributes two electrons to hydrogen. Draw a picture of the electron density distribution for HF and use it to describe how the hydrogen atom can carry a partial positive charge. How can these two models of the HF bond (the electron density and the Lewis structure) seem so different and yet describe the same thing?
The molecular ion S3N3 has the cyclic structure All SN bonds are equivalent. (a) Give six equivalent resonance hybrid Lewis diagrams for this molecular ion. (b) Compute the formal charges on all atoms in the molecular ion in each of the six Lewis diagrams. (c) Determine the charge on each atom in the polyatomic ion, assuming that the true distribution of electrons is the average of the six Lewis diagrams arrived at in parts (a) and (b). (d) An advanced calculation suggests that the actual charge resident on each N atom is 0.375 and on each S atom is +0.041 . Show that this result is consistent with the overall +1 charge on the molecular ion.
a. How many sticks did you need to make the skeleton structure?____________ b. How many sticks are left over? ____________ If your model is to obey the octet rule, each ball must have four sticks in it except for hydrogen atom balls, which need and can only have one. Each atom in an octet rule species is surrounded by four pairs of electrons. c. How many holes remain to be filled? ____________ Fill them with the remaining sticks, which represent nonbonding electron pairs. Draw the complete Lewis structure for NH2Cl using lines for bonds and pairs of dots for nonbonding electrons.
The “Chemistry in Focus" segment Hiding Carbon Dioxide- discusses attempts at sequestering (storing) underground CO2produced at power plants so as to diminish the: greenhouse effect. Draw all resonance structures of the CO2molecule.