A. Complete the table Bond formation according What happened to the Type of compound formed Given to atoms e's? 1, polar bond 2. nonpolar bond 3. ionic bond B. Draw the chemical bond of the following: [Lewis Dot] 1. Phosphorous (P-Group VA) + Chlorine (Cl-Group VIIA) 2. Aluminum (Al 13) + Sulfur (S 16) 3. Nitrogen (N-Group VA) + Flourine (F-Group VIIA) 4. Potassium (K19) + Oxygen (08) 5. Astatine (At-Group VIIA) + Astatine (At-Group VIIA)

A. Complete the table Bond formation according What happened to the Type of compound formed Given to atoms e's? 1, polar bond 2. nonpolar bond 3. ionic bond B. Draw the chemical bond of the following: [Lewis Dot] 1. Phosphorous (P-Group VA) + Chlorine (Cl-Group VIIA) 2. Aluminum (Al 13) + Sulfur (S 16) 3. Nitrogen (N-Group VA) + Flourine (F-Group VIIA) 4. Potassium (K19) + Oxygen (08) 5. Astatine (At-Group VIIA) + Astatine (At-Group VIIA)

Chemistry

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Author:Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste

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Chapter8: Bonding: General Concepts

Section: Chapter Questions

Problem 91E: The most common exceptions to the octet rule are compounds or ions with central atoms having more...

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Predict die molecular structure and bond angles for each molecule or ion in Exercises 88 and 94. a. POCl3, SO42, XeO4, PO43, ClO4 b. NF3, SO32, PO33, ClO3 c.ClO2, SCl2, PCl2 d. Considering your answers to parts a, b, and c. what conclusions can you draw concerning the structures of species containing the same number of atoms and the same number of valence electrons? (O3), sulfur dioxide, and sulfur trioxide.
The most common exceptions to the octet rule are compounds or ions with central atoms having more than eight electrons around them. PF5, SF4, CIF3, and Br3 are examples of this type of exception. Draw the Lewis structure for these compounds or ions. Which elements, when they have to, can have more than eight electrons around them? How is this rationalized?
In Section 12.10 of your text, the term “effective pairs" is used. What does this mean?
7.97 Consider the structure shown below for as well as any other important resonance structures. (a) What is the expected O—N—O bond angle in this structure? (b) The molecule contains N—O bonds of two different lengths. How many sborter N—O bonds would be present?
a Carbonyl fluoride, COF2, is an extremely poisonous gas used in organofluorine synthesis. Give the valence bond description of the carbonyl fluoride molecule. (Both fluorine atoms are attached to the carbon atom.) b Nitrogen, N2, makes up about 80% of the earths atmosphere. Give the valence bond description of this molecule.
Bond Enthalpy When atoms of the hypothetical element X are placed together, they rapidly undergo reaction to form the X2 molecule: X(g)+X(g)X2(g) a Would you predict that this reaction is exothermic or endothermic? Explain. b Is the bond enthalpy of X2 a positive or a negative quantity? Why? c Suppose H for the reaction is 500 kJ/mol. Estimate the bond enthalpy of the X2 molecule. d Another hypothetical molecular compound, Y2(g), has a bond enthalpy of 750 kJ/mol, and the molecular compound XY(g) has a bond enthalpy of 1500 kJ/mol. Using bond enthalpy information, calculate H for the following reaction. X2(g)+Y2(g)2XY(g) e Given the following information, as well as the information previously presented, predict whether or not the hypothetical ionic compound AX is likely to form. In this compound, A forms the A+ cation, and X forms the X anion. Be sure to justify your answer. Reaction: A(g)+12X2(g)AX(s)The first ionization energy of A(g) is 400 kJ/mol. The electron affinity of X(g) is 525 kJ/mol. The lattice energy of AX(s) is 100 kJ/mol. f If you predicted that no ionic compound would form from the reaction in Part e, what minimum amount of AX(s) lattice energy might lead to compound formation?
Consider the following Lewis structure where E is an unknown element: What are some possible identities for element E? Predict the molecular structure (including bond angles) for this ion.
Consider the following compounds: CO2, SO2, KrF2, SO3, NF3, IF3, CF4, SF4, XeF4, PF5, IF5, and SCl6. These 12 compounds arc all examples of different molecular structures. Draw the Lewis structures for each and predict the molecular structure. Predict the bond angles and the polarity of each. (A polar molecule has a net dipole moment, while a nonpolar molecule docs not.) See Exercises 115 and 116 for the molecular structures based on the trigonal bipyramid and the octahedral geometries.
G. N. Lewis developed a model for chemical bonding that you have learned in this chapter. His theory was extremely successful and is used today at all levels of chemistry, from the introductory class to the research laboratory. Why was Lewis theory so successful?
Which of the following statements is/are true? Correct the false statements. a. It is impossible to satisfy the octet rule for all atoms in XeF2. b. Because SF4 exists, OF4 should also exist, because oxygen is in the same family as sulfur. c. The bond in No+ should be stronger than the bond in No. d. As predicted from the two Lewis structures for ozone, one oxygen-oxygen bond is stronger than the other oxygen-oxygen bond.
An important observation supporting the concept of resonance in the localized electron model was that there are only three different structures of dichlorobenzene (C6H4C10). How does this fact support the concept of resonance (see Exercise 89)?
How does the bond energy of HCl(g) differ from the standard enthalpy of formation of HCl(g)?

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