Which of the following molecules have net dipole moments

Explain the electronegativity trends across a row and down a column of the periodic table. Compare these trends with those of ionization energies and atomic radii. How are they related?

The ionic compound AB is formed. The charges on the ions may be 1, 1; 2, 2; 3, 3; or even larger. What are the factors that determine the charge for an ion in an ionic compound? 3. U

Using only the periodic table, predict the most stable ion for Na, Mg, Al, S, Cl, K, Ca, and Ga. Arrange these from largest to smallest radius, and explain why the radius varies as it does. Compare your predictions with Fig. 8.8.

The bond energy for a CH bond is about 413 kJ/mol in CH4 but 380 kJ/mol in CHBr3. Although these values are relatively close in magnitude, they are different. Explain why they are different. Does the fact that the bond energy is lower in CHBr3 make any sense? Why?

Consider the following statement: Because oxygen wants to have a negative two charge, the second electron affinity is more negative than the first. Indicate everything that is correct in this statement. Indicate everything that is incorrect. Correct the incorrect statements and explain.

Which has the greater bond lengths: NO2 or NO3 ? Explain.

The following ions are best described with resonance structures. Draw the resonance structures, and using formal charge arguments, predict the best Lewis structure for each ion. a. NCO b. CNO

Would you expect the electronegativity of titanium to be the same in the species Ti, Ti2, Ti3, and Ti4? Explain. 9

The second electron affinity values for both oxygen and sulfur are unfavorable (endothermic). Explain.

What is meant by a chemical bond? Why do atoms form bonds with each other? Why do some elements exist as molecules in nature instead of as free atoms?

Why are some bonds ionic and some covalent?

How does a bond between Na and Cl differ from a bond between C and O? What about a bond between N and N?

Arrange the following molecules from most to least polar and explain your order: CH4, CF2Cl2, CF2H2, CCl4, and CCl2H2.

Does a Lewis structure tell which electrons come from which atoms? Explain.

Compare and contrast the bonding found in the H2(g) and HF(g) molecules with that found in NaF(s).

Describe the type of bonding that exists in the Cl2(g) molecule. How does this type of bonding differ from that found in the HCl(g) molecule? How is it similar?

Some plant fertilizer compounds are (NH4)2SO4, Ca3(PO4)2, K2O, P2O5, and KCl. Which of these compounds contain both ionic and covalent bonds?

Some of the important properties of ionic compounds are as follows: i. low electrical conductivity as solids and high conductivity in solution or when molten ii. relatively high melting and boiling points iii. brittleness iv. solubility in polar solvents How does the concept of ionic bonding discussed in this chapter account for these properties?

What is the electronegativity trend? Where does hydrogen fit into the electronegativity trend for the other elements in the periodic table?

Give one example of a compound having a linear molecular structure that has an overall dipole moment (is polar) and one example that does not have an overall dipole moment (is nonpolar). Do the same for molecules that have trigonal planar and tetrahedral molecular structures.

When comparing the size of different ions, the general radii trend discussed in Chapter 7 is usually not very useful. What do you concentrate on when comparing sizes of ions to each other or when comparing the size of an ion to its neutral atom?

In general, the higher the charge on the ions in an ionic compound, the more favorable the lattice energy. Why do some stable ionic compounds have 1 charged ions even though 4, 5, and 6 charged ions would have a more favorable lattice energy? 23

Combustion reactions of fossil fuels provide most of the energy needs of the world. Why are combustion reactions of fossil fuels so exothermic?

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 oxygenoxygen bond is stronger than the other oxygen oxygen bond.

Three resonance structures can be drawn for CO2. Which resonance structure is best from a formal charge standpoint?

Which of the following statements is(are) true? Correct the false statements. a. The molecules SeS3, SeS2, PCl5, TeCl4, ICl3, and XeCl2 all exhibit at least one bond angle which is approximately 120. b. The bond angle in SO2 should be similar to the bond angle in CS2 or SCl2. c. Of the compounds CF4, KrF4, and SeF4, only SeF4 exhibits an overall dipole moment (is polar). d. Central atoms in a molecule adopt a geometry of the bonded atoms and lone pairs about the central atom in order to maximize electron repulsions.

Without using Fig. 8.3, predict the order of increasing electronegativity in each of the following groups of elements. a. C, N, O c. Si, Ge, Sn b. S, Se, Cl d. Tl, S, Ge

Without using Fig. 8.3, predict the order of increasing electronegativity in each of the following groups of elements. a. Na, K, Rb c. F, Cl, Br b. B, O, Ga d. S, O, F

Without using Fig. 8.3, predict which bond in each of the following groups will be the most polar. a. CF, SiF, GeF b. PCl or SCl c. SF, SCl, SBr d. TiCl, SiCl, GeCl

Without using Fig. 8.3, predict which bond in each of the following groups will be the most polar. a. CH, SiH, SnH b. AlBr, GaBr, InBr, TlBr c. CO or SiO d. OF or OCl

Repeat Exercises 27 and 29, this time using the values for the electronegativities of the elements given in Fig. 8.3. Are there differences in your answers?

Repeat Exercises 28 and 30, this time using the values for the electronegativities of the elements given in Fig. 8.3. Are there differences in your answers?

Which of the following incorrectly shows the bond polarity? Show the correct bond polarity for those that are incorrect. a. HF d. BrBr b. ClI e. OP c. SiS 34. Ind

Indicate the bond polarity (show the partial positive and partial negative ends) in the following bonds. a. CO d. BrTe b. PH e. SeS c. HCl

Predict the type of bond (ionic, covalent, or polar covalent) one would expect to form between the following pairs of elements. a. Rb and Cl d. Ba and S b. S and S e. N and P c. C and F f. B and H

List all the possible bonds that can occur between the elements P, Cs, O, and H. Predict the type of bond (ionic, covalent, or polar covalent) one would expect to form for each bond.

Hydrogen has an electronegativity value between boron and carbon and identical to phosphorus. With this in mind, rank the following bonds in order of decreasing polarity: PH, OH, NH, FH, CH

Rank the following bonds in order of increasing ionic character: NO, CaO, CF, BrBr, KF

Write electron configurations for the most stable ion formed by each of the elements Rb, Ba, Se, and I (when in stable ionic compounds)

Write electron configurations for the most stable ion formed by each of the elements Te, Cl, Sr, and Li (when in stable ionic compounds).

Predict the empirical formulas of the ionic compounds formed from the following pairs of elements. Name each compound. a. Li and N c. Rb and Cl b. Ga and O d. Ba and S

Predict the empirical formulas of the ionic compounds formed from the following pairs of elements. Name each compound. a. Al and Cl c. Sr and F b. Na and O d. Ca and Se

Write electron configurations for a. the cations Mg2, K, and Al3. b. the anions N3, O2, F, and Te2. 44.

Write electron configurations for a. the cations Sr2, Cs, In, and Pb2. b. the anions P3, S2, and Br. 45. W

Which of the following ions have noble gas electron configurations? a. Fe2, Fe3, Sc3, Co3 b. Tl, Te2, Cr3 c. Pu4, Ce4, Ti4 d. Ba2, Pt2, Mn2 46. What no

What noble gas has the same election configuration as each of the ions in the following compounds? a. cesium sulfide b. strontium fluoride c. calcium nitride d. aluminum bromide

Give the formula of a negative ion that would have the same number of electrons as each of the following positive ions. a. Na c. Al3 b. Ca2 d. Rb 48

Give an example of an ionic compound where both the anion and the cation are isoelectronic with each of the following noble gases. a. Ne c. Kr b. Ar d. Xe

Give three ions that are isoelectronic with neon. Place these ions in order of increasing size.

Consider the ions Sc3, Cl, K, Ca2, and S2. Match these ions to the following pictures that represent the relative sizes of the ions. 51.

For each of the following groups, place the atoms and/or ions in order of decreasing size. a. Cu, Cu, Cu2 b. Ni2, Pd2, Pt2 c. O, O, O2 d. La3, Eu3, Gd3, Yb3 e. Te2, I, Cs, Ba2, La3 52. For each o

For each of the following groups, place the atoms and/or ions in order of decreasing size. a. V, V2, V3, V5 b. Na, K, Rb, Cs c. Te2, I, Cs, Ba2 d. P, P, P2, P3 e. O2, S2, Se2, Te2 1047810_ch08_339

Which compound in each of the following pairs of ionic substances has the most exothermic lattice energy? Justify your answers. a. NaCl, KCl b. LiF, LiCl c. Mg(OH)2, MgO d. Fe(OH)2, Fe(OH)3 e. NaCl, Na2O f. MgO, BaS

Which compound in each of the following pairs of ionic substances has the most exothermic lattice energy? Justify your answers. a. LiF, CsF b. NaBr, NaI c. BaCl2, BaO d. Na2SO4, CaSO4 e. KF, K2O f. Li2O, Na2S

Use the following data to estimate Hf  for potassium chloride. Lattice energy 690. kJ/mol Ionization energy for K 419 kJ/mol Electron affinity of Cl 349 kJ/mol Bond energy of Cl2 239 kJ/mol Enthalpy of sublimation for K 64 kJ/mol

Use the following data to estimate Hf  for magnesium fluoride. Lattice energy 2913 kJ/mol First ionization energy of Mg 735 kJ/mol Second ionization energy of Mg 1445 kJ/mol Electron affinity of F 328 kJ/mol Bond energy of F2 154 kJ/mol Enthalpy of sublimation for Mg 150. kJ/mol

Consider the following energy changes:

Compare the electron affinity of fluorine to the ionization energy of sodium. Is the process of an electron being pulled from the sodium atom to the fluorine atom exothermic or endothermic? Why is NaF a stable compound? Is the overall formation of NaF endothermic or exothermic? How can this be?

LiI(s) has a heat of formation of 272 kJ/mol and a lattice energy of 753 kJ/mol. The ionization energy of Li(g) is 520. kJ/mol, the bond energy of I2(g) is 151 kJ/mol, and the electron affinity of I(g) is 295 kJ/mol. Use these data to determine the heat of sublimation of Li(s).

Use the following data to estimate H for the reaction e . Include an estimate of uncertainty. S S2 1g2 S1g2 H (kJ/mol) 735 Mg 1445 1g2 S Mg2 1g2 e Mg1g2 S Mg1g2 e Magnesium oxide exists as Mg2O2 and not as MgO. Explain. 58. Compare the electron affinity of fluorine to the ionization energy of sodium. Is the process of an electron being pulled from the sodium atom to the fluorine atom exothermic or endothermic? Why is NaF a stable compound? Is the overall formation of NaF endothermic or exothermic? How can this be? 59. LiI(s) has a heat of formation of 272 kJ/mol and a lattice energy of 753 kJ/mol. The ionization energy of Li(g) is 520. kJ/mol, the bond energy of I2(g) is 151 kJ/mol, and the electron affinity of I(g) is 295 kJ/mol. Use these data to determine the heat of sublimation of Li(s). Lattice Hsub Hf Energy I.E. of M of M Na2S 365 2203 495 109 K2S 381 2052 419 90 Rb2S 361 1949 409 82 Cs2S 360 1850 382 78 Assume that all values are known to 1 kJ/mol. 61. Rationalize the fol

Rationalize the following lattice energy values

The lattice energies of FeCl3, FeCl2, and Fe2O3 are (in no particular order) 2631, 5359, and 14,774 kJ/mol. Match the appropriate formula to each lattice energy. Explain. B

Use bond energy values (Table 8.4) to estimate H for each of the following reactions in the gas phase

Use bond energy values (Table 8.4) to estimate H for each of the following reactions

Use bond energies (Table 8.4) to predict H for the isomerization of methyl isocyanide to acetonitrile

Acetic acid is responsible for the sour taste of vinegar. It can be manufactured using the following reaction: Use tabulated values of bond energies (Table 8.4) to estimate H for this reaction

Use bond energies to predict H for the following reaction:

The major industrial source of hydrogen gas is by the following reaction: Use bond energies to predict H for this reaction.

The major industrial source of hydrogen gas is by the following reaction: Use bond energies to predict H for this reaction.

The space shuttle orbiter utilizes the oxidation of methyl hydrazine by dinitrogen tetroxide for propulsion: Use bond energies to estimate H for this reaction. The structures for the reactants are:

Consider the following reaction: Estimate the carbonfluorine bond energy given that the CC bond energy is 347 kJ/mol, the C C bond energy is 614 kJ/mol, and the FF bond energy is 154 kJ/mol.

Consider the following reaction: The bond energy for A2 is one-half the amount of the AB bond energy. The bond energy of B2 432 kJ/mol. What is the bond energy of A2?

Compare your answers from parts a and b of Exercise 63 with H values calculated for each reaction using standard enthalpies of formation in Appendix 4. Do enthalpy changes calculated from bond energies give a reasonable estimate of the actual values?

Compare your answer from Exercise 66 to the H value calculated from standard enthalpies of formation in Appendix 4. Explain any discrepancies.

The standard enthalpies of formation for S(g), F(g), SF4(g), and SF6(g) are 278.8, 79.0, 775, and 1209 kJ/mol, respectively. a. Use these data to estimate the energy of an SF bond. b. Compare your calculated value to the value given in Table 8.4. What conclusions can you draw?

Use the following standard enthalpies of formation to estimate the NH bond energy in ammonia: N(g), 472.7 kJ/mol; H(g), 216.0 kJ/mol; NH3(g), 46.1 kJ/mol. Compare your value to the one in Table 8.4

The standard enthalpy of formation for N2H4(g) is 95.4 kJ/mol. Use this and the data in Exercise 76 to estimate the NN single bond energy. Compare this with the value in Table 8.4

The standard enthalpy of formation for NO(g) is 90. kJ/mol. Use this and the values for the O O and bond energies to estimate the bond strength in NO.

Write Lewis structures that obey the octet rule (duet rule for H) for each of the following molecules. Carbon is the central atom in CH4, nitrogen is the central atom in NH3, and oxygen is the central atom in H2O. a. F2 e. NH3 b. O2 f. H2O c. CO g. HF d. CH4

Write Lewis structures that obey the octet rule (duet rule for H) for each of the following molecules. a. H2CO b. CO2 c. HCN Except for HCN and H2CO, the first atom listed is the central atom. For HCN and H2CO, carbon is the central atom. Carbon is the central atom in all of these molecules

Write Lewis structures that obey the octet rule for each of the following molecules. a. CCl4 c. SeCl2 b. NCl3 d. ICl In each case, the atom listed first is the central atom.

Write Lewis structures that obey the octet rule for each of the following molecules and ions. (In each case the first atom listed is the central atom.) a. POCl3, SO4 2, XeO4, PO4 3, ClO4  b. NF3, SO3 2, PO3 3, 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?

One type of exception to the octet rule are compounds with central atoms having fewer than eight electrons around them. BeH2 and BH3 are examples of this type of exception. Draw the Lewis structures for BeH2 and BH3.

Lewis structures can be used to understand why some molecules react in certain ways. Write the Lewis structures for the reactants and products in the reactions described below. a. Nitrogen dioxide dimerizes to produce dinitrogen tetroxide. b. Boron trihydride accepts a pair of electrons from ammonia, forming BH3NH3. Give a possible explanation for why these two reactions occur.

The most common type of exception to the octet rule are compounds or ions with central atoms having more than eight electrons around them. PF5, SF4, ClF3 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?

SF6, ClF5, and XeF4 are three compounds whose central atoms do not follow the octet rule. Draw Lewis structures for these compounds.

Write Lewis structures for the following. Show all resonance structures where applicable. a. NO2 , NO3 , N2O4 (N2O4 exists as O2NNO2.) b. OCN, SCN, N3 (Carbon is the central atom in OCN and SCN.) 88. S

Some of the important pollutants in the atmosphere are ozone (O3), sulfur dioxide, and sulfur trioxide. Write Lewis structures for these three molecules. Show all resonance structures where applicable.

Benzene (C6H6) consists of a six-membered ring of carbon atoms with one hydrogen bonded to each carbon. Write Lewis structures for benzene, including resonance structures.

Borazine (B3N3H6) has often been called inorganic benzene. Write Lewis structures for borazine. Borazine contains a sixmembered ring of alternating boron and nitrogen atoms with one hydrogen bonded to each boron and nitrogen.

An important observation supporting the concept of resonance in the localized electron model was that there are only three different structures of dichlorobenzene (C6H4Cl2). How does this fact support the concept of resonance (see Exercise 89)?

Consider the following bond lengths: In the CO3 2 ion, all three CO bonds have identical bond lengths of 136 pm. Why?

Order the following species with respect to carbonoxygen bond length (longest to shortest). What is the order from the weakest to the strongest carbon oxygen bond? (CH3OH exists as H3COH.)

Place the species below in order of the shortest to the longest nitrogenoxygen bond. (H2NOH exists as .)

Use the formal charge arguments to rationalize why BF3 would not follow the octet rule

Use formal charge arguments to explain why CO has a much smaller dipole moment than would be expected on the basis of electronegativity.

Write Lewis structures that obey the octet rule for the following species. Assign the formal charge for each central atom. a. POCl3 e. SO2Cl2 b. SO4 2 f. XeO4 c. ClO4 g. ClO3 d. PO4 3 h. NO4 3 98.

Write Lewis structures for the species in Exercise 97 that involve minimum formal charges.

Write the Lewis structure for O2F2 (O2F2 exists as FOOF). Assign oxidation states and formal charges to the atoms in O2F2. This compound is a vigorous and potent oxidizing and fluorinating agent. Are oxidation states or formal charges more useful in accounting for these properties of O2F2?

Oxidation of the cyanide ion produces the stable cyanate ion, OCN. The fulminate ion, CNO, on the other hand, is very unstable. Fulminate salts explode when struck; Hg(CNO)2 is used in blasting caps. Write the Lewis structures and assign formal charges for the cyanate and fulminate ions. Why is the fulminate ion so unstable? (C is the central atom in OCN and N is the central atom in CNO.) 10

When molten sulfur reacts with chlorine gas, a vile-smelling orange liquid forms that has an empirical formula of SCl. The structure of this compound has a formal charge of zero on all elements in the compound. Draw the Lewis structure for the vilesmelling orange liquid.

Nitrous oxide (N2O) has three possible Lewis structures: Given the following bond lengths, NN 167 pm 115 pm 120 pm NO 147 pm 110 pm rationalize the observations that the NN bond length in N2O is 112 pm and that the NO bond length is 119 pm. Assign formal charges to the resonance structures for N2O. Can you eliminate any of the resonance structures on the basis of formal charges? Is this consistent with observation?

Predict the molecular structure and bond angles for each molecule or ion in Exercises 81 and 87

Predict the molecular structure and bond angles for each molecule or ion in Exercises 82 and 88

There are several molecular structures based on the trigonal bipyramid geometry (see Table 8.8). Three such structures are Which of the compounds in Exercises 85 and 86 have these molecular structures?

Two variations of the octahedral geometry (see Table 8.6) are illustrated below. Which of the compounds in Exercises 85 and 86 have these molecular structures

Predict the molecular structure (including bond angles) for each of the following. a. SeO3 b. SeO2

Predict the molecular structure (including bond angles) for each of the following. a. SeO3 b. SeO2

Predict the molecular structure (including bond angles) for each of the following. (See Exercises 105 and 106.) a. XeCl2 b. ICl3 c. TeF4 d. PCl5

Predict the molecular structure (including bond angles) for each of the following. (See Exercises 105 and 106.) a. ICl5 b. XeCl4 c. SeCl6

Which of the molecules in Exercise 107 have net dipole moments (are polar)?

Which of the molecules in Exercise 108 have net dipole moments (are polar)?

Which of the molecules in Exercise 109 have net dipole moments (are polar)?

Which of the molecules in Exercise 110 have net dipole moments (are polar)?

Write Lewis structures and predict the molecular structures of the following. (See Exercises 105 and 106.) a. OCl2, KrF2, BeH2, SO2 b. SO3, NF3, IF3 c. CF4, SeF4, KrF4 d. IF5, AsF5 Which of these compounds are polar?

Write Lewis structures and predict whether each of the following is polar or nonpolar. a. HOCN (exists as HOCN) b. COS c. XeF2 d. CF2Cl2 e. SeF6 f. H2CO (C is the central atom.)

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 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. (See Exercises 105 and 106.)

The molecules BF3, CF4, CO2, PF5, and SF6 are all nonpolar, even though they all contain polar bonds. Why?

Two different compounds have the formula XeF2Cl2. Write Lewis structures for these two compounds, and describe how measurement of dipole moments might be used to distinguish between them.

The alkali metal ions are very important for the proper functioning of biologic systems, such as nerves and muscles, and Na and K ions are present in all body cells and fluids. In human blood plasma, the concentrations are For the fluids inside the cells, the concentrations are reversed: Since the concentrations are so different inside and outside the cells, an elaborate mechanism is needed to transport Na and K ions through the cell membranes. What are the ground-state electron configurations for Na and K? Which ion is smaller in size? Counterions also must be present in blood plasma and inside intracellular fluid. Assume the counterion present to balance the positive charge of Na and K is Cl. What is the ground-state electron configuration for Cl? Rank these three ions in order of increasing size. 122. Pr

Producing ethanol as a biofuel from the sugars in corn (glucose) has become a major industry in the grain belt. This process can be summarized as follows, where glucose is fermented to form ethanol and carbon dioxide. Use bond energies to estimate H for this reaction. C6H12O61s2 2CO21g2 2CH3CH2OH1l2

Most cars in the United States use gasohol for fuel. Gasohol is a mixture consisting of about 10% ethanol and 90% gasoline. The enthalpy of combustion per gram of gasoline is 47.8 kJ/g. Using the bond energies in Table 8.4, estimate the enthalpy of combustion per gram of ethanol. How do the two enthalpies of combustion compare with each other? The combustion reaction for ethanol is

Calcium carbonate (CaCO3) shells are used by many different types of animals to form protective coverings (including mollusks and bivalves, corals, and snails). Draw the Lewis structures for CaCO3. Be sure to include formal charges and any important resonance structures.

Peroxyacetyl nitrate, or PAN, is present in photochemical smog. Draw Lewis structures (including resonance forms) for PAN. The skeletal arrangement is

Cholesterol (C27H46O) has the following structure: In such shorthand structures, each point where lines meet represents a carbon atom, and most H atoms are not shown. Draw the complete structure showing all carbon and hydrogen atoms. (There will be four bonds to each carbon atom.) What are the predicted bond angles exhibited in cholesterol? Is cholesterol a planar molecule as indicated in the structure above?

The study of carbon-containing compounds and their properties is called organic chemistry. Besides carbon atoms, organic compounds also can contain hydrogen, oxygen, and nitrogen atoms (as well as other types of atoms). A common trait of simple organic compounds is to have Lewis structures where all atoms have a formal charge of zero. Consider the following incomplete Lewis structure for an organic compound called histidine (an amino acid), which is one of the building blocks of proteins found in our bodies: Draw a complete Lewis structure for histidine in which all atoms have a formal charge of zero. What should be the approximate bond angles about the carbon atom labeled 1 and the nitrogen atom labeled 2?

Benzoic acid is a food preservative. The space-filling model for benzoic acid is shown below. Draw the Lewis structure for benzoic acid, including all resonance structures in which all atoms have a formal charge of zero. What are the bond angles exhibited in benzoic acid? Is benzoic acid a polar or a nonpolar substance? Explain

Arrange the following in order of increasing radius and increasing ionization energy. a. N, N, N b. Se, Se, Cl, Cl c. Br, Rb, Sr2 130.

For each of the following, write an equation that corresponds to the energy given. a. lattice energy of NaCl b. lattice energy of NH4Br c. lattice energy of MgS d. double bond energy beginning with O2(g) as a reactant

Use bond energies (Table 8.4), values of electron affinities (Table 7.7), and the ionization energy of hydrogen (1312 kJ/mol) to estimate H for each of the following reactions.

Write Lewis structures for CO3 2, HCO3 , and H2CO3. When acid is added to an aqueous solution containing carbonate or bicarbonate ions, carbon dioxide gas is formed. We generally say that carbonic acid (H2CO3) is unstable. Use bond energies to estimate H for the reaction (in the gas phase) Specify a possible cause for the instability of carbonic acid.

Which member of the following pairs would you expect to be more energetically stable? Justify each choice. a. \NaBr\ , or \, NaBr_2\ b. \ClO_4 \, or \, ClO_4\ c. \SO_4 \, or \, XeO_4\ d. \OF_4 \, or \, SeF_4\

What do each of the following sets of compounds/ions have in common with each other? a. SO3, NO3 , CO3 2 b. O3, SO2, NO2 1

What do each of the following sets of compounds/ions have in common with each other? See your Lewis structures for Exercises 107 through 110. a. XeCl4, XeCl2 b. ICl5, TeF4, ICl3, PCl3, SCl2, SeO2

Although both Br3 and I3 ions are known, the F3 ion has not been observed. Explain. 1

Refer back to Exercises 97 and 98. Would you make the same prediction for the molecular structure for each case using the Lewis structure obtained in Exercise 97 as compared with the one obtained in Exercise 98?

Which of the following molecules have net dipole moments? For the molecules that are polar, indicate the polarity of each bond and the direction of the net dipole moment of the molecule. a. CH2Cl2, CHCl3, CCl4 b. CO2, N2O c. PH3, NH3

The structure of TeF5 is Draw a complete Lewis structure for TeF5 , and explain the distortion from the ideal square pyramidal structure. (See Exercise 106.)

Look up the energies for the bonds in CO and N2. Although the bond in CO is stronger, CO is considerably more reactive than N2. Give a possible explanation

Use Coulombs law, to calculate the energy of interaction for the following two arrangements of charges, each having a magnitude equal to the electron charge. a. 1 11010 m 11010 m 1 1 1 V  Q

An alternative definition of electronegativity is where I.E. is the ionization energy and E.A. is the electron affinity using the sign conventions of this book. Use data in Chapter 7 to calculate the (I.E. E.A.) term for F, Cl, Br, and I. Do these values show the same trend as the electronegativity values given in this chapter? The first ionization energies of the halogens are 1678, 1255, 1138, and 1007 kJ/mol, respectively. (Hint: Choose a constant so that the electronegativity of fluorine equals 4.0. Using this constant, calculate relative electronegativities for the other halogens and compare to values given in the text.)

Calculate the standard heat of formation of the compound ICl(g) at 25C, and show your work. (Hint: Use Table 8.4 and Appendix 4.)

Given the following information: Heat of sublimation of Li(s)  166 kJ/mol Bond energy of HCl  427 kJ/mol Ionization energy of Li(g)  520. kJ/mol Electron affinity of Cl(g)  349 kJ/mol Lattice energy of LiCl(s)  829 kJ/mol Bond energy of H2  432 kJ/mol Calculate the net change in energy for the following reaction:

Use data in this chapter (and Chapter 7) to discuss why MgO is an ionic compound but CO is not an ionic compound.

Think of forming an ionic compound as three steps (this is a simplification, as with all models): (1) removing an electron from the metal; (2) adding an electron to the nonmetal; and (3) allowing the metal cation and nonmetal anion to come together. a. What is the sign of the energy change for each of these three processes? b. In general, what is the sign of the sum of the first two processes? Use examples to support your answer. c. What must be the sign of the sum of the three processes? d. Given your answer to part c, why do ionic bonds occur? e. Given your above explanations, why is NaCl stable but not Na2Cl? NaCl2? What about MgO compared to MgO2? Mg2O?

The compound NF3 is quite stable, but NCl3 is very unstable (NCl3 was first synthesized in 1811 by P. L. Dulong, who lost three fingers and an eye studying its properties). The compounds NBr3 and NI3 are unknown, although the explosive compound NI3 NH3 is known. Account for the instability of these halides of nitrogen.

Three processes that have been used for the industrial manufacture of acrylonitrile (CH2CHCN), an important chemical used in the manufacture of plastics, synthetic rubber, and fibers, are shown below. Use bond energy values (Table 8.4) to estimate H for each of the reactions. a. b. The nitrogenoxygen bond energy in nitric oxide (NO) is 630. kJ/mol. c. d. Is the elevated temperature noted in parts b and c needed to provide energy to endothermic reactions?

The compound hexaazaisowurtzitane is the highest-energy explosive known (C & E News, Jan. 17, 1994, p. 26). The compound, also known as CL-20, was first synthesized in 1987. The method of synthesis and detailed performance data are still classified because of CL-20s potential military application in rocket boosters and in warheads of smart weapons. The structure of CL-20 is In such shorthand structures, each point where lines meet represents a carbon atom. In addition, the hydrogens attached to the carbon atoms are omitted; each of the six carbon atoms has one hydrogen atom attached. Finally, assume that the two O atoms in the NO2 groups are attached to N with one single bond and one double bond. Three possible reactions for the explosive decomposition of CL-20 are i. ii. iii. a. Use bond energies to estimate H for these three reactions. b. Which of the above reactions releases the largest amount of energy per kilogram of CL-20?

Many times extra stability is characteristic of a molecule or ion in which resonance is possible. How could this be used to explain the acidities of the following compounds? (The acidic hydrogen is marked by an asterisk.) Part c shows resonance in the C6H5 ring.

A common trait of simple organic compounds is to have Lewis structures where all atoms have a formal charge of zero. Consider the following incomplete Lewis structure for an organic compound called methyl cyanoacrylate, the main ingredient in Super Glue. Draw a complete Lewis structure for methyl cyanoacrylate in which all atoms have a formal charge of zero. What are the approximate bond angles about the carbon atom labeled 1, the carbon atom labeled 2, and the oxygen atom labeled 3?

Draw a Lewis structure for the N,N-dimethylformamide molecule. The skeletal structure is Various types of evidence lead to the conclusion that there is some double bond character to the CN bond. Draw one or more resonance structures that support this observation.

Predict the molecular structure for each of the following. (See Exercises 105 and 106.) a. BrFI2 b. XeO2F2 c. TeF2Cl3 For each formula there are at least two different structures that can be drawn using the same central atom. Draw all possible structures for each formula.

Consider the following computer-generated model of caffeine. Draw a Lewis structure for caffeine in which all atoms have a formal charge of zero.

A compound, XF5, is 42.81% fluorine by mass. Identify the element X. What is the molecular structure of XF5?

A polyatomic ion is composed of C, N, and an unknown element X. The skeletal Lewis structure of this polyatomic ion is [XCN]. The ion X2 has an electron configuration of [Ar]4s 2 3d104p6 . What is element X? Knowing the identity of X, complete the Lewis structure of the polyatomic ion, including all important resonance structures.

Identify the following elements based on their electron configurations and rank them in order of increasing electronegativity: [Ar]4s 1 3d5 ; [Ne]3s 2 3p3 ; [Ar]4s 2 3d104p3 ; [Ne]3s 2 3p5 .

Identify the five compounds of H, N, and O described below. For each compound, write a Lewis structure that is consistent with the information given. a. All the compounds are electrolytes, although not all of them are strong electrolytes. Compounds C and D are ionic and compound B is covalent. b. Nitrogen occurs in its highest possible oxidation state in compounds A and C; nitrogen occurs in its lowest possible oxidation state in compounds C, D, and E. The formal charge on both nitrogens in compound C is 1; the formal charge on the only nitrogen in compound B is 0. c. Compounds A and E exist in solution. Both solutions give off gases. Commercially available concentrated solutions of compound A are normally 16 M. The commercial, concentrated solution of compound E is 15 M. d. Commercial solutions of compound E are labeled with a misnomer that implies that a binary, gaseous compound of nitrogen and hydrogen has reacted with water to produce ammonium ions and hydroxide ions. Actually, this reaction occurs to only a slight extent. e. Compound D is 43.7% N and 50.0% O by mass. If compound D were a gas at STP, it would have a density of 2.86 g/L. f. A formula unit of compound C has one more oxygen than a formula unit of compound D. Compounds C and A have one ion in common when compound A is acting as a strong electrolyte. g. Solutions of compound C are weakly acidic; solutions of compound A are strongly acidic; solutions of compounds B and E are basic. The titration of 0.726 g compound B requires 21.98 mL of 1.000 M HCl for complete neutralization.