Which of the following statements is(are) true

Order the following bonds according to polarity: HOH, OOH, ClOH, SOH, and FOH.

We use differences in electronegativity to account for certain properties of bonds. What if all atoms had the same electronegativity values? How would bonding between atoms be affected? What are some differences we would notice?

For each of the following molecules, show the direction of the bond polarities and indicate which ones have a dipole moment: HCl, Cl2, SO3 (a planar molecule with the oxygen atoms spaced evenly around the central sulfur atom), CH4 [tetrahedral (see Table 8.2) with the carbon atom at the center], and H2S (V-shaped with the sulfur atom at the point).

Ions have different radii than their parent atoms. What if ions stayed the same size as their parent atoms? How would this affect ionic bonding in compounds?

Arrange the ions Se22, Br2, Rb1, and Sr21 in order of decreasing size.

Choose the largest ion in each of the following groups. a. Li1, Na1, K1, Rb1, Cs1 b. Ba21, Cs1, I2, Te22

Using the bond energies listed in Table 8.4, calculate DH for the reaction of methane with chlorine and fluorine to give Freon-12 (CF2Cl2). CH4 1g2 1 2Cl2 1g2 1 2F2 1g2 h CF2Cl2 1g2 1 2HF1g2 1 2HCl1g2

Give the Lewis structure for each of the following. a. HF d. CH4 b. N2 e. CF4 c. NH3 f. NO1

Write the Lewis structure for PCl5.

Write the Lewis structure for each molecule or ion. a. ClF3 b. XeO3 c. RnCl2 d. BeCl2 e. ICl4 2

Describe the electron arrangement in the nitrite anion (NO2 2) using the localized electron model.

Give possible Lewis structures for XeO3, an explosive compound of xenon. Which Lewis structure or structures are most appropriate according to the formal charges?

Describe the molecular structure of the water molecule.

You and a friend are studying for a chemistry exam. What if your friend tells you that all molecules with polar bonds are polar molecules? How would you explain to your friend that this is not correct? Provide two examples to support your

When phosphorus reacts with excess chlorine gas, the compound phosphorus pentachloride (PCl5) is formed. In the gaseous and liquid states, this substance consists of PCl5 molecules, but in the solid state it consists of a 1:1 mixture of PCl4 1 and PCl6 2 ions. Predict the geometric structures of PCl5, PCl4 1, and PCl6 2.

Because the noble gases have filled s and p valence orbitals, they were not expected to be chemically reactive. In fact, for many years these elements were called inert gases because of this supposed inability to form any compounds. However, in the early 1960s several compounds of krypton, xenon, and radon were synthesized. For example, a team at the Argonne National Laboratory produced the stable colorless compound xenon tetrafluoride (XeF4). Predict its structure and whether it has a dipole moment.

Predict the molecular structure of the sulfur dioxide molecule. Is this molecule expected to have a dipole moment?

Distinguish between the terms electronegativity versus electron affinity, covalent bond versus ionic bond, and pure covalent bond versus polar covalent bond. Characterize the types of bonds in terms of electronegativity difference. Energetically, why do ionic and covalent bonds form?

When an element forms an anion, what happens to the radius? When an element forms a cation, what happens to the radius? Why? Define the term isoelectronic. When comparing sizes of ions, which ion has the largest radius and which ion has the smallest radius in an isoelectronic series? Why?

Define the term lattice energy. Why, energetically, do ionic compounds form? Figure 8.11 illustrates the energy changes involved in the formation of MgO(s) and NaF(s). Why is the lattice energy of MgO(s) so different from that of NaF(s)? Magnesium oxide is composed of Mg21 and O22 ions. Energetically, why does Mg21O22 form and not Mg1O2? Why doesnt Mg31O32 form?

Explain how bond energies can be used to estimate DH for a reaction. Why is this an estimate of DH? How do the product bond strengths compare to the reactant bond strengths for an exothermic reaction? For an endothermic reaction? What is the relationship between the number of bonds between two atoms and bond strength? Bond length?

Give a rationale for the octet rule and the duet rule for H in terms of orbitals. Give the steps for drawing a Lewis structure for a molecule or ion. In general, molecules and ions always follow the octet rule unless it is impossible. The three types of exceptions are molecules/ions with too few electrons, molecules/ions with an odd number of electrons, and molecules/ions with too many electrons. Which atoms sometimes have fewer than 8 electrons around them? Give an example. Which atoms sometimes have more than 8 electrons around them? Give some examples. Why are oddelectron species generally very reactive and uncommon? Give an example of an odd-electron molecule

Explain the terms resonance and delocalized electrons. When a substance exhibits resonance, we say that none of the individual Lewis structures accurately portrays the bonding in the substance. Why do we draw resonance structures?

Define formal charge and explain how to calculate it. What is the purpose of the formal charge? Organic compounds are composed mostly of carbon and hydrogen, but also may have oxygen, nitrogen, and/or halogens in the formula. Formal charge arguments work very well for organic compounds when drawing the best Lewis structure. How do C, H, N, O, and Cl satisfy the octet rule in organic compounds so as to have a formula charge of zero?

Explain the main postulate of the VSEPR model. List the five base geometries (along with bond angles) that most molecules or ions adopt to minimize electron-pair repulsions. Why are bond angles sometimes slightly less than predicted in actual molecules as compared to what is predicted by the VSEPR model?

Give two requirements that should be satisfied for a molecule to be polar. Explain why CF4 and XeF4 are nonpolar compounds (have no net dipole moments) while SF4 is polar (has a net dipole moment). Is CO2 polar? What about COS? Explain.

Consider the following compounds: CO2, SO2, KrF2, SO3, NF3, IF3, CF4, SF4, XeF4, PF5, IF5, and SCl6. These 12 compounds are 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 does not.) See Exercises 111 and 112 for the molecular structures based on the trigonal bipyramid and the octahedral geometries.

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 11, 21; 12, 22; 13, 23; or even larger. What are the factors that determine the charge for an ion in an ionic compound?

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 COH 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 2 or NO3 2? 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. NCO2 b. CNO

Would you expect the electronegativity of titanium to be the same in the species Ti, Ti21, Ti31, and Ti41? Explain.

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

The following electrostatic potential diagrams represent H2, HCl, or NaCl. Label each and explain your choices.

Describe the type of bonding that exists in the F2(g) molecule. How does this type of bonding differ from that found in the HF(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 11 charged ions even though 14, 15, and 16 charged ions would have a more favorable lattice energy?

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 NO1 should be stronger than the bond in NO2. 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 1208. 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. COF, SiOF, GeOF b. POCl or SOCl c. SOF, SOCl, SOBr d. TiOCl, SiOCl, GeOCl

Without using Fig. 8.3, predict which bond in each of the following groups will be the most polar. a. COH, SiOH, SnOH b. AlOBr, GaOBr, InOBr, TlOBr c. COO or SiOO d. OOF or OOCl

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. d1HOFd2 d. d1BrOBrd2 b. d1ClOId2 e. d1OOPd2 c. d1SiOSd2

Indicate the bond polarity (show the partial positive and partial negative ends) in the following bonds. a. COO d. BrOTe b. POH e. SeOS c. HOCl

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: POH, OOH, NOH, FOH, COH.

Rank the following bonds in order of increasing ionic character: NOO, CaOO, COF, BrOBr, KOF.

State whether or not each of the following has a permanent dipole moment

The following electrostatic potential diagrams represent CH4, NH3, or H2O. Label each and explain your choices.

Write electron configurations for the most stable ion formed by each of the elements Al, 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 Mg21, K1, and Al31. b. the anions N32, O22, F2, and Te22.

Write electron configurations for a. the cations Sr21, Cs1, In1, and Pb21. b. the anions P32, S22, and Br2

Which of the following ions have noble gas electron configurations? a. Fe21, Fe31, Sc31, Co31 b. Tl1, Te22, Cr31 c. Pu41, Ce41, Ti41 d. Ba21, Pt21, Mn21

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. Na1 c. Al31 b. Ca21 d. Rb1

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 Sc31, Cl2, K1, Ca21, and S22. Match these ions to the following pictures that represent the relative sizes of the ions.

For each of the following groups, place the atoms and/or ions in order of decreasing size. a. Cu, Cu1, Cu21 b. Ni21, Pd21, Pt21 c. O, O2, O22 d. La31, Eu31, Gd31, Yb31 e. Te22, I2, Cs1, Ba21, La31

For each of the following groups, place the atoms and/or ions in order of decreasing size. a. V, V21, V31, V51 b. Na1, K1, Rb1, Cs1 c. Te22, I2, Cs1, Ba21 d. P, P2, P22, P32 e. O22, S22, Se22, Te2

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 DHf 8 for potassium chloride. K1s2 1 1 2Cl2 1g2 h KCl1s2 Lattice energy 2690. kJ/mol Ionization energy for K 419 kJ/mol Electron affinity of Cl 2349 kJ/mol Bond energy of Cl2 239 kJ/mol Enthalpy of sublimation for K 90. kJ/mol

Use the following data to estimate DH8f for magnesium fluoride. Mg1s2 1 F2 1g2 h MgF2 1s2 Lattice energy 22913 kJ/mol First ionization energy of Mg 735 kJ/mol Second ionization energy of Mg 1445 kJ/mol Electron affinity of F 2328 kJ/mol Bond energy of F2 154 kJ/mol Enthalpy of sublimation for Mg 150. kJ/mol

Consider the following energy changes: DH (kJ/mol) Mg1g2 S Mg1 1g2 1 e2 735 Mg1 1g2 S Mg21 1g2 1 e2 1445 O1g2 1 e2 S O2 1g2 2141 O2 1g2 1 e2 S O22 1g2 878 Magnesium oxide exists as Mg21O22 and not as Mg1O2. Explain

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?

Consider the following: Li(s) 1 1 2 I2(g) n LiI(s) DH 5 2292 kJ. LiI(s) has a lattice energy of 2753 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 2295 kJ/mol. Use these data to determine the heat of sublimation of Li(s).

Use the following data (in kJ/mol) to estimate DH for the reaction S2(g) 1 e2 n S22(g). Include an estimate of uncertainty. DHf 8 Lattice Energy Ionization Energy of M DHsub of M Na2S 2365 22203 495 109 K2S 2381 22052 419 90 Rb2S 2361 21949 409 82 Cs2S 2360 21850 382 78 S1s2 h S1g2 DH 5 277 kJ/mol S1g2 1 e2 h S2 1g2 DH 5 2200 kJ/mol Assume that all values are known to 61 kJ/mol.

Rationalize the following lattice energy values: Compound Lattice Energy (kJ/mol) CaSe 22862 Na2Se 22130 CaTe 22721 Na2Te 22095

The lattice energies of FeCl3, FeCl2, and Fe2O3 are (in no particular order) 22631, 25359, and 214,774 kJ/mol. Match the appropriate formula to each lattice energy. Explain.

Use bond energy values (Table 8.4) to estimate DH for each of the following reactions in the gas phase. a. H2 1 Cl2 S 2HCl b. N{N 1 3H2 S 2NH3

Use bond energy values (Table 8.4) to estimate DH for each of the following reactions. a. H C N(g) 2H2(g) H H H H H + C N(g) b. N N (g) + 2F2(g) N N(g) + 4HF(g

Use bond energies (Table 8.4) to predict DH for the isomerization of methyl isocyanide to acetonitrile: CH3N{C1g2 h CH3C{N1g2

Acetic acid is responsible for the sour taste of vinegar. It can be manufactured using the following reaction: CH3OH(g) + C O(g) CH OH(l) 3C O Use tabulated values of bond energies (Table 8.4) to estimate DH for this reaction.

Use bond energies to predict DH for the following reaction: H2S1g2 1 3F2 1g2 h SF4 1g2 1 2HF1g

The major industrial source of hydrogen gas is by the following reaction: CH4 1g2 1 H2O1g2 h CO1g2 1 3H2 1g2 Use bond energies to predict DH for this reaction.

Use bond energies to estimate DH for the combustion of one mole of acetylene: C2H2 1g2 1 5 2O2 1g2 h 2CO2 1g2 1 H2O1g2

Use data from Table 8.4 to estimate DH for the combustion of methane (CH4), as shown below:

Consider the following reaction: H H H H C C (g) + F2(g) H H(g) H = 549 kJ F F H H C C Estimate the carbonfluorine bond energy given that the COC bond energy is 347 kJ/mol, the CPC bond energy is 614 kJ/mol, and the FOF bond energy is 154 kJ/mol.

Consider the following reaction: A2 1 B2 h 2AB DH 5 2285 kJ The bond energy for A2 is one-half the amount of the AB bond energy. The bond energy of B2 5 432 kJ/mol. What is the bond energy of A2?

Compare your answers from parts a and b of Exercise 65 with DH 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 68 to the DH 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 1278.8, 179.0, 2775, and 21209 kJ/mol, respectively. a. Use these data to estimate the energy of an SOF bond. b. Compare your calculated value to the value given in Table 8.4. What conclusions can you draw? c. Why are the DHf 8 values for S(g) and F(g) not equal to zero, since sulfur and fluorine are elements?

Use the following standard enthalpies of formation to estimate the NOH bond energy in ammonia: N(g), 472.7 kJ/mol; H(g), 216.0 kJ/mol; NH3(g), 246.1 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 78 to estimate the NON 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 OPO and NqN 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 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 22, XeO4, PO4 32, ClO4 2 b. NF3, SO3 22, PO3 32, ClO3 2 c. ClO2 2, SCl2, PCl2 2 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 2 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 2, NO3 2, N2O4 (N2O4 exists as O2NONO2.) b. OCN2, SCN2, N3 2 (Carbon is the central atom in OCN2 and SCN2.)

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 91.)

Consider the following bond lengths: CiO 143 pm CwO 123 pm C{O 109 pm In the CO3 22 ion, all three COO bonds have identical bond lengths of 136 pm. Why?

A toxic cloud covered Bhopal, India, in December 1984 when water leaked into a tank of methyl isocyanate, and the product escaped into the atmosphere. Methyl isocyanate is used in the production of many pesticides. Draw the Lewis structures for methyl isocyanate, CH3NCO, including resonance forms. The skeletal structure is H C H H C N O

Peroxyacetyl nitrate, or PAN, is present in photochemical smog. Draw Lewis structures (including resonance forms) for PAN. The skeletal structure is H O H C C O O O O H N

Order the following species with respect to carbonoxygen bond length (longest to shortest). CO, CO2, CO3 22, CH3OH What is the order from the weakest to the strongest carbon oxygen bond? (CH3OH exists as H3COOH.)

Place the species below in order of the shortest to the longest nitrogenoxygen bond. H2NOH, N2O, NO1, NO2 2, NO3 2 (H2NOH exists as H2NiOH.)

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 22 f. XeO4 c. ClO4 2 g. ClO3 2 d. PO4 32 h. NO4

Write Lewis structures for the species in Exercise 101 that involve minimum formal charges.

Write the Lewis structure for O2F2 (O2F2 exists as FOOOOOF). 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, OCN2. The fulminate ion, CNO2, 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 OCN2 and N is the central atom in CNO2.)

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 vile-smelling orange liquid.

Nitrous oxide (N2O) has three possible Lewis structures: N N O N O N N O N Given the following bond lengths, NiN 167 pm NwO 115 pm NwN 120 pm NiO 147 pm N{N 110 pm rationalize the observations that the NON bond length in N2O is 112 pm and that the NOO 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?

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. H C O H C C N O C H H C H Draw a complete Lewis structure for methyl cyanoacrylate in which all atoms have a formal charge of zero

Benzoic acid is a food preservative. The space-filling model for benzoic acid is shown below. Benzoic acid (C6H5CO2H) C O H Draw the Lewis structure for benzoic acid, including all resonance structures in which all atoms have a formal charge of zero.

Predict the molecular structure and bond angles for each molecule or ion in Exercises 83 and 89.

Predict the molecular structure and bond angles for each molecule or ion in Exercises 84 and 90.

There are several molecular structures based on the trigonal bipyramid geometry (see Table 8.8). Three such structures are A B A 180 A B A A A 120 A B A A 90 90 90 90 Linear T-shaped See-saw Which of the compounds in Exercises 87 and 88 have these molecular structures?

Two variations of the octahedral geometry (see Table 8.6) are illustrated below. AA AA B 90 90 AA AA A B 90 90 90 Square planar Square pyramid Which of the compounds in Exercises 87 and 88 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. PCl3 b. SCl2 c. SiF4

Predict the molecular structure (including bond angles) for each of the following. (See Exercises 111 and 112.) a. XeCl2 b. ICl3 c. TeF4 d. PCl

Predict the molecular structure (including bond angles) for each of the following. (See Exercises 111 and 112.) a. ICl5 b. XeCl4 c. SeCl6

Which of the molecules in Exercise 113 have net dipole moments (are polar)?

Which of the molecules in Exercise 114 have net dipole moments (are polar)?

Which of the molecules in Exercise 115 have net dipole moments (are polar)?

hich of the molecules in Exercise 116 have net dipole moments (are polar)?

Write Lewis structures and predict the molecular structures of the following. (See Exercises 111 and 112.) 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 HOOCN) 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: O E O O 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: 2 F E O F What are some possible identities for element E? Predict the molecular structure (including bond angles) for this ion. (See Exercises 111 and 112.)

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.

Arrange the following in order of increasing radius and increasing ionization energy. a. N1, N, N2 b. Se, Se2, Cl, Cl1 c. Br2, Rb1, Sr21

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. OwO 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 DH for each of the following reactions. a. HF1g2 S H1 1g2 1 F2 1g2 b. HCl1g2 S H1 1g2 1 Cl2 1g2 c. HI1g2 S H1 1g2 1 I 2 1g2 d. H2O1g2 S H1 1g2 1 OH2 1g2 (Electron affinity of OH(g) 5 2180. kJ/mol.)

Write Lewis structures for CO3 22, HCO3 2, 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 DH for the reaction (in the gas phase) H2CO3 h CO2 1 H2O 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 NaBr2 b. ClO4 or ClO4 2 c. SO4 or XeO4 d. OF4 or SeF

What do each of the following sets of compounds/ions have in common with each other? a. SO3, NO3 2, CO3 22 b. O3, SO2, NO2 2

What do each of the following sets of compounds/ions have in common with each other? See your Lewis structures for Exercises 113 through 116. a. XeCl4, XeCl2 b. ICl5, TeF4, ICl3, PCl3, SCl2, SeO2

Although both Br3 2 and I3 2 ions are known, the F3 2 ion has not been observed. Explain

Refer back to Exercises 101 and 102. Would you make the same prediction for the molecular structure for each case using the Lewis structure obtained in Exercise 101 as compared with the one obtained in Exercise 102?

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 2 is F F F F F Te 79  Draw a complete Lewis structure for TeF5 2, and explain the distortion from the ideal square pyramidal structure. (See Exercise 112.)

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.

Classify the bonding in each of the following molecules as ionic, polar covalent, or nonpolar covalent. a. H2 e. HF b. K3P f. CCl4 c. NaI g. CF4 d. SO2 h. K2S

List the bonds POCl, POF, OOF, and SiOF from least polar to most polar

Arrange the atoms and/or ions in the following groups in order of decreasing size. a. O, O2, O22 b. Fe21, Ni21, Zn21 c. Ca21, K1, Cl2

Use the following data to estimate DHf 8 for barium bromide. Ba1s2 1 Br2 1g2 h BaBr2 1s2 Lattice energy 21985 kJ/mol First ionization energy of Ba 503 kJ/mol Second ionization energy of Ba 965 kJ/mol Electron affinity of Br 2325 kJ/mol Bond energy of Br2 193 kJ/mol Enthalpy of sublimation of Ba 178 kJ/mol

Use bond energy values to estimate DH for the following gas phase reaction: C2H4 1 H2O2 h CH2OHCH2OH

Which of the following compounds or ions exhibit resonance? a. O3 d. CO3 22 b. CNO2 d. AsF3 c. AsI3

The formulas of several chemical substances are given in the table below. For each substance in the table, give its chemical name and predict its molecular structure. Formula Compound Name Molecular Structure CO2 NH3 SO3 H2O ClO4 2

Predict the molecular structure, bond angles, and polarity (has a net dipole moment or has no net dipole moment) for each of the following compounds. a. SeCl4 d. CBr4 b. SO2 e. IF3 c. KrF4 f. ClF5

Use Coulombs law, V 5 Q1Q2 4pP0r 5 2.31 3 10219 J # nm a Q1Q2 r b to calculate the energy of interaction for the following two arrangements of charges, each having a magnitude equal to the electron charge. a. +1 1 +1 1 1 1010 m 1 1010 m b. +1 +1 1 1010 m 1 1010 m 1 1010 m 1 1010 m

An alternative definition of electronegativity is Electronegativity 5 constant 1I.E. 2 E.A.2 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. 2 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 258C. (Hint: Use Table 8.4 and Appendix 4 data.)

Given the following information: Heat of sublimation of Li(s) 5 166 kJ/mol Bond energy of HCl 5 427 kJ/mol Ionization energy of Li(g) 5 520. kJ/mol Electron affinity of Cl(g) 5 2349 kJ/mol Lattice energy of LiCl(s) 5 2829 kJ/mol Bond energy of H2 5 432 kJ/mol Calculate the net change in energy for the following reaction: 2Li1s2 1 2HCl1g2 h 2LiCl1s2 1 H2 1g2

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 DH for each of the reactions. a. CH2 CH2 HCN O HOC C N H H H H + C C N + H H H HOCH2CH2CN C C H2O b. 4CH2wCHCH3 1 6NO h700C Ag 4CH2wCHCN 1 6H2O 1 N2 The nitrogenoxygen bond energy in nitric oxide (NO) is 630. kJ/mol. c. 2CH2wCHCH3 1 2NH3 1 3O2 88888n 2CH2wCHCN 1 6H2O d. Is the elevated temperature noted in parts b and c needed to provide energy to endothermic reactions?

The compound hexaazaisowurtzitane is one of the highestenergy explosives 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 O2N O2N N O2N NO2 NO2 NO2 N N CL-20 N N N 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. C6H6N12O12 1s2 S 6CO1g2 1 6N2 1g2 1 3H2O1g2 1 3 2O2 1g2 ii. C6H6N12O12 1s2 S3CO1g2 1 3CO2 1g2 1 6N2 1g2 1 3H2O1g2 iii. C6H6N12O12 1s2 S 6CO2 1g2 1 6N2 1g2 1 3H2 1g a. Use bond energies to estimate DH 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. H c. O C OH* OH* OH* b. O CH3 C CH CH C 3

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: H H H N H H N C H H C H C C N C O C O H 2 1 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?

Draw a Lewis structure for the N,N-dimethylformamide molecule. The skeletal structure is H CH3 CH3 O C N Various types of evidence lead to the conclusion that there is some double bond character to the CON bond. Draw one or more resonance structures that support this observation

Predict the molecular structure for each of the following. (See Exercises 111 and 112.) a. BrFI2 b. XeO2F2 c. TeF2Cl3 2 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. H O N C 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 [XOCON]2. The ion X22 has an electron configuration of [Ar]4s2 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]4s1 3d5 ; [Ne]3s2 3p3 ; [Ar]4s2 3d104p3 ; [Ne]3s2 3p5 .

Identify the five compounds of H, N, and O described as follows. 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 11; 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