Although !3- is known, f:J- is not. Using Lewis structures, explain why F3- does not

For each of these Lewis symbols, indicate the group in the periodic table in which the element X belongs: (a) (b) X (c) :

lilustrated at right are four ions-A, B, X, and Yshowing their relative ionic radii. The ions shown in red carry positive charges: a 2+ charge for A and a 1 + charge for B. Ions shown in blue carry negative charges: a 1 - charge for X and a 2- charge for Y. (a) Which combinations of these ions produce ionic compounds where there is a 1: 1 ratio of cations and anions? (b) Among those compounds, which combination of ions leads to the ionic compound having the largest lattice energy? (c) Which combination of ions leads to the ionic cornpound having the smallest lattice energy? [Section 8.2]

The orbital diagram below shows the valence electrons for a 2+ ion of an element. (a) What is the element? (b) What is the electron configuration of an atom of this element? [Section 8.2]

In the Lewis structure shown below, A, D, E, Q, X, and Z represent elements in the first two rows of the periodic table (H-Ne). Identify all six elements so that the formal charges of all atoms are zero. [Section 8.3]

The partial Lewis structure below is for a hydrocarbon molecule. In the full Lewis structure, each carbon atom satisfies the octet rule, and there are no unshared electron pairs in the molecule. The carbon-carbon bonds are labeled 1, 2, and 3. (a) Determine where the hydrogen atoms are in the molecule. (b) Rank the carbon-carbon bonds in order of increasing bond length. (c) Rank the carbon-carbon bonds in order of increasing bond enthalpy. [Section 8.3 and 8.8] cc2..c c

Consider the Lewis structure for the polyatomic oxyanion shown below, where X is an element from the 3rd period (Na-Ar). By changing the overall charge, n, from 1 - to 2- to 3- we get three ctifferent polyatomic ions. For each of these ions (a) Identify the central atom, X. (b) Determine the formal charge of the central atom, X . (c) Draw a Lewis structure that makes the formal charge on the central atom equal to zero. (d) If the Lewis structure you drew in part (c) differs from the Lewis structure shown below, which one do you think is the best one (if you were able to choose only a single Lewis structure)? [Sections 8.5, 8.6, and 8.7]

(a) What are valence electrons? (b) How many valence electrons does a nitrogen atom possess? (c) An atom has the electron configuration ls2 2s2 2p6 3s2 3p2 . How many valence electrons does the atom have?

(a) What is the octet rule? (b) How many electrons must a sulfur atom gain to achieve an octet in its valence shell? (c) If an atom has the electron configuration ls2 2s2 2p3 , how many electrons must it gain to achieve an octet?

Write the electron configuration for phosphorus. Identify the valence electrons in this configuration and the nonvalence electrons. From the standpoint of chemical reactivity, what is the important difference between them?

Write the Lewis symbol for atoms of each of the following elements: (a) AI, (b) Br, (c) Ar, (d) Sr.

What is the Lewis symbol for each of the following atoms or ions: (a) Ca, (b) P, (c) Mg2+, (d) s 2-?

Using Lewis symbols, diagram the reaction between magnesium and oxygen atoms to give the ionic substance MgO.

Use Lewis symbols to represent the reaction that occurs between Ca and F atoms.

Use Lewis symbols to represent the reaction that occurs between Ca and F atoms.

Which ionic compound is expected to form from combining the following pairs of elements: (a) barium and fluorine, (b) cesium and chlorine, (c) lithium and nitrogen, (d) aluminum and oxygen?

Write the electron configuration for each of the following ions, and determine which ones possess noble-gas configurations: (a) Sr2 +, (b) Ti2 +, (c) Se2-, (d) Ni2 +, (e) Br-, (f) Mn3+.

Write electron configurations for the following ions, and determine which have noble-gas configurations: (a) zn 2 +, (b) Te2-, (c) Sc3+, (d) Rh3+, (e) n+, (f) Bi3+

(a) Define the term lattice energy. (b) Which factors govern the magnitude of the lattice energy of an ionic compound?

The ionic substances KF, CaO, and SeN are isoelectronic (they have the same number of electrons). Examine the lattice energies for these substances in Table 8.2, and account for the trends you observe.

(a) Does the lattice energy of an ionic solid increase or decrease (i) as the charges of the ions increase, (ii) as the sizes of the ions increase? (b) Using a periodic table, arrange the following substances according to their expected lattice energies, listing them from lowest lattice energy to the highest: SeN, KBr, MgO, NaF. Compare your list with the data in Table 8.2.

The lattice energies of KBr and CsCI are nearly equal (Table 8.2). What can you conclude from this observation?

Explain the following trends in lattice energy: (a) CaF2 > BaF2; (b) NaCI > RbBr > CsBr; (c) BaO > KF.

Energy is required to remove two electrons from Ca to form Ca 2+ and is required to add two electrons to 0 to form 0 2-. Why, then, is CaO stable relative to the free elements?

Construct a Born-Haber cycle for the formation of the hypothetical compound NaClz, where the sodium ion has a 2+ charge (the 2nd ionization energy for sodium is given in Table 7.2). (a) How large would the lattice energy need to be for the formation of N aC12 to be exothermic? (b) If we were to estimate the lattice energy of NaC12 to be roughly equal to that of MgCI2 (2326 kJ/mol ___ from Table 8.2), what value would you obtain for the standard enthalpy of formation, !lHJ, of NaC12?

Use data from Appendix C, Figure 7.12, and Figure 7.14 to calculate the lattice energy of RbCI. Is this value greater than or less than the lattice energy of NaCI? Explain.

(a) Based on the lattice energies of MgCI2 and SrCI2 given in Table 8.2, what is the range of values that you would expect for the lattice energy of CaCI2? (b) Using data from Appendix C, Figure 7.12, and Figure 7.14 and the value of the second ionization energy for Ca, 1145 kJ/mol, calculate the lattice energy of CaCI2.

(a) What is meant by the term covalent bo11d? (b) Give three examples of covalent bonding. (c) A substance XY, formed from two different elements, boils at -33 oc. Is XY likely to be a covalent or an ionic substance? Explain.

Using Lewis symbols and Lewis structures, diagram the formation of SiC14 from Si and Cl atoms.

Use Lewis symbols and Lewis structures to diagram the formation of PF3 from P and F atoms.

(a) Construct a Lewis structure for 02 in which each atom achieves an octet of electrons. (b) Explain why it is necessary to form a double bond in the Lewis structure. (c) The bond in 02 is shorter than the 0-0 bond in compounds that contain an 0-0 single bond. Explain this observation.

(a) Construct a Lewis structure for hydrogen peroxide, H20z, in which each atom achieves an octet of electrons. (b) Do you expect the 0-0 bond in H202 to be longer or shorter than the 0-0 bond in 02?

(a) What is meant by the term electronegativity? (b) On the Pauling scale what is the range of electronegativity values for the elements? (c) Which element has the greatest electronegativity? (d) Which element has the smallest electronegativity?

(a) What is the trend in electronegativity going from left to right in a row of the periodic table? (b) How do electronegativity values generally vary going down a column in the periodic table? (c) How do periodic trends in ___ electronegativity relate to those for ionization energy and electron affinity?

Using only the periodic table as your guide, select the most electronegative atom in each of the following sets: (a) Se, Rb, 0, In; (b) Al, Ca, C, Si; (c) Ge, As, P, Sn; (d) Li, Rb, Be, Sr.

By referring only to the periodic table, select (a) the most electronegative element in group 6A; (b) the least electronegative element in the group AI, Si, P; (c) the most electronegative element in the group Ga, P, Cl, Na; (d) the element in the group K, C, Zn, F, that is most likely to form an ionic compound with Ba.

Which of the following bonds are polar: (a) B-F, (b) Cl-Cl, (c) Se-0, (d) H-I? Which is the more electronegative atom in each polar bond?

The dipole moment and bond distance measured for the highly reactive gas phase OH molecule are 1 .78 D and 0.98 A, respectively. (a) Given these values calculate the effective charges on the H and 0 atoms of the OH molecule in units of the electronic charges e. (b) Is this bond more or less polar than the H- Cl bond in an HCI molecule? (c) Is that what you would have expected based on electronegativities?

The iodine monobromide molecule, IBr, has a bond length of 2.49 A and a dipole moment of 1.21 D. (a) Which atom of the molecule is expected to have a negative charge? Explain. (b) Calculate the effective charges on the I and Br atoms in IBr, in units of the electronic charge e.

In the following pairs of binary compounds determine which one is a molecular substance and which one is an ionic substance. Use the appropriate naming convention (for ionic or molecular substances) to assign a name to each compound: (a) SiF4 and LaF3, (b) FeC12 and ReCI6, (c) PbCI4 and RbCI.

In the following pairs of binary compounds determine which one is a molecular substance and which one is an ionic substance. Use the appropriate naming convention (for ionic or molecular substances) to assign a name to each compound: (a) TICI4 and CaFz, (b) ClF3 and VF3, (c) SbCls and AlF3.

Draw Lewis structures for the following: (a) SiH4, (b) CO, (c) SF:u (d) H2S04 (H is bonded to 0), (e) Cl02- (f) NH20H.

Write Lewis structures for the following: (a) H2CO (both H atoms are bonded to C), (b) Hp2, (c) C2F6 (contains a ___ C-C bond), (d) Asol-, (e) H2S03 (H is bonded to 0), (f) C2H2.

(a) When talking about atoms in a Lewis structure, what is meant by the term formal charge? (b) Does the formal charge of an atom represent the actual charge on that atom? Explain. (c) How does the formal charge of an atom in a Lewis structure differ from the oxidation number of the atom?

(a) Write a Lewis structure for the phosphorus trifluoride molecule, PF3. Is the octet rule satisfied for all the atoms in your structure? (b) Determine the oxidation numbers of the P and F atoms. (c) Determine the formal charges of the P and F atoms. (d) Is the oxidation num ___ ber for the P atom the same as its formal charge? Explain why or why not.

Write Lewis structures that obey the octet rule for each of the following, and assign oxidation numbers and formal charges to each atom: (a) NO+, (b) POC13 (P is bonded to the three Cl atoms and to the 0), (c) Cl04 -, (d) HC103 (H is bonded to 0).

(a) Write one or more appropriate Lewis structures for the nitrite ion, N02 -. (b) With what allotrope of oxygen is it isoelectronic? (c) What would you predict for the lengths of the bonds in N02- relative to N-O single bonds?

Consider the nitryl cation, N02 +. (a) Write one or more appropriate Lewis structures for this ion. (b) Are resonance structures needed to describe the structure? (c) With what familiar molecule is it isoelectronic?

Predict the ordering of the C-0 bond lengths in CO, CO:u and co/-.

Based on Lewis structures, predict the ordering of _ _ N -0 bond lengths in NO+, N02 -, and N03 -.

(a) Use the concept of resonance to explain why all six C -C bonds in benzene are equal in length. (b) Are the c-c bond lengths in benzene shorter than c-c single bonds? Are they shorter than C=C double bonds?

Mothballs are composed of naphthalene, C10H8, a molecule of which consists of two six-membered rings of carbon fused along an edge, as shown in this incomplete Lewis structure: (a) Write two complete Lewis structures for naphthalene. (b) The observed C-C bond lengths in the molecule are intermediate between C-C single and C=C double bonds. Explain. (c) Represent the resonance in naphthalene in a way analogous to that used to represent it in benzene.

(a) State the octet rule. (b) Does the octet rule apply to ionic as well as to covalent compounds? Explain, using examples as appropriate.

Considering the nonmetals, what is the relationship between the group number for an element (carbon, for example, belongs to group 4A; see the periodic table on the inside front cover) and the number of single covalent bonds that element needs to form to conform to the octet rule?

What is the most common exception to the octet rule? Give two examples.

Draw the Lewis structures for each of the following ions or molecules. Identify those that do not obey the octet rule, and explain why they do not. (a) S032 -, (b) AlH3, (c) N3 -, (d) CHzCl:u (e) SbFs.

Draw the Lewis structures for each of the following molecules or ions. Which do not obey the octet rule? (a) NH4 +, (b) SCN-, (c) PC13, (d) TeF4, (e) XeF2.

In the vapor phase, BeC12 exists as a discrete molecule. (a) Draw the Lewis structure of this molecule, using only single bonds. Does this Lewis structure satisfy the octet rule? (b) What other resonance forms are possible that satisfy the octet rule? (c) Using formal charges, select the resonance form from among all the Lewis structures that is most important in describing BeC12.

(a) Describe the molecule chlorine dioxide, ClO:u using three possible resonance structures. (b) Do any of these resonance structures satisfy the octet rule for every atom in the molecule? Why or why not? (c) Using formal charges, select the resonance structure(s) that is (are) most important.

Using the bond enthalpies tabulated in Table 8.4, estimate t.H for each of the following gas-phase reactions:

Using bond enthalpies (Table 8.4), estimate t.H for the following gas-phase reactions:

Using bond enthalpies (Table 8.4), estimate t.H for each of the following reactions: (a) 2 CH4(g) + 02(g) ----> 2 CH30H(g) (b) Hz(g) + Brz(g) ----> 2 HBr(g) (c) 2 Hz02(g) ----> 2 H20(g) + 02(g)

Use bond enthalpies (Table 8.4) to estimate the enthalpy change for each of the following reactions: (a) C3H8(g) + 5 02(g) ----> 3 C02(g) + 4 H20(g) (b) C2H50H(g) + 3 02(g) ----> 2 C02(g) + 3 H20(g) (c) 8 H2S(g) ----> 8 Hz(g) + Ss(s)

Ammonia is produced directly from nitrogen and hydrogen by using the Haber process. The chemical reaction is (a) Use bond enthalpies (Table 8.4) to estimate the enthalpy change for the reaction, and tell whether this reaction is exothermic or endothermic. (b) Compare the enthalpy change you calculate in (a) to the true enthalpy change as obtained using llHJ values.

Given the following bond-dissociation energies, calculate the average bond enthalpy for the Ti -Cl bond.

(a) Using average bond enthalpies, predict which of the following reactions will be most exothermic: (i) C(g) + 2 F2(g) ----> CF4(g) (ii) CO(g) + 3 F2 ----> CF4(g) + OFz(g) (iii) C02(g) + 4 F2 ----> CF4(g) + 2 OF2(g) (b) Explain the trend, if any, that exists between reaction exothermicity and the extent to which the carbon atom is bonded to oxygen.

How many elements in the periodic table are represented by a Lewis symbol with a single dot? Are all these elements in the same group? Explain.

(a) Explain the following trend in lattice energy: BeH2, 3205 kJ/mol; MgH2, 2791 kJ/mol; CaHz, 2410 kJ/mol; SrHz, 2250 kJ/mol; BaHz, 2121 k)/mol. (b) The lattice energy of ZnH2 is 2870 kJ/mol. Based on the data given in part (a), the radius of the Zn2+ ion is expected to be closest to that of which group 2A element?

Based on data in Table 8.2, estimate (within 30 kJ/mol) the lattice energy for (a) LiBr, (b) CsBr, (c) CaC12.

Would you expect AlN to have a lattice energy that is larger or smaller than SeN? Explain.

From the ionic radii given in Figure 7.8, calculate the potential energy of a Ca2 + and 0 2 - ion pair that are just touching (the magnitude of the electronic charge is given on the back inside cover). Calculate the energy of a mole of such pairs. How does this value compare with the lattice energy of CaO (fable 8.2)? Explain the dHference.

From Equation 8.4 and the ionic radii given in Figure 7.8, calculate the potential energy of the following pairs of ions. Assume that the ions are separated by a ctistance equal to the sum of their ionic radii: (a) Na+, Br-; (b) Rb+, Br-; (c) Sr2 +, s 2 -

(a) How does a polar molecule diller from a nonpolar one? (b) Atoms X and Y have dHferent electronegativities. Will the diatomic molecule X -Y necessarily be polar? Explain. (c) What factors affect the size of the dipole moment of a diatomic molecule?

To address energy and environmental issues, there is great interest in powering vehicles with hydrogen rather than gasoline. One of the most attractive aspects of the "hydrogen economy" is the fact that in principle the only emission would be water. However, two daunting obstacles must be overcome before this vision can become a reality. First, an economical method of producing hydrogen must be found. Second, a safe, lightweight, and compact way of storing hydrogen must be found. The hydrides of light metals are attractive for hydrogen storage because they can store a high weight percentage of hydrogen in a small volume. One of the most attractive hydrides is NaAlH4, which can release 5.6% of its mass as H2 upon decomposing to NaH(s), AJ(s), and H2(g). NaAJH4 possesses both covalent bonds, which hold polyatomic anions together, and ionic bonds. (a) Write a balanced equation for the decomposition of Na. (b) Which element in Na is the most electronegative? Which one is the least electronegative? (c) Based on electronegativity differences, what do you think is the identity of the polyatomic anion? Draw a Lewis structure for this ion.

For the following collection of nonmetallic elements, 0, P, Te, I, B, (a) which two would form the most polar single bond? (b) Whlch two would form the longest single bond? (c) Which two would be likely to form a compound of formula XY2? (d) Which combinations of elements would likely yield a compound of empirical formula XzY3? In each case explain your answer.

You and a partner are asked to complete a lab entitled "Oxides of Ruthenium" that is scheduled to extend over two lab periods. The first lab, which is to be completed by your partner, is devoted to carrying out compositional analysis. In the second lab, you are to determine melting points. Upon going to lab you find two unlabeled vials, one containing a soft yellow substance and the other a black powder. You also find the following notes in your partner's notebook-Compound 1: 76.0% Ru and 24.0% 0 (by mass), Compound 2: 61.2% Ru and 38.8% 0 (by mass). (a) What is the empirical formula for Compound 1? (b) What is the empirical formula for Compound 2? (c) Upon determining the melting points of these two compounds, you find that the yellow compound melts at 25 oc, while the black powder does not melt up to the maximum temperature of your apparatus, 1200 oc. What is the identity of the yellow compound? What is the identity of the black compound? Be sure to use the appropriate naming convention depending upon whether the compound is better described as a molecular or ionic compound.

You and a partner are asked to complete a lab entitled "Fluorides of Group 6B Metals" that is scheduled to extend over two lab periods. The first lab, which is to be completed by your partner, is devoted to carrying out compositional analysis. In the second lab, you are to determine melting points. Upon going to lab you find two unlabeled vials, one containing a colorless liquid and the other a green powder. You also find the following notes in your partner's notebook-Compound 1: 47.7% Cr and 52.3% F (by mass), Compound 2: 45.7% Mo and 54.3% F (by mass). (a) What is the empirical formula for Compound 1? (b) What is the empirical formula for Compound 2? (c) Upon determining the melting points of these two compounds you find that the colorless liquid solidifies at 18 C, while the green powder does not melt up to the maximum temperature of your apparatus, 1200 oc. What is the identity of the colorless liquid? What is the identity of the green powder? Be sure to use the appropriate naming convention depending upon whether the compound is better described as a molecular or ionic compound.

(a) Triazine, C3H3N3, is like benzene except that in triazine every other C -H group is replaced by a nitrogen atom. Draw the Lewis structure(s) for the triazine molecule. (b) Estimate the carbon-nitrogen bond distances in the ring.

Using the electronegativities of Br and Cl, estimate the partial charges on the atoms in the Br-Cl molecule. Using these partial charges and the atomic radii given in Figure 7.7, estimate the dipole moment of the molecule. The measured dipole moment is 0.57 D.

Although !3- is known, f:J- is not. Using Lewis structures, explain why F3- does not form.

Calculate the formal charge on the indicated atom in each of the following molecules or ions: (a) the central oxygen atom in 03, (b) phosphorus in PF6 -, (c) nitrogen in NOz, (d) iodine in JCJ3, (e) chlorine in HCI04 (hydrogen is bonded to 0).

(a) Determine the formal charge on the chlorine atom in the hypochlorite ion, ClO-, and the perchlorate ion, Cl04 -, using resonance structures where the Cl atom has an octet. (b) What are the oxidation numbers of chlorine in oo- and in Cl04 -? (c) Is it uncommon for the formal charge and the oxidation state to be different? Explain. (d) Perchlorate is a much stronger oxidizing agent than hypochlorite. Would you expect there to be any relationship between the oxidizing power of the oxyanion and either the oxidation state or the formal charge of chlorine?

An important reaction for the conversion of natural gas to other useful hydrocarbons is the conversion of methane to ethane. 2 CH4(g) ----> CzH6(g) + Hz(g) In practice, this reaction is carried out in the presence of oxygen, which converts the hydrogen produced to water. 2 CH4(g) + ! Oz(g) ----> CzH6(g) + HzO(g) Use bond enthalpies (Table 8.4) to estimate !l.H for these two reactions. Why is the conversion of methane to ethane more favorable when oxygen is used?

Two compounds are isomers if they have the same chemical formula but a different arrangement of atoms. Use bond enthalpies (Table 8.4) to estimate !l.H for each of the following gas-phase isomerization reactions, and indicate which isomer has the lower enthalpy:

With reference to the "Chemistry Put to Work" box on explosives, (a) use bond enthalpies to estimate the enthalpy change for the explosion of 1.00 g of nitroglycerin. (b) Write a balanced equation for the decomposition of TNT. Assume that, upon explosion, TNT decomposes into Nz(g), COz(g), HzO(g), and C(s).

The "plastic" explosive C-4, often used in action movies, contains the molecule cyclotrimethylenetrinitramine, which is often called RDX (for Royal Demolition eXplosive): o _.....o N H I H H-c

The bond lengths of carbon-carbon, carbon-nitrogen, carbon-oxygen, and nitrogen-nitrogen single, double, and triple bonds are listed in Table 8.5. Plot bond enthalpy (Table 8.4) versus bond length for these bonds. What do you conclude about the relationship between bond length and bond enthalpy? What do you conclude about the relative strengths of C -C, C-N, C-0, and N- N bonds?

Use the data in Table 8.5 and the following data: S-S distance in Ss = 2.05 A; s-o distance in SOz = 1.43 A to answer the following questions: (a) Predict the bond length in a S-N single bond. (b) Predict the bond length in a s-o single bond. (c) Why is the s-o bond length in S02 considerably shorter than your predicted value for the S-O single bond? (d) When elemental sulfur, 5s, is carefully oxidized, a compound SsO is formed, in which one of the sulfur atoms in the S8 ring is bonded to an oxygen atom. The S-O bond length in this compound is 1.48 A. In light of this information, write Lewis structures that can account for the observed S-0 bond length. Does the sulfur bearing the oxygen in this compound obey the octet rule?

The Ti2+ ion is isoelectronic with the Ca atom. (a) Are there any differences in the electron configurations of Ti2+ and Ca? (b) With reference to Figure 6.25, comment on the changes in the ordering of the 4s and 3d subshells in Ca and Ti2+ (c) Will Ca and Ti2+ have the same number of unpaired electrons? Explain.

(a) Write the chemical equations that are used in calculating the lattice energy of SrC12(s) via a Bom-Haber cycle. (b) The second ionization energy of Sr(g) is 1064 kJ/mol. Use this fact along with data in Appendix C, Figure 7.12, Figure 7.14, and Table 8.2 to calculate !l.HJ for SrC!2(s).

The electron affinity of oxygen is -141 kJ/mol, corresponding to the reaction O(g) + e- --> o-(g) The lattice energy of K20(s) is 2238 kJ/mol. Use these data along with data in Appendix C and Figure 7.12 to calculate the "second electron affinity" of oxygen, corresponding to the reaction o-(g) + e- - o2 -(g)

One scale for electronegativity is based on the concept that the electronegativity of any atom is proportional to the ionization energy of the atom minus its electron affinity: electronegativity = k(IE - EA), where k is a proportionality constant. (a) How does this definition explain why the electronegativity of F is greater than that of Cl even though Cl has the greater electron affinity? (b) Why are both ionization energy and electron affinity relevant to the notion of electronegativity? (c) By using data in Chapter 7, determine the value of k that would lead to an electronegativity of 4.0 for F under this definition. (d) Use your result from part (c) to determine the electronegativities of Cl and 0 using this scale. Do these values follow the trend shown in Figure 8.6?

The compound chloral hydrate, known in detective stories as knockout drops, is composed of 14.52% C, 1.83% H, 64.30% Cl, and 19.35% 0 by mass and has a molar mass of 165.4 g/mol. (a) What is the empirical formula of this substance? (b) What is the molecular formula of this substance? (c) Draw the Lewis structure of the molecule, assuming that the Cl atoms bond to a single C atom and that there are a C-C bond and two c-o bonds in the compound.

Barium azide is 62.04% Ba and 37.96% N. Each azide ion has a net charge of 1-. (a) Determine the chemical formula of the azide ion. (b) Write three resonance structures for the azide ion. (c) Which structure is most important? (d) Predict the bond lengths in the ion.

Acetylene (C2H2) and nitrogen (N2) both contain a triple bond, but they differ greatly in their chemical properties. (a) Write the Lewis structures for the two substances. (b) By referring to Appendix C, look up the enthalpies of formation of acetylene and nitrogen and compare their reactivities. (c) Write balanced chemical equations for the complete oxidation of N2 to form N205(g) and of acetylene to form C02(g) and H20(g). (d) Calculate the enthalpy of oxidation per mole of N2 and C2H2 (the enthalpy of formation of N205(g) is 11.30 kJ/mol). How do these comparative values relate to your response to part (b)? Both N2 and C2H2 possess triple bonds with quite high bond enthalpies (Table 8.4). What aspect of chemical bonding in these molecules or in the oxidation products seems to account for the difference in chemical reactivities?

Under special conditions, sulfur reacts with anhydrous liquid ammonia to form a binary compound of sulfur and nitrogen. The compound is found to consist of 69.6% S and 30.4% N. Measurements of its molecular mass yield a value of 184.3 g mol-1 . The compound occasionally detonates on being struck or when heated rapidly. The sulfur and nitrogen atoms of the molecule are joined in a ring. All the bonds in the ring are of the same length. (a) Calculate the empirical and molecular formulas for the substance. (b) Write Lewis structures for the molecule, based on the information you are given. (Hint: You should find a relatively small number of dominant Lewis structures.) (c) Predict the bond distances between the atoms in the ring. (Note: The S-S distance in the 58 ring is 2.05A.) (d) The enthalpy of formation of the compound is estimated to be 480 kJ mol-1 . t!.HJ of S(g) is 222.8 kJ mol-1 Estimate the average bond enthalpy in the compound.

A common form of elemental phosphorus is the tetrahedral P 4 molecule, where all four phosphorus atoms are equivalent: At room temperature phosphorus is a solid. (a) Do you think there are any unshared pairs of electrons in the P 4 molecule? (b) How many P-P bonds are there in the molecule? (c) Can you draw a Lewis structure for a linear P 4 molecule that satisfies the octet rule? (d) Using formal charges, what can you say about the stability of the linear molecule vs. that of the tetrahedral molecule?

Consider benzene (C6H6) in the gas phase. (a) Write the reaction for breaking all the bonds in C6H6(g), and use data in Appendix C to determine the enthalpy change for this reaction. (b) Write a reaction that corresponds to breaking all the carbon--<:arbon bonds in C6H6(g). (c) By combining your answers to parts (a) and (b) and using the average bond enthalpy for C -H from Table 8.4, calculate the average bond enthalpy for the carbon--<:arbon bonds in C6H6(g). (d) Comment on your answer from part (c) as compared to the values for C-C single bonds and C=C double bonds in Table 8.4.

Average bond enthalpies are generally defined for gasphase molecules. Many substances are liquids in their standard state. =(Section 5.7) By using appropriate thermochemical data from Appendix C, calculate average bond enthalpies in the liquid state for the following bonds, and compare these values to the gas-phase values given in Table 8.4: (a) Br-Br, from Br2(1); (b) C -Cl, from CC14(1); (c) 0-0, from Hz02(1) (assume that the 0-H bond enthalpy is the same as in the gas phase). (d) What can you conclude about the process of breaking bonds in the liquid as compared to the gas phase? Explain the difference in the llH values between the two phases.