22.1 (a) One of these structures is a stable compound; the other is not. Identify the stable compound, and explain why it is stable. Explain why the other compound is not stable. (b) What is the geometry around the central atoms of the stable compound? [Section 22.1]

22.2 (a) Identify the type of chemical reaction represented by the following diagram. (b) Place appropriate charges on the species on both sides of the equation. (c) Write the chemical equation for the reaction. [Section 22.1]

22.3 Which of the following species (there may be more than one) is/are likely to have the structure shown here: (a) XeF4, (b) BrF4+(c) SiF4(d) TeCl4(e) HClO4? (The colors do not reflect atom identities.) [Sections 22.3, 22.4, 22.6, and 22.10]

22.4 You have two glass bottles, one containing oxygen and one filled with ozone. How could you determine which one is which? [Section 22.5]

22.5 Write the molecular formula and Lewis structure for each of the following oxides of nitrogen: [Section 22.7]

22.6 Which property of the group 6A elements might be the one depicted in the graph shown here: (a) electronegativity, (b) first ionization energy, (c) density, (d) X — X single-bond enthalpy, (e) electron affinity? Explain your answer. [Sections 22.5 and 22.6]

22.7 The atomic and ionic radii of the first three group 6A elements are

(a) Explain why the atomic radius increases in moving downward in the group. (b) Explain why the ionic radii are larger than the atomic radii. (c) Which of the three anions would you expect to be the strongest base in water? Explain. [Sections 22.5 and 22.6]

22.8 Which property of the third-row nonmetallic elements might be the one depicted in the graph on the next page: (a) first ionization energy, (b) atomic radius, (c) electronegativity, (d) melting point, (e) X — X single-bond enthalpy? Explain both your choice and why the other choices would not be correct. [Sections 22.3, 22.4, 22.6, 22.8, and 22.10]

22.9 Which of the following compounds would you expect to be the most generally reactive, and why? [Section 22.8]

22.10 (a) Draw the Lewis structures for at least four species that have the general formula

where X and Y may be the same or different, and n may have a value from +1 to –2. (b) Which of the compounds is likely to be the strongest Brønsted base? Explain. [Sections 22.1, 22.7, and 22.9]


22.11 Identify each of the following elements as a metal, nonmetal, or metalloid: (a) phosphorus, (b) strontium, (c) manganese, (d) selenium, (e) sodium, (f) krypton.

22.12 Identify each of the following elements as a metal, nonmetal, or metalloid: (a) gallium, (b) molybdenum, (c) tellurium, (d) arsenic, (e) xenon, (f) ruthenium.

22.13 Consider the elements O, Ba, Co, Be, Br, and Se. From this list select the element that (a) is most electronegative, (b) exhibits a maximum oxidation state of +7, (c) loses an electron most readily, (d) forms π bonds most readily, (e) is a transition metal, (f) is a liquid at room temperature and pressure.

22.14 Consider the elements Li, K, Cl, C, Ne, and Ar. From this list select the element that (a) is most electronegative, (b) has the greatest metallic character, (c) most readily forms a positive ion, (d) has the smallest atomic radius, (e) forms π bonds most readily, (f) has multiple allotropes.

22.15 Explain the following observations: (a) The highest fluoride compound formed by nitrogen is NF3, whereas phosphorus readily forms PF5(b) Although CO is a well-known compound, SiO does not exist under ordinary conditions. (c) AsH3 is a stronger reducing agent than NH3.

22.16 Explain the following observations: (a) HNO3 is a stronger oxidizing agent than H3PO4(b) Silicon can form an ion with six fluorine atoms, SiF62–, whereas carbon is able to bond to a maximum of four, CF4. (c) There are three compounds formed by carbon and hydrogen that contain two carbon atoms each (C2H2, C2H4, and C2H6), whereas silicon forms only one analogous compound (Si2H6).

22.17 Complete and balance the following equations:

22.18 Complete and balance the following equations:


22.19 (a) Give the names and chemical symbols for the three isotopes of hydrogen. (b) List the isotopes in order of decreasing natural abundance. (c) Which hydrogen isotope is radioactive? (d) Write the nuclear equation for the radioactive decay of this isotope.

22.20 Are the physical properties of H2O different from D2O? Explain.

22.21 Give a reason why hydrogen might be placed along with the group 1A elements of the periodic table.

22.22 What does hydrogen have in common with the halogens? Explain.

22.23 Write a balanced equation for the preparation of H2 using (a) Mg and an acid, (b) carbon and steam, (c) methane and steam.

22.24 List (a) three commercial means of producing H2(b) three industrial uses of H2.

22.25 Complete and balance the following equations:

22.26 Write balanced equations for each of the following reactions (some of these are analogous to reactions shown in the chapter). (a) Aluminum metal reacts with acids to form hydrogen gas. (b) Steam reacts with magnesium metal to give magnesium oxide and hydrogen. (c)Manganese(IV) oxide is reduced to manganese(II) oxide by hydrogen gas. (d) Calcium hydride reacts with water to generate hydrogen gas.

22.27 Identify the following hydrides as ionic, metallic, or molecular: (a) BaH2(b) H2Te, (c) TiH1.7.

22.28 Identify the following hydrides as ionic, metallic, or molecular: (a) B2H6(b) RbH, (c) TH4H1.5.

22.29 Describe two characteristics of hydrogen that are favorable for its use as a general energy source in vehicles.

22.30 The H2/O2 fuel cell converts elemental hydrogen and oxygen into water, producing, theoretically, 1.23 V of energy. What is the most sustainable way to obtain hydrogen to run a large number of fuel cells? Explain.

22.31 Why does xenon form stable compounds with fluorine, whereas argon does not?

22.32 A friend tells you that the “neon” in neon signs is a compound of neon and aluminum. Can your friend be correct? Explain.

22.33 Write the chemical formula for each of the following, and indicate the oxidation state of the halogen or noble-gas atom in each: (a) calcium hypobromite, (b) bromic acid, (c) xenon trioxide, (d) perchlorate ion, (e) iodous acid, (f) iodine pentafluoride.

22.34 Write the chemical formula for each of the following compounds, and indicate the oxidation state of the halogen or noble-gas atom in each: (a) chlorate ion, (b) hydroiodic acid, (c) iodine trichloride, (d) sodium hypochlorite, (e) perchloric acid, (f) xenon tetrafluoride.

22.35 Name the following compounds and assign oxidation states to the halogens in them: (a) Fe(ClO3)3(b) HClO2(c) XeF6(d) BrF5(e) XeOF4(f) HIO3.

22.36 Name the following compounds and assign oxidation states to the halogens in them: (a) KC lO3(b) Ca(IO3)2(c) AlCl3(d) HBrO3(e) H5IO6(f) XeF4.

22.37 Explain each of the following observations: (a) At room temperature I2 is a solid, Br2 is a liquid, and Cl2 and F2 are both gases. (b) F2 cannot be prepared by electrolytic oxidation of aqueous F solutions. (c) The boiling point of HF is much higher than those of the other hydrogen halides. (d) The halogens decrease in oxidizing power in the order F2 > Cl2 > Br2 > I2.

22.38 Explain the following observations: (a) For a given oxidation state, the acid strength of the oxyacid in aqueous solution decreases in the order chlorine > bromine > iodine. (b) Hydrofluoric acid cannot be stored in glass bottles. (c) HI cannot be prepared by treating NaI with sulfuric acid. (d) The inter-halogen ICl3 is known, but BrCl3 is not.

OXYGEN AND THE OTHER GROUP 6A ELEMENTS (sections 22.5 and 22.6)

22.39 Write balanced equations for each of the following reactions. (a) When mercury(II) oxide is heated, it decomposes to form O2 and mercury metal. (b) When copper(II) nitrate is heated strongly, it decomposes to form copper(II) oxide, nitrogen dioxide, and oxygen. (c) Lead(II) sulfide, PbS(s), reacts with ozone to form PbSO4(s) and O2(g). (d) When heated in air, ZnS(s) is converted to ZnO. (e) Potassium peroxide reacts with CO2(g) to give potassium carbonate and O2(f) Oxygen is converted to ozone in the upper atmosphere.

22.40 Complete and balance the following equations:

22.41 Predict whether each of the following oxides is acidic, basic, amphoteric, or neutral: (a) NO2(b) CO2(c) Al2O3(d) CaO.

22.42 Select the more acidic member of each of the following pairs: (a) MN2O7 and MnO2(b) SnO and SnO2(c) SO2 and SO3(d) SiO2 and SO2(e) Ga2O3 and IN2O3(f) SO2 and SeO2.

22.43 Write the chemical formula for each of the following compounds, and indicate the oxidation state of the group 6A element in each: (a) selenous acid, (b) potassium hydrogen sulfite, (c) hydrogen telluride, (d) carbon disulfide, (e) calcium sulfate, (f) cadmium sulfide, (g) zinc telluride.

22.44 Write the chemical formula for each of the following compounds, and indicate the oxidation state of the group 6A element in each: (a) sulfur tetrachloride, (b) selenium trioxide, (c) sodium thiosulfate, (d) hydrogen sulfide, (e) sulfuric acid, (f) sulfur dioxide, (g) mercury telluride.

22.45 In aqueous solution, hydrogen sulfide reduces (a) Fe3+ to Fe2+(b) Br2 to Br(c) MnO4 to Mn2+(d) HNO3 to NO2. In all cases, under appropriate conditions, the product is elemental sulfur. Write a balanced net ionic equation for each reaction.

22.46 An aqueous solution of SO2 reduces (a) aqueous KMnO4 to MnSO4(aq), (b) acidic aqueous K2Cr2O7 to aqueous Cr3+(c) aqueous Hg2(NO3)2 to mercury metal. Write balanced equations for these reactions.

22.47 Write the Lewis structure for each of the following species, and indicate the structure of each: (a) SeO32–(b) S2Cl2(c) chlorosulfonic acid, HSO3Cl (chlorine is bonded to sulfur).

22.48 The SF5 ion is formed when SF4(g) reacts with fluoride salts containing large cations, such as CsF(s). Draw the Lewis structures for SF4 and SF5, and predict the molecular structure of each.

22.49 Write a balanced equation for each of the following reactions: (a) Sulfur dioxide reacts with water. (b) Solid zinc sulfide reacts with hydrochloric acid. (c) Elemental sulfur reacts with sulfite ion to form thiosulfate. (d) Sulfur trioxide is dissolved in sulfuric acid.

22.50 Write a balanced equation for each of the following reactions. (You may have to guess at one or more of the reaction products, but you should be able to make a reasonable guess, based on your study of this chapter.) (a) Hydrogen selenide can be prepared by reaction of an aqueous acid solution on aluminum selenide. (b) Sodium thiosulfate is used to remove excess Cl2 from chlorine-bleached fabrics. The thiosulfate ion forms SO42– and elemental sulfur, while Cl2 is reduced to Cl.


22.51 Write the chemical formula for each of the following compounds, and indicate the oxidation state of nitrogen in each: (a) sodium nitrite, (b) ammonia, (c) nitrous oxide, (d) sodium cyanide, (e) nitric acid, (f) nitrogen dioxide, (g) nitrogen, (h) boron nitride.

22.52 Write the chemical formula for each of the following compounds, and indicate the oxidation state of nitrogen in each: (a) nitric oxide, (b) hydrazine, (c) potassium cyanide, (d) sodium nitrite, (e) ammonium chloride, (f) lithium nitride.

22.53 Write the Lewis structure for each of the following species, describe its geometry, and indicate the oxidation state of the nitrogen: (a) HNO2(b) N3(c) N2H5+(d) NO3.

22.54 Write the Lewis structure for each of the following species, describe its geometry, and indicate the oxidation state of the nitrogen: (a) NH4+(b) NO2(c) N2O, (d) NO2.

22.55 Complete and balance the following equations:

Which ones of these are redox reactions?

22.56 Write a balanced net ionic equation for each of the following reactions: (a) Dilute nitric acid reacts with zinc metal with formation of nitrous oxide. (b) Concentrated nitric acid reacts with sulfur with formation of nitrogen dioxide. (c) Concentrated nitric acid oxidizes sulfur dioxide with formation of nitric oxide. (d) Hydrazine is burned in excess fluorine gas, forming NF3(e) Hydrazine reduces CrO42– to Cr(OH)4 in base (hydrazine is oxidized to N2).

22.57 Write complete balanced half-reactions for (a) oxidation of nitrous acid to nitrate ion in acidic solution, (b) oxidation of N2 to N2O in acidic solution.

22.58 Write complete balanced half-reactions for (a) reduction of nitrate ion to NO in acidic solution, (b) oxidation of HNO2 to NO2 in acidic solution.

22.59 Write a molecular formula for each compound, and indicate the oxidation state of the group 5A element in each formula: (a) phosphorous acid, (b) pyrophosphoric acid, (c) antimony trichloride, (d) magnesium arsenide, (e) diphosphorus pen-toxide, (f) sodium phosphate.

22.60 Write a chemical formula for each compound or ion, and indicate the oxidation state of the group 5A element in each formula: (a) phosphate ion, (b) arsenous acid, (c) antimony(III) sulfide, (d) calcium dihydrogen phosphate, (e) potassium phosphide, (f) gallium arsenide.

22.61 Account for the following observations: (a) Phosphorus forms a pentachloride, but nitrogen does not. (b) H3PO2 is a monoprotic acid. (c) Phosphonium salts, such as PH4Cl, can be formed under anhydrous conditions, but they can't be made in aqueous solution. (d) White phosphorus is more reactive than red phosphorus.

22.62 Account for the following observations: (a) H3PO3 is a diprotic acid. (b) Nitric acid is a strong acid, whereas phosphoric acid is weak. (c) Phosphate rock is ineffective as a phosphate fertilizer. (d) Phosphorus does not exist at room temperature as diatomic molecules, but nitrogen does. (e) Solutions of Na3PO4 are quite basic.

22.63 Write a balanced equation for each of the following reactions:

(a) preparation of white phosphorus from calcium phosphate,

(b) hydrolysis of PBr3(c) reduction of PBr3 to P4 in the gas phase, using H2.

22.64 Write a balanced equation for each of the following reactions: (a) hydrolysis of PCl5(b) dehydration of phosphoric acid (also called orthophosphoric acid) to form pyrophosphoric acid, (c) reaction of P4O10 with water.

CARBON, THE OTHER GROUP 4A ELEMENTS, AND BORON (sections 22.922.1022.11)

22.65 Give the chemical formula for (a) hydrocyanic acid, (b) nickel tetracarbonyl, (c) barium bicarbonate, (d) calcium acetylide (e) potassium carbonate.

22.66 Give the chemical formula for (a) carbonic acid, (b) sodium cyanide, (c) potassium hydrogen carbonate, (d) acetylene, (e) iron pentacarbonyl.

22.67 Complete and balance the following equations:

22.68 Complete and balance the following equations:

22.69 Write a balanced equation for each of the following reactions: (a) Hydrogen cyanide is formed commercially by passing a mixture of methane, ammonia, and air over a catalyst at 800 °C. Water is a by-product of the reaction. (b) Baking soda reacts with acids to produce carbon dioxide gas. (c) When barium carbonate reacts in air with sulfur dioxide, barium sulfate and carbon dioxide form.

22.70 Write a balanced equation for each of the following reactions: (a) Burning magnesium metal in a carbon dioxide atmosphere reduces the CO2 to carbon. (b) In photosynthesis, solar energy is used to produce glucose (C6H12O6) and O2 from carbon dioxide and water. (c) When carbonate salts dissolve in water, they produce basic solutions.

22.71 Write the formulas for the following compounds, and indicate the oxidation state of the group 4A element or of boron in each: (a) boric acid, (b) silicon tetrabromide, (c) lead (II) chloride, (d) sodium tetraborate decahydrate (borax), (e) boric oxide, (f) germanium dioxide.

22.72 Write the formulas for the following compounds, and indicate the oxidation state of the group 4A element or of boron in each: (a) silicon dioxide, (b) germanium tetrachloride, (c) sodium borohydride, (d) stannous chloride, (e) diborane, (f) boron trichloride.

22.73 Select the member of group 4A that best fits each description: (a) has the lowest first ionization energy, (b) is found in oxidation states ranging from –4 to +4, (c) is most abundant in Earth's crust.

22.74 Select the member of group 4A that best fits each description: (a) forms chains to the greatest extent, (b) forms the most basic oxide, (c) is a metalloid that can form 2+ ions.

22.75 (a) What is the characteristic geometry about silicon in all silicate minerals? (b) Metasilicic acid has the empirical formula H2SiO3. Which of the structures shown in Figure 22.34 would you expect metasilicic acid to have?

22.76 Speculate as to why carbon forms carbonate rather than silicate analogs.

22.77 (a) How does the structure of diborane (B2H6) differ from that of ethane (C2H6)? (b) Explain why diborane adopts the geometry that it does. (c) What is the significance of the statement that the hydrogen atoms in diborane are described as “hydridic"?

22.78 Write a balanced equation for each of the following reactions: (a) Diborane reacts with water to form boric acid and molecular hydrogen. (b) Upon heating, boric acid undergoes a condensation reaction to form tetraboric acid. (c) Boron oxide dissolves in water to give a solution of boric acid.


22.79 In your own words, define the following terms: (a) allotrope, (b) disproportionation, (c) interhalogen, (d) acidic anhydride, (e) condensation reaction, (f) protium.

22.80 Although the ClO4 and IO4 ions have been known for a long time, BrO4 was not synthesized until 1965. The ion was synthesized by oxidizing the bromate ion with xenon difluoride, producing xenon, hydrofluoric acid, and the perbromate ion. (a) Write the balanced equation for this reaction. (b) What are the oxidation states of Br in the Br-containing species in this reaction?

22.81 Write a balanced equation for the reaction of each of the following compounds with water: (a) SO2(g), (b) Cl2O7(g), (c) Na2O2(s), (d) BaC2(s), (e) RbO2(s), (f) Mg3N2(s), (g) NaH(s).

22.82 What is the anhydride for each of the following acids: (a) H2SO4(b) HClO3(c) HNO2(d) H2CO3(e) H3PO4?

22.83 Explain why SO2 can be used as a reducing agent but SO3 cannot.

22.84 A sulfuric acid plant produces a considerable amount of heat. This heat is used to generate electricity, which helps reduce operating costs. The synthesis of H2SO4 consists of three main chemical processes: (1) oxidation of S to SO2, (2) oxidation of SO2 to SO3, (3) the dissolving of SO3 in H2SO4 and its reaction with water to form H2SO4. If the third process produces 130 kJ/mol, how much heat is produced in preparing a mole of H2SO4 from a mole of S? How much heat is produced in preparing 5000 pounds of H2SO4?

22.85 (a) What is the oxidation state of P in PO43– and of N in NO3(b) Why doesn't N form a stable NO43– ion analogous to P?

22.86 (a) The P4, P4O6, and P4O10 molecules have a common structural feature of four P atoms arranged in a tetrahedron (Figures 22.27 and 22.28). Does this mean that the bonding between the P atoms is the same in all these cases? Explain. (b) Sodium trimetaphosphate (Na3P3O9) and sodium tetra-metaphosphate (Na4P4O12) are used as water-softening agents. They contain cyclic P3O93– and P4O124– ions, respectively. Propose reasonable structures for these ions.

22.87 Ultrapure germanium, like silicon, is used in semiconductors. Germanium of “ordinary” purity is prepared by the high-temperature reduction of GeO2 with carbon. The Ge is converted to GeCl4 by treatment with Cl2 and then purified by distillation; GeCl4 is then hydrolyzed in water to GeO2 and reduced to the elemental form with H2. The element is then zone refined. Write a balanced chemical equation for each of the chemical transformations in the course of forming ultra-pure Ge from GeO2.

22.88 Hydrogen peroxide is capable of oxidizing (a) hydrazine to N2 and H2O, (b) SO2 to SO42–(c) NO2 to NO3(d) H2S(g) to S(s), (e) Fe2+ to Fe3+. Write a balanced net ionic equation for each of these redox reactions.


[22.89] (a) How many grams of H2 can be stored in 100.0 kg of the alloy FeTi if the hydride FeTiH2 is formed? (b) What volume does this quantity of H2 occupy at STP? (c) If this quantity of hydrogen was combusted in air to produce liquid water, how much energy could be produced?

[22.90] Using the thermochemical data in Table 22.1 and Appendix C, calculate the average Xe — F bond enthalpies in XeF2, XeF4, and XeF6, respectively. What is the significance of the trend in these quantities?

22.91 Hydrogen gas has a higher fuel value than natural gas on a mass basis but not on a volume basis. Thus, hydrogen is not competitive with natural gas as a fuel transported long distances through pipelines. Calculate the heats of combustion of H2 and CH4 (the principal component of natural gas) (a) per mole of each, (b) per gram of each, (c) per cubic meter of each at STP. Assume H2O(l) as a product.

22.92 The solubility of Cl2 in 100 g of water at STP is 310 cm3. Assume that this quantity of Cl2 is dissolved and equilibrated as follows:

(a) If the equilibrium constant for this reaction is 4.7 × 10-4, calculate the equilibrium concentration of HClO formed. (b) What is the pH of the final solution?

[22.93] When ammonium perchlorate decomposes thermally, the products of the reaction are N2(g), O2(g), H2O(g), and HCl(g). (a) Write a balanced equation for the reaction. [Hint: You might find it easier to use fractional coefficients for the products.] (b) Calculate the enthalpy change in the reaction per mole of NH4ClO4. The standard enthalpy of formation of NH4ClO4(s) is –295.8 kJ. (c) When NH4ClO4(s) is employed in solid-fuel booster rockets, it is packed with powdered aluminum. Given the high temperature needed for NH4ClO4(s) decomposition and what the products of the reaction are, what role does the aluminum play? (d) Calculate the volume of all the gases that would be produced at STP, assuming complete reaction of one pound of ammonium perchlorate.

22.94 The dissolved oxygen present in any highly pressurized, high-temperature steam boiler can be extremely corrosive to its metal parts. Hydrazine, which is completely miscible with water, can be added to remove oxygen by reacting with it to form nitrogen and water. (a) Write the balanced equation for the reaction between gaseous hydrazine and oxygen. (b) Calculate the enthalpy change accompanying this reaction. (c) Oxygen in air dissolves in water to the extent of 9.1 ppm at 20 °C at sea level. How many grams of hydrazine are required to react with all the oxygen in 3.0 × 104 L (the volume of a small swimming pool) under these conditions?

22.95 One method proposed for removing SO2 from the flue gases of power plants involves reaction with aqueous H2S. Elemental sulfur is the product. (a) Write a balanced chemical equation for the reaction. (b) What volume of H2S at 27 °C and 760 torr would be required to remove the SO2 formed by burning 2.0 tons of coal containing 3.5% S by mass? (c) What mass of elemental sulfur is produced? Assume that all reactions are 100% efficient.

22.96 The maximum allowable concentration of H2S(g) in air is 20 mg per kilogram of air (20 ppm by mass). How many grams of FeS would be required to react with hydrochloric acid to produce this concentration at 1.00 atm and 25 °C in an average room measuring 12 ft × 20 ft × 8 ft? (Under these conditions, the average molar mass of air is 29.0 g/mol.)

22.97 The standard heats of formation of H2O(g), H2S(g), H2Se(g), and H2Te (g) are –241.8, –20.17, +29.7, and +99.6 kJ/mol, respectively. The enthalpies necessary to convert the elements in their standard states to one mole of gaseous atoms are 248, 277, 227, and 197 kJ/mol of atoms for O, S, Se, and Te, respectively. The enthalpy for dissociation of H2 is 436 kJ/mol. Calculate the average H — O, H — S, H — Se, and H — Te bond enthalpies, and comment on their trend.

22.98 Manganese silicide has the empirical formula MnSi and melts at 1280 °C. It is insoluble in water but does dissolve in aqueous HF. (a) What type of compound do you expect MnSi to be: metallic, molecular, covalent-network, or ionic? (b) Write a likely balanced chemical equation for the reaction of MnSi with concentrated aqueous HF.

[22.99] Chemists tried for a long time to make molecular compounds containing silicon–silicon double bonds; they finally succeed in 1981. The trick is having large, bulky R groups on the silicon atoms to make R2Si ═ SiR2 compounds. What experiments could you do to prove that a new compound has a silicon–silicon double bond rather than a silicon–silicon single bond?

22.100 Hydrazine has been employed as a reducing agent for metals. Using standard reduction potentials, predict whether the following metals can be reduced to the metallic state by hydrazine under standard conditions in acidic solution: (a) Fe2+(b) Sn2+(c) Cu2+(d) Ag+(e) Cr3+,(f) Co3+.

22.101 Both dimethylhydrazine, (CH3)2NNH2, and methylhydrazine, CH3NHNH2, have been used as rocket fuels. When dinitrogen tetroxide (N2O4) is used as the oxidizer, the products are H2O, CO2, and N2. If the thrust of the rocket depends on the volume of the products produced, which of the substituted hydrazines produces a greater thrust per gram total mass of oxidizer plus fuel? [Assume that both fuels generate the same temperature and that H2O(g) is formed.]

22.102 Carbon forms an unusual unstable oxide of formula C3O2, called carbon suboxide. Carbon suboxide is made by using P2O5 to dehydrate the dicarboxylic acid called malonic acid, which has the formula HOOC — CH2 — COOH. (a) Write a balanced reaction for the production of carbon suboxide from malonic acid. (b) How many grams of carbon suboxide could be made from 20.00 g of malonic acid? (c) Suggest a Lewis structure for C3O2. (Hint: The Lewis structure of malonic acid suggests which atoms are connected to which.) (d) By using the information in Table 8.5, predict the C — C and C — O bond lengths in C3O2(e) Sketch the Lewis structure of a product that could result by the addition of 2 mol of H2 to 1 mol of C3O2.

22.103 Borazine, (BH)3(NH)3, is an analog of C6H6, benzene. It can be prepared from the reaction of diborane with ammonia, with hydrogen as another product; or from lithium borohydride and ammonium chloride, with lithium chloride and hydrogen as the other products. (a) Write balanced chemical equations for the production of borazine using both synthetic methods. (b) Draw the Lewis dot structure of borazine. (c) How many grams of borazine can be prepared from 2.00 L of ammonia at STP, assuming diborane is in excess?

22.104 Throughout history, arsenic(III) oxide, known simply to the general public as “arsenic,” has been a poison favored by murderers: It is tasteless, colorless, can be easily added to food or drink, produces symptoms that are similar to several diseases, and until the mid-1800s, was undetectable in the body. James Marsh developed the famous “Marsh test” for arsenic that was instrumental in convicting the murderer in a famous poisoning case in France in 1840. The Marsh test relies on the reaction of arsenic(III) oxide in a sample with elemental zinc and sulfuric acid to produce arsine (an analog of ammonia), zinc sulfate, and water. Upon igniting the final mixture, arsine is oxidized to elemental arsenic, and if captured on a ceramic bowl, a characteristic silvery-black powder would appear. (a) Write the balanced chemical equations of the Marsh test. (b) Antimony is the only potential interferent, as it reacts similarly to arsenic and produces a similar silvery-black film; however, antimony does not dissolve in a solution of sodium hypochlorite, but arsenic does. Therefore, the completion of the Marsh test is to add a solution of sodium hypochlorite to the elemental arsenic and see if the silvery-black film dissolves. Write the balanced chemical equation for this final reaction. (c) Today, commercial kits for arsenic testing rely on a different reaction. The sample is reacted with hydrogen sulfide in the presence of hydrochloric acid; if arsenic is present, As2S3, which is a bright yellow precipitate, forms. Write the balanced equation for this reaction.