?40EOxygen and the Other Group 6A Elements

(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]

(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]

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]

You have two glass bottles, one containing oxygen and one filled with nitrogen. How could you determine which one is which? [Sections 22.5 and 22.7]

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

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) XX single-bond enthalpy, (e) electron affinity? Explain your answer. [Sections 22.5 and 22.6]

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]

Which property of the third-row nonmetallic elements might be the one depicted below: (a) first ionization energy, (b) atomic radius, (c) electronegativity, (d) melting point, (e) XX 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]

Which of the following compounds would you expect to be the most generally reactive, and why? (Each corner in these structures represents a CH2 group.) [Section 22.8]

(a) Draw the Lewis structures for at least four species that have the general formula X Y n 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 Brnsted base? Explain. [Sections 22.1, 22.7, and 22.9]

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

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

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 p bonds most readily, (e) is a transition metal, (f) is a liquid at room temperature and pressure.

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 p bonds most readily, (f) has multiple allotropes.

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

Explain the following observations: (a) HNO3 is a stronger oxidizing agent than H3PO4. b) Silicon can form an ion with six fluorine atoms, SiF6 2 - , 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 1Si2H62.

Complete and balance the following equations:

Complete and balance the following equations:

(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

Are the physical properties of H2O different from D2O? Explain

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

What does hydrogen have in common with the halogens? Explain

Complete and balance the following equations:

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

Complete and balance the following equations

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.

Identify the following hydrides as ionic, metallic, or molecu

Identify the following hydrides as ionic, metallic, or molecular:

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

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

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

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

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

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

Name the following compounds and assign oxidation states to the halogens in them: (a) Fe1ClO323, (b) HClO2, (c) XeF6, (d) BrF5, (e) XeOF4, (f) HIO3.

Name the following compounds and assign oxidation states to the halogens in them: (a) KClO3, (b) Ca1IO322, (c) AlCl3, (d) HBrO3, (e) H5IO6, (f) XeF4.

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 7 Cl2 7 Br2 7 I2.

Explain the following observations: (a) For a given oxidation state, the acid strength of the oxyacid in aqueous solution decreases in the order chlorine 7 bromine 7 iodine. (b) Hydrofluoric acid cannot be stored in glass bottles. (c) HI cannot be prepared by treating NaI with sulfuric acid. (d) The interhalogen ICl3 is known, but BrCl3 is not.

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 PbSO41s2 and O21g2. (d) When heated in air, ZnS(s) is converted to ZnO. (e) Potassium peroxide reacts with CO21g2 to give potassium carbonate and O2. (f) Oxygen is converted to ozone in the upper atmosphere.

Complete and balance the following equations:

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

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

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.

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

In aqueous solution, hydrogen sulfide reduces (a) Fe 3 + to Fe 2 + , (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.

An aqueous solution of SO2 reduces (a) aqueous KMnO4 to MnSO41aq2, (b) acidic aqueous K2Cr2O7 to aqueous Cr 3 + , (c) aqueous Hg21NO322 to mercury metal. Write balanced equations for these reactions.

Write the Lewis structure for each of the following species, and indicate the structure of each: (a) SeO3 2 - ; (b) S2Cl2; (c) chlorosulfonic acid, HSO3Cl (chlorine is bonded to sulfur).

The SF5 - ion is formed when SF41g2 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

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.

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 SO4 2 - and elemental sulfur, while Cl2 is reduced to Cl - .

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.

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

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 - .

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

Complete and balance the following equations:

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 CrO4 2 - to Cr1OH24 - in base (hydrazine is oxidized to N2).

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.

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.

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 pentoxide, (f) sodium phosphate

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.

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 cannot be made in aqueous solution. (d) White phosphorus is more reactive than red phosphorus

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.

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.

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.

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

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

Complete and balance the following equations:

Complete and balance the following equations:

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.

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 1C6H12O62 and O2 from carbon dioxide and water. (c) When carbonate salts dissolve in water, they produce basic solutions

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.

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

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 Earths crust.

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.

(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?

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

(a) Determine the number of calcium ions in the chemical formula of the mineral hardystonite, CaxZn1Si2O72. (b) Determine the number of hydroxide ions in the chemical formula of the mineral pyrophyllite , Al21Si2O5221OH2x.

(a) Determine the number of sodium ions in the chemical formula of albite, NaxAlSi3O8. (b) Determine the number of hydroxide ions in the chemical formula of tremolite, Ca2Mg51Si4O11221OH2x.

(a) How does the structure of diborane 1B2H62 differ from that of ethane 1C2H62? (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?

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.

ndicate whether each of the following statements is true or false (a) H21g2 and D21g2 are allotropic forms of hydrogen. (b) ClF3 is an interhalogen compound. (c) MgO(s) is an acidic anhydride. (d) SO21g2 is an acidic anhydride. (e) 2 H3PO41l2 S H4P2O71l2+H2O1g2 is an example of a condensation reaction. (f) Tritium is an isotope of the element hydrogen. (g) 2SO21g2+O21g2 S 2SO31g2 is an example of a disproportionation reaction.

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?

Write a balanced equation for the reaction of each of the following compounds with water: (a) SO21g2, (b) Cl2O71g2, (c) Na2O21s2, (d) BaC21s2, (e) RbO21s2 (f) Mg3N21s2, (g) NaH(s).

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

Hydrogen peroxide is capable of oxidizing (a) hydrazine to N2 and H2O, (b) SO2 to SO4 2 - , (c) NO2 - to NO3 - , (d) H2S1g2 to S(s), (e) Fe 2 + to Fe 3 + . Write a balanced net ionic equation for each of these redox reactions.

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

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: (a) oxidation of S to SO2, (b) oxidation of SO2 to SO3, (c) 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?

(a) What is the oxidation state of P in PO4 3 - and of N in NO3 - ? (b) Why doesnt N form a stable NO4 3 - . ion analogous to P?

(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 1Na3P3O92 and sodium tetrametaphosphate 1Na4P4O122 are used as water-softening agents. They contain cyclic P3O9 3 - and P4O12 4 - ions, respectively. Propose reasonable structures for these ions.

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 ultrapure Ge from GeO2

1 (a) Determine the charge of the aluminosilicate ion whose composition is AlSi3O10. (b) Using Figure 22.33, propose a reasonable description of the structure of this aluminosilicate.

(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?

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

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 H2O1l2 as a product

Using G f for ozone from Appendix C, calculate the equilibrium constant for Equation 22.24 at 298.0 K, assuming no electrical input

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: Cl21aq2 + H2O Cl - 1aq2 + HClO1aq2 + H+1aq2 (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?

When ammonium perchlorate decomposes thermally, the products of the reaction are N21g2, O21g2, H2O1g2, 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 NH4ClO41s2 is -295.8 kJ. (c) When NH4ClO41s2 is employed in solid-fuel booster rockets, it is packed with powdered aluminum. Given the high temperature needed for NH4ClO41s2 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.

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?

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.

The maximum allowable concentration of H2S1g2 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.)

The standard heats of formation of H2O1g2, H2S1g2, H2Se1g2, and H2Te1g2 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 HO, HS, HSe, and HTe bond enthalpies, and comment on their trend.

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

Chemists tried for a long time to make molecular compounds containing siliconsilicon double bonds; they finally succeeded in 1981. The trick is having large, bulky R groups on the silicon atoms to make R2SiSiR2 compounds. What experiments could you do to prove that a new compound has a siliconsilicon double bond rather than a siliconsilicon single bond?

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) Fe 2 + , (b) Sn2 + , (c) Cu2 + , (d) Ag+, (e) Cr 3 + , (f) Co3 +

Both dimethylhydrazine, 1CH322NNH2, and methylhydrazine, CH3NHNH2, have been used as rocket fuels. When dinitrogen tetroxide 1N2O42 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 H2O1g2 is formed.)

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 HOOCCH2COOH. (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 CC and CO 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

Borazine, 1BH231NH23, 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?

These exercises are not divided by category, although they are roughly in the order of the topics in the chapter. They are not paired. The element lead (Pb) consists of four naturally occurring isotopes with atomic masses 203.97302, 205.97444, 206.97587, and 207.97663 amu. The relative abundances of these four isotopes are 1.4, 24.1, 22.1, and 52.4% respectively. From these data, calculate the atomic weight of lead.

(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]

(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]

Which of the following species (there may be more than one) is/are likely to have the structure shown here: (a) \(\mathrm{XeF}_{4}\), (b) \(B r F_{4}^{+}\), (c) \(S i F_{4}\), (d) \(\mathrm{TeCl}_{4}\), (e) \(\mathrm{HClO}_{4}\)? The colors do not reflect atom identities.) [Sections 22.3, 22.4, 22.6, and 22.10] Equation transcription: Text transcription: {XeF}_{4} B r F_{4}^{+} S i F_{4} {TeCl}_{4} {HClO}_{4}

These exercises are not divided by category, although they are roughly in the order of the topics in the chapter. They are not paired. There are two different isotopes of bromine atoms. Under normal conditions, elemental bromine consists of molecules, and the mass of a molecule is the sum of the masses of the two atoms in the molecule. The mass spectrum of consists of three peaks: (a) What is the origin of each peak (of what isotopes does each consist)? (b) What is the mass of each isotope? (c) Determine the average molecular mass of a molecule, (d) Determine the average atomic mass of a bromine atom, (e) Calculate the abundances of the two isotopes.

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

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]

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]

Which property of the third-row nonmetallic elements might be the one depicted below: (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]

Which of the following compounds would you expect to be the most generally reactive, and why? (Each corner in these structures represents a \(\mathrm{CH}_{2}\) group.) [Section 22.8] Equation transcription: Text transcription: {CH}_{2}

10E Visualizing Concepts (a) Draw the Lewis structures for at least four species that have the general formula (:X=Y:)n 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 Bronsted base? Explain. (Sections 22.1.22.7, and 22.9]

11E Periodic Trends and Chemical Reactions (Section) Identify each of the following elements as a metal, nonmetal, or metalloid: (a) phosphorus, (b) strontium, (c) manganese, (d) selenium, (e) sodium, (f) krypton.

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

13E Periodic Trends and Chemical Reactions (Section) Consider the elements O. Ba. Co. Be. Br. and Se. From this bst. select the element that (a) is most electronegative, (b) exhibits a maximum oxidation state of +7, (c) loses an election most readily, (d) forms tt bonds most readily, (e) is a transition metal, (f) is a liquid at room temperature and pressure.

Periodic Trends and Chemical Reactions (Section) Consider the elements Li. K. Cl. C. Ne. and Ar. From this 1st. select the element that (a) is most electronegative, (b) has the greatest metalb: character, (c) most readily forms a positive ion, (d) has the smallest atomic radius, (e) forms p bonds most readily, (f) has multiple allotropes

Periodic Trends and Chemical Reactions (Section) Explain the following observations: (a) The highest fluoride compound formed by nitrogen is . whereas phosphorus readily forms . (b) Although CO is a well-known compound. SiO does not exist under ordinary conditions, (c) is a stronger reducing agent than .

Explain the following observations: (a) HN03 is a stronger oxidizing agent than . (b) Sic on can form an ion with six fluorine atoms. . whereas carbon is able to bond to a maximum of four. . (c) There are three compounds formed by carbon and hydrogen that contain two carbon atoms each . and , whereas sic on forms only one analogous compound .

Periodic Trends and Chemical Reactions (Section) Complete and balance the following equations: (a) (b) (C) (d) (e)

Periodic Trends and Chemical Reactions (Section) Complete and balance the following equations: (a) (b) (c) (d) (e)

Hydrogen, the Noble Gases, and the Halogens (Sections 22.2. 22.3. and 22.4) (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.

Hydrogen, the Noble Gases, and the Halogens (Sections 22.2. 22.3. and 22.4) Are the physical properties of different from ? Explain.

21E Hydrogen, the Noble Gases, and the Halogens (Sections 22.2. 22.3. and 22.4) Gwe a reason why hydrogen might be placed along with the group 1A elements of the periodic table.

22E Hydrogen, the Noble Gases, and the Halogens (Sections 22.2. 22.3. and 22.4) What does hydrogen have in common with the halogens’ Explain.

23E Hydrogen, the Noble Gases, and the Halogens (Sections 22.2. 22.3. and 22.4) Write a balanced equation for the preparation of H2 using (a) Mg and an acid, (b) carbon and steam, (c) methane and steam.

24E Hydrogen, the Noble Gases, and the Halogens (Sections 22.2. 22.3. and 22.4) List (a) three commercial means of producing H2, (b) three industrial uses of H2.

25E Hydrogen, the Noble Gases, and the Halogens (Sections 22.2. 22.3. and 22.4) Complete and balance the following equations: (a) NaH(s) + H20(/) - (b) Fe(s) + H2S04(aq) - (c) H2(g) - Br2(g) - (d) Na(/) + H2(g) - (e) PbO(s) + H2(g) -

26E Hydrogen, the Noble Gases, and the Halogens (Sections 22.2. 22.3. and 22.4) 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(ll) oxide by hydrogen gas. (d) Calcium hydride reacts with water to generate hydrogen gas.

27E Hydrogen, the Noble Gases, and the Halogens (Sections 22.2. 22.3. and 22.4) Identify the following hydrides as ionic. metallic. or molecular: (a) BaH2, (b) H2Te, (c) TH1.7.

28E Hydrogen, the Noble Gases, and the Halogens (Sections 22.2. 22.3. and 22.4) Identify the following hydrides as ionic. metallic. or molecular: (a) B2H6. (b) RbH. (c) Th4H1.5.

29E Hydrogen, the Noble Gases, and the Halogens (Sections 22.2. 22.3. and 22.4) Describe two characteristics of hydrogen that are favorable for its use as a general energy source in vehicles.

30E Hydrogen, the Noble Gases, and the Halogens (Sections 22.2. 22.3. and 22.4) The H2/02 fuel cell converts elemental hydrogen and oxygen into water, producing, theoretically. 1.23 V. What is the most sustainable way to obtain hydrogen to run a large number of fuel cells-7 Explain.

31E Hydrogen, the Noble Gases, and the Halogens (Sections 22.2. 22.3. and 22.4) Why does xenon form stable compounds with fluorine, whereas argon does not?

32E Hydrogen, the Noble Gases, and the Halogens (Sections 22.2. 22.3. and 22.4) A friend tells you that the -neon' in neon signs is a compound of neon and aluminum. Can your friend be correct? Explain.

33E Hydrogen, the Noble Gases, and the Halogens (Sections 22.2. 22.3. and 22.4) Write the c hemic al formula for eac h of the following, and indie ate the oxidation state of the halogen or noble-gas atom in each: (a) calcium hypobromite. (b) bromic acid, (c) xenon trioxide, (d) perchbrate bn. (e) bdous acid, (f) iodine pentafluoride.

34E Hydrogen, the Noble Gases, and the Halogens (Sections 22.2. 22.3. and 22.4) 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.

35E Hydrogen, the Noble Gases, and the Halogens (Sections 22.2, 22.3, and 22.4) Name the following compounds and assign oxidation states to the halogens in them: (a) Fe(CI03)3, (b) HCI02, (c) XeF6, (d) BrF5, (e) XeOF4, (f) HI03.

36E Hydrogen, the Noble Gases, and the Halogens (Sections 22.2. 22.3. and 22.4) Name the following compounds and assign oxidation states to the halogens in them: (a) KCI03, (b) Ca(l03)2, (c) AICI3, (d) HBr03, (e) H5I06. (f) XeF4.

37E Hydrogen, the Noble Gases, and the Halogens (Sections 22.2. 22.3. and 22.4) Explain each of the following observations: (a) At room temperature 12 is a solid. Br2 is a liquid, and CC and F2 are both gases, (b) F2 cannot be prepared by electrolytic oxidation of aqueous F-sohitions. (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 > CC > Br2 > 12.

38E Hydrogen, the Noble Gases, and the Halogens (Sections 22.2. 22.3. and 22.4) 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 Nal with sulfuric acid, (d) The interhalogen ICI3 is known, but BrCI3 is not.

39E Oxygen and the Other Group 6A Elements (Sections) Write balanced equations for each of the following reactions, (a) When mercury(ll) oxide is heated, it decomposes to form 02 and mercury metal, (b) When copper(ll) nitrate is heated strongly, it decomposes to form copper(ll) oxide, nitrogen dioxide, and oxygen, (c) Lead(ll) sulfide. PbS(s), reacts with ozone to form PbS04(s) and 02(g). (d) When heated in air. ZnS(s) is converted to ZnO. (e) Potassium peroxide reacts with 002(g) to give potassium carbonate and 02. (f) Oxygen is converted to ozone in the upper atmosphere.

40E Oxygen and the Other Group 6A Elements (Sections) Complete and balance the following equations: (a) CaO(s) + H20(/) - (b) AE03(s) + H+<ag) - (c) Na202(s) + H20(/) - (d) N203(g) + H20(/) - (e) K02(s) + H20(/) - (f) NO(g) + 03(g) —*

41E Oxygen and the Other Group 6A Elements (Sections) Predict whether each of the following oxides is acidic, basic, amphoteric, or neutral: (a) N02, (b) C02, (c) AE03, (d) CaO.

42E Oxygen and the Other Group 6A Elements (Sections) Select the more acidic member of each of the following pairs: (a) Mn207 and Mn02, (b) SnO and Sn02, (c) S02 and SO3. (d) Si02 and S02, (e) Ga202 and In203. (f) S02 and Se02.

43E Oxygen and the Other Group 6A Elements (Sections) Write the chemical formula for each of the following compounds, and indicate the oxidation state of the group 6A element in each: (a) seienous acid, (b) potassium hydrogen sulfite, (c) hydrogen teDuride, (d) carbon disulfide, (e) calcium sulfate, (f) cadmium sulfide, (g) zinc teDuride.

44E Oxygen and the Other Group 6A Elements (Sections) 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 teDuride.

45E Oxygen and the Other Group 6A Elements (Sections) In aqueous solution, hydrogen sulfide reduces (a) Fe3 + to Fe2 +. (b) Br2 to Br-. (c) Mn04 - to Mn2 +, (d) HN03 to N02. In all cases, under appropriate conditions, the product is elemental sulfur. Write a balanced net ionic equation for each reaction.

46E Oxygen and the Other Group 6A Elements (Sections) An aqueous solution of S02 reduces (a) aqueous KMn04 to MnS04(aq), (b) acidic aqueous K2Cr207 to aqueous Cr3 +, (c) aqueous Hg2(N03)2 to mercury metal. Write balanced equations for these reactions.

47E Oxygen and the Other Group 6A Elements (Sections) Write the Lewis structure for each of the following species, and indicate the structure of each: (a) Se032 (b) S2CI2: (c) chlorosulfonic acid. HSOSCI (chlorine is bonded to sulfur)

48E Oxygen and the Other Group 6A Elements (Sections) 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.

49E Oxygen and the Other Group 6A Elements (Sections) 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.

50E Oxygen and the Other Group 6A Elements (Sections) 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 CC from chlorine-bleached fabrics. The thiosulfate ion forms SQ42- and elemental sulfur, while CC is reduced to Ch

51E Nitrogen and the Other Group 5A Elements (Sections) 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.

52E Nitrogen and the Other Group 5A Elements (Sections) 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) ithium nitride

53E Nitrogen and the Other Group 5A Elements (Sections) Write the Lewis structure for each of the folowing species, describe its geometry, and indicate the oxidation state of the nitrogen: (a) HN02, (b) N3-. (c) N2H5+. (d) N03-.

54E Nitrogen and the Other Group 5A Elements (Sections) Write the Lewis structure for each of the folowing species, describe its geometry, and indicate the oxidation state of the nitrogen: (a) NH4 +, (b) N02 (c) N20. (d) N02.

55E Nitrogen and the Other Group 5A Elements (Sections) Complete and balance the following equations: (a) Mg3N2(s) + H20(/) — (b) NO(g) + 02(g) - (c) N205(g) + H20(/) - (d) NH3(ag) + H+(sg) - (e) N2H4(/) + 02(g) - Which ones of these are redox reactions?

56E Nitrogen and the Other Group 5A Elements (Sections) Write a balanced net ionic equation for each of the following reactions: (a) Dlute 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 Cr042- to Cr(OH)4- in base (hydrazine is oxidized to N2).

57E Nitrogen and the Other Group 5A Elements (Sections) Write complete balanced half-reactions for (a) oxidatbn of nitrous acid to nitrate ion in acidic solution, (b) oxidation of N2 to N20 in acidic solution.

58E Nitrogen and the Other Group 5A Elements (Sections) Write complete balanced half-reactions for (a) reduction of nitrate ion to NO in acidic solution, (b) oxidation of HN02 to N02 in acidic solution.

59E Nitrogen and the Other Group 5A Elements (Sections) 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 pentoxide. (f) sodium phosphate.

60E Nitrogen and the Other Group 5A Elements (Sections) Write a chemical formula for each compound or ion. and indicate the oxidation state of the group 5A element in each formula: (a) phosphate bn, (b) arsenous acid, (c) antimony(lll) sulfide, (d) calcium diiydrogen phosphate, (e) potassium phosphide, (f) gallium arsenide.

61E Nitrogen and the Other Group 5A Elements (Sections) Account for the following observations: (a) Phosphorus forms a pentachloride. but nitrogen does not. (b) H3P02 is a monoprotic acid, (c) Phosphonium salts, such as PH4CI. can be formed under anhydrous conditions, but they cannot be made in aqueous solution, (d) White phosphorus is more reactive than red phosphorus.

62E Nitrogen and the Other Group 5A Elements (Sections) Account for the following observations: (a) H3P03 is a diprotic acid, (b) Nitric acid is a strong acid, whereas phosphor*: 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 Na3P04 are quite basic.

63E Nitrogen and the Other Group 5A Elements (Sections) 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.

64E Nitrogen and the Other Group 5A Elements (Sections) Write a balanced equation for each of the following reactions: (a) hydrolysis of PCI5. (b) dehydration of phosphoric acid (also called orthophosphoric acid) to form pyrophosphoric acid, (c) reaction of P4O10 with water.

65E Carbon, the Other Group 4A Elements, and Boron (Sections 22.9. 22.10. and 22.11) Gh/e the chemical formula for (a) hydrocyanic acid, (b) nickel tetracarbonyl. (c) barium bicarbonate, (d) calcium acety&de. (e) potassium carbonate.

66E Carbon, the Other Group 4A Elements, and Boron (Sections 22.9. 22.10. and 22.11) Gwe the chemical formula for (a) carbonic acid, (b) sodium cyanide, (c) potassium hydrogen carbonate, (d) acetylene, (e) iron pentacarbonyl.

67E Carbon, the Other Group 4A Elements, and Boron (Sections 22.9. 22.10. and 22.11) Complete and balance the following equations: (a) ZnCO,(j) — (b) BaC2(s) + H20(/) - (c) C2H2(g) + 02(g) - (d) CS2(g) * 02(g) - (e) Ca(CN)2(s) + HBr(aq) -

Complete and balance the following equations: 1. \(\mathrm{CO}_{2}(g)+\mathrm{OH}^{-}(a q) \rightarrow\) 2. \(\mathrm{NaHCO}_{3}(s)+\mathrm{H}^{+}(a q) \rightarrow\) 3. \(C a O(s)+C(s) \rightarrow\) 4. \(\mathrm{C}(s)+\mathrm{H}_{2} \mathrm{O}(g) \rightarrow\) 5. \(C u O(s)+C O(g) \rightarrow\) Equation transcription: Text transcription: {CO}_{2}(g)+{OH}^{-}(a q) rightarrow {NaHCO}_{3}(s)+{H}^{+}(a q) rightarrow C a O(s)+C(s) rightarrow {C}(s)+{H}_{2}{O}(g) rightarrow C u O(s)+C O(g) rightarrow

69E Carbon, the Other Group 4A Elements, and Boron (Sections 22.9. 22.10. and 22.11) 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 SC. 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.

70E Carbon, the Other Group 4A Elements, and Boron (Sections 22.9. 22.10. and 22.11) Write a balanced equation for each of the following reactions: (a) Burning magnesium metal in a carbon dioxide atmosphere reduces the C02 to carbon, (b) In photosynthesis, solar energy is used to produce glucose (C6H1206) and 02 from carbon dioxide and water, (c) When carbonate salts dissolve in water, they produce basic solutions.

71E Carbon, the Other Group 4A Elements, and Boron (Sections 22.9. 22.10. and 22.11) 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(ll) chloride, (d) sodium tetraborate decahydrate (borax), (e) boric oxide, (f) germanium dioxide.

72E Carbon, the Other Group 4A Elements, and Boron (Sections 22.9. 22.10. and 22.11) 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.

73E Carbon, the Other Group 4A Elements, and Boron (Sections 22.9. 22.10. and 22.11) 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 Earths crust.

74E Carbon, the Other Group 4A Elements, and Boron (Sections 22.9. 22.10. and 22.11) 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 metaltoid that can form 2+ ions.

(a) What is the characteristic geometry about silicon in all silicate minerals? (b) Metasilicic acid has the empirical formula \(\mathrm{H}_{2} \mathrm{SiO}_{3}\). Which of the structures shown in Figure 22.34 would you expect metasilicic acid to have? Equation transcription: Text transcription: {H}_{2}{SiO}_{3}

76E Carbon, the Other Group 4A Elements, and Boron (Sections 22.9. 22.10. and 22.11) Speculate as to why carbon forms carbonate rather than silicate analogs.

77E Carbon, the Other Group 4A Elements, and Boron (Sections 22.9. 22.10. and 22.11) (a) Determine the number of calcium tons in the chemical formula of the mineral hardystonite, CaxZn(Si207). (b) Determine the number of hydroxide ions in the c hemic al formula of the mineral pyrophylBte. AE(Si205)2(OH)x.

78E Carbon, the Other Group 4A Elements, and Boron (Sections 22.9. 22.10. and 22.11) (a) Determine the number of sodium ions in the chemical formula of abite. NaxAlSi308. (b) Determine the number of hydroxide ions in the chemical formula of tremoBte. Ca2Mg5(Si4011)2(0 H)x.

79E Carbon, the Other Group 4A Elements, and Boron (Sections 22.9. 22.10. and 22.11) (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'?

80E Carbon, the Other Group 4A Elements, and Boron (Sections 22.9. 22.10. and 22.11) Write a balanced equation for each of the following reactions: (a) Dborane 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 dissoh/es in water to give a solution of boric acid.

81AE Indicate whether each of the following statements is true or false (a) H2(g) and D2(g) are allotropic forms of hydrogen, (b) CIF3 is an interhalogen compound, (c) MgO(s) is an acidic anhydride, (d) S02(g) is an acidic anhydride, (e) 2 H3P04(/) — H4P207(/)+H20(g) is an example of a condensation reaction, (f) Tritium is an isotope of the element hydrogen, (g) 2S02(g)+02(g) — 2S03(g) is an example of a disproportionation reaction.

82AE Although the CI04- and 104- ions have been known for a long time. Br04- 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?

83AE Write a balanced equation for the reaction of each of the following compounds with water: (a) S02(g), (b) C1207(g), (c) Na202(s), (d) BaC2(s), (e) Rb02(s) (f) Mg3N2(s), (g) NaH(s).

84AE What is the anhydride for each of the following acids: (a) H2S04, (b) HCI03. (c) HN02, (d) H2C03, (e) H3P04?

85AE Hydrogen peroxide is capable of oxidizing (a) hydrazine to N2 and H20. (b) S02 to S042-, (c) N02- to N03-. (d) H2S(g) to S(s), (e) Fe2+ to Fe3+. Write a balanced net ionic equation for each of these redox reactions.

86AE Explain why S02 can be used as a reducing agent but S03 cannot.

87AE 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 H2S04 consists of three main chemical processes: (a) oxidation of S to S02, (b) oxidation of S02 to S03. (c) the dissolving of SO3 in H2S04 and its reaction with water to form H2S04. If the third process produces 130 kJ/mol. how much heat is produced in preparing a mole of H2S04 from a mole of S? How much heat is produced in preparing 5000 pounds of H2S04?

88AE (a) What is the oxidation state of P in P043- and of N in N03-? (b) Why doesnl N form a stable N043-. ion analogous to P7

(a) The \(P_{4}, P_{4} O\) and \(P_{4} O_{10}\) 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 \(\left(\mathrm{Na}_{3} \mathrm{P}_{3} \mathrm{O}_{9}\right)\) and sodium tetrametaphosphate \(\left(\mathrm{Na}_{4} P_{4} O_{12}\right)\) are used as water-softening agents. They contain cyclic \(P_{3} O_{9}^{3-}\) and \(P_{4} O_{12}^{4-}\) ions, respectively. Propose reasonable structures for these ions. How do the electron domains about P in \(P_{4} O_{6}\) differ from those about P in \(P_{4} O_{10}\)? ? Figure 22.27 Structures of \(P_{4} O_{6}\) (top) and \(P_{4} O_{10}\) (bottom). ? Figure 22.28 Structures of \(\mathrm{H}_{3} \mathrm{PO}_{4}\) (top) and \(\mathrm{H}_{3} \mathrm{PO}_{3}\) (bottom). Equation transcription: Text transcription: P_{4}, P_{4} O P_{4} O_{10} ({Na}_{3}{P}_{3}{O}_{9}) ({Na}_{4} P_{4} O_{12}) P_{3} O_{9}^{3-} P_{4} O_{12}^{4-} P_{4} O_{6} {H}_{3}{PO}_{4} {H}_{3}{PO}_{3}

90AE Ultrapure germanium. l<e silicon, is used in semiconductors. Germanium of "ordinary’ purity is prepared by the high-temperature reduction of Ge02 with carbon. The Ge is converted to GeCI4 by treatment with Ct2 and then purified by distillation: GeCI4 is then hydrolyzed in water to Ge02 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 ultrapure Ge from Ge02.

(a) Determine the charge of the aluminosilicate ion whose composition is \(A I S \dot{i}_{3} O_{10}\). (b) Using Figure 22.33, propose a reasonable description of the structure of this aluminosilicate. ? Figure 22.33 Silicate chains and sheets. Equation transcription: Text transcription: A I S dot{i}_{3} O_{10}

92IE (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

93IE Using the thermochemical data in Table 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? Table Properties of Xenon Compounds Compound Oxidation State of Xe Melting Point (SC) AH°f kJ mol i[a] XeF2 +2 129 -1091g2 XeF4 +4 117 -2181g2 XeF6 +6 49 -2981g2 XeOF4 +6 -41 to-28 +146112 Xe03 +6 _[b] +4021s2 Xe02F2 +6 31 +1451s2 Xe04 +8 Jc] -

94IE 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 H20(/) as a product.

Using \(\Delta G_{f}^{0}\) for ozone from Appendix C, calculate the equilibrium constant for Equation 22.24 at 298.0 K, assuming no electrical input. Equation transcription: Text transcription: Delta G_{f}^{0}

The solubility of \(C I_{2}\) in 100 g of water at STP is \(310 \mathrm{~cm}^{3}\). Assume that this quantity of \(C I_{2}\) is dissolved and equilibrated as follows: \(C L_{2}(a q)+H_{2} O \leftrightarrow C l^{-}(a q)+H C l O(a q)+H^{+}(a q)\) (a) If the equilibrium constant for this reaction is \(4.7 \times 10^{-4}\), calculate the equilibrium concentration of HClO formed. (b) What is the pH of the final solution? Equation transcription: Text transcription: C I_{2} 310 {~cm}^{3} C L_{2}(a q)+H_{2} O leftrightarrow C l^{-}(a q)+H C l O(a q)+H^{+}(a q) 4.7 times 10^{-4}

When ammonium perchlorate decomposes thermally, the products of the reaction are \(N_{2}(g), O_{2}(g), H_{2} O(g)\), and \(H C I(g)\). (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 \(\mathrm{NH}_{4} \mathrm{C} I O_{4}\). The standard enthalpy of formation of \(\mathrm{NH}_{4} \mathrm{C} I O_{4}\)(s) is -295.8 kJ. (c) When \(\mathrm{NH}_{4} \mathrm{C} I O_{4}\)(s) is employed in solid-fuel booster rockets, it is packed with powdered aluminum. Given the high temperature needed for \(\mathrm{NH}_{4} \mathrm{C} I O_{4}\)(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. Equation transcription: Text transcription: N_{2}(g), O_{2}(g), H_{2} O(g) H C I(g) {NH}_{4}{C} I O_{4}

Problem 98IE 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 SC at sea level. How many grams of hydrazine are required to react with all the oxygen in 3.0 x 104 L (the volume of a small swimming pool) under these conditions?

Problem 99IE One method proposed for removing S02 from the flue gases of power plants invokes 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 SC and 760 torr would be required to remove the S02 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.

100IE The maximum allowable concentration of H2S(g) in air is 20 mg per kiogram 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 SC in an average room measuring 12 ft x 20 ft x 8 ft? (Under these conditions, the average molar mass of air is 29.0 g/mol.)

101IE The standard heats of formation of H20(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.

102IE Manganese silk:»de 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 MnSito be: metallic, molecular, covalent-network, or ionic? (b) Write a likely balanced chemical equation for the reaction of MnSi with concentrated aqueous HF.

103IE Chemists tried for a long time to make molecular compounds containing silicon-silicon double bonds: they finally succeeded 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?

104IE Hydrazine has been employed as a reducing agent for metals. Using standard reduction potentials, predict whether the following metals can be reduced to the metalbc state by hydrazine under standard conditions in acidic solution: (a) Fe2+ , (b) Sn2+ , (c) Cu2+ , (d) Ag+. (e) Cr3+ , (f) Co3+

105IE Both dimethyliydrazine. (CH3)2NNH2. and methyliydrazine. CH3NHNH2. have been used as rocket fuels. When dinitrogen tetroxide (N204) is used as the oxidizer, the products are H20. C02. 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 H20(g) is formed.) Step-by-step solution

106 IE 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, predict the C—C and C—O bond lengths in C3O2. Table Average Bond Lengths for Some Single, Double, and Triple Bonds Bond Bond Length (Å) Bond Bond Length (Å) C—C 1.54 N—N 1.47 C=C 1.34 N=N 1.24 C?C 1.20 N?N 1.10 C—N 1.43 N—O 1.36 C=N 1.38 N=O 1.22 C?N 1.16 O—O 1.48 C—O 1.43 O=O 1.21 C=O 1.23 C?O 1.13 (e) Sketch the Lewis structure of a product that could result by the addition of 2 mol of H2 to 1 mol of C3O2.

107IE Borazine, (BH)3(NH2), 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?