chemistry 134 problem set introduction

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Chemistry 134 Problem Set

Introduction The following problem set includes all assigned problems for the Chemistry 134 Sections taught by R.H. Langley.

The coding is to T.L. Brown, H. E. LeMay, Jr., B. E. Bursten, C. J. Murphy, and P. M. Woodward, Chemistry the

Central Science, 12th ed., Prentice Hall, 2012. Reference to this text will be “BLBMWS.”

Selected answers in Blue.

Chapter 14 Corresponds to BLBMWS Chapter 22

14.1 Corresponds to BLBMWS 22.1

14.1 Label each of the following as a metal, nonmetal, or metalloid. (a) silicon (b) titanium (c) radon (d) gallium

(e) rubidium (f) chlorine

14.2 Of the elements Li, Rb, Br, C, and Kr, select which is (a) most electronegative (b) most metallic (c) most easily

oxidized (d) smallest in size (e) best able to form bonds

14.3 What oxidation states are possible for each of the following? (a) lithium (b) aluminum (c) tin (d) strontium

(e) nitrogen

14.4 Which compound is each of the following pairs is the more covalent? (a) magnesium chloride or beryllium

chloride (b) tin(IV) chloride or tin(II) chloride (c) aluminum chloride or rubidium chloride (d) sodium chloride or

sodium fluoride (e) potassium chloride or bromine trifluoride

14.5 Complete and balance the following equations based upon information from Chapters 4 and 8.

(a) Na2O(s) + H2O(l) (b) C3H7OH(l) + O2(g) (c) NiCl2(s) + Mg(s)

(d) Li3N(s) + H2O(l) (e) Na2SO3(s) + HBr(aq)


14.6 (a) Why is hydrogen an unusual element? (b) What other element has this characteristic? (c) What is the

isotope effect, and why is it more important to hydrogen than any other element?

14.7 Determine the number of electrons, protons, and neutrons in each of the isotopes of hydrogen.

14.8 What are the electron configurations of hydrogen and the two ions it may form?

14.9 Draw a molecular orbital energy level diagram for a hydrogen molecule. From this diagram, predict the bond

order of the hydrogen molecule.

14.10 Zinc is above hydrogen in the activity series, and copper is below hydrogen in the activity series. How does

their placement in the activity series lead to differences in the reactions of zinc and copper with hydrochloric acid?

14.11 Much of the hydrogen gas produced in the United States is utilized in the synthesis of ammonia. The other

reactant in this process is nitrogen gas. (a) Write a balanced chemical equation for the reaction. (b) Using bond

energies, determine the enthalpy of reaction. (c) Why is the production of ammonia so expensive?

14.12 Classify each of the following hydrides as ionic, metallic, or molecular. (a) PH3 (b) KH (c) PdH0.7

14.13 Complete and balance the following equations: (a) LiH(s) + H2O(l) (b) Fe(s) + H3PO4(aq)

(c) H2(g) + F2(g) (d) K(l) + H2(g) (e) WO3(s) + H2(g)


14.14 Using information from this chapter write balanced chemical equations for each of the following. (a) Zinc

metal will react with acids to produce hydrogen gas. (b) Steam will react with magnesium metal to produce solid

magnesium oxide and hydrogen gas. (c) Hydrogen gas reduces solid iron(III) oxide to solid iron(II) oxide. (d)

Nitrogen gas reacts with hydrogen gas in the presence of a catalyst to produce ammonia gas. (e) Solid calcium

hydride forms when heating calcium metal in the presence of hydrogen gas.

14.15 Propylene (CH3CH=CH2) is a major industrial chemical. Large amounts of this substance are necessary for

the preparation of polypropylene. A simplified reaction for the preparation of propylene is

CH3CH2CH3(g) CH3CH=CH2(g) + H2(g)

If a particular plant produces 1.7 × 107 kg of polypropylene per year, how many kilograms of hydrogen gas form as

a by-product?

14.16 It is possible to use a titration to produce a particular product. An example is the reaction of barium hydroxide

with phosphoric acid to produce barium phosphate. How many milliliters of 0.2275 M barium hydroxide are

necessary to react completely with 50.00 mL of 0.1250 M phosphoric acid? 41.21 mL

14.2 Corresponds to BLBMWS

14.17 (a) Which of the alkali metals are among the more abundant elements in the Earth's crust? (b) Which alkali

metal has no known minerals? Where does this element occur? (c) Why are very few of the properties of francium

known?

14.18 Use molecular orbital theory to show the development of a conduction band in lithium metal.

14.19 (a) What is the trend in the melting points of the alkali metals? (b) Which alkali metals are less dense than

water? (c) Which metal has the lowest density of all metals?

14.20 (a) Which of the important metal structures do the alkali metals adopt? (b) Sketch the structure of the alkali

metals. (c) How many atoms are present in a unit cell of an alkali metal?

14.21 Write balanced chemical equations for the reaction of each of the alkali metals with oxygen gas.

14.22 Write balanced chemical equations for the following reactions. (a) lithium metal with nitrogen gas (b) sodium

metal with solid phosphorus (c) potassium metal with solid iodine (d) rubidium metal with solid sulfur (e) cesium

metal with graphite

14.23 Why are most sodium and potassium compounds strong electrolytes?

14.24 Complete and balance the following equations. (a) Li(s) + O2(g) (b) Na(s) + O2(g)

(c) K(s) + O2(g) (d) Rb(s) + H2O(l) (e) KOH(s) + CO2(g) (f) Cs2CO3(aq) + Ca(OH)2(aq)

14.3 Corresponds to BLBMWS

14.25 (a) Which alkaline earth metals are among the ten most abundant elements in the Earth's crust? (b) What are

the third and fourth most abundant alkaline earth metals?

14.26 The analysis of a calcium hydroxide solution found that 35.75 mL of 0.1230 M hydrochloric acid were

necessary for titration of 50.00 mL of calcium hydroxide solution. What was the molarity of the calcium hydroxide

solution?

14.27 (a) Unlike the compounds of other alkaline earth metals, many beryllium compounds are covalent. What

property of beryllium explains this observation? (b) Write balanced chemical reactions to illustrate the amphoteric

properties of beryllium.

14.28 Write balanced chemical equations for the steps in the preparation of magnesium.


14.29 What is the trend in reactivity of the alkaline earth metals with respect to their position in the periodic table?

14.30 Write balanced chemical equations for each of the following reactions. (a) beryllium metal with chlorine gas

(b) magnesium metal with steam (c) calcium metal with solid phosphorus (d) strontium metal with oxygen gas

(e) barium metal with solid sulfur

14.31 What are the names of the three minerals with the chemical formula CaCO3?

14.32 Write balanced chemical equations showing two methods of preparing calcium oxide.

14.33 What is the difference between temporary and permanent hardness?

14.34 In a particular water supply, the concentration of Ca2+ is 4.7 × 10–3 M, and the concentration of bicarbonate

ion is 1.5 × 10–3 M. What masses of Ca(OH)2 and Na2CO3 are required to reduce the level of Ca2+ to one-fourth of

its original level if 2.0 × 108 L of water must be treated? The reaction is Ca2+(aq) + 2 HCO3–(aq) +

Ca(OH)2(s) 2 CaCO3(s) + 2 H2O(l)

14.35 The major source of magnesium is seawater. The first step in recovering magnesium from seawater is to

precipitate magnesium hydroxide with calcium oxide:

Mg2+(aq) + CaO(s) + H2O(l) Mg(OH)2(s) + Ca2+(aq)

How many kilograms of calcium oxide are required to precipitate 1.0 × 109 g of magnesium hydroxide?

14.36 It is possible to form calcium chloride by the reaction of solid calcium oxide with gaseous hydrogen chloride.

(a) Write a balanced chemical equation for this reaction. (b) How many grams of calcium chloride will form when

100.0 g of calcium oxide react with 15.0 L of hydrogen chloride gas? The pressure of the gas is 775 torr at 40.0 °C.

(c) What is the percent yield if only 13.5 g of calcium chloride form?

(a) CaO(s) + 2 HCl(g) CaCl2(s) + H2O(l) (b) 33.0 g CaCl2 (c) 40.9 %


14.37 (a) Which of the elements in the boron family is the most abundant? (b) How abundant is this element in the

Earth’s crust?

14.38 (a) What is the difference between a sapphire and a ruby? (b) Why might aluminum be present with silicon in

many minerals?

14.39 (a) List the stable oxidation states for each member of the boron family. (b) For any element that may have

more than one stable oxidation state, identify the more stable state.

14.40 (a) What is the inert pair effect? (b) What is an allotrope?

14.41 The tetrahydroaluminate ion (AlH4–) and the borohydride ion (BH4

–) are important reducing agents used in

organic chemistry. Each of these ions can form through a Lewis acid base reaction between a Lewis acid (AlH3 or

BH3) and a Lewis base (H–). Use Lewis structures to illustrate the reaction forming the tetrahydroaluminate ion and

the borohydride.

14.42 (a) Give a balanced chemical equation for the preparation of aluminum metal using sodium metal. (b) What

other metals might replace the sodium metal?

14.43 (a) What is an important use of boron? (b) What is the formula of the simplest borane? (c) What is unusual

about the melting point of gallium?

14.44 Aluminum hydroxide is soluble in hydrochloric acid. How many milliliters of 6.0 M hydrochloric acid are

necessary to dissolve 25.0 g of aluminum hydroxide?


14.45 Alum is an important aluminum compound. It is possible to prepare small quantities of alum by the following

two reactions:

6 H2O(l) + 2 Al(s) + 2 KOH(aq) 2 KAl(OH)4(aq) + 3 H2(g)

KAl(OH)4(aq) + 2 H2SO4(aq) KAl(SO4)2(aq) + 4 H2O(l)

How many grams of KAl(SO4)2 will be produced in the reaction of 10.0 g of aluminum metal with an excess of the

other reactants?

14.46 Using Lewis structures illustrate why the reaction of water with boric acid is a Lewis acid-base reaction.

14.47 Write balanced chemical equations illustrating the amphoteric nature of aluminum hydroxide.

14.48 Gallium hydroxide, like aluminum hydroxide, is amphoteric. Write balanced chemical equations, similar to

the equations for aluminum hydroxide in this chapter, to illustrate the reaction of gallium hydroxide with sulfuric

acid, and the reaction of gallium hydroxide with hydroxide ion.

14.5 Corresponds to BLBMWS 22.9 & 22.10

14.49 (a) Label each of the elements in the carbon family as a nonmetal, a metalloid, or a metal. (b) Which of the

members of the carbon family is the most abundant? How abundant is this element in the Earth's crust? (c) List one

important use of each of the elements in the carbon family. (d) What is the difference between charcoal and coke?

14.50 Which element or elements in the carbon family illustrate the inert pair effect?

14.51 Draw the resonance structures for the cyanate ion (OCN–) and determine the formal charge of each atom in

each resonance structure.

14.52 Which member of the carbon family was predicted to exist by Mendeleev before the element was actually

discovered?

14.53 Which tin and lead compounds tend to be ionic and which tend to be covalent?

14.54 What is catenation, and which element best illustrates this property?

14.55 (a) Draw the band structure for germanium. (b) Show the band structure for germanium containing a small

amount of gallium. Is this an n-type or a p-type semiconductor? (c) Show the band structure for germanium

containing a small amount of arsenic. Is this an n-type or a p-type semiconductor?

14.56 (a) Which elements in the carbon family may adopt the diamond structure? (b) Describe the structures of the

following allotropes of carbon: diamond, graphite, and buckminsterfullerene.

14.57 Silicon dioxide is the main ingredient in glass. (a) What substances lower the melting point of silicon

dioxide? (b) What substance makes glass brighter? (c) What additive produces blue glass? (d) What substance

produces red glass?

14.58 (a) Why is hydrofluoric acid the only common acid to react with glass? (b) Write a balanced chemical

equation for the reaction of hydrofluoric acid with silicon dioxide.

14.59 (a) Write a balanced chemical equation for the reaction of beryllium carbide with water. (b) What is the

maximum amount of gas that could form upon mixing 60.0 g of beryllium carbide with 60.0 g of water? (c) If 12.5

g of gas form, what is the percent yield?

14.60 What is the basic unit for most silicate minerals? (b) How do these units produce the different categories of

silicates?


14.61 (a) Write the electron configuration for a silicon atom. (b) Draw an orbital diagram for a silicon atom.

(c) Assign a possible set of four quantum numbers for each of the electrons in the valence shell of a silicon atom.

14.62 Ethane (C2H6) is a flammable gas. It will burn in oxygen gas to produce carbon dioxide and water vapor.

(a) Write a balanced chemical reaction for ethane burning in oxygen gas. (b) Use bond energies to estimate the

enthalpy change for this reaction. (c) Use standard heats of reaction to calculate the enthalpy change for this

reaction.

14.63 The most abundant isotope of carbon is 12C. About 1.1% of natural carbon is 13C, and a trace amount is 14C.

Carbon-14, though very rare, is the key to carbon dating, a way to measure the age of ancient artifacts. Determine

the number of protons, neutrons, and electrons in each of these isotopes of carbon.

14.64 Boron nitride (BN) occurs in two forms. One is very soft and has a layered structure. The other is very hard

and has a three-dimensional-network structure. Locate boron and nitrogen in the periodic table. Which element has

the same electron configuration as the average electron configuration of boron and nitrogen? Based on the average

electron configurations of boron and nitrogen, how do you think the structures of boron nitride might relate to those

of other materials discussed in this chapter?

14.65 It is necessary to make extremely pure silicon (Si) for computer chips. In one of the steps in the purification

process, impure silicon is heated in chlorine gas to produce volatile silicon tetrachloride (SiCl4). Most of the

impurities remain behind in the solid phase, as they are not volatile. How many grams of chlorine gas are necessary

to react with 2.500 kg of silicon?

14.66 It requires about 331 kJ/mol to break the C–Cl bond in a chlorofluorocarbon. (a) Calculate the wavelength, in

nanometers, of a photon with sufficient energy to break a C–Cl bond. (b) What is the frequency of this photon?

(a) 3.61 × 10–7m (b) 8.30 × 1014 s–1

14.67 For each of the following ions or molecules, draw the Lewis structure. If resonance is possible, show all

resonance structures. (a) carbon monoxide, CO (b) carbon dioxide, CO2 (c) thiocyanate ion, SCN– (d) bicarbonate

ion, HCO3– (e) carbonate ion, CO3

2–

14.68 Complete and balance the following equations: (a) BaCO3(s) + heat

(b) C12(g) + Ge(s) (c) C6H14(g) + O2(g) (d) CH3CH2OH(l) + O2(g)

(e) BaH2(s) + H2O(l)

14.69 Write balanced chemical equations for each of the following reactions. (a) It is possible to produce hydrogen

cyanide along with water vapor when methane, ammonia, and air are heated in the presence of a catalyst.

(b) Baking soda (sodium bicarbonate) reacts with the acid in vinegar (acetic acid) to produce carbon dioxide gas,

water, and sodium acetate. (c) Lime (calcium oxide) can remove sulfur dioxide, a potential air pollutant, from air by

forming calcium sulfite.

14.70 Determine the formula of each of the following compounds, and determine the oxidation number for each

element present. (a) silicon tetrachloride (b) germanium dioxide (c) sodium tetraborate (d) stannous fluoride

(e) diborane

14.71 Determine the formula of each of the following compounds, and determine the oxidation number for each

element present. (a) orthoboric acid (b) carbon monoxide (c) lead(IV) fluoride (d) potassium cyanide (e) boric oxide

14.72 Name or give the formula for each of the following. (a) potassium superoxide (b) sodium orthoborate (c)

sodium tetraborate (d) silicon carbide (e) hexafluorosilicic acid (f) D (g) NaBH4 (h) B2H6 (i) C60 (j) C22– (k) barium

peroxide (l) boron nitride (m) CH4 (n) barium acetylide (o) graphite

14.73 Name or give the formula for each of the following. (a) tritium (b) calcium borohydride (c) boron nitride

(d) diamond (e) calcium acetylide (f) RbO2 (g) BO33– (h) B4O7

2– (i) CH4 (j) Na2SiF6 (k) calcium orthoborate

(l) deuterium (m) Na2O2 (n) calcium tetraborate (o) H2B4O7



14.74 (a) List each of the elements in the nitrogen family and classify it as a metal, metalloid, or nonmetal.

(b) Which member of the nitrogen family is the most abundant element in the Earth's atmosphere? (c) List the

highest and lowest oxidation state for each member of the nitrogen family.

14.75 Draw a molecular orbital energy level diagram for the cyanide ion. Using this diagram, predict the bond order

of the ion and determine if it is paramagnetic or diamagnetic.

14.76 (a) Draw a Lewis structure for white phosphorus. (b) What is the bond angle predicted by VSEPR? (c) Why

is white phosphorus so unstable?

14.77 (a) Draw resonance structures for the azide ion (N3–) and the thiocyanate ion (SCN–). (b) For each of the

resonance structures, determine the formal charge on each atom.

14.78 Dimethylamine is a weak base related to ammonia. The titration of dimethylamine with hydrochloric acid

gives dimethylammonium chloride. (The analogous reaction of ammonia with hydrochloric acid produces

ammonium chloride.) (a) What is the molarity of a dimethylamine solution if 50.00 mL of this base reacts with

45.25 mL of 0.01275 M hydrochloric acid? The reaction is (CH3)2NH(aq) + HCl(aq) [(CH3)2NH2]+(aq) + Cl–(aq)

(b) Draw the Lewis structures for dimethylamine and the dimethylammonium ion.

14.79 (a) Draw the Lewis structures for water and ammonia. (b) What are the bond angles in water and ammonia?

(c) What are the intermolecular forces in each of these compounds?

14.80 Write balanced chemical equations for the formation of each of the following compounds. (a) N2O (b) NO

(c) N2O2 (d) NO2 (e) N2O3

14.81 Dinitrogen trioxide and dinitrogen pentoxide are both acid anhydrides. Write balanced chemical equations for

the reaction of each of these oxides to produce an acid.

14.82 Write a balanced chemical equation for each step of the Ostwald process.

14.83 Write balanced chemical equations for the reaction of sodium hydroxide with each of the following. Assume

that all the acid hydrogen atoms react. (a) phosphoric acid (b) phosphorous acid (c) pyrophosphoric acid

(d) tripolyphosphoric acid (e) trimetaphosphoric acid

14.84 (a) Define a condensation reaction. (b) Define hydrolysis.

14.85 For each of the following compounds, give the chemical formula and determine the oxidation state of each

element. (a) nitric acid (b) methylhydrazine (c) potassium cyanate (d) sodium nitrite (e) hydroxylamine

(f) magnesium nitride


element. (a) calcium nitrate (b) ammonium hydrogen sulfate (c) dinitrogen oxide (d) sodium thiocyanate (e) nitrous

acid (f) nitrogen oxide

14.87 Draw Lewis structures for each of the following and predict their molecular geometries. Determine the formal

charge of nitrogen in each substance. (a) NH4+ (b) HNO2 (c) N2O (one N is central) (d) NO2 (e) NaNO3

14.88 Complete and balance the following equations: (a) N2O3(g) + H2O(l) (b) Mg3N2(s) + H2O(l)

(c) NH3(aq) + HMnO4(aq) (d) (CH3)2N2H2(l) + O2(g) (e) NH4NO3(s) + heat


element. (a) tripolyphosphoric acid (b) arsenious acid (c) antimony(V) fluoride (d) magnesium dihydrogen

phosphate (e) potassium hydrogen phosphate


14.90 For each of the following compounds give the chemical formula and determine the oxidation state of each

element. (a) phosphine (b) pyrophosphoric acid (c) diantimony trioxide (d) copper(II) hydrogen arsenate

(e) diphosphorus trioxide

14.91 Write a balanced chemical equation for each of the following reactions. (a) It is possible to produce

phosphorus pentafluoride by a halogen exchange reaction of phosphorus pentachloride with arsenic trifluoride.

(b) Diarsenic trioxide reacts with water to produce arsenious acid. (c) Pyrophosphoric acid forms when heating

orthophosphoric acid until dehydration occurs. (d) Chlorine gas will oxidize antimony to antimony(III) chloride.

(e) Bismuth(V) fluoride may be produced by the direct reaction of bismuth metal with fluorine gas.

14.92 Nitrogen oxide gas is a pollutant in urban environments. In a particular city, the NO concentration is

determined to be 0.35 ppm. The air temperature the day of the measurement was 33°C and the air pressure was 755

torr. Determine the partial pressure of NO and the number of NO molecules present in a room measuring 5.0 m ×

7.5 m × 2.0 m.

14.93 Write the chemical formula for each of the following compounds and indicate the oxidation state of nitrogen

in each:

Formula Ox. State

(a) Nitrous acid ____________________ _____

(b) Calcium cyanide ____________________ _____

(c) Barium nitrate ____________________ _____

(d) Ammonium chloride ____________________ _____

(e) Magnesium nitride ____________________ _____

14.94 Write the chemical formula for each of the following and indicate the oxidation state of nitrogen in each of the

following: (a) Nitrous acid (b) Potassium cyanide (c) Sodium nitrate (d) Ammonium chloride (e) Lithium nitride

14.95 Describe the geometry of each of the following: (a) NH3 (b) HNO3 (c) N2O (d) HNO2

14.96 Write Lewis structures for each of the following species: (a) NH4+ (b) NO (c) N2O (d) NO2

14.97 For each of the following draw the Lewis structures and predict the relative N–O bond lengths: NO, NO2,

HNO2, NO3–, and NO4

3–.

14.98 Give the hybridization on the nitrogen in each of the following compounds: (a) NH3 (b) HNO3 (c) N2O

(d) HNO2 (a) sp3 (b) sp2 (c) sp (d) sp2


14.99 (a) Draw a molecular orbital energy level diagram for the oxygen molecule. Using this diagram, predict the

bond order of the molecule and determine if it is paramagnetic or diamagnetic. (b) Repeat the procedure in part (a)

for the peroxide ion and the superoxide ion.

14.100 (a) Draw Lewis structures for the two allotropes of oxygen. (b) Approximately how much of the Earth's

crust is oxygen? Approximately how much of the Earth’s atmosphere is oxygen?

14.101 Write balanced chemical equations for the formation of a base from a base anhydride, such as sodium oxide,

and for the formation of an acid from an acid anhydride, such as carbon dioxide.

14.102 (a) What is an important use for sodium thiosulfate? (b) Draw the Lewis structure for sodium thiosulfate.

14.103 Write balanced chemical equations for each step in the preparation of sulfuric acid from sulfur.


14.104 Draw Lewis structures for the oxyacids of sulfur, selenium, and tellurium where the element is in a +6

oxidation state.

14.105 Write the formulas for the products, and then balance each of the following.

(a) Cs2O(s) + H2O(l) (b) ZnO(s) + HCl(aq) (c) Li2O2(s) + H2O(l)

(d) N2O5(g) + H2O(l) (e) RbO2(s) + H2O(l) (f) NO2(g) + O3(g)

14.106 Write the balanced chemical equation for each of following reactions. (a) Solid mercury(II) oxide

decomposes to the elements when heated. (b) Heating solid copper(II) nitrate produces solid copper(II) oxide,

nitrogen dioxide gas, and oxygen gas. (c) Ozone will oxidize magnesium sulfide to magnesium sulfate. Dioxygen

is the other product. (d) Heating solid lead(II) sulfide in air produces sulfur dioxide gas and solid lead(II) oxide.

(e) Oxygen gas may form through the reaction of carbon dioxide gas with solid potassium peroxide. Potassium

carbonate is the other product. (f) Silver metal will partially react with ozone gas to produce solid silver(II) oxide

and oxygen gas.

14.107 Each of the following substances has polar bonds. Predict which member of each pair has the more polar

bonds. (a) SiO2 and PbO2 (b) P2O5 and V2O5 (c) CO2 and SO2 (d) SiF4 and SiBr4

14.108 Label each of the following oxides as acidic, basic, amphoteric, or neutral. (a) NO (b) SO2 (c) BaO (d) ZnO

(e) BeO

14.109 For each of the following compounds, give the chemical formula and determine the oxidation number of

each element. (a) sulfurous acid (b) potassium hydrogen selenite (c) hydrogen telluride (d) carbon diselenide

(e) barium sulfite (f) calcium thiosulfate

14.110 Draw the Lewis structure and predict the molecular geometry for each of the following. (a) the thiosulfate

ion, S2O32– (b) orthotelluric acid, H6TeO6 (c) gaseous selenium dioxide, SeO2 (d) disulfur dichloride, S2Cl2 (there is

a sulfur-sulfur bond) (e) methylsulfonic acid, HSO3CH3 (CH3– is bonded to the central sulfur)

(a)

(b) (c) (d) (e)

14.111 Write the balanced chemical equation for each of the following reactions. (a) Gaseous sulfur dioxide

dissolves in water. (b) Solid iron(II) sulfide reacts with hydrochloric acid to form hydrogen sulfide gas and aqueous

iron(II) chloride. (c) Aqueous thiosulfate ion forms when heating an aqueous solution containing the sulfite ion with

solid sulfur. (d) Gaseous sulfur dioxide undergoes air oxidation in the present of a catalyst to produce gaseous

sulfur trioxide. (e) Gaseous sulfur trioxide reacts with liquid sulfuric acid to produce liquid pyrosulfuric acid.

14.112 Ozone depletion in the upper atmosphere may occur through the following cycle of reactions. Calculate the

enthalpy change for each step and the overall enthalpy change for this cycle.

NO(g) + O3(g) NO2(g) + O2(g)

NO2(g) + O(g) O2(g) + NO(g)

14.113 Each of the following substances will react with water. In each case, write a balanced chemical equation.

(a) CO2(g) (b) Cl2O7(g) (c) K2O(s) (d) CaC2(s) (e) KO2(s) (f) Li3N(s) (g) Na2O2(s) (h) CaH2(s) (i) P4O10(s) (j) Ba(s).

14.114 Give the formula for the acid anhydride of each of the following. (a) H2SO3 (b) HClO4 (c) HNO3 (d) H2CO3

(e) H3PO4 (a) SO2 (b) Cl2O7 (c) N2O5 (d) CO2 (e) P2O5


14.115 Name or give the formula for each of the following. (a) cesium superoxide (b) sodium orthoborate (c)

tetraborate ion (d) silicon carbide (e) potassium hexafluorosilicate (f) 1H (g) KBH4 (h) B2O3 (i) C(dia) (j) C22–

(k) CH4 (l) graphite (m) B2H6 (n) CaC2 (o) borohydride ion

14.116 Name or give the formula for each of the following. (a) deuterium (b) sodium borohydride (c) boron nitride

(d) calcium acetylide (e) buckminsterfullerene (f) BaO2 (g) K3BO3 (h) K2B4O7 (i) CH4 (j) H2SiF6 (k) CsO2 (l) T

(m) protium (n) boron oxide (o) K2SiF6

14.117 Name or give the formula for each of the following. (a) potassium borohydride (b) diethylammonium ion

(c) methane (d) trimetaphosphoric acid (e) ammonium chloride (f) H3PO4 (g) NH2OH (h) P2O74– (i) B2O3 (j)

CH3NH3+ (k) chloramine (l) dimethylhydrazine (m) (C2H5)2NH (n) CH3NH2

14.118 Name or give the formula for each of the following. (a) tritium (b) tetraboric acid (c) ammonium

pyrophosphate (d) methylammonium chloride (e) ammonia (f) Ca(BH4)2 (g) H5P3O10 (h) (NH4)2SO4 (i) N2H4 (j) SiC

(k) sodium tripolyphosphate (l) NH2Cl (m) C60 (n) H3PO3 (o) orthoboric acid


14.119 (a) Why are many of the properties of astatine unknown? (b) What is the only nonmetal that is a liquid at

room temperature? (c) List the highest and lowest oxidation states for each of the halogens. (d) What color is each

of the halogens?

14.120 (a) Estimate the percent ionization of a 0.168 m aqueous solution of acetic acid (CH3CO2H) if the solution

freezes at –0.320°C. (b) Estimate the percent ionization of a 0.125 m aqueous solution of trichloroacetic acid

(CCl3CO2H) if the solution freezes at –0.420°C. (c) Draw the Lewis structures for acetic acid and trichloroacetic

acid. (d) These are both weak acids, with the one of lower percent ionization being the weaker of the two. Based

upon the Lewis structures, why is the percent ionization of one acid higher than percent ionization of the other?

(a) 2 % (b) 80.8 %

14.121 (a) Dichloromethane (CH2Cl2) is a nonelectrolyte that is very slightly soluble in water. Approximately 2.7 g

of dichloromethane will dissolve in 100.0 mL of water at 25.0 °C. Determine the osmotic pressure of this solution.

(b) What is the osmotic pressure of a solution made by dissolving 2.7 g of sodium chloride in 100.0 mL of water at

25.0 °C?

14.122 Give balanced chemical equations for the formation of chlorine, bromine, and iodine.

14.123 Define disproportionation.

14.124 (a) Bromine, like chlorine, undergoes disproportionation in cold water. Write a balanced chemical equation

for this reaction. (b) Bromine, like chlorine, undergoes disproportionation in hot base. Write a balanced chemical

equation for this reaction.

14.125 (a) Give the name and formula for each of the stable oxides of the halogens that contains the halogen in a

positive oxidation state. (b) Show the reaction of each of the halogen acid anhydrides with water to produce an acid.

14.126 Chlorine forms four oxyacids: hypochlorous acid, chlorous acid, chloric acid, and perchloric acid. The

strengths of these acids increase in the order HClO < HClO2 < HClO3 < HClO4. Draw the Lewis structure for each

of these acids and use these structures to explain the order of increase in acid strength.

14.127 (a) Draw the Lewis structure of the triiodide ion. (b) What is the molecular structure of the triiodide ion?

(c) What is the hybridization of the central iodine in the triiodide ion?

14.128 (a) What is an interhalogen? (b) Name and draw Lewis structures for the following compounds: ClF3, BrF5,

and IF7. (c) What is the molecular geometry of each of these compounds? (d) What is the hybridization of the

central halogen in each of these compounds? (e) Which of these compounds is polar?


14.129 Provide a short explanation of the following properties of the halogens. (a) The ability to oxidize other

substances decreases with increasing atomic number. (b) The melting points of the hydrogen halides increase in the

order HCl < HBr < HI < HF. (c) It is possible to produce all the halogens except fluorine by the electrolysis of an

aqueous solution of one of their salts. (d) At room temperature, both fluorine and chlorine are gases, whereas

bromine is a liquid and iodine is a solid. (e) The compounds KClO4, KBrO4, and KIO4 exist, but KFO4 does not

exist.

14.130 Fluorine and iodine are monoisotopic elements, meaning elements consisting of only one isotope. The

isotope of fluorine is fluorine-19, and the isotope of iodine is iodine-127. The halogens chlorine and bromine each

have two naturally occurring isotopes. The two isotopes of chlorine are chlorine-35 and chlorine-37. The two

isotopes of bromine are bromine-79 and bromine-81. Determine the number of protons and the number of neutrons

present in each of the halogen isotopes.

14.131 For each of the following reactions, write a balanced chemical equation. (a) Bromine reacts with water to

produce small amounts of hypobromite ion and bromide ion. (b) Bromine will react with an aqueous solution of

potassium iodide. (c) Bromine reacts with a hot aqueous solution of potassium hydroxide to produce a mixture of

potassium bromide and potassium bromate. (d) Potassium bromide reacts with phosphoric acid. (e) Hydrobromic

acid reacts with manganese(IV) oxide.

14.132 Bromine consists of two naturally occurring isotopes. Bromine-79, with a mass of 78.918338 amu, has an

abundance of 50.69 %, and bromine-81, with a mass of 80.916291 amu, has an abundance of 49.31%. Using these

data, calculate the atomic weight of bromine. 79.90 amu

14.133 The atomic weight of chlorine is 35.453 amu. Natural chlorine contains two isotopes. Chlorine-35 has a

mass of 34.96885271 amu, and chlorine-37 has a mass of 36.96590260 amu. Calculate the percent abundance of

each isotope.

14.134 For each of the following compounds give the chemical formula and determine the oxidation number of each

element. (a) iodite ion (b) bromous acid (c) chlorine trifluoride (d) sodium hypofluorite (e) paraperiodic acid


14.135 Which noble gas was discovered on the Sun before it was discovered on the Earth?

14.136 Write balanced chemical reactions for the formation of xenon difluoride, xenon tetrafluoride, and xenon

hexafluoride.

14.137 (a) Draw Lewis structures for xenon difluoride and xenon tetrafluoride. (b) What is the molecular geometry

of each of these compounds? (c) What is the hybridization of xenon in each of these compounds?

(a)

(b) XeF2 is linear and XeF4 is square planar. (c) XeF2 is sp3d and XeF4 is sp3d2.

14.138 (a) Draw Lewis structures for xenon trioxide, xenon tetroxide, and the perxenate ion. (b) What is the

molecular geometry of each of these compounds? (c) What is the hybridization of xenon in each of these

compounds?


14.139 Air contains the following mole fractions of the noble gases: 5.24 × 10–6 He, 1.82 × 10–5 Ne, 9.34 × 10–3 Ar,

1.14 × 10–6 Kr, and 8.70 × 10–8 Xe. (a) Determine the partial pressure of helium and xenon when the total pressure

is 775 mmHg. (b) How many moles of neon atoms are in 10.0 m3 of air at 25°C? (c) Calculate the moles of krypton

atoms in a room measuring 15.0 ft × 10.0 ft × 8.0 ft at 25°C. (d) How many argon atoms are in the room given in

part (c)? (a) 6.74 × 10–5 mmHg (b) 0.00759 mole Ne (c) 0.0016 mole Kr (d) 8.0 × 1024 Ar Atoms

14.140 The mole fraction of xenon in the air is 8.70 × 10–8. Calculate the number of xenon atoms in 1.0 L of air at a

pressure of 0.94 atm and a temperature of 300.0 K.

Summary

14.141 Complete and balance equations for the reaction of each of the following combinations. If no reaction

occurs, label the reaction with NR. If there is a reaction, write a net ionic equation. (a) NaOH(aq) + Mg(NO3)2(aq)

(b) Ca(OH)2(aq) + NH4Cl(aq) (c) BaCl2(aq) + (NH4)3PO4(aq) (d) KNO3(aq) + SrCl2(aq) (e) Pb(C2H3O2)2(aq) +

HI(aq)

14.142 Write balanced equations for each of the following, completing the reaction if necessary. (a) Potassium plus

water (b) Aluminum nitrate decomposes to aluminum oxide, nitrogen dioxide, and oxygen gas.

14.143 Write balanced chemical equations for each of the following. Complete the reactions where necessary.

(a) sodium plus water (b) calcium carbonate plus heat (c) magnesium plus oxygen (d) lithium plus hydrogen

(e) aluminum plus aqueous sodium nitrate (basic) produces ammonia plus NaAlO2

14.144 Complete and balance the following equations. (a) Sn(s) + HF(aq) (b) Tl3+(aq) + OH–(aq)

(c) Fe2O3(s) + Al(s) (d) Pb3O4(s) + C(s) + heat (e) Al(OH)3(s) + NaOH(aq)

14.145 Draw Lewis structures and predict the molecular geometry for each of the following. (a) I3– (b) BrF4

–

(c) ClO2– (d) H5IO6 (e) XeF4

14.146 Name or give the formula for each of the following. (a) deuterium (b) sodium borohydride (c) diborane

(d) acetylide ion (e) buckminsterfullerene (f) KO2 (g) Na3BO3 (h) Na2B4O7 (i) SiC (j) H2SiF6 (k) protium (l) boron

oxide (m) hexafluorosilicic acid (n) graphite (o) CsO2

14.147 Name or give the formula for each of the following. (a) rubidium superoxide (b) orthoborate ion

(c) tetraborate ion (d) methane (e) sodium hexafluorosilicate (f) T (g) Ca(BH4)2 (h) BN (i) C(dia) (j) CaC2 (k) C60

(l) barium peroxide (m) D (n) potassium tetraborate (o) potassium orthoborate

14.148 Name or give the formula for each of the following. (a) ammonium chloride (b) ammonium sulfate

(c) barium peroxide (d) calcium tetraborate (e) phosphorous acid (f) T (g) H3PO4 (h) (NH4)3PO4 (i) (C2H5)2NH2+

(j) H2SiF6 (k) chloramine (l) hydroxylamine (m) methylammonium ion (n) N2H3CH3 (o) Na5P3O10

14.149 Name or give the formula for each of the following. (a) calcium orthoborate (b) trimetaphosphoric acid

(c) hydrazine (d) sodium hexafluorosilicate (e) ammonium pyrophosphate (f) Na2O2 (g) NH2Cl (h) NH2OH

(i) CH3NH2 (j) P2O74– (k) CH3NH3Cl (l) NH3 (m) diamond (n) K2SiF6 (o) methylamine

14.150 Name or give the formula for each of the following. (a) pyrosulfuric acid (b) paraperiodic acid (c)

tetraborate ion (d) iron(II) disulfide (e) ammonium sulfate (f) O3 (g) KI3 (h) NaBH4 (i) H6TeO6 (j) P2O74– (k) XeF6

(l) orthotelluric acid (m) H3BO3 (n) potassium tetraborate (o) metaperiodic acid

14.151 Name or give the formula for each of the following. (a) xenon tetrafluoride (b) ammonium pyrophosphate

(c) selenous acid (d) potassium borohydride (e) calcium triiodide (f) KrF2 (g) H5IO6 (h) CH3NH2 (i) Na2B4O7 (j) S22–

(k) H2SeO4 (l) H2S2O7 (m) ozone

14.152 (a) What are two chemical properties that distinguish metals from nonmetals? (b) What are two physical

properties that distinguish metals from nonmetals?


14.153 Some periodic tables place hydrogen in both column 1 (1A) and column 17 (7A). What is the justification

for placing hydrogen in each of these two positions?

14.154 Both hydrogen (H2) and methane (CH4) are useful fuels. (a) Write balanced chemical equations for the

combustion of each of these gases to produce only gaseous products. (b) Calculate the heat of reaction for each of

your equations in part (a). (c) Compare the amount of heat produced by the combustion of 10.0 g of hydrogen to the

amount of heat produced by the combustion of 10.0 g of methane. (d) Compare the amount of heat produced by the

combustion of 10.0 L of hydrogen gas at 25°C and 760 torr to the amount of heat produced by the combustion of

10.0 L of methane at the same temperature and pressure.

14.155 A magnesium boride reacts with acid to form boric acid and a boron hydride. The boron hydride is 84.3

percent boron and 15.7 percent hydrogen. The density of a sample of this gas is 2.55 g/L at 27°C and 745 torr.

What is the molecular formula for the boron hydride? B5H10

14.156 There is a compound known with the formula NF5. The properties of this compound are not like those of

PF5. Draw a Lewis structure for NF5. (Hint: All atoms in the structure obey the octet rule.)

14.157 Elemental bromine and the interhalogen iodine chloride have about the same molar mass and certain similar

physical properties. However, some of the physical properties, such as the boiling point, are different. Account for

the differences in the physical properties.

14.158 Indium forms a diamagnetic solid with the empirical formula InCl2. Write the electron configuration for an

indium atom and for the indium +1, +2, and +3 ions. Use the electron configuration of the ions to describe the

diamagnetic nature of InCl2 and to predict the “molecular” formula.

14.159 (a) A 0.250 g sample of an unknown liquid produced 215 mL of vapor at 25°C and a pressure of 674 torr.

What is the molecular weight of the unknown liquid? (b) Combustion of the sample in pure oxygen produced 225

mL of nitrogen gas at 25°C and 644 torr. In addition, a small quantity of water formed. Heating the water produced

475 mL of vapor at 125°C and 815 torr. The unknown liquid only contained nitrogen and hydrogen. Determine the

empirical and molecular formula of the unknown liquid. (c) Draw the Lewis structure of the unknown liquid.

14.160 The halogens will react with rhenium metal (Re) to produce the following compounds: ReF7, ReCl6, ReBr5,

and ReI4. (a) Write balanced chemical equations for the formation of each of these rhenium compounds. (b) Name

each of these compounds. (c) Explain why the rhenium compounds formed with the halogens do not all have the

same general formula.

Chapter 15 Corresponds to BLBMWS Chapter 14 15.1 Corresponds to BLBMWS 14.1

15.2 Explain why the reaction of water vapor with sodium vapor is faster than the reaction of liquid water with solid

sodium.

15.3 In addition to the properties of reactants, what factors may influence the rate of a chemical reaction?

15.4 In general, how does an increase in temperature affect the rate of a reaction?

15.5 In general, how will a decrease in concentration affect the rate of a reaction?


15.6 Define reaction rate. What are the usual units for the reaction rate?

15.7 (a) Why do reactants have negative values for their reaction rates? (b) Why do products have positive values

for their reaction rates?


15.8 (a) Sketch a graph representing the change in concentration of a reactant verses time. (b) On your sketch,

indicate how you would determine the reaction rate.

15.9 (a) Sketch a graph representing the change in concentration of a product verses time. (b) On your sketch,

indicate how you would determine the reaction rate.

15.10 (a) What is a rate law? (b) Write an equation indicating the general form of a rate law.

15.11 Define the following terms as they apply to a rate law: (a) rate constant (b) order (c) overall order

15.12 Show how the rates of change in concentration for each substance in each of the following reactions relate to

each of the other substances in the reaction.

(a) 2 H2(g) + O2(g) 2 H2O(g) (b) Fe3+(aq) + 6 CN–(aq) Fe(CN)63–(aq)

(c) N2H4(g) + 2 F2(g) N2(g) + 4 HF(g) (d) 2 C4H10(g) + 13 O2(g) 8 CO2(s) + 10 H2O(g)

(e) 14 H+(aq) + 6 Fe2+(aq) + Cr2O72–(aq) 6 Fe3+(aq) + 2 Cr3+(aq) + 7 H2O(l)

15.13 Indicate how the rates of change for each of the substances in the following reactions relate to each of the

other substances: (a) C3H6(g) + Br2(g) C3H6Br2(g)

(b) CH4(g) + 2 O2(g) CO2(g) + 2 H2O(g)

(c) 2 CH3OH(g) + 8 F2(g) 2 CF4(g) + 8 HF(g) + O2(g)

(a) −∆[𝑪𝟑𝑯𝟔]

∆𝒕 = −

∆[𝑩𝒓𝟐]

∆𝒕 =

∆[𝑪𝟑𝑯𝟔𝑩𝒓𝟐]

∆𝒕 (b) −

∆[𝑪𝑯𝟒]

∆𝒕 = −

𝟏

𝟐

∆[𝑶𝟐]

∆𝒕 =

∆[𝑪𝑶𝟐]

∆𝒕 =

𝟏

𝟐

∆[𝑯𝟐𝑶]

∆𝒕

(c) −𝟏

𝟐

∆[𝑪𝑯𝟑𝑶𝑯]

∆𝒕 = −

𝟏

𝟖

∆[𝑭𝟐]

∆𝒕 =

𝟏

𝟐

∆[𝑪𝑭𝟒]

∆𝒕 =

𝟏

𝟖

∆[𝑯𝑭]

∆𝒕 =

∆[𝑶𝟐]

∆𝒕

15.14 The concentration of H+ was determined at various times for the following reaction:

CH3OH(aq) + H+(aq) + Cl–(aq) CH3Cl(aq) + H2O(l)

Time (min) [H+] (M)

0.00 2.00

50.0 1.86

150.0 1.64

300.0 1.39

650.0 1.02

(a) Determine the average rate from the initial and final times. (b) From each of the time intervals between

successive measurements, determine the average rate of disappearance for the hydrogen ion. (c) Construct a graph

of the hydrogen ion concentration versus time. (d) Draw a tangent at 100 minutes and determine the rate at that

time. (e) Draw a tangent at 325 minutes and determine the rate at that time.

15.16 The following reaction is second order in NO and first order in H2:

2 NO(g) + 2 H2(g) N2(g) + 2 H2O(g) at 730°C

(a) Write the rate law for this reaction. (b) The rate constant (k) is 6.0 × 104 M–2s–1. Calculate the rate of this

reaction if [NO] = 0.10 M and [H2] = 0.10 M. (c) Calculate the reaction rate after doubling the hydrogen

concentration to 0.20 M while the nitrogen oxide concentration remains at 0.10 M. (d) Calculate the reaction rate if

the nitrogen oxide concentration is doubled to 0.20 M while the hydrogen concentration remains at 0.10 M.

(a) Rate = k [NO]2 [H2] (b) 6.0 × 101 M/s (c) 1.2 × 102 M/s (d) 2.4 × 102 M/s

15.17 An investigation of the following reaction gave the data in the table for the rate of appearance of NO2:

2 NO(g) + O2(g) 2 NO2(g)

____Initial concentration___ Initial rate

Experiment [NO] (M) [O2] (M) (M/s)_____

1 0.0150 0.0150 2.40 × 10–2

2 0.0300 0.0300 1.92 × 10–1

3 0.0300 0.0150 9.59 × 10–2

(a) What is the rate law for this reaction? (b) What is the overall order of this reaction? (c) Calculate the rate

constant for each of the experiments. (d) What is the average value of the rate constant?


15.18 The rate of disappearance of ammonia (NH3) was determined for the following reaction:

BF3(g) + NH3(g) H3NBF3(g)

__Initial concentration__

Initial rate

Experiment [BF3] (M) [NH3] (M) (M/s)_____

1 0.125 0.125 5.32 × 10–2

2 0.125 0.250 1.065 × 10–1

3 0.250 0.100 8.52 × 10–2

4 0.450 0.100 1.534 × 10–1

5 0.225 0.100 7.67 × 10–2

(a) Determine the rate law for this reaction. (b) Calculate the value of the rate constant for this reaction.

(c) Calculate the rate of reaction when the boron trifluoride (BF3) concentration is 0.300 M and the ammonia (NH3)

concentration is 0.300 M.

15.19 The peroxydisulfate ion (S2O82–) will oxidize iodide ion (I–) according to the following equation:

S2O82–(aq) + 3 I–(aq) 2 SO4

2–(aq) + I3–(aq)

The rate of appearance of the triiodide ion (I3–) was measured in the following series of experiments:

Initial concentration

∆[I3–]/∆t

Experiment [S2O82–] (M) [I–] (M) (M/s)___

1 0.075 0.080 3.9 × 10–5

2 0.150 0.080 7.9 × 10–5

3 0.075 0.040 2.0 × 10–5

(a) Determine the rate law for this reaction. (b) Calculate the value of the rate constant for the formation of the

triiodide ion. (c) Determine the rate of appearance of the triiodide ion if the peroxydisulfate ion concentration is

0.100 M and the iodide ion concentration is 0.100 M. (d) Calculate the rate of disappearance of the iodide ion if the

peroxydisulfate ion concentration is 0.080 M and the iodide ion concentration is 0.120 M.

15.20 Determine the value for the rate constant for the following reaction:

HI

2 H2O2(aq) 2 H2O(l) + O2(g)

A series of experiments yielded the following experimental data.

Concentration (M)______ Initial rate of

[H2O2] [I–] [H+] formation of O2 (M/s)

0.40 0.010 0.020 3.5 × 10–4

0.80 0.010 0.020 7.0 × 10–4

0.40 0.020 0.020 1.4 × 10–3

0.40 0.020 0.040 2.8 × 10–3

15.21 It is possible to remove the pollutant hydrogen sulfide (H2S) from industrial wastewater by oxidation with

chlorine in the reaction

H2S(aq) + Cl2(aq) S(s) + 2 H+(aq) + 2 Cl–(aq)

This reaction is first-order in each of the reactants. At 25°C, the disappearance of hydrogen sulfide has a rate

constant of 3.5 × 10–2 /M s. If a wastewater sample has [Cl2] = 0.50 M and [H2S] = 1.5 × 10–3 M, calculate the rate

of formation of H+.

15.22 Molecules of cyclopentadiene (C5H6) will react with other molecules of this substance to form

dicyclopentadiene (C10H12). This process is an example of dimerization reaction. The reaction is:

2 C5H6 C10H12

Varying the concentration of cyclopentadiene in a solution gave the following data.

Time (s) [C5H6] (M)

0.00 0.0600

75.0 0.0347

150.0 0.0244

225.0 0.0188

300.0 0.0153


(a) Plot the above data as concentration versus time. (b) Plot the above data as ln [concentration] versus time.

(c) Plot the above data as 1 / [concentration] versus time. (d) Based on your graphs, what is the order of this

reaction? (e) Calculate the value of the rate constant.

15.23 Gaseous chloromethane (CH3Cl) reacts with water vapor,( H2O) to produce hydrogen chloride (HCl) and

methanol (CH3OH) vapor. The reaction is

CH3Cl(g) + H2O(g) CH3OH(g) + HCl(g)

Initial concentration

[CH3Cl] [H2O] Initial rate of HCl formation

Experiment (M) (M) (M/s)____________________

1 0.250 0.250 2.838

2 0.375 0.250 4.256

3 0.250 0.375 6.384

4 0.250 0.125 0.709

5 0.375 0.0625 0.266

(a) Determine the rate law for this reaction. (b) Determine the overall order of this reaction. (c) Calculate the rate

constant for the reaction.

15.24 The reaction of iodide ion (I–) with hypochlorite ion (ClO–) will take place in basic solution. The overall

reaction is

I–(aq) + ClO–(aq) IO–(aq) + Cl–(aq)

The following data were obtained at a certain temperature:

______Initial concentration______

Experiment [I–] [ClO–] [OH–] Initial rate of IO– formation

(M) (M) (M) (M/s)___________________

1 0.0015 0.020 0.20 9.0 × 10–3

2 0.0030 0.020 0.20 1.8 × 10–2

3 0.0015 0.010 0.20 4.5 × 10–3

4 0.0015 0.030 0.20 1.36 × 10–2

5 0.0015 0.020 0.10 1.81 × 10–2

6 0.0015 0.020 0.40 4.5 × 10–3

7 0.0015 0.030 0.40 6.8 × 10–3

(a) What is the rate law for this reaction? (b) What is the overall order of this reaction? (c) Calculate the rate

constant for each of the experiments. (d) What is the average value of the rate constant?


15.25 What is an integrated rate law, and why is it useful?

15.26 Write the integrated rate law for each of the following: (a) a reaction that is first order in A (b) a reaction that

is second order in A

15.27 Define each of the terms in the following zero-order integrated rate law.

[A]0 – [A]t = kt

15.28 It is possible to determine the order of a reaction graphically. Plotting the correct terms will produce a

straight-line plot. The x-axis for such a graph is always time and the y-axis will be related to the concentration.

(a) What concentration term should you use for the y-axis for a first-order reaction? (b) What concentration term

should you use for the y-axis for a second-order reaction? (c) What concentration term should you use for the y-axis

for a zero-order reaction?

15.29 Define half-life.

15.30 Write the different forms of the half-life equation for first-, second-, and zero-order reactions.


15.31 The first-order decomposition of gaseous dinitrogen pentoxide (N2O5) to nitrogen dioxide (NO2) and oxygen

(O2) has a rate constant of 4.9 × 10–4 s–1 at a certain temperature. Calculate the half-life of this reaction. 1.4 × 103 s

15.32 Molecules of butadiene (C4H6) can dimerize to C8H12. If the rate equation is second order in butadiene with a

rate constant of 0.014 L/mol s, how many seconds, will it take for the concentration of butadiene to drop from 0.020

M to 0.0020 M?

15.33 The dimerization of butadiene (C4H6) is second order in the concentration of butadiene (see Problem 15.32).

During one particular experiment, the butadiene concentration changed from 0.100 M to one-half of that value

(0.0500 M) in 714 s. How long would it have taken the concentration to change from an original concentration of

0.200 M to 0.0200 M?

15.34 The following reaction is second order in NO2 with a rate constant of 0.54 M–1s–1:

2 NO2(g) 2 NO(g) + O2(g)

If the initial concentration of nitrogen dioxide is 2.50 M, how long will it take the concentration to drop to 30.0

percent of the initial concentration?

15.35 In solution, dinitrogen pentoxide (N2O5) decomposes to nitrogen dioxide (NO2) and oxygen (O2). The

following data were obtained at a certain temperature:

Time (min) [N2O5] (M)

0.0 1.000

10.0 0.700

20.0 0.490

40.0 0.240

80.0 0.057


(c) Plot the above data as 1/[concentration] versus time. (d) Based on your graphs, what is the order of this reaction?

(e) Calculate the value of the rate constant.

15.36 Nitrogen dioxide gas (NO2) decomposes to nitrogen oxide (NO) and oxygen (O2). The following data were

obtained at 650 K:

Time (s) [NO2] (M)

0.00 2.00 × 10–1

10.0 9.1 × 10–3

20.0 4.7 × 10–3

30.0 3.1 × 10–3

40.0 2.4 × 10–3


(c) Plot the above data as 1 / [concentration] versus time. (d) Based on your graphs, what is the order of this

reaction? (e) Calculate the value of the rate constant.

15.37 (a) A certain first-order reaction is 35.5 percent complete in 4.90 min at 25°C. What is the rate constant?

What is the half-life? (b) A certain first-order reaction is 42.5 percent complete in 5.14 min at 35°C. Calculate the

value of the rate constant. What is the half-life?

15.38 The decomposition of hydrogen iodide to hydrogen gas and iodine vapor will take place on a heated gold

wire. The reaction is zero order with a rate constant of 1.20 × 10–4 M/s at 150°C. If the initial concentration of

hydrogen iodide is 0.135 M, how long will it take the concentration to drop to 75.0% of this value? What is the

half-life of 0.135 M hydrogen iodide in this reaction? 562 s


15.39 Define activation energy and reactive site.

15.40 Why is it possible for two reactant molecules to collide and not react?


15.41 (a) By what factor does a 10°C increase in temperature change the rate of a reaction? (b) By what factor does

a 20°C increase in temperature change the rate of a reaction? (c) By what factor does a 10°C decrease in

temperature change the rate of a reaction?

15.42 In terms of kinetics, why does refrigerating food retard spoilage?

15.43 What two conditions must two molecules meet to have an effective collision?


15.44 Define an activated complex or transition state.

15.45 (a) Sketch a general energy diagram for an exothermic reaction. (b) Sketch a general energy diagram for an

endothermic reaction. (c) On each of your diagrams, indicate the reactants, products, transition state, heat of

reaction, and activation energy.

15.46 (a) How does the activation energy of a reverse reaction compare to the activation energy for the forward

reaction? Assume the forward reaction is exothermic. (b) In this case, will the rate of the reverse reaction be greater

than that of the forward reaction? Why or why not?

15.47 The reaction CH4(g) + 2 S2(g) CS2(g) + 2 H2S(g)

This reaction has an activation energy, Ea, of 140 kJ/mol. The ∆H for the reaction is –106.5 kJ/mol. (a) Draw an

energy profile for this reaction. (b) Determine the activation energy of the reverse reaction.


15.48 (a) Why is the Arrhenius equation useful? (b) Write the Arrhenius equation. (c) Define each of the terms in

the Arrhenius equation, including the value and units for R.

15.49 Why might it be difficult to measure the frequency factor for a reaction?

15.50 What are the two general characteristics of the activation energy?

15.51 It is difficult to solve the Arrhenius equation directly for the activation energy. In order to determine the

activation energy, it is convenient to determine the rate constant at two different temperatures. If the temperatures

are T1 and T2, write the Arrhenius equation in the form used to determine the activation energy.

15.52 What are the two reasons why an increase in temperature will increase the rate of a reaction?

15.53 The following reaction has an activation energy, Ea, of 140 kJ/mol.

CH4(g) + 2 S2(g) CS2(g) + 2 H2S(g)

The rate constant for this reaction is 6.4 M–1 s–1 at 625°C. (a) Calculate the value of the rate constant, k, at 575°C.

(b) Calculate the value of the rate constant, k, at 775°C.

15.54 A certain reaction has a rate constant, k, of 2.5 × 10–1 s–1 at 25.0°C. Determine the value of the rate constant

at 50.0°C for this reaction if the value of the activation energy is (a) = 53.5 kJ/mol (b) 145 kJ/mol.


15.55 Determine the activation energy for the following reaction using the data given below.

NO(g) + O3(g) NO2(g) + O2(g)

Temperature Rate constant

(°C) (M–1 s–1)

–78.0 1.08 × 109

–50.0 2.74 × 109

0.00 8.99 × 109

25.0 1.40 × 1010

100.0 3.70 × 1010

15.56 When an egg cooks, the protein albumin in it denatures, and the time required to achieve a particular degree of

denaturation is related to the inverse of the rate constant of the process. The activation energy of this process is 41.8

kJ/mol. If it takes 3.00 min to cook an egg at sea level (T = 100.°C, implying k = 1/3.00 min), how long would it

take to cook the same egg on top of Mt. McKinley in Alaska, where the water boils at 80.0°C?


15.57 (a) What is the definition of a reaction mechanism? (b) What is an elementary step? (c) What is the rate-

determining step in a mechanism?

15.58 (a) What are the definitions of a unimolecular step and a bimolecular step? (b) Why are termolecular steps

(three molecules colliding) unlikely to occur in a mechanism?

15.59 Why is the following an unlikely step in a mechanism?

2 H2(g) + O2(g) 2 H2O(g)

15.60 What are the three rules that every mechanism must follow?

15.61 Define an intermediate.

15.62 The reaction of nitrogen dioxide (NO2) with fluorine (F2) is thought to proceed by the following mechanism:

NO2(g) + F2(g) NO2F(g) + F(g)

NO2(g) + F(g) NO2F(g)

(a) Determine the overall reaction from these elementary steps. (b) Determine the rate law for each step in the

mechanism. (c) What intermediates, if any, are present? (d) The rate law for this reaction is Rate = k [NO2] [F2].

Which step in the mechanism is the rate-determining step? (a) 2 NO2(g) + F2(g) 2 NO2F(g)

(b) Rate = k [NO2] [F2] – first step, Rate = [NO2] [F2]1/2 – second step (c) F(g) (d) The first step is rate-

determining.

15.63 Nitrogen oxide (NO) reacts with bromine (Br2) according to the following reaction:

Br2(g) + 2 NO(g) 2 NOBr(g)

The observed rate law for this reaction is Rate = k [Br2][NO]2. Which of the following mechanisms would be

consistent with the observed rate law?

(a) Br2(g) + 2 NO(g) 2 NOBr(g)

(b) Br2(g) + NO(g) NOBr2(g) (slow)

NOBr2(g) + NO(g) 2 NOBr(g) (fast)

(c) Br2(g) + NO(g) NOBr2(g) (fast)

NOBr2(g) + NO(g) 2 NOBr(g) (slow)

(d) Br2(g) 2 Br(g) (slow)

2 [Br(g) + NO(g) NOBr(g)] (fast)

(e) 2 NO(g) N2O2(g) (fast)

N2O2(g) + Br2(g) 2 NOBr(g) (slow)


15.64 The following mechanism has been proposed for the gas-phase reaction of methane (CH4) and chlorine (Cl2):

Step 1: Cl2(g) 2 Cl(g) (fast) k1 and k–1

Step 2: Cl(g) + CH4(g) HCl(g) + CH3(g) (slow) k2

Step 3: Cl(g) + CH3(g) CH3Cl(g) (fast) k3

(a) Determine the overall reaction. (b) Identify the intermediates, if any, in this mechanism. (c) Determine the

molecularity of each step. (d) Identify the rate-determining step. (e) Predict the rate law for this mechanism.

15.65 The reaction 2 ICl(g) + H2(g) I2(g) + 2 HCl(g) at temperatures above 200°C is first order in H2 and first

order in ICl. Suggest a mechanism of two steps with the first step rate determining to account for the rate law.

15.66 A reaction has the following mechanism:

1. 2 NO(g) N2O2(g)

2. N2O2(g) + H2(g) N2O(g) + H2O(g)

3. N2O(g) + H2(g) N2(g) + H2O(g)

(a) Write the net reaction. (b) Write the rate equation if step 1 is the slow step. (c) Write the rate equation if step 2

is the slow step. (a) 2 NO(g) + 2 H2(g) N2(g) + 2 H2O(g) (b) Rate = k [NO]2 (c) Rate = k [NO]2 [H2]

15.68 A reaction has the following proposed mechanism.

1. 2 A B

2. B + C D + E

The activation energies for each step are 35 kJ/mol (step 1) and 25 kJ/mol (step 2). The overall reaction is

exothermic by 25 kJ/mol. (a) Sketch an energy profile corresponding to this mechanism. (b) Which is the rate-

determining step?

15.69 A reaction has the following proposed mechanism.

1. 2 A B

2. B + C D + E

3. D + C F + E

The activation energies for each step are 25.0 kJ/mol (step 1), 15.0 kJ/mol (step 2), and 35.0 kJ/mol (step 3).

(a) Sketch an energy profile corresponding to this mechanism. (b) Which is the rate-determining step?


15.70 Define a catalyst.

15.71 (a) What is a homogeneous catalyst? (b) What is a heterogeneous catalyst?

15.72 (a) Sketch an energy diagram for an exothermic reaction. (b) Sketch an energy diagram for the same reaction

if a catalyst is present.

15.73 Is it possible for a catalyst to slow down a reaction? Why or why not?

15.74 A catalyst lowers the activation energy of a reaction. How does it change the activation energy of the reverse

reaction?

15.75 In the upper atmosphere, chlorofluorocarbons may decompose to produce, among other things, chlorine

atoms. Chlorine atoms interact with, and destroy, ozone by the following proposed mechanism:

Cl(g) + O3(g) ClO(g) + O2(g)

ClO(g) + O(g) Cl(g) + O2(g)

(a) Combine the two steps of this mechanism to produce the overall reaction. (b) Identify each substance in the

mechanism as a reactant, a product, an intermediate, or a catalyst.

Corresponds to BLBMWS 14.7

15.76 (a) Define enzyme. (b) Is an enzyme likely to be a homogeneous catalyst or a heterogeneous catalyst?


15.77 What is an enzyme inhibitor?

15.78 For the following reaction the rate of the uncatalyzed decomposition is slow and −∆[H2O2]

∆𝑡 = 1 × 10–8 M/s.

2 H2O2(aq) 2 H2O(aq) + O2(g)

In the presence of hemoglobin, which acts as a catalyst, the rate increases from 1 × 10–8 M/s to 1 × 10–1 M/s?

(a) What is the rate of oxygen formation for the uncatalyzed reaction? (b) What is the rate of oxygen production for

the catalyzed reaction?

Summary

15.79 The bromate ion will oxidize bromide ions to elemental bromine. It is necessary to add an acid to catalyze the

reaction. Using the data in the following table, determine the rate law.

______Initial concentration_______

Experiment [BrO3–](M) [Br–](M) [H+](M) Relative rate

1 0.05 0.25 0.30 1

2 0.05 0.25 0.60 4

3 0.10 0.25 0.60 8

4 0.05 0.25 0.60 4

5 0.05 0.50 0.30 2

15.80 Calculate the value of the rate constant for the following reaction:

CH3COCH3 + Br2 AcidCH3COCH2Br + H+ + Br–

Use the following experimental data from a series of experiments investigating this reaction.

Concentrations (M) Initial Rate

[CH3COCH3] [Br2] [H+] [mol/L·s]__

0.30 0.05 0.05 5.7 × 10–5

0.30 0.10 0.05 5.7 × 10–5

0.30 0.05 0.10 12.0 × 10–5

0.40 0.05 0.20 31.0 × 10–5

0.40 0.05 0.05 7.6 × 10–5

15.81 Cyclobutane (C4H8) decomposes to ethene (C2H4). Determine the half-life of this first-order reaction from the

following data.

Initial

[C4H8] Rate (L/mol·s)

0.0100 9.20 × 10–5

0.0200 1.84 × 10–4

0.0300 2.76 × 10–4

15.82 The following reaction occurs in basic solution:

I–(aq) + OCl–(aq) OI–(aq) + Cl–(aq)

Determine the rate equation for this reaction from the following data.

[ClO–] [I–] [OH–] Rate (mol/ L ·s)

4.00 × 10–3 2.00 × 10–3 1.00 4.80 × 10–4

2.00 × 10–3 4.00 × 10–3 1.00 4.80 × 10–4

2.00 × 10–3 2.00 × 10–3 1.00 2.40 × 10–4

2.00 × 10–3 2.00 × 10–3 0.500 4.80 × 10–4

2.00 × 10–3 2.00 × 10–3 0.250 9.60 × 10–4

Rate = k [ClO–] [I–] [OH–]–1


15.83 A series of experiments investigating the reaction of iodate ion with iodide ion in an acid solution yielded the

following data.

[IO3–] [I–] [H+] Rate (mol/L·s)

2.25 × 10–3 1.8 × 10–3 1.4 × 10–4 2.38 × 10–7

2.25 × 10–3 3.6 × 10–3 1.4 × 10–4 9.07 × 10–7

4.50 × 10–3 1.8 × 10–3 1.4 × 10–4 4.77 × 10–7

2.25 × 10–3 1.8 × 10–3 0.7 × 10–4 0.60 × 10–7

2.25 × 10–3 1.8 × 10–3 2.8 × 10–4 8.95 × 10–7

The reaction was IO3–(aq) + 8 I–(aq) + H+(aq) 3 I3

–(aq) + 3 H2O(l)

Find the rate law and calculate the rate constant. (Note: round off orders to the nearest whole numbers.)

15.84 The dimerization reaction of butadiene (C4H6) to produce C8H12 is second order in butadiene. In one

experiment, butadiene concentration changed from 0.100 M to one-half of that value (0.0500 M) in 714 s. How

much time would it take the concentration to change from the original 0.100 M to 0.0100 M?

15.85 A first-order reaction is 28.2 percent complete after 3.57 min at 25°C. Calculate the value of the rate constant

for this reaction.

15.86 The following reaction illustrates the decomposition of acetaldehyde:

CH3CHO(g) CH4(g) + CO(g)

This reaction is second order with a rate of 0.18 mol L–1 s–1 at a certain temperature. Calculate the rate constant at

this temperature if the initial concentration of acetaldehyde is 0.10 M.

15.87 The rate constant, k, for a certain reaction is 2.5 × 10–2 s–1 at 0.0°C. Determine the value of the rate constant

at 65.0°C if Ea = 135 kJ/mol.

15.88 The rate law for the decomposition of dinitrogen oxide,

2 N2O(g) 2 N2(g) + O2(g)

is second order in N2O. When this reaction occurs at 900 K with an initial N2O concentration of 2.0 × 10–2 M, it

took 4500 s for the N2O concentration to fall to half its initial concentration. Calculate the value of the rate constant

for this reaction.

15.89 For each of the following reactions, indicate the relationship between the rates of disappearance of each

reactant to the rate of appearance of each product:

(a) N2(g) + 3 F2(g) 2 NF3(g)

(b) 2 NO(g) + Br2(g) 2 NOBr(g)

(c) 2 N2O(g) 2 N2(g) + O2(g)

15.90 For each of the following reactions, express the rate in terms of the change in concentration with time of each

reactant, and relate this to the rate of formation of each product.

(a) 2 C4H10(g) + 13 O2(g) 10 H2O(g) + 8 CO2(g)

(b) 5 Br–(aq) + BrO3–(aq) + 6 H+(aq) 3 Br2(aq) + 3 H2O(l)

(a) −𝟏

𝟐

∆[𝑪𝟒𝑯𝟏𝟎]

∆𝒕 = −

𝟏

𝟏𝟑

∆[𝑶𝟐]

∆𝒕 =

𝟏

𝟖

∆[𝑪𝑶𝟐]

∆𝒕 =

𝟏

𝟏𝟎

∆[𝑯𝟐𝑶]

∆𝒕 (b) −

𝟏

𝟔

∆[𝑯+]

∆𝒕 = −

𝟏

𝟓

∆[𝑩𝒓−]

∆𝒕 = −

∆[𝑩𝒓𝑶𝟑−]

∆𝒕 =

𝟏

𝟑

∆[𝑩𝒓𝟐]

∆𝒕 =

𝟏

𝟑

∆[𝑯𝟐𝑶]

∆𝒕

15.91 The following reaction has a rate constant of 1.63 × 10–1 M–1s–1 at a certain temperature:

2 ICl(g) + H2(g) I2(g) + 2 HCl(g)

This reaction is first order in each of the reactants. Calculate the rate of reaction if [ICl]= 0.33 M and [H2] = 0.20 M.

15.92 An experiment investigating the oxidation of iodide ion by permanganate ion in acid solution, shown below,

found the rate of disappearance of the permanganate ion to be 4.56 × 10–3M/s. Determine the rate of appearance of

iodine under these conditions.

10 I–(aq) + 2 MnO4–(aq) + 16 H+(aq) 2 Mn2+(aq) + 5 I2(aq) + 8 H2O(l) 0.0114 M/s


15.93 For the following reaction the rate law is Rate = k [NH4+] [NO2

–].

NH4+(aq) + NO2

–(aq) N2(g) + 2 H2O(l)

At room temperature, the rate constant is 3.0 × 10–4/M·s. Determine the rate of reaction, at this temperature, if

[NH4+] = 0.26 M and [NO2

–] = 0.080 M.

15.94 The following reaction occurs in the gas phase.

2 NO(g) + O2(g) 2 NO2(g)

This reaction has the following rate law: rate = k [NO]2[O2]. At 25°C, the rate constant is 7.1 × 109 L2·mol–2·s–1.

Determine the reaction rate when [NO] = 0.0020 M and [O2] = 0.027 M.

15.95 The following first-order reaction illustrates the decomposition of N2O5 in a carbon tetrachloride solution.

2 N2O5(sol) 4 NO2(sol) + O2(g)

The rate constant is 6.08 × 10–4 s–1 at 45°C. Determine the rate of reaction when [N2O5] = 0.295 M. (The term "sol"

refers to molecules in carbon tetrachloride solution.)

15.96 The following reaction has an activation energy of 102 kJ/mol:

2 N2O5(g) 2 NO2(g) + O2(g)

At 45°C, the rate constant for this reaction is 5.0 × 10–4 s–1. Determine the value of the rate constant at 75°C.

0.014 s–1

15.97 Studies of a certain first-order reaction yielded the following data.

Temperature Rate constant

(K) (s–1)_________

300. 1.00 × 10–5

320. 5.00 × 10–5

340. 2.00 × 10–4

355. 5.00 × 10–4

Determine the value of the activation energy for this reaction (in kJ/mol).

15.98 Using the data in the following table determine the activation energy, in kJ/mol, for the following reaction.

[Mn(CO)5(CH3CN)]+(aq) + NC5H5(aq) [Mn(CO)5(NC5H5)]+(aq) + CH3CN(aq)

T (°C) k, min–1

25 0.0409

35 0.0818

45 0.157

15.99 The following reaction occurs in basic solution:

3 BrO–(aq) BrO3–(aq) + 2 Br–(aq)

This reaction is second order in BrO– and has a rate constant of 0.056 M–1s–1 at 80.°C. In one experiment, the initial

concentration of BrO– was 0.225 M. In this experiment, what will be the concentration of BrO– after 1.00 minute?

15.100 The following equation illustrates the decomposition of hydrogen iodide to the elements:

2 HI(g) H2(g) + I2(g)

This is a second-order reaction that has a rate constant equal to 1.6 × 10–3 L/mol·s at 973 K. In one experiment, the

hydrogen iodide concentration was 2.9 × 10–4 M after 3.0 × 103 minutes. What was the initial molarity of HI?

15.101 The decomposition of aqueous hydrogen peroxide at 70°C is first order in hydrogen peroxide. This reaction

has a half-life of 20.0 min. The reaction is

2 H2O2(aq) 2 H2O(l) + O2(g)

In one experiment, the initial concentration of hydrogen peroxide was 0.100 M. What was the hydrogen peroxide

concentration after 25.0 min?

15.102 The rate constant for the disappearance of NOBr in the following second-order reaction is 0.80/M·s at 10°C:

2 NOBr(g) 2 NO(g) + Br2(g)

If the initial concentration of NOBr was 0.186 M, calculate the concentrations of each substance after 1.50 s.


15.103 Molecules of butadiene (C4H6) can dimerize to C8H12. If the rate equation is second order in butadiene with

a rate constant of 0.014 L/mol·s, calculate the half-life if the initial concentration of butadiene is 0.018 M.

15.104 The following reaction has a rate constant of 6.2 × 10–4 min−1 when dissolved in carbon tetrachloride at

45°C. 2 N2O5 4 NO2 + O2

This reaction is first order. Calculate the half-life of this reaction in seconds.

15.105 The following reaction occurs in the gas phase.

2 NO(g) + Br2(g)2 NOBr(g)

This reaction follows the rate law k [NO]2[Br2]. Derive a possible mechanism.

15.106 A reaction has the following mechanism:

1. 2 NON2O2

2. N2O2 + H2 N2O + H2O

3. N2O + H2N2 + H2O

(a) Write the net reaction. (b) Write the rate equation if step 1 is the slow step. (c) Write the rate equation if step 2

is the slow step.

15.107 Write molecular and net ionic equations for each of the following reactions. Assume that all the reactions

take place in aqueous solution. (a) Potassium hydroxide plus hydrochloric acid (b) Barium hydroxide plus nitric

acid (c) Sodium bicarbonate plus chloric acid (d) Rubidium sulfite plus hydrobromic acid (e) Ammonium carbonate

plus hydrochloric acid

15.108 For the following reaction

NO2(g) + CO(g) NO(g) + CO2(g)

The proposed mechanism is

NO2(g) + NO2(g) NO3(g) + NO(g) slow

NO3(g) + CO(g) NO2(g) + CO2(g) fast

What is the rate equation for this reaction?

15.109 How many grams of calcium chloride are necessary to prepare 775 mL of a 0.150 M calcium chloride

solution if the solid calcium chloride is only 95.0% pure?

15.110 Circle the more ionic compound in each of the following pairs.

(a) Magnesium chloride Beryllium chloride

(b) Tin(IV) chloride Tin(II) chloride

(c) Aluminum chloride Rubidium chloride

(d) Sodium chloride Bromine trichloride

(e) Lead(II) chloride Lead(IV) chloride

15.111 Write the balanced chemical equations for each of the following reactions.

(a) Aluminum metal reacts with hydrochloric acid to form hydrogen gas.

(b) Steam reacts with magnesium metal to give solid magnesium oxide and hydrogen gas.

(c) Hydrogen gas will reduce manganese(IV) oxide to manganese(II) oxide.

(d) Lithium hydride reacts with water to generate hydrogen gas.

15.112 Give the name or formula for each of the following:

(a) Manganese(IV) oxide (f) HNO2

(b) Calcium hydride (g) KO2

(c) Ammonium chloride (h) NO2

(d) Hydrocyanic acid (i) CH4

(e) Nitric acid (j) Li3N

15.113 A 1.5060 g sample of bronze, an alloy of copper and tin, dissolves in nitric acid to yield 0.3450 g of SnO2.

What is the mass percent of tin in the sample?


15.114 What is the expected electron configuration for each of the following ions?

(a) Co2+ (b) Pd2+ (c) Tc4+

15.115 Give the oxidation number of xenon in the following compounds: (a) XeF4 (b) XeF6 (c) XeOF4 (d) XeO2F2

(e) XeO4.

15.116 Draw Lewis structures and predict the geometries for each of the following.

(a) ClF3 (b) IF5 (c) CH2O

15.117 Complete and balance the following equations. (a) Barium metal plus chloric acid

(b) Aqueous strontium hydroxide plus chlorous acid (c) Solid lithium hydride plus water

(d) Solid tetraphosphorus hexaoxide plus water (e) Aqueous rubidium carbonate plus sulfuric acid

15.118 Write the chemical formula for each of the following and indicate the oxidation state of nitrogen in each.

(a) Nitrous acid (b) Potassium cyanide (c) Sodium nitrate (d) Ammonium chloride (e) Lithium nitride

15.119 Cryosurgical procedures lower the body temperature of a patient prior to surgery. Given that the activation

energy for the beating of heart muscle is about 30 kJ/mol, estimate the pulse rate at 72.0°F (22.2°C) if the pulse rate

at 98.6°F (37.0°C) is 75 beats/min. (Assume that 75 beats/min is the value of k at this temperature.)


16.1 Why is chemical equilibrium so very important?

16.2 How does chemical equilibrium apply to the weak electrolytes discussed in Chapter 4?

16.3 How do the rates of the forward and reverse reactions compare when a system is in equilibrium?

16.4 If no concentrations are changing, why is an equilibrium considered to be a dynamic process?


16.6 State the law of mass action.

16.7 What is the general form of the expression that produces a reaction quotient, Q?

16.8 When does the value of the reaction quotient become the same as that of the equilibrium constant?

16.9 What is the difference between K and k?

16.10 In mass action expressions, what symbolism is often useful to indicate molarity?

16.11 Where do the coefficients from a balanced chemical equation appear in a mass action expression?

16.12 (a) How does the value of the equilibrium constant relate to the value of the equilibrium constant for the

reverse reaction? (b) How does the equilibrium constant change when the coefficients in a reaction are doubled?

16.13 What changes are necessary in a mass action expression when partial pressures replace concentrations?

16.14 Given the following equilibrium, how would the expressions for Kc and Kp differ?

COCl2(g) CO(g) + Cl2(g)


16.15 Write a mass action expression describing Kc for each of the following equilibria. Determine which of these

equilibria are homogeneous and which are heterogeneous.

(a) H2(g) + Br2(g) 2 HBr(g) (b) Fe2O3(s) + 3 CO(g) 2 Fe(s) + 3 CO2 (g)

(c) N2(g) + 3 H2(g) 2 NH3(g) (d) CH4(g) + 2 O2(g) CO2(g) + 2 H2O(l)

(e) HNO2(aq) H+(aq) + NO2–(aq) (f) Cu2+(aq) + 4 NH3(aq) Cu(NH3)4

2+(aq)

(g) ZnCO3(s) Zn2+(aq) + CO32–(aq)

16.16 Write mass action expressions describing Kc and Kp for each of the following equilibria. Determine which of

these equilibria are homogeneous and which are heterogeneous.

(a) 2 NO2(g) N2(g) + 2 O2(g) (b) H2SO3(aq) H+(aq) + HSO3–(aq)

(c) C(s) + H2O(g) CO(g) + H2(g) (d) Fe3+(aq) + 3 C2O42–(aq) Fe(C2O4)3

3–(aq)

(e) 2 Zn(s) + O2(g) 2 ZnO(s) (f) Mg3(AsO4)2(s) 3 Mg2+(aq) + 2 AsO43–(aq)

(g) 2 C4H10(g) + 13 O2(g) 8 CO2(g) + 10 H2O(l) (h) MnCO3(s) CO2(g) + MnO(s)

16.17 Write the equilibrium constant expressions (Kc) for each of the following reactions:

(a) Solid calcium carbonate in equilibrium with solid calcium oxide and gaseous carbon dioxide

(b) Solid carbon plus gaseous carbon dioxide in equilibrium with carbon monoxide gas

(c) Sulfur dioxide gas plus oxygen gas in equilibrium with sulfur trioxide gas.

(d) Solid sodium hydrogen carbonate in equilibrium with solid sodium carbonate, gaseous carbon dioxide,

and liquid water

(e) Nitrogen gas plus hydrogen gas in equilibrium with ammonia gas

16.18 Write mass action expressions for each of the following equilibria:

(a) 2 O3(g) 3 O2(g) (b) CaCO3(s) CaO(s) + CO2(g)

(c) PCl5(s) + H2O(l) 2 HCl(g) + POCl3(g) (d) Pb3(PO4)2(s) 3 Pb2+(aq) + 2 PO43–(aq)

(e) MnC2O4(s) + C2O42–(aq) [Mn(C2O4)2]2–(aq)

16.19 Write the equilibrium constant (Kc) for each of the following reactions:

(a) 2 NOBr(g) 2 NO(g) + Br2(g) (b) Fe2O3(s) + 2 CO(g) 2 Fe(s) + 3 CO2(g)

(c) N2O(g) + 4 H2(g) 2 NH3(g) + H2O(g) (d) 2 KNO3(s) 2 KNO2(s) + O2(g)

(e) 2 Pb(NO3)2(s) 2 PbO(s) + 2 NO2(g) + O2(g)

16.3 Corresponds to BLBMWS 15.3 & 15.4 &15.5

16.20 Define a homogeneous equilibrium.

16.21 Write the equation for relating Kp to Kc. Define each term in this equation.

16.22 Determine ∆ng for each of the following equilibria: (a) H2(g) + Br2(g) 2 HBr(g)

(b) Fe2O3(s) + 3 CO(g) 2 Fe(s) + 3 CO2 (g) (c) N2(g) + 3 H2(g) 2 NH3(g)

(d) CH4(g) + 2 O2(g) CO2(g) + 2 H2O(g) (e) 2 NO(g) + Cl2(g) 2 NOCl(g)

16.23 Determine ∆ng for each of the following equilibria: (a) 2 NO2(g) N2(g) + 2 O2(g)

(b) C(s) + H2O(g) CO(g) + H2(g) (c) 2 Zn(s) + O2(g) 2 ZnO(s)

(d) Mg3(AsO4)2(s) 3 Mg2+(aq) + 2 AsO43–(aq) (e) 2 C4H10(g) + 13 O2(g) 8 CO2(g) + 10 H2O(l)

(f) MnCO3(s) CO2(g) + MnO(s)

16.24 The following equilibrium was studied at 973 K: 2 SO3(g) 2 SO2(g) + O2(g). At this temperature, the

value of Kc is 2.4 × 10–3. (a) What is the value of Kc for the following reaction at 973 K? 4SO3(g)4SO2(g)+2O2(g)

(b) What is the value of Kc for the following reaction at 973 K? 2 SO2(g) + O2(g) 2 SO3(g). (c) What is the value

of Kc for the following reaction at 973 K? SO2(g) + 1/2 O2(g) SO3(g).

(a) 5.8 × 10–6 (b) 4.2 × 102 (c) 2.0 × 102

16.25 (a) The reaction 2 NOCl(g) 2 NO(g) + Cl2(g) has Kc = 1.6 × 10–5 at 35°C. Determine the value of Kp at this

temperature. (b) The following 2 Mo(s) + CH4(g) Mo2C(s)+ 2 H2(g) has Kp = 3.55 at 700.°C. Determine the

value of Kc at this temperature.


16.26 The compound nitrosyl bromide, NOBr, decomposes according to the equilibrium 2 NOBr(g)2 NO(g) +

Br2(g). After reaching equilibrium, a 3.50L vessel at 373K was found to contain 2.25g of NOBr, 2.16g of NO, and

2.93g of Br2. (a) Calculate the value of Kc. (b) Calculate the value of Kp by two different methods.

16.27 When heated, hydrogen iodide (HI) decomposes by the following equilibrium: 2 HI(g) H2(g) + I2(g). An

equilibrium mixture of these substances was analyzed and found to contain [HI] = 3.98 × 10–3 M; [H2] = 5.29 × 10–4

M; and [I2] = 5.29 × 10–4 M. Calculate the value of Kc for this mixture.

16.28 At 1000 K, methane (CH4) and water (H2O) react in the presence of a nickel catalyst. After equilibrium was

established, a 5.00 L sample of the gaseous mixture was analyzed and found to contain 0.308 mol of carbon

monoxide (CO), 1.29 mol of hydrogen (H2), 2.69 mol of methane, and 2.69 mol of water vapor. Calculate Kc for the

reaction CH4(g) + H2O(g) CO(g) + 3 H2(g)

16.29 Nitrogen oxide (NO) reacts with hydrogen gas (H2) to establish the following equilibrium:

2 NO(g) + 2 H2(g) N2(g) + 2 H2O(g)

In an experiment, 0.250 mol of NO, 0.100 mol of H2, and 0.200 mol of H2O were placed in a 1.00 L container and

allowed to come to equilibrium. After equilibrium was established, the concentration of N2 in the container was

0.0322 M. (a) Determine the equilibrium concentrations of the other three gases. (b) Calculate the value of Kc.

16.30 At 25°C, the value of Kc for the following equilibrium is 8.6 × 1035. If 0.200 mol of bromine trifluoride, 0.200

mol of fluorine and 0.200 mol of bromine pentafluoride are mixed in a 5.00 L container at 25°C, what will be the

equilibrium concentrations of the three gases? BrF3(g) + F2(g) BrF5(g)

16.31 The equilibrium constant (Kc), for the following reaction, is 5.29 × 10–10 at 1454°C.

2 H2O(g) 2 H2(g) + O2(g)

What are the equilibrium concentrations of these three gases if 36.0 g of H2O, 4.03 g of H2, and 32.0 g O2 are sealed

in a 10.0 L container at 1454°C? 0.399 M H2O 5.53 × 10–4 M H2 2.77 × 10–4 M O2

16.32 The equilibrium constant (Kc), for the following reaction, is 5.29 × 10–10 at 1454°C.

2 H2O(g) 2 H2(g) + O2(g)

What are the equilibrium concentrations of these three gases if 72.0 g of H2O, 8.06 g of H2, and 64.0 g O2 are sealed

in a 20.0 L container at 1454°C?

16.33 Determine Kc for the equilibrium 2 CH4(g) C2H2(g) + 3 H2(g). In order to determine this value, a

mixture initially 0.0300 M in CH4 was allowed to come to equilibrium. At equilibrium, the concentration of C2H2

was 0.01375 M.

16.34 Like all weak acids, chloroacetic acid (CH2ClCOOH) is a weak electrolyte. A 0.100 M solution of

chloroacetic acid is 11.8 percent ionized [percent ionized = (ionized/initial) × 100%]. Using this information,

calculate K for the following equilibrium: CH2ClCOOH(aq) H+(aq) + CH2ClCOO–(aq)

16.35 A sample of 0.0493 mol of phosgene gas (COCl2) was placed in a 2.50 L container at 800.0 K. After

establishing equilibrium, the partial pressure of COCl2 was 0.800 atm. Determine Kp for the reaction

CO(g) + Cl2(g) COCl2(g)

16.36 Calculate Kp for the reaction below, given the following partial pressures: C2H5OH = 0.00250 atm, O2 = 0.200

atm, CO2 = 0.00452 atm, and H2O = 0.00750 atm. C2H5OH(g) + 3 O2(g) 2 CO2(g) + 3 H2O(g)

16.37 Determine Kc for the following equilibrium: 2 CH4(g) C2H2(g) + 3 H2(g)

In order to determine this value, a mixture initially 0.0300 M in CH4 was allowed to come to equilibrium. At

equilibrium, the concentration of C2H2 was 0.01375 M. 0.15

16.38 Calculate Kc for the following reaction: CO(g) + 2 H2(g) CH3OH(g)

The equilibrium amounts of the gases are 0.500 mol of carbon monoxide, 3.00 mol of hydrogen, and 0.563 mol of

CH3OH in a 2.00 L container.


16.39 For the following reaction calculate the value of K if the initial concentration of N2 and H2 were both 1.000 M

and the equilibrium concentration of N2 was 0.922 M. N2(g) + 3 H2(g) 2 NH3(g)

16.40 For the following equilibrium, Kc = 4.6 × 109. If 0.15 mol of carbon monoxide, 0.30 mol of chlorine, and 0.15

mol of COCl2 are mixed in a 1.00 L container and allowed to come to equilibrium, what will be the equilibrium

concentrations? CO(g) + Cl2(g) COCl2(g)

16.41 For the equilibrium H2(g) + I2(g) 2 HI(g)

Kc is 54.4 at 698 K. If 0.100 mol of H2 and 0.150 mol of I2 are placed in a 10.0 L container at 698 K and allowed to

come to equilibrium, what will be the equilibrium concentrations of each of the gases?

16.42 For the following equilibrium, Kc = 8.2 × 10–44 at 25°C. Calculate the equilibrium concentrations of the three

gases if the initial concentrations are [H2O] = 2.00 M, [H2] = 2.00 M, and [O2] = 1.00 M.

2 H2O(g) 2 H2(g) + O2(g)

6.9 × 10–15 M = [O2] 1.4 × 10–14 M = [H2] 4.00 M = [H2O]

16.43 At 1727°C, the equilibrium constant for the following reaction is 5.29 × 10–10.

2 H2O(g) 2 H2(g) + O2(g)

Determine the equilibrium concentration of each of these gases if 0.200 mol of water, 0.400 mol of hydrogen, and

0.200 mol of oxygen are mixed in a 10.00 L container at this temperature.

5.96 × 10–5 M = [O2] 1.19 × 10–4 M = [H2] 0.0399 M = [H2O]

16.44 At a temperature of 700°C, hydrogen sulfide gas dissociates according to the following equilibrium:

2 H2S(g) 2 H2(g) + S2(g)

If the value of Kc for this reaction is 9.1 × 10–8, what will be the concentrations of the three gases if 0.800 mol of

hydrogen sulfide, 0.800 mol of hydrogen, and 0.400 mol of sulfur are placed in a 4.00 L container at 700°C and

allowed to come to equilibrium?

16.45 The value of Kp is 1.32 at 700.0 K for the reaction C(s) + 2 Cl2(g) CCl4(g)

Determine Kc for this reaction at the given temperature.

16.46 Determine Kp for the following equilibrium if Kc is 2.24 × 1022 at 1273°C.

2 CO(g) + O2(g) 2 CO2(g)


16.47 How does a heterogeneous equilibrium differ from a homogeneous equilibrium?

16.48 How does the mass action expression for a heterogeneous equilibrium differ from that of a homogeneous

equilibrium?

16.49 The following reaction produces water gas, an important industrial fuel consisting of carbon monoxide and

hydrogen. The reaction takes place near 800°C, when water vapor (steam) passes over hot carbon (coke, a form of

carbon produced from heating coal). C(s) + H2O(g) CO(g) + H2(g)

Calculate Kc for this reaction if at equilibrium the following concentrations are: [H2] = 3.5 × 10–2 M,

[CO] = 3.5 × 10–2 M, and [H2O] = 7.7 × 10–3 M.

16.50 A sample of ammonium hydrogen sulfide (NH4HS) is sealed in a flask. The sample decomposes according to

the following equilibrium NH4HS(s) NH3(g) + H2S(g)

This equilibrium has Kc = 1.2 × 10–4 at this temperature. Determine the equilibrium concentrations of ammonia and

hydrogen sulfide in the flask. 0.011 M = [NH3] = [H2S]

16.51 Two molecules may join to form an adduct. The compound PH3BCl3 is an adduct formed by the reaction of

PH3(g) with BCl3(g). The adduct PH3BCl3, when heated, decomposes according to the following equilibrium.

PH3BCl3(s) PH3(g) + BCl3(g) Kp = 1.60


Determine the equilibrium partial pressures of phosphine (PH3) and boron trichloride (BCl3) when a sample of

PH3BCl3 is placed in a sealed flask and allowed to come to equilibrium.

16.52 For the following equilibrium, Kc = 1.58 × 10–8 at 523 K: NH2COONH4(s) 2 NH3(g) + CO2(g)

(a) Calculate Kp at 523 K. (b) A sample of solid NH2COONH4 is sealed in a 275 mL container and heated to 523 K.

Calculate the partial pressures of ammonia and carbon dioxide after equilibrium has been established. (c) In a

separate experiment, a sample of solid NH2COONH4 and 0.0100 mol of ammonia are sealed in a 275 mL container

and heated to 523 K. Calculate the partial pressures of ammonia and carbon dioxide after equilibrium has been

established.

16.53 The following equilibrium is established at 500 K: S2(g) + C(s) CS2(g).

The initial pressure of sulfur vapor was 0.431 atm. After equilibrium has been established, the S2 pressure is 0.414

atm. Determine the value of Kp at 500 K.

16.54 For the following reaction, Kp = 3.55 at 700°C. 2 Mo(s) + CH4(g) Mo2C(s) + 2 H2(g)

What is the equilibrium partial pressure of methane (CH4) when the equilibrium partial pressure of hydrogen (H2) is

0.100 atm?

16.55 The addition of sparingly soluble silver(I) carbonate to water at 25°C gives a carbonate ion concentration of

1.3 × 10–4mol/L. Calculate K for the following equilibrium: Ag2CO3(s) 2 Ag+(aq) + CO32–(aq)

16.56 A total of 250.0 g of NH4CO2NH2 was placed in a 0.5000 L container at 40.0°C. After the following

equilibrium was established, the total pressure in the container was 0.363 atm. Calculate the value of Kp.

NH4CO2NH2(s) 2 NH3(g) + CO2(g) 0.00709

16.57 A sample of NO2 having a partial pressure of 0.750 atm at 1000.0 K was sealed in a container and the

equilibrium shown below was established. At equilibrium, the total pressure in the system was 1.098 atm.

Calculate Kp. 2 NO2(g) 2 NO(g) + O2(g)

16.58 The following equilibrium produces water gas (see Problem 16.49):

C(s) + H2O(g) CO(g) + H2(g)

An equilibrium mixture of these substances was collected in a 1.00 L flask and found to contain 5.25 g of carbon,

[H2O] = 3.1 × 10–2 M, [CO] = 7.0 × 10–2 M, and [H2] = 7.0 × 10–2 M. Determine the value of Kc for this mixture.

16.59 For the following equilibrium, Kp = 6.66 × 10–3 at 20.0°C. Calculate Kc at this temperature.

CuSO4•4NH3(s) CuSO4•2NH3(s) + 2 NH3(g) 1.15 × 10–5

16.60 For the following reaction, Kc = 0.238 at 900°C. Calculate Kp. C(s) + CO2(g) 2 CO(g)


16.61 (a) What information about the equilibrium a very large value of K provides? (b) What information about the

equilibrium does a very small value of K provides?

16.62 Carbon monoxide (CO) appears in both of the following equilibria:

CO(g) + Cl2(g) COCl2(g) Kp = 6.0 × 10–3

C(s) + CO2(g) 2 CO(g) K'p = 1.3 × 1014

In which equilibrium is the carbon monoxide concentration expected to be higher. Why?

16.63 (a) Which direction will a reaction shift if K > Q? (b) Which direction will a reaction shift if K < Q? (c) What

changes, if any, will take place when Q = K?


16.65 Decide which direction the following reaction will shift if Kc = 3.59 at 900.0°C, and the initial quantities of

materials given below are mixed.

CH4(g) + 2 H2S(g) CS2(g) + 2 H2(g)

(a) [CH4] = 1.15 M, [H2S] = 1.20 M, [CS2] = 1.51 M, [H2] = 1.08 M

(b) [CH4] = 1.07 M, [H2S] = 1.20 M, [CS2] = 0.90 M, [H2] = 1.78 M

(c) [CH4] = 1.10 M, [H2S] = 1.49 M, [CS2] = 1.10 M, [H2] = 1.68 M

(d) [CH4] = 1.45 M, [H2S] = 1.29 M, [CS2] = 1.25 M, [H2] = 1.75 M

16.66 The following reaction has Kc = 2.19 × 10–10 at 100°C:


Determine which, if any, of the following mixtures are at equilibrium. If the mixture is not at equilibrium, predict

which direction the reaction will shift to achieve equilibrium.

(a) [COCl2] = 5.00 × 10–2 M, [CO] = 3.31 × 10–6 M, [Cl2] = 3.31 × 10–6 M

(b) [COCl2] = 3.50 × 10–3 M, [CO] = 1.11 × 10–5 M, [Cl2] = 3.25 × 10–6 M

(c) [COCl2] = 1.45 M, [CO] = [Cl2] = 1.56 × 10–6 M

(a) No Change (b) Left (c) Right

16.67 For the following equilibrium, Kp is 1.05 at 250.0 °C.

PCl5(g) PCl3(g) + Cl2(g)

The initial pressures for a particular experiment are PCl5 = 0.177 atm, PCl3 = 0.223 atm, and Cl2 = 0.111 atm. When

this mixture comes to equilibrium which of the pressures will increase, decrease, or remain unchanged?


16.68 State Le Châtelier's principle.

16.69 (a) How does a reaction at equilibrium change if the concentration of a reactant increases? (b) How does a

gaseous reaction at equilibrium change if the concentration of a product increases? (c) How does an endothermic

equilibrium change if the temperature increases? (d) How does an exothermic equilibrium change if the temperature

increases? (e) How does a reaction at equilibrium change upon the addition of a catalyst?

16.70 (a) How does a gaseous reaction at equilibrium change if the concentration of a reactant decreases? (b) How

does a gaseous reaction at equilibrium change if the concentration of a product decreases? (c) How does an

endothermic equilibrium change if the temperature decreases? (d) How does an exothermic equilibrium change if

the temperature decreases? (e) How does a reaction at equilibrium change upon the addition of an inert gas?

16.71 (a) How does the value of the homogeneous equilibrium constant change if the concentration of a reactant

increases? (b) How does the value of a homogeneous equilibrium constant change if the concentration of a product

increases? (c) How does the value of a homogeneous equilibrium constant change upon the addition of a catalyst?

16.72 Which way will this equilibrium shift in response to each of the given changes?

B(s) + 6 H2O(g) 2 B2H6(g) + 3 O2(g) ∆H = +1524 kJ

(a) Adding water (b) Adding boron (c) Increasing pressure (d) Increasing temperature (e) Adding helium

16.73 Predict which direction (left, right, unchanged) each of the following equilibrium systems would be shifted by

the stress indicated.

(a) Br2(l) + Cl2(g) 2 BrCl(g) ∆H = 29.37 kJ/mol Increase the temperature

(b) 3 O2(g) 2 O3(g) Increase pressure

(c) BaCl2·2H2O(s) BaCl2(s) + 2 H2O(g) Remove H2O(g)

(d) BaCl2·2H2O(s) BaCl2(s) + 2 H2O(g) Remove BaCl2(s)

(e) CaCO3(s) CaO(s) + CO2(g) Add CO2(g)

(a) Right (b) Right (c) Right (d) No change (e) Left

16.74 Consider the following equilibrium system: C(s) + CO2(g) 2 CO(g) ∆H° = 119.8 kJ

If the reaction is at equilibrium, what would be the effect of: (a) Adding CO(g) (b) Adding C(s) (c) Lowering the

temperature (d) Increasing the pressure on the system by decreasing the volume (e) Adding a catalyst


16.75 In which direction will the equilibrium shift when the following changes are made?

C(s) + 2 H2(g) CH4(g) H = –75 kJ

(a) Increasing the temperature (b) Increasing the volume (c) Adding hydrogen (d) Adding carbon.

16.76 For the reaction C(s) + O2(g) CO2(g), ∆H° = −393.5 kJ. How will the amount of CO2 present at

equilibrium be affected by: (a) Removing O2(g) (b) Adding C(s) (c) Increasing the volume of the container

(d) Adding a catalyst (e) Increasing the temperature.

16.77 The following equilibrium was established in a 1.00 L container at 298 K. How will the equilibrium

concentration of oxygen change after the adjustments listed in each part?

C(s) + O2(g) CO2(g) ∆H = −395.5 kJ

(a) Adding carbon (b) Adding carbon dioxide (c) Removing carbon dioxide (d) Lowering the temperature

(e) Adding a catalyst

16.78 The value of Kp for the following equilibrium is 0.26 at 425°C.

2 NO(g) + Cl2(g) 2 NOCl(g)

Predict the direction the equilibrium will shift in each of the following situations.

Initial Initial Initial

NO Cl2 NOCl

(atm) (atm) (atm)

(a) 0.12 0.30 0.15

(b) 0.30 0.45 0.050

(c) 0.050 0.75 0.022

16.79 For the following equilibrium, Kc = 0.0108 at 1173 K.

CaCO3(s) CaO(s) + CO2(g)

Samples of the compositions described below are placed in a 25.00 L container and heated to 1173 K. Predict how

the amount of calcium oxide (CaO) will change (increase, decrease, remain the same) in each case.


CaCO3 CaO CO2

(a) 20.0 g 20.0 g 5.00 g

(b) 10.0 g 30.0 g 0.123 g

(c) 30.0 g 10.0 g 0.100 g

16.80 At a certain temperature, Kc for the following equilibrium is 2.20 × 10–10.


Predict the direction the equilibrium will shift in each of the following situations.


COCl2 CO Cl2

(mol/L) (mol/L) (mol/L)

(a) 4.50 × 10–3 2.11 × 10–5 4.20 × 10–6

(b) 4.95 × 10–2 3.30 × 10–6 3.30 × 10–6

(c) 2.45 × 10–1 1.00 × 10–6 1.50 × 10–6

16.81 Determine the effect of the stresses listed below on the following equilibrium:

2 SO2(g) + O2(g) 2 SO3(g) ∆H = −198 kJ

(a) doubling the volume of the reaction chamber (b) increasing the temperature (c) adding O2 (d) adding helium

16.82 In each of the following conditions, what will happen to the number of moles of SO3 in equilibrium with SO2

and O2 in the reaction? 2 SO3(g) 2 SO2(g) + O2(g)

The standard heat of reaction is 198 kJ/mol.

(a) Remove oxygen gas. (b) Decrease the pressure by increasing the volume. (c) Add argon gas to increase the

pressure. (d) Lower the temperature. (e) Add a catalyst.


16.83 Indicate how the partial pressure of hydrogen will change if the following stresses are applied to the

equilibrium below: H2(g) + CO2(g) H2O(l) + CO(g) ∆H = +41 kJ

(a) Adding carbon dioxide (b) Adding water (c) Adding a catalyst (d) Increasing the temperature (e) Decreasing

the volume

16.84 Given the following equilibria:

I. H2(g) + I2(g) 2 HI(g) ∆H° = −2.26 kcal/mol

II. 2 Cl2(g) + 2 H2O(g) 4 HCl(g) + O2(g) ∆H° = 113 kJ/mol

III. C(s) + H2O(g) CO(g) + H2(g) ∆H° = 131 kJ/mol

IV. 2 PbS(s) + 3 O2(g) 2 PbO(s) + 2 SO2(g) ∆H° = −197.76 kJ/mol

Tell how each of the changes listed below would affect these equilibria (by a shift right, a shift left, or no change).

(a) Increase the temperature in I. (b) Decrease the pressure in I. (c) Add Cl2(g) to II. (d) Increase the pressure in II.

(e) Increase the temperature in II. (f) Add CO(g) to III. (g) Add C(s) to III. (h) Increase the pressure in IV.

(i) Increase the temperature in IV. (j) Add PbS(s) to IV.


16.85 What is the relationship between the equilibrium constants K1 and K2 for an equilibrium that is the sum of two

separate equilibria?

16.86 Oxalic acid solutions contain the following simultaneously occurring equilibria:

H2C2O4(aq) H+(aq) + HC2O4–(aq) K1

HC2O4–(aq) H+(aq) + C2O4

2–(aq) K2

(a) Write the mass action expression for K1. (b) Write the mass action expression for K2. (c) Add the two equilibria

together to produce an overall equilibrium. (d) Write a mass action expression for the overall equilibrium.

(e) Multiply the K1 expression by the K2 expression and compare the result to your answer for part (d).

16.87 Given the following equilibrium constants (determined at 1000 K)

Fe3O4(s) + 4 CO(g) 3 Fe(s) + 4 CO2(g) K1 = 1.67 × 10–2

C(s) + CO2(g) 2 CO(g) K2 = 2.03 × 10–2

Calculate K3 for

Fe3O4(s) + 2 C(s) 3 Fe(s) + 2 CO2(g) K3 = ?

16.88 Write the equilibrium constant Kp, and calculate the equilibrium constant at 1123 K for

C(s) + CO2(g) + 2 Cl2(g) 2 COCl2(g)

Using the following data:

CO(g) + Cl2(g) COCl2(g) K"p = 6.0 × 10–3

C(s) + CO2(g) 2 CO(g) K'p = 1.3 × 1014

Summary

16.99 Given the following equilibria,

I. H2(g) + I2(g) 2 HI(g) H° = −2.26 kcal/mol

II. 2 Cl2(g) + 2 H2O(g) 4 HCl(g) + O2(g) H° = 113 kJ/mol

III. C(s) + H2O(g) CO(g) + H2(g) H° = 131 kJ/mol

IV. 2 PbS(s) + 3 O2(g) 2 PbO(s) + 2 SO2(g) H° = −197.76 kJ/mol

Tell how each of the following changes would affect these equilibria (by a shift right, a shift left, or no change).

(a) Increase the temperature in I. (b) Decrease the pressure in II. (c) Add Cl2(g) to II. (d) Increase the pressure in

II. (e) Increase the temperature in II. (f) Add CO(g) to III. (g) Add C(s) to III. (h) Increase the pressure in IV.

(i) Increase the temperature in IV. (j) Add PbS(s) to IV.


16.100 What will happen to the number of moles of SO3 in equilibrium with SO2 and O2 in the reaction in each of

the following cases? (The standard heat of reaction is +198 kJ/mol.)

2 SO3(g) 2 SO2(g) + O2(g)

(a) Remove oxygen gas. (b) Decrease the pressure by increasing the volume. (c) Add argon gas to increase the

pressure. (d) Lower the temperature. (e) Add a catalyst.

16.101 A sample containing sulfur trioxide gas was prepared from sulfur dioxide gas and oxygen gas. The

equilibrium mixture of the three gases had a sulfur dioxide concentration of 0.010 M, an oxygen concentration of

0.20 M, and a sulfur trioxide concentration of 0.100 M. Calculate the value of K for this equilibrium.

16.102 At 60.0 °C, the value of Kc for the following equilibrium is 7.0 × 10–5: PH3BCl3(s) PH3(g) + BCl3(g)

Calculate Kp at this temperature.

16.103 The equilibrium constant for the formation of ammonia gas from hydrogen gas and nitrogen gas is Kp = 4.51

× 10–5 at 450°C. Calculate Kc at this temperature.

16.104 Given the following equilibrium constants (determined at 750°C),

SnO2(s) + 2 H2(g) Sn(s) + 2 H2O(g) Kp = 8.12

H2(g) + CO2(g) CO(g) + H2O(g) K’p = 0.771

Determine the value of K”p for SnO2(s) + 2 CO(g) Sn(s) + 2 CO2(g) 13.6

16.105 For the following reaction, Kp is 14.1 at 850°C: C(s) + CO2(g) 2 CO(g)

A 5.00 L reaction vessel (at 850°C) initially contains 25.2 g of carbon, 3.32 atm of carbon monoxide, and 1.02 atm

of carbon dioxide gas. Determine which direction the reaction will go in order to reach equilibrium.

16.106 The following equilibrium has K = 9.0 × 10–8 at 700°C: 2 H2S(g) 2 H2(g) + S2(g)

The initial concentrations of the three gases were 0.300 M H2S, 0.300 M H2, and 0.150 M S2. Determine the

equilibrium concentration of these gases at 700°C.

16.107 Antimony pentachloride decomposes in a gas phase reaction at 448°C as follows:

SbCl5(g) SbCl3(g) + Cl2(g)

(a) An equilibrium mixture in a 5.00 L vessel has 3.84 g of SbCl5, 9.14 g of SbCl3, and 2.84 g of Cl2. Evaluate Kc at

448°C. (b) Draw Lewis structures for the reactant and both products of this reaction.

16.108 For the following reaction, the value of K is 1.6 × 10–5. 2 NOCl(g) 2 NO(g) + Cl2(g)

(a) Calculate the equilibrium concentration of each species if the initial concentrations were 0.00 M NOCl, 2.50 ×

10–8 M NO, and 1.75 × 10–8 M Cl2. (b) Draw Lewis structures for the reactant and both products of this reaction.

16.109 If Kc for the following reaction is 7.7 × 10–17 (at 25°C), what is the value of Kp?

NH4Cl(s) NH3(g) + HCl(g)

16.110 Calculate Kp for the following reaction (Kc = 2.2 × 1059) at 300°C.

2 NO(g) + 2 CO(g) N2(g) + 2 CO2(g)

16.112 Weak electrolytes (Chapter 4) are present in a variety of equilibria. Write appropriate mass action

expressions (describing Kc) for each of the following weak electrolytes.

(a) HNO2(aq) H+(aq) + NO2–(aq) (b) HSCN(aq) H+(aq) + SCN–(aq)

(c) NH3(aq) + H2O(l) NH4+(aq) + OH–(aq) (d) CH3NH2(aq) + H2O(l) CH3NH3

+(aq) + OH–(aq)

(e) NH4+(aq) NH3(aq) + H+(aq)

Chapter 17 Corresponds to BLBMWS Chapter 16 17.1 Corresponds to BLBMWS 16.1 & 16.2 & 16.10

17.1 (a) Give the definitions of an Arrhenius acid and an Arrhenius base. (b) When an Arrhenius acid reacts with an

Arrhenius base, what two products always form?


17.2 (a) Define a strong acid. (b) Give the name and formula for each of the seven strong acids.

17.3 (a) Define a weak acid. (b) Does a Ka apply to a strong acid or a weak acid?

17.4 (a) Define a strong base. (b) Give the name and formula for each of the eight strong bases.

17.5 (a) Define a monoprotic acid. (b) List the strong acids that are monoprotic acids.

17.6 (a) Define a diprotic acid. (b) List the strong acids that are diprotic acids.

17.7 What equilibrium constant is present whenever liquid water is present?

17.8 Using balanced chemical equations show how each of the following serves as an Arrhenius acid or base when

dissolved in water. (a) HCl (b) HNO2 (c) NaOH (d) Ca(OH)2 (e) NH3

17.9 Write balanced molecular and net ionic equations for each of the following. Assume that all reactions go to

completion and are in aqueous solution. (a) Hydrochloric acid plus sodium hydroxide (b) Ammonia plus nitric acid

(c) Oxalic acid plus potassium hydroxide (d) Ammonium chloride plus lithium hydroxide (e) Phosphoric acid plus

strontium hydroxide

17.10 Write balanced molecular and net ionic equations for each of the following. Assume that all of the reactions

go to completion and are in aqueous solution. (a) Perchloric acid plus sodium carbonate (b) Ammonia plus sulfuric

acid (c) Carbonic acid plus calcium hydroxide (d) Potassium hydroxide plus acetic acid (e) Strontium hydroxide

plus nitric acid

17.12 Define each of the following (a) binary acid, (b) ternary acid, and (c) oxyacid.

17.13 (a) What is the general trend in the strength of binary acids for the elements within a column (family) on the

periodic table? (b) What is the general trend in the strength of binary acids within a period (row) on the periodic

table?

17.14 What factor or factors are important in predicting the relative strengths of two oxyacids with similar formulas,

such as H3PO4 and H3AsO4?

17.15 Explain why the strengths of the iodine oxyacids are in the order HIO < HIO2 < HIO3 < HIO4.

17.16 Explain why HSO4– is a weaker acid than H2SO4.

17.17 Explain why an iron(III) chloride (FeCl3) solution is acidic.

17.18 (a) Write the equilibrium reaction to illustrate why ammonia (NH3) is a base in water. (b) Write a similar

equilibrium reaction to illustrate why methylamine (CH3NH2) is a base. (c) Write a similar equilibrium reaction to

illustrate why hydroxylamine (NH2OH) is a base. (d) Predict the relative strengths of these three bases: NH3,

CH3NH2, and NH2OH. (e) Explain your predictions for part (d).

17.19 Choose the member of each pair that is the stronger acid. (a) H3PO3 or H3AsO3; (b) H2SO4 or H2SO3;

(c) H2SO3 or H2CO3 (a) H3PO3 (b) H2SO4 (c) H2SO3

17.20 Identify the strongest acid in each of the following sets.

(a) H2S H2SO3 H2SO4 (b) HBrO4 HBrO HBrO3

(c) HClO HCl NCl3 (d) HN3 HNO3 HNO2

(e)


17.21 Identify the strongest acid in each of the following groups:

(a) HClO3 HClO HClO4 HClO2

(b) H3PO2 H3PO4 H3PO3

(c) H2SO3 H2SO4

(d) HNO3 H2N2O2 HNO2

(e) HAsO42– H2AsO4

– H3AsO4


17.22 (a) Define an acidic oxide (acid anhydride) and give two examples. (b) Give the formula for two compounds

that are not acidic oxides even though their formulas indicate they should be.

17.23 (a) Define a basic oxide (base anhydride) and give two examples. (b) Give two examples of compounds that

are not basic oxides but that have a formula indicating they should be.

17.24 Dinitrogen pentoxide (N2O5) is an acid anhydride. (a) What is the expected oxidation state of the nitrogen in

the acid formed from this compound? (b) What are the name and formula of the acid formed when this compound

reacts with water? (c) Write a balanced chemical equation for the reaction of dinitrogen pentoxide with water.

17.25 Write balanced chemical equations showing the reaction of the following compounds with water to produce

an acid. (a) carbon dioxide (CO2) (b) sulfur trioxide (SO3) (c) dinitrogen trioxide (N2O3) (d) tetraphosphorus

decaoxide (P4O10) (e) chromium(VI) oxide (CrO3)

17.26 Write balanced chemical equations showing the reaction of the following compounds with water to produce a

base. (a) potassium oxide (K2O) (b) sodium oxide (Na2O) (c) calcium oxide (CaO) (d) barium oxide (BaO)

17.27 Complete and balance the following chemical equations. Write NR (no reaction) for any combination that

will not react. (a) Li2O(s) + H2O(l) (b) SrO(s) + H2O(l) (c) SO2(g) + H2O(l) (d) CO(g) + H2O(l)

(e) Cl2O7(s) + H2O(l)

17.28 Complete and balance the following chemical equations. Write NR (no reaction) for any combination that

will not react. (a) Cs2O(s) + H2O(l) (b) BaO(s) + H2O(l) (c) CuO(s) + H2O(l) (d) MgO(s) + HCl(aq)

(e) SO2(g) + HBr(g)


17.29 Define a Brønsted-Lowry acid and give two examples.

17.30 Define a Brønsted-Lowry base and give two examples.

17.31 Nitric acid (HNO3) will react with potassium hydroxide (KOH) to form the salt potassium nitrate (KNO3) and

water. Why is potassium nitrate a salt?

17.32 (a) Define a conjugate acid-base pair. (b) What are the members of the conjugate acid-base pair for

hypochlorous acid (HClO)? (c) What are the members of the conjugate acid-base pair for oxalic acid (H2C2O4)?

17.33 (a) Why can we assume that H+(aq) and H3O+ are analogous? (b) Hydrogen ions are present in what form in

aqueous solution?

17.34 For the following questions, use CA to represent the conjugate acid, and CB to represent the conjugate base.

Remember, a CA always gives a CB and a CB always gives a CA. (a) Write the balanced equilibrium reaction

equation for a weak acid. (b) What variations are allowed in this reaction? (c) Write the balanced equilibrium

reaction equation for a weak base. (d) What variations are allowed in this reaction?


17.35 (a) Write the equilibrium constant, Ka, expression for a weak acid. Use CA to represent the conjugate acid,

and use CB to represent the conjugate base. (b) Other than entering the actual formulas of the CA and CB, what

variations are allowable in this expression?

17.36 (a) Write the equilibrium constant, Kb, expression for a weak base. Use CA to represent the conjugate acid,

and use CB to represent the conjugate base. (b) Other than entering the actual formulas of the CA and CB, what

variations are allowable in this expression? (a) Kb = [𝐇+] [𝐂𝐀]

[𝐂𝐁] (b) No other variations are allowable.

17.37 (a) Define amphoteric. (b) Write balanced chemical equations to demonstrate the amphiprotic nature of

dihydrogen phosphate ion, H2PO4–. (c) Write balanced chemical equations to demonstrate the amphiprotic nature of

aluminum hydroxide, Al(OH)3.

17.38 Beryllium hydroxide, Be(OH)2, zinc hydroxide, Zn(OH)2, and gallium hydroxide, Ga(OH)3, are amphiprotic

like aluminum hydroxide. (a) Write balanced chemical equations illustrating the reaction of each of these

hydroxides with nitric acid, HNO3. (b) Write balanced chemical equations illustrating the reaction of each of these

hydroxides with potassium hydroxide, KOH. All these metals form anions with the general formula M(OH)4n–.

17.39 List the formulas of the conjugate base for each of the following Brønsted-Lowry acids: (a) HC2H3O2; (b)

HClO3; (c) H2CO3; (d) HCO3–; (e) PH4

+

17.40 Label the acid and conjugate base or the base and conjugate acid (there are two pairs in each equation) in each

of the following equilibria:

(a) H2S(aq) + CN–(aq) HCN(aq) + HS–(aq) (b) HClO(aq) + H2O(l) ClO–(aq) + H3O+(aq)

(c) NH4+(aq) + S2–(aq) NH3(aq) + HS–(aq) (d) CO3

2–(aq) + HC2H3O2(aq) HCO3–(aq) + C2H3O2

–(aq)

(e) SO32–(aq) + H2C2O4(aq) HSO3

–(aq) + HC2O4–(aq)

17.41 Predict whether K for each of the following equilibria is large or small:

(a) HCl(aq) + NaC2H3O2(aq) HC2H3O2(aq) + NaCl(aq) (b) CN–(aq) + HNO2(aq) HCN(aq) + NO2–(aq)

(c) CO32–(aq) + H2O(l) HCO3

–(aq) + OH–(aq) (d) NH4+(aq) + OH–(aq) NH3(aq) + H2O(l)

(e) HBrO(aq) + HC2O4–(aq) BrO–(aq) + H2C2O4(aq)

17.42 Predict which member of the following pairs would have a stronger Brønsted-Lowry conjugate acid. Explain

your predictions. (a) HCO3– or Br– (b) ClO3

–or CO32– (c) H2PO4

– or HSO4– (d) HSO3

– or SO32– (e) NH3 or NH2

–

17.43 Predict which of the following acids is the weakest, and which is the strongest: HNO3, HClO2, HIO, and

HC2H3O2.

17.44 List the following bases in order of increasing base strength: NH3, H2O, OH–, NO2–, and Br–.

17.45 (a) Write the balanced equilibrium reaction equation for acetic acid (HC2H3O2). (b) Write the balanced

equilibrium reaction equation for the hydrolysis of the acetate ion (C2H3O2–) in water. (c) Write equilibrium

constant expressions for parts (a) and (b). (d) Add the two reactions from parts (a) and (b), and then write the

equilibrium constant expression for the resultant reaction. (e) If Ka for acetic acid is 2 × 10–5 and Kb for the acetate

ion is 5 × 10–10, what is the value of the equilibrium constant from part (d)?

17.46 If the Ka for an acid is 2.5 × 10–3, what is the Kb for its conjugate base?

17.47 (a) Write the Ka expression for HClO(aq) H+(aq) + ClO–(aq). (b) Write the Kb expression for ClO–(aq) +

H2O(l) OH–(aq) + HClO(aq). (c) Write the equilibrium constant expression resulting when you multiply your

answer from part (a) by your answer from part (b).


17.48 (a) Define a Lewis acid. (b) Define a Lewis base.


17.49 (a) Draw the Lewis structure for ammonia (NH3). (b) Draw the Lewis structure for the cyanide ion (CN–).

(c) Why may both ammonia and the cyanide ion serve as a Lewis base?

17.51 (a) Why can the hydrogen ion be an Arrhenius acid, a Brønsted-Lowry acid, and a Lewis acid? (b) Why can

the hydrogen ion not be a Lewis base?

17.52 (a) Is the zinc ion (Zn2+) more likely to be a Lewis acid or a Lewis base? (b) Why?

(a) It is more likely to be a Lewis acid. (b) The positive charge indicates that it has lost electrons;

therefore, it is more likely to accept electrons (Lewis acid) than to donate electrons (Lewis base).

17.53 (a) Draw the Lewis structure for water (H2O), for the nitrite ion (NO2–), and for nitrous acid (HNO2). (b) Use

your Lewis structures from part (a) to show how water behaves as a Lewis acid and the nitrite ion behaves as a

Lewis base during the formation of nitrous acid.

17.54 In each of the following reactions, which of the reactants is the Lewis acid and which is the Lewis base?

(a) FeCl3(s) + Cl–(aq) FeCl4–(aq) (b) H2O(l) + NO2

–(aq) HNO2(aq) + OH–(aq)

(c) NH3(g) + BF3(g) H3NBF3(s) (d) BrF3(l) + HF(l) H+(sol) + BrF4–(sol) (sol = solvated)

(e) CaO(s) + H2O(l) Ca(OH)2(s)

17.55 In each of the following reactions, which of the reactants is the Lewis acid and which is the Lewis base?

(a) SiF4(g) + 2 F–(aq) SiF62–(aq) (b) H2O(l) + NH2

–(aq) OH–(aq) + NH3(aq)

(c) Cu2+(aq) + 4 NH3(aq) Cu(NH3)42+(aq) (d) OH–(aq) + NH4

+(aq) NH3(g) + H2O(l)

(e) CO2(aq) + H2O(l) H2CO3(aq)


17.56 Define autoionization and write a balanced equilibrium reaction equation to illustrate the autoionization of

water.

17.57 Why is there no term in the denominator of the Kw expression?

17.58 Ammonia, like water, undergoes autoionization. The conjugate acid of ammonia is the ammonium ion

(NH4+), and the conjugate base is the amide ion (NH2

–). (a) Write a balanced equilibrium reaction equation to

illustrate the autoionization of ammonia. (b) Write a mass action expression, which is analogous to Kw, for this

equilibrium.

17.59 The value of Kw is 1.0 × 10–14 near room temperature. However, like all equilibrium constants, this value will

vary with temperature. The value of Kw increases with temperature. (a) How does the hydrogen ion concentration

in a neutral solution change as the temperature increases? (b) How does the pH of a neutral solution change as the

temperature increases?

17.60 Define the leveling effect.

17.61 Liquid ammonia, because it is a polar solvent capable of hydrogen bonding, often serves as an alternative to

water as a solvent. Ammonia, like water, is capable of autoionization to form the ammonium ion (NH4+) and the

amide ion (NH2–). (a) What is the strongest acid that can exist in liquid ammonia? (b) What is the strongest base

that can exist in liquid ammonia?

17.62 Label each of the following solutions as acidic, basic, or neutral (a) [H+] = [OH–]; (b) [H+] = 0.000000075 M;

(c) [OH–] = 0.00000800 M (d) 3.0 × 10–8 M OH–; (e) 7.5 × 10–6 M H+.

(a) Neutral (b) Basic (c) Basic (d) Acidic (e) Acidic

17.63 Determine the hydroxide ion concentration in each of the following solutions: (a) 0.0150 M H+(aq) (b) 1.7 ×

10–10 M H+(aq) (c) 1.0 × 10–14 M H+(aq) (d) 1.0 M H+(aq) (e) a solution with a [OH–] that is 50.0 times that of the

[H+].


17.64 Determine the hydrogen ion concentration in each of the following solutions: (a) 5.5 × 10–12 M OH– (b) 4.25

M OH– (c) 1.0 × 10–7 M OH– (d) 3.25 × 10–4 M OH– (e) a solution with a [H+] that is 25.0 times that of the [OH–].

(a) 1.8 × 10–3 M H+ (b) 2.35 × 10–15 M H+ (c) 1.0 × 10–7 M H+ (d) 3.08 × 10–11 M H+ (e) 5.00 × 10–7 M H+

17.65 Determine the pH of each of the following strong acid solutions. (a) 0.040 M nitric acid (HNO3) (b) A

solution prepared by dissolving 1.32 g of hydrogen iodide (HI) in sufficient water to produce 275 mL of solution.

(c) A solution prepared by dissolving 36.45 g of hydrogen chloride (HCl) in sufficient water to produce 1.000 L of

solution. (d) A solution prepared by diluting 15.0 mL of 2.50 M hydrobromic acid (HBr) to 500.0 mL. (e) A

solution prepared by mixing 175 mL of 0.100 M chloric acid (HClO3) with 225 mL of 0.30 M perchloric acid

(HClO4).

17.66 Determine the hydroxide ion concentration and the pH of the following strong base solutions. (a) 0.150 M

cesium hydroxide (CsOH) (b) A solution prepared by dissolving 40.0 g of sodium hydroxide (NaOH) in sufficient

water to produce 1000.0 mL of solution. (c) A solution prepared by diluting 25.0 mL of 0.200 M barium hydroxide

(Ba(OH)2), to 750.0 mL. (d) A solution prepare by mixing 25.0 mL of 0.100 M lithium hydroxide (LiOH) with 75.0

mL of 0.250 M rubidium hydroxide (RbOH). (e) A solution prepared by mixing 175 mL of 0.500 M potassium

hydroxide (KOH) with 225 mL of 0.500 M strontium hydroxide (Sr(OH)2).

17.67 Determine the hydroxide ion concentration and the pH of the following strong base solutions. (a) 2.75 × 10–3

M barium hydroxide (Ba(OH)2). (b) A solution made to contain 112 g of potassium hydroxide (KOH) in 1000.0 mL

of solution. (c) A solution made by diluting 175 mL of 0.00175 M lithium hydroxide (LiOH) to 3.00 L. (d) A

solution prepared by mixing 25.00 mL of 0.0105 M cesium hydroxide (CsOH), with 175.00 mL of 2.50 × 10–2 M

barium hydroxide, (Ba(OH)2).

17.68 Determine the pH for each of the following solutions. (a) 10–4 M H+ (b) 0.147 M H+ (c) pOH = 4.75

(d) 3.8 × 10–3 M OH– (e) 2.0 M OH–

17.69 Determine the hydrogen ion concentration in each of the following solutions. (a) pH = 3.78 (b) pH = 10.32

(c) pH = −1.20 (d) pH = 6.8 (e) pH = 0.00

17.70 Determine both the hydrogen ion and the hydroxide concentrations in each of the following solutions. (a) pH

= 4.21 (b) pH = 12.00 (c) pOH = 2.88 (d) pOH = 7.5 (e) pOH = 14.0


17.71 (a) Define percent ionization. (b) What do the percent ionizations of all strong acids have in common?

(c) What do the percent ionizations of all weak acids have in common?

17.72 (a) Write a generic equilibrium chemical reaction for the ionization of a weak acid. You may use CA to

represent the conjugate acid, and CB to represent the conjugate base. (b) Other then substituting the actual formulas

of the CA and CB into the reaction, what other changes are possible in this equation?

17.73 (a) Define a polyprotic acid and give two examples. (b) Can a polyprotic acid be a strong acid? If so, give an

example.

17.74 (a) Write a generic mass action expression, Ka, for a weak acid. You may use CA to represent the acid, and

CB to represent the conjugate base. (b) Other then substituting the actual formulas of the CA and CB into the

reaction, what other changes are possible in this equation?

17.75 Oxalic acid (H2C2O4) is a diprotic acid. As a diprotic acid, it has two Ka values: Ka1 and Ka2. (a) Write an

equilibrium chemical reaction for Ka1 and a separate reaction for Ka2. (b) Write the mass action expressions for Ka1

and Ka2. (c) Add your two answers for part (a) to give one overall equation, and show how to determine Kcombined for

the new equilibrium.

17.76 Arsenic acid (H3AsO4) is a triprotic acid. Like all triprotic acids, arsenic acid has three Ka's (Ka1, Ka2, and

Ka3). How do the three Ka values for arsenic acid relate to each other?


17.77 A 0.100 M solution of hydrazoic acid (HN3) has a pH of 2.86; calculate the Ka of hydrazoic acid.

Ka = 1.9 × 10–5

17.78 Calculate the pH for each of the following solutions (see the Appendix for the Ka or Kb values). (a) 0.145 M

cyanic acid (HOCN) (b) 0.018 M phenol (C6H5OH) (c) 0.180 M methylamine

17.79 Calculate the percent ionization of salicylic acid, C6H4(OH)COOH, in each of the following cases. The pKa of

salicylic acid is 2.97. (a) 0.500 M; (b) 0.0500 M; (c) 0.00500 M

17.80 A 0.200 M solution of 3-nitrobenzoic acid (NO2C6H4COOH) is 4.18 percent ionized. Calculate the Ka for this

acid, and find the following concentrations in the 0.200 M solution: [H+], [OH–], [NO2C6H4COOH], and

[NO2C6H4COO–].

17.81 Calculate the concentrations of H3PO4, H+(aq), H2PO4–, HPO4

2–, and PO43– at equilibrium in a 0.200 M

solution of phosphoric acid (H3PO4).

0.164 M H3PO4 0.0355 M H+ 0.0355 M H2PO4– 6.2 × 10–8 M HPO4

2– 7.3 × 10–19 M PO43–

17.82 Citric acid (H3C6H5O7) is present in citrus fruits. Citric acid is a triprotic acid with the following Ka values:

Ka1 = 7.4 × 10–4, Ka2 = 1.8 × 10–5, and Ka3 = 4.0 × 10–7. What is the pH of a 0.100 M solution of citric acid? What

approximations or assumptions, if any, can you make in your calculations?


17.83 What do weak bases and weak acids have in common?

17.84 Ammonia is a weak base and exists in aqueous solution as part of the following equilibrium:

NH3(aq) + H2O(l) OH–(aq) + NH4+(aq)

(a) Use Lewis structures to illustrate this reaction. (b) Why does ammonia qualify as an Arrhenius base? (c) Why

does ammonia qualify as a Brønsted-Lowry base? (d) Why does ammonia qualify as a Lewis base?

17.85 Methylamine is a weak base. In aqueous solution, methylamine is part of the following equilibrium:

CH3NH2(aq) + H2O(l) OH–(aq) + CH3NH3+(aq)

(a) Use Lewis structures to illustrate this reaction. (b) Why does methylamine qualify as an Arrhenius base?

(c) Why does methylamine qualify as a Brønsted-Lowry base? (d) Why does methylamine qualify as a Lewis base?

17.86 (a) Write a generic equilibrium chemical reaction for the ionization of a weak base. You may use CA to

represent the conjugate acid, and CB to represent the conjugate base. (b) Other then substituting the actual formulas

of the CA and CB into the reaction, what other changes are possible in this equation?

17.87 (a) Write a generic mass action expression defining Ka for a weak acid. You may use CA to represent the

conjugate acid, and CB to represent the conjugate base. (b) Other then substituting the actual formulas of the CA

and CB into the reaction, what other changes are possible in this equation?

17.88 Codeine (C18H21NO3), an alkaloid narcotic, is a weak base. The pH of a 1.0 × 10–2 M solution of codeine is

9.98. Determine the Kb of codeine.

17.89 (a) The pH of a 0.100 M ethylamine (C2H5NH2) solution is 11.80. Determine the Kb of ethylamine. (b) What

is the pH of a 0.10 M ethylammonium chloride (C2H5NH3+Cl–) solution? (a) 4.2 × 10–4 (b) 5.81

17.90 (a) Pyridine (C5H5N) is a weak organic base. A 5.0 × 10–3 M solution of pyridine has a pH of 8.44; calculate

the value of Kb for this substance. (b) What is the pH of a 0.15 M solution of pyridinium chloride (C5H5NH+Cl–)

solution?



17.91 When an acid reacts with a base, the two substances "neutralize" each other. Why is the solution resulting

from the complete neutralization of a base by an acid not necessarily neutral?

17.92 Identify the "parent" acid and base for each of these salts. (a) potassium nitrate (KNO3) (b) calcium chloride

(CaCl2) (c) ammonium bromide (NH4Br) (d) sodium bisulfate (NaHSO4) (e) sodium sulfate (Na2SO4)

17.93 Identify the ions that will undergo hydrolysis. Cl–, BrO2–, SO3

2–, NO3–, SCN–, NH4

+, Rb+, Fe3+, Ca2+, Cs+

17.94 (a) Write the chemical equation showing the Ka equilibrium for hydrofluoric acid (HF). (b) Write the

chemical equation showing the Kb equilibrium for the hydrolysis of the fluoride ion (F–). (c) What equation results

when your answers to parts (a) and (b) are summed? (d) Write equilibrium constant expressions for parts (a) and

(b). (e) What is the product of your two answers for part (d)?

17.95 Determine the Kb values for each of the following bases using the Ka values in the Appendix. (a) the chlorite

ion (ClO2–); (b) the hypoiodite ion (IO–); (c) the hydrogen carbonate ion (HCO3

–); (d) the sulfide ion (S2–).

(a) 9.1 × 10–13 (b) 4.3 × 10–4 (c) 2.1 × 10–4 (d) 7.7 × 10–2

17.96 Determine the pOH, pH, [H+], and [OH–] in each of the following. (a) A 0.10 M potassium hypochlorite

(KClO) solution (b) A 0.20 M sodium nitrite (NaNO2) solution (c) A 0.15 M rubidium carbonate (Rb2CO3) solution

(d) A solution prepared by mixing 275 mL of 0.15 M potassium cyanate (KOCN) with 225 mL of 0.15 M barium

cyanate (Ba(OCN)2). (e) A solution made by mixing 225.0 mL of 0.1000 M acetic acid (HC2H3O2) with 275.0 mL

of 0.1500 M sodium acetate (NaC2H3O2).

(a) 1.4 × 10–4 M OH– pOH = 3.85 pH = 10.15 [H+] = 7.1 × 10–11 M H+

(b) 2.0 × 10–6 M OH– pOH = 5.70 pH = 8.30 5.0 × 10–9 M H+

(c) 5.6 × 10–3 M OH– pOH = 2.25 pH = 11.75 1.8 × 10–12 M H+

(d) 2.5 × 10–6 M OH– pOH = 5.60 pH = 8.40 4.0 × 10–9 M H+

(e) 1.05 × 10–9 M OH– pOH = 8.977 pH = 5.023 9.491 × 10–6 M H+

17.97 Predict whether an acidic, basic, or neutral solution will result when dissolving each of the following in water.

(a) potassium cyanide (KCN) (b) sodium hydrogen arsenate (Na2HAsO4) (c) methylammonium chloride

(CH3NH3Cl) (d) potassium carbonate (K2CO3) (e) ammonium nitrate (NH4NO3)

17.98 Benzoic acid (HC7H5O2) has Ka = 6.46 × 10–5. One of its salts (sodium benzoate) is added to food as a

preservative. Calculate the pH of a solution made by dissolving 5.15 g of sodium benzoate in sufficient water to

produce 0.750 L of solution.

Summary

17.110 Compare the following thermochemical equations:

NaOH(aq) + HNO3(aq) NaNO3(aq) + H2O(l) ∆H = −55.8 kJ

KOH(aq) + HCl(aq) KCl(aq) + H2O(l) ∆H = −55.8 kJ

(a) Estimate the heat of reaction for one of the following reactions:

CsOH(aq) + HClO4(aq) CsClO4(aq) + H2O(l)

RbOH(aq) + HC2H3O2(aq) CsC2H3O2(aq) + H2O(l)

(b) Why do the first two reactions presented in this exercise have the same enthalpy change?

17.111 A 0.1321 M solution of barium hydroxide was titrated with an arsenic acid solution. A total of 43.95 mL of

acid was required to titrate (see Chapter 4) 50.00 mL of the base. What was the concentration of the acid?

17.112 Sulfuric acid (H2SO4), like water, undergoes autoionization. (a) Write a balanced equilibrium reaction

equation to illustrate the autoionization of sulfuric acid. (b) Write a mass action expression, which is analogous to

Kw, for this equilibrium.


17.113 Cyanic acid (HOCN) and fulminic acid (HCNO) are weak acids. (a) Draw the Lewis structure for each of

the acids. (b) Determine the formal charges for each atom in each structure. (c) Based on the formal charges, which

is the stronger acid?

17.114 What is the pH of a 1.00 × 10–8 M solution of HClO3? pH = 6.9784

17.115 The amide ion (NH2–) is a stronger base than the hydroxide ion (OH–). It is such a strong base that it cannot

exist in aqueous solution. Write a balanced chemical equation illustrating the probable fate of the amide ion in

water.

17.116 (a) What volume (in mL) of 0.1650 M HCl is required to titrate 150.0 mL of a solution that contains 5.65 g

of NaOH per liter? (b) Calculate the number of milliliters of 0.1000 M sulfuric acid required to titrate 25.00 mL of

0.2000 M sodium hydroxide to the endpoint. (a) 128 mL (b) 25.00 mL

17.117 Ammonia (NH3), methylamine (CH3NH2), dimethylamine ((CH3)2NH), and trimethylamine ((CH3)3N) are all

weak bases. (a) Draw Lewis structures for each of these compounds. (b) What do the Lewis structures of each of

these compounds have in common? (c) Which of these compounds can serve as a Lewis base? (d) Which of these

compounds can form hydrogen bonds? (e) Use Lewis structures to illustrate how each of these compounds may

serve as a Brønsted-Lowry base.

17.118 Ammonia gas is very soluble in water. As much as 700 mL of the gas will dissolve in 1.0 mL of water.

Determine the pH of a solution obtained by dissolving 1.00 L of ammonia gas, at 15 °C and 755 torr, in 275 mL of

water.

17.119 Determine the pH of a solution formed by dissolving 775 mL of hydrogen chloride gas, at 27 °C and 785

mmHg, in 225 mL of water.


18.1 Based upon information from Chapter 4, define a weak acid and a weak base.

18.2 (a) Write the mass action expression for a Ka. (b) Write the mass action expression for a Kb.

18.3 (a) What is the definition of pH? (b) What is the definition of pOH? (c) What is the relationship between the

pH and the pOH of a solution?

18.4 How does the Ka of a weak acid relate to the Kb of its conjugate base?

18.5 (a) What is the definition of pKa? (b) What is the definition of pKb? (c) How does the pKa of a weak acid relate

to the pKb of its conjugate base?

18.6 (a) What is the common ion effect? (b) Define a common ion.

18.7 What chemical principle does the common ion effect illustrate?

18.8 Given the following equilibrium, what are two ways to increase the acetate ion concentration?

HC2H3O2(aq) H+(aq) + C2H3O2–(aq)


18.9 Give the formula of the missing member of each of the following conjugate acid-base pairs.

Acid Conjugate base

H2O _____

_____ H2O

NH3 _____

_____ NH3

HSO4– _____

18.10 Calculate the pH of each of the following solutions using the Ka or Kb values from the Appendix. (a) 0.15 M

in cyanic acid (HOCN) and 0.25 M in sodium cyanate (NaOCN); (b) 0.25 M in methylamine (CH3NH2) and 0.15 M

in methylammonium chloride (CH3NH3Cl).

18.11 Determine the pH of each of the following solutions. (a) 0.100 M in sodium bicarbonate (NaHCO3) and 0.175

M in sodium carbonate (Na2CO3); (b) 0.085 M in pyridine (C5H5N) and 0.075 M in pyridinium bromide

(C5H5NHBr).

18.12 What is the pH of a solution made by adding 0.040 mol of KOH to 825 mL of 0.150 M HC2H3O2?

18.13 Calculate the pH of a solution that is made by adding 185.0 g of barium acetate, Ba(C2H3O2)2, to 1.000 L of a

0.3500 M hydrochloric acid solution. pH = 5.256

18.14 Calculate the pH of a solution that initially contained 1.22 g/L benzoic acid (C6H5COOH) and 2.88 g/L of

sodium benzoate (NaC6H5COO).

18.15 Calculate the pH of a solution containing 6.50 g of ammonium sulfate, (NH4)2SO4, in 100.0 mL of 1.000 M

ammonia, NH3.

18.16 The acetate ion undergoes hydrolysis in aqueous solution. The equilibrium expression for the reaction is

C2H3O2–(aq) + H2O(l) OH–(aq) + HC2H3O2(aq)

Predict how the equilibrium will change (left, right, or no change), with the application of each of the following

stresses. (a) the addition of sodium acetate (NaC2H3O2) (b) the addition of acetic acid (HC2H3O2) (c) the addition of

sodium hydroxide (NaOH) (d) the addition of water (e) the addition of hydrochloric acid (HCl)


18.17 Define a buffer solution.

18.18 (a) Write the Henderson-Hasselbalch equation in the form used to find the pH directly. (b) Write the

Henderson-Hasselbalch equation in the form used to find the pOH directly. (c) What relationship(s) allow indirect

determination of pH from the pOH?

18.19 What does the term buffer capacity mean?

18.20 Identify which of the following combinations will produce a buffer solution. (a) acetic acid (HC2H3O2) with

sodium acetate (NaC2H3O2) (b) ammonium chloride (NH4Cl) with ammonia (NH3) (c) hydrochloric acid (HCl) with

sodium chloride (NaCl) (d) benzoic acid (C6H5COOH) with sodium benzoate (NaC6H5COO) (e) nitrous acid

(HNO2) with sodium nitrate (NaNO3)

18.21 A buffer contains 1.0 mol of formic acid (HCHO2) and 0.75 mol of sodium formate (NaCHO2). (a) What is

the buffer capacity with respect to the number of moles of sodium hydroxide (NaOH) that could be added?

(b) What is the buffer capacity with respect to the number of moles of hydrochloric acid (HCl) that could be added?

(c) What is the buffer capacity with respect to the number of moles of calcium hydroxide, Ca(OH)2, that could be

added? (a) 1.0 mol (b) 0.75 mol (c) 0.50 mol


18.22 Give directions on how to prepare a buffer solution, with the pH equal to the pKa, in each of the following

cases. If any reactions occur, give the appropriate chemical equations. (a) You have an acetic acid (HC2H3O2)

solution and solid sodium acetate (NaC2H3O2). (b) You have an acetic acid (HC2H3O2) solution and solid sodium

hydroxide (NaOH).

18.23 (a) What is the pH of a solution prepared by adding 125.5 g of acetic acid (HC2H3O2) and 125.0 g of sodium

acetate (NaC2H3O2) to enough water to form 3.50 L of solution? (b) Calculate the pH of a buffer containing 1.00

mol of calcium acetate (Ca(C2H3O2)2) and 1.00 mol of acetic acid in 2.00 L of solution. (c) Calculate the pH of a

buffer containing 1.00 mol of ammonium sulfate ((NH4)2SO4) and 2.00 mol of ammonia (NH3) in 2.00 L of solution.

18.24 (a) Calculate the pH of a solution made by adding 135.20 g of sodium acetate (NaC2H3O2) to 500.0 mL of

1.000 M hydrochloric acid (HCl). (b) Calculate the pH of a solution made by adding 130.2 g of calcium acetate

(Ca(C2H3O2)2) to 500.0 mL of 1.000 M hydrochloric acid.

18.25 A 0.146 g sample of an unknown monoprotic acid (molar mass = 83.0 g/mol) was dissolved in 50.0 mL of

water and titrated with a 0.101 M NaOH solution. After the addition of 14.0 mL of base, the pH was 5.12. What

was the Ka for the unknown acid? Ka = 3.1 × 10–5

18.26 A mixture of 0.250 mol of a weak base (XOH) and 0.075 mol of hydrochloric acid (HCl) was diluted to 750.0

mL. After dilution, the pH of the solution was 10.15. Determine the Kb of the base.

18.27 A solution is prepared by dissolving 35.0 g butanoic acid (HC4H7O2) and 40.0 g of potassium butanoate

(KC4H7O2) in a little water and then diluting to 1.00 L. The Ka for butanoic acid is 1.5 × 10–5. (a) Calculate the pH

of this buffer solution. (b) Write the molecular, complete ionic, and net ionic equations illustrating the reaction that

occurs when a small quantity of nitric acid is added to this buffer. (c) Write the molecular, complete ionic, and net

ionic equations illustrating the reaction that occurs when a small quantity of potassium hydroxide is added to this

buffer. (d) What is the maximum number of moles of nitric acid that may be added to this buffer before it ceases to

be a buffer? (e) What is the maximum number of moles of potassium hydroxide that may be added to this buffer

before it ceases to be a buffer?

18.28 A buffer solution is prepared by dissolving 35.0 g ammonium nitrate (NH4NO3) in 100.0 mL of 2.75 M

ammonia (NH3) and then diluting the mixture to 1.000 L. (a) Calculate the pH of this buffer solution. (b) Write the

molecular, complete ionic, and net ionic equations illustrating the reaction that occurs upon the addition of a small

quantity of nitric acid to this buffer. (c) Write the molecular, complete ionic, and net ionic equations illustrating the

reaction that occurs upon the addition of a small quantity of sodium hydroxide to this buffer. (d) What is the

maximum number of moles of nitric acid that may be added to this buffer before it ceases to be a buffer? (e) What

is the maximum number of moles of sodium hydroxide that may be added to this buffer before it ceases to be a

buffer?

18.29 A sample of 500.0 mL of a buffer solution is prepared that contains 0.150 mol of acetic acid (HC2H3O2) and

0.150 mol of potassium acetate (KC2H3O2). (a) Determine the pH of this buffer solution. (b) Determine the pH of

the buffer after the addition of 0.020 mol of solid potassium hydroxide. Assume there is no change in volume.

(c) A second sample of the same buffer solution is mixed with 50.0 mL of 0.500 M hydrochloric acid. Determine

the pH of the final solution. Assume the volumes are additive. (a) pH = 4.759 (b) pH = 4.876 (c) pH = 4.613

18.30 A 50.00 mL sample of 0.450 M NaOH solution is added to 1.00 L of a buffer solution containing 0.115 mol of

propanoic acid (HC3H5O2) and 0.125 mol of sodium propanoate (NaC3H5O2). What is the pH after the addition?


18.31 What is an acid-base indicator?

The following two questions refer to BLBMWS Figure 16.7

18.32 Which acid-base indicator in Figure 18.XXX would be the best choice in each of the following cases? (a) The

pH is 7. (b) The pH is 5.5. (c) The pH is 10.5.


18.33 (a) Which of the indicators in Figure 18.XXX would probably not be useful in the titration of a strong base

with a strong acid? (b) Which of the indicators in Figure 18.XXX would probably not be useful in the titration of a

weak base with a strong acid? (c) Which of the indicators in Figure 18.XXX would probably not be useful in the

titration of a strong base with a weak acid?


18.34 What are the four types of calculations that may be necessary during a titration?

18.35 (a) What are the four possible types of acid-base titrations? (b) Which of the four types of acid-base titrations

is usually not practical?

18.36 Consider the following possible titrations: (1) nitric acid (HNO3) with ammonia, (NH3) (2) nitric acid with

sodium hydroxide (NaOH) (3) nitrous acid (HNO2) with sodium hydroxide (4) phosphoric acid (H3PO4) with

sodium hydroxide. (a) Which of these four titrations, if any, will have pH = 7 at the equivalence point? (b) Which

of these four titrations, if any, will have pH > 7 at the equivalence point? (c) Which of these four titrations, if any,

will have pH < 7 at the equivalence point? (a) (2) (b) (3) and (4) (c) (1)

18.37 (a) Sketch a rough titration curve for the addition of a strong acid to a strong base. (b) Indicate the point on

your titration curve corresponding to the equivalence point. (c) Indicate the point on your titration curve

corresponding to pH = 7. (d) Indicate any portion(s) of your titration curve where the solution behaves as a buffer.

18.38 (a) Sketch a rough titration curve for the addition of a strong base to a weak acid. (b) Indicate the point on

your titration curve corresponding to the equivalence point. (c) Indicate the point on your titration curve

corresponding to pH = 7. (d) Indicate any portion(s) of your titration curve where the solution behaves as a buffer.

18.39 Calculate the volume, in milliliters, of 0.450 M KOH required in the titration of each of the following to the

equivalence point corresponding to the removal of all acidic hydrogen atoms. (a) 45.0 mL of 0.350 M HClO3;

(b) 25.0 mL of 0.186 M HF; (c) a 100.0 mL sample of a solution containing 3.45 g of HNO3 per liter of solution;

(d) 50.0 mL of 0.225 M sulfuric acid; (e) a solution containing 0.257 g of phosphoric acid in 65.0 mL of solution.

(a) 35.0 mL (b) 10.3 mL (c) 12.2 mL (d) 50.0 mL (e) 17.5 mL

18.40 Calculate the volume, in milliliters, of 0.450 M HNO3 required in the titration of each of the following to the

equivalence point. (a) 25.0 mL of 0.175 M KOH; (b) 40.0 mL of 0.378 M LiOH; (c) 20.0 mL of 0.425 M Ba(OH)2;

(d) a solution containing 0.328 mol of sodium hydroxide in a total volume of 100.0 mL; (e) a solution containing

0.135 g of strontium hydroxide in 50.00 mL

18.41 A scientist wishes to determine the concentration of an acetic acid solution. She removes three 25.00 mL

portions of liquid from the solution and places each in a flask containing phenolphthalein as an indicator. Each of

these samples is titrated with 0.09985 M sodium hydroxide. The volumes of sodium hydroxide solution required for

each titration are, respectively, 45.02 mL, 44.85 mL, and 45.16 mL. (a) Calculate the concentration of acetic acid

found in each of the three titrations. (b) What is the average concentration of the acetic acid solution?

18.42 A chemist in a company producing barium compounds wishes to analyze some samples of barium hydroxide

solution by titrating them with 0.03800 M hydrochloric acid. The volume of each of the barium hydroxide samples

is 50.00 mL. One sample requires 45.02 mL of hydrochloric acid, and a second sample requires 25.48 mL of

hydrochloric acid. (a) Calculate the concentration of barium hydroxide in each of the samples. (b) Calculate the

original pH and pOH of each of the barium hydroxide samples.

(a) 0.01711 M and 0.009682 M (b) pH = 12.534, pOH = 1.446; pH = 12.287, pOH = 1.713

18.43 The pH of a 50.00 mL sample of 0.1100 M ammonia (NH3) is measured as it is being titrated with 0.1100 M

hydrochloric acid (HCl). Determine the pH of the solution after the addition of the following total volumes of

hydrochloric acid. (a) 0.00 mL (b) 25.00 mL (c) 49.00 mL (d) 50.00 mL (e) 51.00 mL (f) 75.00 mL.


18.44 The pH of a 50.00 mL sample of 0.06000 M strontium hydroxide is measured as it is titrated with a 0.1200 M

hydrochloric acid. Determine the pH of the solution after the addition of the following total volumes of

hydrochloric acid. (a) 0.00 mL (b) 25.00 mL (c) 49.00 mL (d) 50.00 mL (e) 51.00 mL (f) 75.00 mL.

18.45 Calculate the pH at the equivalence point for titrating 0.200 M solutions of each of the following bases with

0.200 M HBr: (a) sodium hydroxide (NaOH); (b) hydroxylamine (NH2OH); (c) aniline (C6H5NH2).

18.46 Determine the pH at the equivalence point for each of the following titrations. (a) 50.00 mL of 0.1000 M

hydrochloric acid with 0.09985 M sodium hydroxide (b) 50.0 mL of 0.1000 M benzoic acid with 0.09825 M

potassium hydroxide (c) 50.00 mL of 0.1000 M hydrochloric acid with 0.09755 M ammonia (d) 50.00 mL of 0.1000

M sodium hydrogen carbonate with 0.09550 M sodium hydroxide (e) 50.00 mL of 0.1000 M acetic acid with

0.05525 M strontium hydroxide (a) pH = 7 (b) pH = 8.44 (c) pH = 5.276 (d) pH = 11.51 (e) pH = 8.740

18.47 A 1.2450 g sample of an unknown weak monoprotic acid is dissolved in water. This solution is titrated with

0.2000 M sodium hydroxide. The titration requires 48.96 mL to reach the endpoint. After 24.48 mL of base have

been added, the pH of the solution is 3.75. (a) Determine the molar mass of the unknown acid. (b) Calculate the Ka

of the unknown acid. (c) What indicator should be used to determine the endpoint? (a) 127.1 g/mol (b) 1.8 × 10–4

18.48 A mixture containing 225 g of nitrogen gas (N2) and 40.0 g of hydrogen gas (H2) is heated with a catalyst in a

100.0 L container. After the reaction, the contents are cooled to room temperature. The following reaction occurs in

the container:

N2(g) + 3 H2(g) 2 NH3(g)

(a) Determine the theoretical yield for this experiment. (b) A 1.00 L sample of the contents is bubbled through

500.0 mL of a 0.1000 M hydrochloric acid solution. Titration of the unreacted acid requires 43.25 mL of 0.2302 M

sodium hydroxide (NaOH) to reach the endpoint. What is the percent yield for this experiment?


18.54 One of the important buffers in blood is the HCO3–/H2CO3 system. (a) Determine the HCO3

–-to-H2CO3 ratio

in a blood sample with a pH of 7.4. (b) Determine the HCO3–-to-H2CO3 ratio in a blood sample from a patient

suffering from acidosis. This sample has a pH of 7.1. (c) Determine the HCO3–-to-H2CO3 ratio in a blood sample

from a patient suffering from alkalosis. This sample has a pH of 7.6.

18.62 (a) What will be the pH at the equivalence point in the titration of 50.00 mL of a solution that has a codeine

monohydrate concentration of 0.01000 M with a 0.005000 M hydrochloric acid solution? (b) What would be a good

indicator for this titration?

18.63 The typical dose of morphine sulfate, (C17H20NO3)2SO4, is the intravenous injection of between 8 and 15 mg.

If a patient is to receive 12 mg of morphine sulfate via a 5.0 mL injection, what is the pH of the solution being

injected? pH = 5.01

18.64 Codeine (C18H21NO3) is seldom administered directly. It is most commonly given as codeine sulfate

monohydrate. A typical oral dose of codeine sulfate monohydrate is between 15 and 60 mg. (a) Write the formula

of codeine sulfate monohydrate. (b) What is the pKa of codeine sulfate monohydrate? (c) One gram of codeine

sulfate monohydrate will dissolve in 30 mL of water. What is the molarity of this solution? (d) What is the pH of

the solution in part (c)?

Summary

18.65 A saturated solution of morphine sulfate monohydrate contains 1 g of this compound in 15.5 mL of water.

What is the pH of this solution? Assume the volume of the solution is the same as the volume of the water present.


18.66 The generic name of the analgesic narcotic Demerol is meperidine hydrochloride. Demerol is 63.48% C,

7.81% H, 11.28% O, 12.49% Cl, and 4.94% N. The typical dosage of Demerol is between 25 and 150 mg. (a) What

is the empirical formula of Demerol? (b) What is the molarity of a solution containing 75 mg of Demerol in 25 mL

of solution? Assume the answer from part (a) is the molecular formula. (a) C15H22O2ClN (b) 0.011 M

18.67 (a) Sketch the structures of morphine, codeine, heroin, cocaine, caffeine, and nicotine given in the chapter.

(b) Indicate which, if any, of the carbon atoms in each of these compounds is not sp3 hybridized. (c) The presence

of sp3-hybridized nitrogen leads to the basicity of most alkaloids. Indicate the sp3-hybridized nitrogen atoms, if any,

in each of your structures.

18.68 A solution was prepared containing 200.0 g of calcium acetate and 120.0 g of acetic acid in a total volume of

500.0 mL. Determine the pH of this solution after the addition of 10.0 mL of 6.0 M hydrochloric acid.

18.69 An ammonia-ammonium sulfate buffer has a pH of 9.90 and is 0.05000 M in ammonium sulfate. Calculate

the pH after 25.00 mL of 0.1000 M sulfuric acid has been added to 250.00 mL of this buffer.

18.70 A buffer solution with a pH of 4.62 was made by mixing 0.30 mol of an unknown acid (HA) with 0.060 mol

of calcium hydroxide and diluting to a final volume of 1.00 L. Determine the Ka for the unknown acid.

18.71 Determine the pH of a buffer solution produced by dissolving 10.0 L of ammonia gas (at 25°C and 855 torr)

in 750.0 mL of a solution containing 36.0 g of ammonium nitrate (NH4NO3).

18.72 (a) What is the pH at the equivalence point in the titration of 25.00 mL of 0.1000 M barium hydroxide with

0.05000 M acetic acid? (b) What would be an appropriate indicator?

18.73 Determine the pH at the equivalence point for each of the following titrations:

(a) 25.00 mL of 0.100 M chloric acid with 0.100 M sodium hydroxide. (b) 25.00 mL of 0.100 M nitrous acid with

0.100 M potassium hydroxide. (c) What would be an appropriate indicator in each case?

18.74 Calculate the pH of the following solutions. (a) a solution containing 0.21 M HNO2 and 0.12 M NaNO2.

(b) 65.0 g of acetic acid and 64.0 g of sodium acetate diluted to 500.0 mL with water.

18.75 Calculate the pH in a titration of a 50.00 mL sample of 0.2500 M acetic acid (HC2H3O2) after the addition of

25.00 mL of a 0.2500 M sodium hydroxide (NaOH).

18.76 Calculate the original molarity of a solution of nitrous acid if the pH of the solution is 2.00 at equilibrium.

18.77 (a) Calculate Kb for the oxalate ion. (b) Calculate Kb for the phosphate ion. (c) Calculate Kb for the sulfide

ion.

18.78 Calculate the pH of each of the following solutions. (a) 0.100 M sodium hypochlorite (b) 0.200 M barium

cyanide

18.79 How many grams of potassium cyanide must be added to 250.0 mL of water to produce a solution with pH =

11.00?

18.80 Are solutions of the following acidic, basic, or neutral? (a) NH4Cl (b) Ca(ClO4)2 (c) NaC2H3O2 (d) KNO3

(e) LiCN

18.81 Calculate the pH of a solution made by adding 135 g of cyanic acid and 95.0 g of sodium hydroxide to

sufficient water to produce 2.00 L of solution.

18.82 Calculate the pH of a solution made by adding 248.23 g of calcium cyanate to 500.0 mL of a 2.000 M nitric

acid solution.


18.83 Write molecular equations (balanced) for each of the following reactions. (a) potassium hydroxide +

hypochlorous acid (b) calcium hydroxide + chlorous acid (c) sodium hydroxide + chloric acid (d) magnesium

hydroxide + perchloric acid (e) barium carbonate + nitric acid

18.84 Write balanced molecular and net ionic equations for each of the following reactions. Assume that all of the

reactions go to completion. (a) Perchloric acid plus sodium carbonate (b) Ammonia plus sulfuric acid (c) Carbonic

acid plus calcium hydroxide (d) Potassium hydroxide plus acetic acid (e) Strontium hydroxide plus nitric acid

18.85 (a) A 22.15 mL volume of 0.5000 M H2SO4 neutralizes 50.00 mL of NaOH. What is the concentration of the

NaOH solution? (b) A 25.27 mL volume of 0.5000 M sulfuric acid neutralizes 50.00 mL of sodium hydroxide.

What is the concentration of the base?

18.86 A 60.00 mL sample of a sodium hydroxide solution was titrated with a 0.1000 M phosphoric acid solution.

The titration required 20.00 mL of the acid to reach the endpoint corresponding to the removal of all acidic

hydrogen ions from the acid. What was the concentration of the base?

18.87 Calculate the molarity of a hydrochloric acid solution if 50.00 mL of the solution reacted with 0.4185 g of

sodium carbonate in the reaction

2 HCl + Na2CO3 2 NaCl + CO2 + H2O

18.88 What volume, in milliliters, of 0.04815 M NaOH is required to reach the equivalence point when it is titrated

with 25.00 mL of 0.01805 M H2SO4?

18.89 What volume (mL) of 0.1650 M HCl is required to titrate 150.0 mL of a solution that contains 5.65 g of

NaOH per liter?

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

18.91 Name each of the following acids and draw its Lewis structure. (a) H2SO4 (b) HClO3 (c) HNO2 (d) H2CO3

(e) H3PO4

18.92 Name each of the following acids and draw its Lewis structure. (a) HNO2 (b) CH3COOH (c) H3PO4

(d) H2CO3 (e) HF

18.93 Write the chemical formula for each of the following compounds and indicate the oxidation state of nitrogen

in each:

Formula Oxidation state

(a) Nitrous acid ____________________ _____

(b) Calcium cyanide ____________________ _____

(c) Barium nitrate ____________________ _____

(d) Ammonium chloride ____________________ _____

(e) Magnesium nitride ____________________ _____

18.94 Predict the molecular geometries of the following. (a) H2SO4 (b) H6TeO6

18.95 Use appropriate equations to indicate why each of the following aqueous solutions is basic. (a) KOH (b) NH3

(c) Ca(C2H3O2)2 (d) Ba(CN)2 (e) Na2O

18.96 Use appropriate equations to indicate why each of the following aqueous solutions is acidic. (a) HF

(b) HC2H3O2 (c) NH4NO3 (d) FeCl3 (e) N2O5

18.97 A solution was prepared containing 55.0 g of ammonia and 85.0 g of ammonium sulfate in a total volume of

500.0 mL. Determine the pH of this solution after the addition of 10.0 mL of 6.0 M chloric acid. (Ignore any effect

that the sulfate might have on the pH.)


18.98 What is the pH of a solution made by dissolving 286 g of calcium chlorite in 1.000 L of a 3.00 M chlorous

acid solution?

18.99 A 0.0350 M solution of a monoprotic acid in water is 13.5 percent ionized. Calculate the Ka of the acid.

18.100 A 5.0 × 10–3 M solution of the weak organic base codeine (C18H21NO3) has a pH of 9.95. Calculate the value

of Kb for this substance.

18.101 Name or give the formula for each of the following. (a) Tritium (b) Potassium borohydride (c) Diborane

(d) Diamond (e) Calcium acetylide (f) RbO2 (g) BO33– (h) B4O7

2– (i) SiC (j) Na2SiF6 (k) Buckminsterfullerene

(l) Tetraboric acid (m) BaO2 (n) BN (o) H2SiF6

18.102 Name or give the formula for each of the following. (a) Sodium orthoborate (b) Pyrophosphate ion

(c) Methylammonium ion (d) Barium acetylide (e) Ammonia (f) NH2Cl (g) Na2B4O7 (h) H5P3O10 (i) (NH4)2SO4

(j) NH4Cl (k) Orthophosphoric acid (l) Hydroxylamine (m) N2H2(CH3)2 (n) Sodium borohydride

(o) Diethylammonium ion

18.103 Name or give the formula for each of the following. (a) Cesium superoxide (b) Potassium orthoborate

(c) Potassium tetraborate (d) Methane (e) Potassium hexafluorosilicate (f) 1H (g) BH4– (h) B2O3 (i) C(gr) (j) C2

2–

(k) Deuterium (l) CaB4O7 (m) Ca3(BO3)2 (n) T (o) Sodium peroxide

18.104 Write molecular and net ionic equations (balanced) for each of the following reactions. (a) sodium

hydroxide + hydrobromic acid (b) barium hydroxide + chloric acid (c) calcium carbonate + hydrochloric acid

(d) magnesium hydroxide + perbromic acid (e) rubidium hydroxide + chlorous acid

18.105 Heroin is diacetylmorphine. The compound is 68.28% C, 6.28% H, 3.79% N, and 21.66% O. What is the

empirical formula of heroin?

18.106 (a) What is the percent ionization of a 0.0100 molal solution of acetic acid (HC2H3O2) if the solution freezes

at −0.0194°C? (b) Calculate the ionization constant, Ka, for acetic acid at this temperature. Assume that, in this

case, the molality (m) = molarity (M). The freezing-point depression constant for water is 1.86 °C/m.

18.107 The freezing point of a 0.9775 M hydrofluoric acid solution (HF) is −1.858°C. (a) Estimate the percent

ionization of the hydrofluoric acid solution at this temperature. Assume the molality of this solution is

approximately equal to the molarity. (b) Calculate the Ka for hydrofluoric acid at this temperature.

(a) 2.30 % (b) Ka = 5.30 × 10–4

18.108 When 0.30 mol of an unknown monoprotic acid (HA) are added to 100.0 mL of a 1.20 M potassium

hydroxide (KOH) solution, the resultant pH is 4.85. What is the Ka of the unknown acid? Ka = 9.4 × 10–6

18.109 The following reaction will produce dinitrogen trioxide (N2O3) at temperatures below −30°C:

4 NO(g) + O2(g) 2 N2O3(l)

Dinitrogen trioxide is the acid anhydride of nitrous acid. An evacuated 10.00 L container is cooled to −40.0°C. The

introduction of some nitrogen oxide (NO) increases the pressure to 8.00 atm. The introduction of some oxygen gas

quickly raises the pressure to 9.80 atm before the reaction begins to lower the pressure. After a period of time, the

unreacted NO and O2 are removed and the contents dissolved in 1.525 L of a 1.00 M sodium nitrite solution to give

a solution with a pH of 3.20. (a) What was the theoretical yield of the reaction? (b) What was the percent yield of

the reaction?



19.1 Which of the two new types of equilibria presented in this chapter is heterogeneous, and which is

homogeneous?

19.3 List the solubility rules from Chapter 4.

19.4 How are heterogeneous equilibria different from homogeneous equilibria?

19.6 (a) Define molar solubility. (b) Define solubility.

19.7 What term(s) go into the denominator of a Ksp mass action expression?

19.8 In Ksp, what does the subscript "sp" mean?

19.10 Zinc sulfide (ZnS) is an insoluble compound. Answer the following questions as they apply to the results of

adding a quantity of zinc sulfide to water. (a) Is this a homogeneous or heterogeneous equilibrium? (b) Where

(reactant or product side) does the zinc sulfide appear in the equilibrium reaction equation? (c) Where (reactant or

product side) does the water appear in the equilibrium reaction equation? (d) Where (reactant or product side) do

the zinc ions appear in the equilibrium reaction equation? (e) Where (reactant or product side) do the sulfide ions

appear in the equilibrium reaction equation?

19.11 Could a precipitate form when the following pairs of solutions are mixed? Answer yes or no for each. If a

precipitate forms gives its name and formula. (a) sodium bromide + iron(III) sulfate (b) ammonium sulfide +

iron(II) sulfate (c) rubidium hydroxide + cobalt(II) chloride (d) sodium phosphate + strontium iodide (e) potassium

iodide + silver nitrate

19.12 (a) How does the equilibrium constant expression for a Ksp differ from a homogeneous equilibrium constant

expression? (b) Write Ksp expressions for each of the following solids in water: AlPO4, ZnS, Tl2S, AlF3, Ag2CrO4,

and Cu3(AsO4)2

19.13 Determine which member of the following pairs has the greater molar solubility. (a) ZnS or MnS; (b) CuBr or

AgBr; (c) Cu(OH)2 or Ga(OH)3; (d) Ca3(PO4)2 or NiS; (e) zirconium(IV) phosphate or copper(II) sulfide.

19.14 Calculate the molar solubility of calcium carbonate (CaCO3) that will dissolve in each of the following. The

Ksp for calcium carbonate is in the Appendix. (a) in pure water (b) in 4.5 × 10–2 M Ca(NO3)2 solution (c) in 0.25 M

Na2CO3 solution. (a) 9.3 × 10–5 M (b) 1.9 × 10–7 M (c) 3.5 × 10–8 M

19.15 Solid sodium sulfate is added slowly to a solution that is 0.115 M in lead(II) nitrate and 0.115 M in barium

nitrate. Determine the percentage of the barium that will precipitate just before the lead(II) sulfate begins to

precipitate. The Ksp of barium sulfate is 1.1 × 10–10, and the Ksp of lead sulfate is 1.6 × 10–8.

19.16 (a) Determine the Ksp of lead(II) bromide if the molar solubility of PbBr2 at 25°C is 1.0 × 10–2 mol/L. (b) If

0.0386 g of lead(II) sulfate dissolves per liter of solution, calculate the solubility product. (c) Lead(II) iodate

(Pb(IO3)2) has a molar solubility of 2.4 × 10–11 mol/L in a 0.100 M NaIO3 solution. Calculate the Ksp for Pb(IO3)2.

19.17 A slightly soluble metal hydroxide (MOH) produces a saturated solution with a pH of 9.72. Calculate the Ksp

of the compound. Ksp = 2.8 × 10–9

19.18 Given that the Ksp of Fe(OH)2 is 7.9 × 10–16, determine the solubility in grams per liter (a) at pH 7.0; (b) at pH

8.9; (c) at pH 12.1. The Ksp of Fe(OH)2 is 7.9 × 10–16.


19.19 (a) Will cadmium hydroxide (Ksp = 2.2 × 10–14) precipitate from solution if the pH of a 0.050 M solution of

cadmium nitrate is adjusted to 8.0? (b) Will lead(II) chloride (Ksp = 1.6 × 10–5) precipitate when 100.0 mL of 1.00

M lead(II) nitrate is mixed with 20.0 mL of 5.0 × 10–2 M calcium chloride solution?

19.20 Zinc ions may be precipitated as zinc hydroxide (Ksp = 4.5 × 10–17). What is the minimum pH required to

reduce the zinc ion concentration to 1 part per billion (ppb)? A concentration of 1 ppb is equal to 1 g/L.

19.21 Some substances are more soluble in acidic solution than in pure water. Predict which of the following are

more soluble in an acidic solution: (a) CaCO3 (b) AlF3 (c) BaI2 (d) Hg2(CN)2 (e) Ca3(PO4)2

19.22 A solution buffered at pH = 2.50 is saturated with hydrogen sulfide (0.10 M). What is the solubility (in g/L)

of CuS? 1.8 × 10–30 g/L


19.23 Is a ligand a Lewis acid or a Lewis base?

19.24 Ammonia is a common ligand for many transition metals. Draw the Lewis structure of ammonia and show

how it might interact with a transition metal ion to form a complex. What is the molecular geometry of ammonia?

19.25 Describe what each of the following means: Kf, Kstab, Kd, and Kinstab.

19.26 How does the value of a Kf relate to the Kd for the same complex?

19.27 Which concentrations, if any, often do not have an exponent of 1 in a Kf mass action expression?

19.28 From the Kf for [Zn(NH3)4]2+ of 5.0 × 108, calculate the concentration of zinc ion in 1.0 L of a solution that

contains a total of 1.0 × 10–3 mol of zinc ion and that is 0.10 M in ammonia. 2.0 10–8 M Zn2+

19.29 (a) Determine the Ag+ concentration present at equilibrium in a 1.000 L solution containing a total of 2.00 ×

10–4 mol of metal ion and an ammonia concentration of 0.300 M. (b) Determine the Co3+ concentration present at

equilibrium in a 1.000 L solution containing a total of 1.00 × 10–4 mol of metal ion and an ammonia concentration of

0.200 M. (c) Determine the Zn2+ concentration present at equilibrium in a 1.000 L solution containing a total of 1.50

× 10–4 mol of metal ion and an ammonia concentration of 0.220 M.

(a) 1.3 × 10–10 M Ag (b) 9.8 10–36 M Co3+ (c) 1.3 10–10 M Zn2+

19.30 (a) What is the uncomplexed cobalt(III) concentration in a 0.200 M solution of [Co(NH3)6]Cl3? (b) What is

the uncomplexed iron(III) concentration in a 0.100 M solution of K3[Fe(CN)6]? (c) Calculate the uncomplexed

mercury(II) ion concentration present in a 0.1245 M solution of [HgCl4]2–.

(a) 1.6 10–6 M (b) 5.8 10–6 M (c) 1.2 10–4 M

19.31 Determine the uncomplexed mercury(II) ion concentration in 500.0 mL of a solution with an initial Hg2+

concentration of 0.0100 M that has had 63 g of potassium iodide added to it. The formation constant of HgI42– is 1.0

× 1030.

19.32 Calculate the pH of a solution prepared by dissolving 0.100 mol of tetraamminecopper(II) chloride in 2.00 L

of solution. pH = 10.35


19.33 Why is it easier to combine the Ksp and Kf constants, then to treat them separately, when investigating the

interaction of an insoluble substance with a solution containing a complexing agent?


19.34 (a) Write the balanced equilibrium reaction equation for the Ksp of CuCO3. (b) Write the mass action

expression for the Ksp of CuCO3. (c) Write the balanced equilibrium reaction equation for the Kf of [Cu(NH3)4]2+.

(d) Write the mass action expression for the Kf of [Cu(NH3)4]2+. (e) Write the balanced equilibrium reaction

equation that results when you add your answer for part (a) to your answer for part (c). (f) Write the mass action

expression for part (e). (g) Multiply your answer for part (b) by your answer for part (d) and compare the result to

your answer for part (f).

19.35 (a) Show how to combine the equilibrium reactions and equilibrium constants for the Ksp of silver chloride

(AgCl) and the Kf of [Ag(NH3)2]+ to get an overall expression for adding silver chloride to an ammonia solution.

(b) Show how to combine the equilibrium reactions and equilibrium constants for the Ksp of copper(I) carbonate

(Cu2CO3) and the Kf of [Cu(NH3)2]+ to get an overall expression for adding copper(I) carbonate to an ammonia

solution. (c) What procedural difference is there between your solutions to parts (a) and (b)?

19.36 In order to cause a significant amount of an insoluble solid to dissolve in a solution containing a complexing

agent, what must be the relationship between the Ksp and Kf values?

19.37 Calculate the total molarity of copper in a solution made by adding solid copper(I) iodide (CuI), with a Ksp of

5.1 × 10–12, to a 0.25 M solution of ammonia (NH3). The Kf of [Cu(NH3)2]+ is 7.2 × 1010. 0.056 M Cu2+

19.38 Calculate the molar solubility of AgCl in a 0.750 M solution of ammonia.

19.39 Calculate how many grams of copper(II) arsenate will dissolve in 3.50 L of 1.50 M ammonia.

205 g Cu3(AsO4)2

19.40 What is the minimum number of moles of solid sodium hydroxide, NaOH, necessary to dissolve 0.10 mole of

solid zinc hydroxide, Zn(OH)2, to make 1.00 L of a solution with a free zinc ion concentration of 1.00 × 10–3 M?

Most of the zinc in the solution will be present as the [Zn(OH)4]2– ion? Kf [Zn(OH)4]2– = 4.6 × 1014

19.41 A total of 1.20 g of solid zinc hydroxide (Zn(OH)2) is to be dissolved in sufficient water to produce 1.00 L of

solution. To do this the zinc ions are to be complexed as Zn(OH)42–. Potassium hydroxide (KOH) will be the source

of hydroxide ions. At what pH will all of the solid dissolve and have a free zinc ion concentration of 1.00 × 10–3 M?


19.43 The addition of an aqueous hydrochloric acid (HCl) solution to solid calcium carbonate (CaCO3) results in the

formation of carbon dioxide (CO2) gas. (a) Write a balanced molecular chemical equation for the reaction of

hydrochloric acid with calcium carbonate. (b) Write the balanced net ionic equation for the reaction of hydrochloric

acid with calcium carbonate.

19.44 The addition of an aqueous acetic acid (HC2H3O2) solution to solid calcium carbonate (CaCO3) results in the

formation of carbon dioxide (CO2) gas. (a) Write a balanced molecular chemical equation for the reaction of acetic

acid with calcium carbonate. (b) Write the balanced net ionic equation for the reaction of acetic acid with calcium

carbonate. (c) Write equilibrium reactions for each of the equilibria occurring while acetic acid reacts with calcium

carbonate.

19.45 Many sources list the compound AgCN (silver cyanide) as "not soluble" in water. The Ksp for AgCN is 1.4 ×

10–16. (a) Write the balanced chemical equation and the mass action expression for the Ksp of AgCN in water. (b)

Write the balanced chemical equations and the mass action expressions for two other equilibria present when you

add AgCN to water.

19.46 The compound Ag3PO4 (silver phosphate) is listed as "not soluble" in water. The Ksp for Ag3PO4 is

1.3 × 10–20. (a) Write the balanced chemical equation and the mass action expression for the Ksp of Ag3PO4 in water.

(b) Write the balanced chemical equations and the mass action expressions for four other equilibria present when

you add Ag3PO4 to water.



19.47 Define qualitative analysis.

19.48 A solution contains the following cations: Ag+, Cu2+, and Fe3+. Identify a reagent that will precipitate one of

these ions but not the other two.

19.49 A solution contains the following cations: K+, Ba2+, and Fe3+. Identify a reagent that will precipitate one of

these ions but not the other two.

19.50 A solution contains the following cations: K+, Pb2+, and Fe3+. (a) Identify a reagent that will precipitate one of

these ions but not the other two. (b) Identify a reagent that will precipitate one of the ions remaining in solution

after the treatment in part (a) and not the other.

19.51 A solution contains the following cations: NH4+, Ca2+, and Fe3+. (a) Identify a reagent that will precipitate one

of these ions but not the other two. (b) Identify a reagent that will precipitate one of the ions remaining in solution

after the treatment in part (a) and not the other.

19.52 Suggest how the ions in each pair may be separated in a qualitative analysis experiment. (a) K+ and Cu2+

(b) Zn2+ and Ba2+ (c) Cd2+ and Cr3+ (d) Hg22+ and Hg2+

19.53 A solution buffered at pH = 2.60 is saturated with hydrogen sulfide (0.10 M). If the solution is 0.010 M in

each of the ions La3+, Mn2+, and Sb3+, which, if any, will precipitate? (La = lanthanum.) Sb2S3

19.54 A solution buffered at pH = 2.50 is saturated with hydrogen sulfide (0.10 M). If the solution is 0.010 M in

each of the ions Ag+, Bi3+, and Tl+, which, if any, will precipitate? (Tl = thallium)

19.55 A solution is saturated with hydrogen sulfide (0.10 M). The solution is 0.010 M in Fe2+ and 0.010 M in Mn2+.

Determine the highest and lowest pH at which only one of the ions will precipitate.

Fe2+ only will precipitate between pH = 2.79 and pH = 5.14

19.56 A student wishes to test a sample for the presence of silver ions. According to the qualitative analysis scheme,

she should be able to test for silver ions by adding a substance, such as hydrochloric acid, that will supply chloride

ions. If silver ion is present, a white precipitate of silver chloride (AgCl) should form. She begins to add

hydrochloric acid to her sample, and a white precipitate immediately forms. However, as she continues to add

hydrochloric acid, the white precipitate disappears. Explain why the precipitate formed and then disappeared. You

may use any appropriate equilibrium constants in your explanation.


19.57 How many grams of silver nitrate (AgNO3) and how many grams of potassium bromide (KBr) are necessary

to precipitate 2.50 g of silver bromide to produce photographic film?

19.59 The following chemical reaction takes place during the development stage of film processing:

2 AgBr(s) + SO32–(aq) + H2O(l) 2 Ag(s) + SO4

2–(aq) + 2 H+(aq) + 2 Br–(aq)

Assign oxidation numbers to each of the elements in this reaction.

19.60 The following reaction occurs when a piece of photographic film is "fixed”:

AgBr(s) + 2 S2O32–(aq) [Ag(S2O3)2]3–(aq) + Br–(aq)

What is the value of the equilibrium constant for this reaction? The Ksp for AgBr is 5.3 × 10–13, and the Kf for

[Ag(S2O3)2]3– is 1.6 × 107.

19.61 How many grams of silver bromide (AgBr) will dissolve in 1.00 L of a 0.200 M sodium thiosulfate (Na2S2O3)

solution?


19.65 BAL (British anti-Lewisite) is the compound 2,3-dimercapto-1-propanol (C3H8S2O). The compound has three

carbon atoms in a row. The first carbon atom also has two hydrogen atoms and a mercapto group (–SH) attached.

The central carbon also has one hydrogen atom and a mercapto group attached. The remaining carbon atom has two

hydrogen atoms and a hydroxyl group (–OH) attached. Draw the structure of BAL. What is the hybridization of

each carbon atom?

Summary

19.69 A sample of an alloy containing chromium was dissolved in acid. Oxidation of the chromium produced the

chromate ion and the solution was separated from other contaminants. Next, the chromate ion was reduced by 3.17

g of sodium sulfite in a basic solution. This converted the chromium to chromium(III) hydroxide, and the sulfite

became sulfate. If the original sample weighed 2.971 g, what was the percent chromium?

19.70 Calculate the number of milliliters of 0.1500 M sulfuric acid required to titrate 25.00 mL of 0.3000 M sodium

hydroxide to the endpoint.

19.71 The acetic acid in a sample of vinegar was extracted and then was diluted with water to a total volume of

200.00 mL. This final solution was titrated with 0.1200 M barium hydroxide solution. The titration required 40.00

mL of the base to reach the endpoint. What was the concentration of the extracted acetic acid solution?

19.72 A 20.00 mL sample of a phosphoric acid solution was titrated with a 0.1000 M barium hydroxide solution.

What was the concentration of the acid if 30.00 mL of base were required to reach the endpoint?

19.73 Draw the Lewis dot formula for each species and label the Lewis base in each of the following. (a) 2 NH3(aq)

+ Ag+(aq) [Ag(NH3)2]+ (aq) (b) NH3(aq) + H+(aq) NH4+(aq) (c) Cl–(aq) + AlCl3(aq) AlCl4

–(aq)

(d) SnCl2(aq) + 2 Cl–(aq) SnCl42–(aq) (e) Cu+(aq) + 2 CN–(aq) [Cu(CN)2]– (aq)

19.74 Name or give the formula for each of the following. (a) Sodium peroxide (b) Calcium orthoborate (c)

Calcium tetraborate (d) Silicon carbide (e) Sodium hexafluorosilicate (f) T (g) KBH4 (h) B2H6 (i) C(dia) (j) BaC2

(k) Deuterium (l) Buckminsterfullerene (m) Borohydride ion (n) CH4 (o) KO2

19.75 Name or give the formula for each of the following. (a) Potassium superoxide (b) Orthoboric acid

(c) Tetraboric acid (d) Methane (e) Hexafluorosilicic acid (f) D (g) BH4– (h) BN (i) C60 (j) C2

2– (k) Barium acetylide

(l) Ca3(BO3)2 (m) Potassium borohydride (n) SiC (o) Diborane

19.76 Name or give the formula for each of the following. (a) Protium (b) Diethylamine (c) Orthophosphoric acid

(d) B4O72– (e) Ammonia (f) H3BO3 (g) Na5P3O10 (h) (C2H5)2NH2

+ (i) NH2Cl (j) N2H2(CH3)2 (k) Methylamine

(l) NH4Cl (m) (HPO3)3 (n) Hydroxylamine (o) P2O74–

19.77 Name or give the formula for each of the following. (a) Na3BO3 (b) NH3 (c) Ammonium pyrophosphate

(d) Potassium orthoborate (e) Dimethylhydrazine (f) Phosphorous acid (g) Tripolyphosphoric acid (h) CsO2

(i) N2H3CH3 (j) K2B4O7 (k) CH3NH3+ (l) Hydrazine (m) Rubidium superoxide (n) Orthoborate ion (o) Ammonium

sulfate

19.78 Name or give the formula for each of the following. (a) Xenon tetrafluoride (b) Orthotelluric acid (c) Calcium

triiodide (d) Chloramine (e) Boron nitride (f) KrF2 (g) S22– (h) HIO4 (i) N2H4 (j) CaC2 (k) Selenous acid (l) Boron

oxide (m) Na2SiF6 (n) H2B4O7 (o) Ozone

19.79 Name or give the formula for each of the following. (a) Pyrosulfuric acid (b) Iron(II) disulfide

(c) Paraperiodic acid (d) Methylhydrazine (e) XeF6 (f) H2SeO3 (g) I3– (h) NH2OH (i)Trioxygen (j) Selenic acid

(k) Potassium triiodide

19.80 Titanium(III) will react with water and acidified permanganate as follows:

5 Ti3+ + MnO4– + 6 H2O 5 TiO2(s) + Mn2+ + 12 H+

What is the percentage of titanium in a 2.468 g sample if 43.96 mL of a 0.02317 M potassium permanganate

solution are required to react with it?


19.81 Write balanced molecular and net ionic equations for each of the following. Assume all reactions go to

completion in aqueous solution. (a) Chloric acid plus potassium carbonate (b) Ammonium chloride plus sodium

hydroxide (c) Carbonic acid plus calcium hydroxide (d) Acetic acid plus strontium hydroxide (e) Nitric acid plus

barium hydroxide


completion in aqueous solution. (a) Hydrochloric acid plus ammonia (b) Sodium hydroxide plus nitric acid

(c) Oxalic acid plus lithium hydroxide (d) Ammonium chloride plus potassium hydroxide (e) Calcium hydroxide

plus phosphoric acid

19.83 The following equilibrium is established in a 2.00 L container:

CS2(g) + 2 H2(g) CH4(g) + 2 S(s) H° = −162 kJ

Determine if each of the following will result in a shift to the right, to the left, or no change. (a) Add CH4(g)

(b) Add CS2(g) (c) Add S(s) (d) Increase the temperature (e) Increase the volume to 4.0 L

19.84 At 25°C, hydrosulfuric acid has the following equilibria:

H2S(aq) H+(aq) + HS–(aq) Ka1 = 1.0 × 10–7

HS–(aq) H+(aq) + S2–(aq) Ka2 = 1.3 × 10–13

Find the value of the equilibrium constant for the following reaction at the same temperature:

H2S(aq) 2 H+(aq) + S2–(aq)

19.85 Write the Ka expressions for arsenic acid (H3AsO4); it is a triprotic acid.

19.86 Give the conjugate base of each of the following Brønsted-Lowry acids. (a) H2Se (b) HNO2 (c) HClO2

(d) HCO3– (e) NH4

+

19.87 Calculate the pH of a solution made by mixing 79.085 g of calcium acetate with 250.0 mL of 1.000 M

perchloric acid.

19.88 Which of the following ions will undergo hydrolysis? (a) Cl– (b) BrO2– (c) SO3

2– (d) NO3– (e) SCN– (f) NH4

+

(g) Rb+ (h) Fe3+ (i) Ca2+ (j) Cs+

19.89 How many grams of NaCN would you need to dissolve in enough water to make exactly 500.0 mL of a

solution whose pH is 9.95?

19.90 Calculate the pH of a 0.150 M solution of ammonium sulfate (ignore hydrolysis of the sulfate ion).

19.91 Calculate the pH of a solution that is 0.350 M in strontium cyanate and 0.100 M in cyanic acid.

19.92 Determine the pH at the equivalence point for the titration of 25.00 mL of 0.1000 M barium hydroxide with

0.1000 M chlorous acid.

19.93 Write balanced equilibrium equations and solubility product mass action expressions for the solubility

equilibria of the following. (a) AgBr (b) Hg2Br2 (c) Co3(PO4)2

19.94 The molar solubility of silver(I) carbonate at 25°C is 1.3 × 10–4 mol/L. Calculate Ksp.

19.95 A solution containing arsenate, carbonate, fluoride, hydroxide, sulfite, and sulfate (all at 0.00010 M) gradually

has calcium ion added. In what order will the ions precipitate?

19.96 A solution buffered at pH = 2.00 is saturated with hydrogen sulfide (0.10 M). What is the solubility (in g/L)

of Cu2S?


19.97 Does a precipitate form when the following pairs of solutions are mixed? If a precipitate forms, what is its

name and chemical formula? (a) sodium bromide + iron(III) sulfate (b) ammonium sulfide + iron(II) sulfate

(c) rubidium hydroxide + cobalt(II) chloride (d) sodium phosphate + strontium iodide (e) potassium iodide + silver

nitrate

19.98 Will lead(II) chloride (Ksp = 1.6 × 10–5) precipitate when 100.0 mL of 1.00 M lead(II) nitrate is mixed with

20.0 mL of 5.0 × 10–2 M calcium chloride solution?

19.99 Two solutions are mixed and a precipitate forms. The first solution was 80.0 mL of 0.10 M silver nitrate, and

the second solution was 20.0 mL of 0.15 M sodium bromide. Calculate the equilibrium concentrations of silver ion

and bromide ion if the solubility product constant for the precipitate is 5.0 × 10–13.

19.100 What is the copper ion concentration in a saturated solution of copper(I) sulfide?

19.101 Which is more soluble (i.e., which yields the larger molarity in solution), chromium(III) hydroxide or

chromium(III) phosphate?

19.102 What is the equilibrium lead ion concentration resulting when 50.0 mL of 0.100 M lead(II) nitrate is mixed

with 50.0 mL of 0.250 M calcium chloride? 7.1 × 10–5 M Pb2+

19.103 What is the solubility (in grams per liter) of mercury(I) sulfide?

19.104 What is the iron(II) concentration in a 0.200 M solution of (NH4)4[Fe(CN)6]?

19.105 Aqueous solutions of copper(II) ions have a characteristic blue color due to the complex formed by

copper(II) ions with water. The addition of ammonia changes the color to a significantly more intense blue as

ammonia replaces the water as the ligand about the copper(II) ions. It is possible to estimate the formula of the

complex by observing the color change due to ligand displacement. A solution containing 2.49 g of CuSO4•5H2O in

100.0 mL of solution was exposed to ammonia gas. The following change occurred:

n NH3(g) + Cu2+(aq) [Cu(NH3)n]2+(aq)

colorless light blue intense blue

The blue color of the solution intensified until 978 mL of ammonia gas (at 25°C and 1.00 atm) dissolved. Estimate

the value of n.


20.1 Which of the three laws of thermodynamics appears both in Chapter 6 and in this chapter?

20.2 What do the symbols H, S, and G represent?

20.3 (a) What is a spontaneous process? (b) What is a nonspontaneous process?

20.4 Is a process that takes place very slowly a spontaneous or a nonspontaneous process? Explain your answer is

either case.

20.5 Predict whether each of the following processes is spontaneous or nonspontaneous. (a) Charcoal ignites. (b) A

bottle of household ammonia sits open, and the odor spreads throughout the room. (c) The N2 in the air all moves

toward the ceiling of the room. (d) A "dead" battery in a calculator is replaced with a new one. (e) Mixtures of

gasoline vapor and air ignite when exposed to an electrical spark. (f) Table salt dissolves in hot water.

(a) Nonspontaneous (b) Spontaneous (c) Nonspontaneous (d) Nonspontaneous (e) Spontaneous (f) Spontaneous


20.6 Predict whether each of the following processes is spontaneous or nonspontaneous. (a) Ice cubes melt at −15°C

under standard pressure. (b) Milk mixes in a cup of hot coffee. (c) Chlorine molecules separate into chlorine atoms

under standard conditions. (d) When shaken, salad oil mixes with vinegar. (e) Carbon dioxide gas reacts with water

vapor under standard conditions to produce oxygen gas and natural gas (methane).

20.7 (a) Thermodynamics is the study of what aspect of chemistry? (b) Initially, the study of thermodynamics

centered on what type of energy?

20.8 What does the symbol "∆" represent?

20.9 What is the state of a system? (It may be necessary to review Chapter 6.)

20.10 State the first law of thermodynamics.

20.11 What is the relationship between changes in internal energy, heat, and work?

20.12 (a) What is the amount of work produced in a process occurring under constant volume conditions? (b) How

do heat and work compare in a process occurring under constant pressure conditions?

20.13 (a) What is the sign of the enthalpy change for an exothermic process? (b) What is the sign of the enthalpy

change for an endothermic process?

20.14 Determine how much the internal energy of the system changes in each of the following cases. (a) A balloon

pops, and the released gas expands, doing 145 J of work on the surroundings, without any significant heat exchange.

(b) A pan containing water is heated on a stove and absorbs 2250 J of heat. (c) A sample of argon gas is compressed

by 475 J of work; at the same time, 225 J of heat energy are removed from the sample.

(a) ΔE = –145 J (b) ΔE = +2250 J (c) ΔE = +250. J

20.2 Corresponds to BLBMWS 19.2 & 19.3 & 19.4

20.15 What is the key to the second law of thermodynamics?

20.16 (a) Describe the thought experiment developed by James Clerk Maxwell in 1867. (b) What was the error that

Leó Szilárd identified in Maxwell's thought experiment? (c) Give an example of where Maxwell's demon is

apparently working (although nonspontaneously) in our body.

20.17 What is the IUPAC definition of entropy?

20.18 What is a microstate?

20.19 (a) What is the relationship between the entropy values of solids, liquids, and gases? (b) When water freezes,

is the entropy change positive or negative? (c) When water boils, is the entropy change positive or negative?

20.20 In Chapter 6, heating curves show how the temperature of a substance increases with increasing amount of

heat added. Construct a graph similar to a heating curve using entropy instead of temperature as the vertical axis.

Plot how the entropy changes when a sample of ice is heated from −30 °C to steam at 130 °C.

20.21 Predict whether the entropy change of the system will be positive or negative in each of the following cases.

Explain your reasoning in each one. (a) ice melts (b) liquid ammonia vaporizes (c) solid sugar dissolves in water

(d) carbon dioxide gas is converted to dry ice (solid carbon dioxide) (e) dew forms on grass

20.24 Predict which member of the following pairs will have the greater entropy: (a) 2.0 mol of oxygen gas at 298

K and 1.0 atm pressure, or 2.0 mol of oxygen gas at 500 K and 1.0 atm pressure. (b) 1.0 mol of C2H5OH(g) at 78.5

°C and 1.0 atm pressure, or 1.0 mol of C2H5OH(l) at 78.5 °C and 1.0 atm. (c) 2.0 mol of chlorine gas at 298 K in a

30.0 L volume or 4.0 mol of gaseous chlorine atoms at 298 K in a 60.0 L volume. (d) 2.0 mol of solid calcium

chloride at 25 °C, or 2.0 mol of aqueous calcium chloride at 25 °C.



20.25 (a) State the second law of thermodynamics. (b) State the third law of thermodynamics.

20.26 When water freezes, the entropy change of the system is negative. How can water freeze spontaneously if the

entropy change is negative?

20.27 What is the relationship between the entropy change of the surroundings and the enthalpy change of the

system?

20.28 The third law of thermodynamics states that the entropy of a pure crystalline solid at absolute zero is zero.

Why is this true?

20.29 Why do thermodynamic tables, such as the one in Appendix X, not normally contain ∆Sf° values?

20.30 Predict if the entropy change in each of the following cases will be positive or negative.

(a) C(s) + 2 Cl2(g) CCl4(g) (b) 3 Mg(s) + N2(g) Mg3N2(s)

(c) 2 KClO3(s) 2 KCl(s) + 3 O2(g) (d) 3 O2(g) 2 O3(g)

(e) 2 C8H18(l) + 25 O2(g) 16 CO2(g) + 18 H2O(g)

20.31 Calculate the entropy change in each of the following. (a) 2 HCl(g) + Br2(g) 2 HBr(g) + Cl2(g)

(b) N2(g) + 3 H2(g) 2 NH3(g) (c) 2 CH3OH(l) + 3 O2(g) 2 CO2(g) + 4 H2O(l)

(d) 2 MnO(s) + O2(g) 2 MnO2(s) (e) H2O(l) H2O(g)

20.32 Calculate ∆S° for each of the following reactions: (a) C2H4(g) + 2 O2(g) 2 CO(g) + 2 H2O(g)

(b) H2SO4(l) SO3(g) + H2O(l) (c) BaCO3(s) + H2SO4(l) H2O(l) + CO2(g) + BaSO4(s)

(d) C2H5OH(l) + 3 O2(g) 2 CO2(g) + 3 H2O(l) (e) 2 Fe2O3(s) + 3 Cl2(g) 4 FeCl3(s) + 3 O2(g)

20.33 Calculate ∆S° for each of the following reactions. (a) Al2O3(s) + 3 H2SO4(l) Al2(SO4)3(s) + 3 H2O(l)

(b) 2 C2H6(g) + 7 O2(g) 4 CO2(g) + 6 H2O(g) (c) Ca(OH)2(s) + 2 HI(g) CaI2(s) + 2 H2O(l)

20.34 Calculate ∆S° for each of the following reactions: (a) CH3OH(l) + O2(g) CO(g) + 2 H2O(l)

(b) 4 FeCl3(s) + 3 O2(g) 2 Fe2O3(s) + 3 Cl2(g) (c) BaO(s) + H2SO4(l)H2O(l) + BaSO4(s)


20.35 Place a plus or minus sign in each empty space in the following table:

∆H (= ∆Ssurroundings) ∆S (system) Implication

______________ _________ always spontaneous

______________ _________ never spontaneous

______________ _________ spontaneous at low T

______________ _________ spontaneous at high T

20.36 How can an endothermic process be spontaneous?

20.37 What does "reversible" mean?

20.38 What do each of the following mean in relation to whether a process is spontaneous or not? (a) ∆G° > 0

(b) ∆G° < 0 (c) ∆G° = 0.

20.39 What are two equations for calculating the ∆G° of a process?


20.40 Determine the value of ∆G° for the following reaction, given that ∆H° = −120.0 kJ and ∆S° = −356 J/K:

CO2(g) + 2 NH3(g) H2O(l) + (NH2)2CO(aq)

20.41 Calculate ∆G° for each of the following reactions: (a) Al2O3(s) + 3 H2SO4(l) Al2(SO4)3(s) + 3 H2O(l)


20.42 Calculate ∆H°, ∆S°, and ∆G° at 25 °C for each of the following, using data from Appendix X. Compare the

value of ∆G° calculated directly from the Appendix values to the values calculated from ∆G° = ∆H° − T ∆S°.

(a) MgO(s) + CO2(g) MgCO3(s) (b) TiO2(s) + 2 C(s) + 2 Cl2(g) TiCl4(g) + 2 CO(g)

(c) C2H5OH(l) + 3 O2(g) 2 CO2(g) + 3 H2O(g) (d) C2H5OH(g) + 3 O2(g) 2 CO2(g) + 3 H2O(l)

(e) Sr2+(aq) + SO42–(aq) SrSO4(s)

(a) ΔH° = –117.6 kJ/mol ΔS° = –174.8 J/mol K ΔG° = –65.3 kJ/mol ΔG° = ΔH° – T ΔS° = –65.5 kJ/mol

(b) ΔH° = –39.5 kJ/mol ΔS° = 242.3 J/mol K ΔG° = –111.8 kJ/mol ΔG° = ΔH° – T ΔS° = –111.7 kJ/mol

(c) ΔH° = –1234.8 kJ/mol ΔS° = 217 J/mol K ΔG° = –1299.8 kJ/mol ΔG° = ΔH° – T ΔS° = –1299.6 kJ/mol

(d) ΔH° = –1409.3 kJ/mol ΔS° = –260. J/mol K ΔG° = –1331.8 kJ/mol ΔG° = ΔH° – T ΔS° = –1331.7 kJ/mol

(e) ΔH° = 8 kJ/mol ΔS° = 144 J/mol K ΔG° = –35 kJ/mol ΔG° = ΔH° – T ΔS° = –35 kJ/mol

20.43 (a) Predict how the ∆G° for the following reaction will change with temperature: BaCO3(s) BaO(s) +

CO2(g) (b) Assuming that ∆H° and ∆S° at 1727°C are approximately the same as at 25°C, calculate ∆G° at 727°C.

20.44 Each of the reactions below fits into one of these categories: always spontaneous, never spontaneous,

spontaneous at high temperature, or spontaneous at low temperature. Decide to which category each belongs.

(a) 4 KClO3(s) 3 KClO4(s) + KCl(s) ∆H° = −140 kJ ∆S° = −45 J/K

(b) 2 H2O2(l) 2 H2O(l) + O2(g) ∆H° = −196 kJ ∆S° = 125 J/K

(c) 3 O2(g) 2 O3(g) ∆H° = 285 kJ ∆S° = −140 J/K

(d) 2 N2O(g) + O2(g) 4 NO(g) ∆H° = 197 kJ ∆S° = 198 J/K

(e) N2(g) + 3 H2(g) 2 NH3(g) ∆H° = −92 kJ ∆S° = –199 J/K

20.45 For each of the following processes, indicate whether ∆S and ∆H are expected to be positive or negative. (a)

Dry ice sublimes (b) The temperature of water is lowered by 55°C (c) Methyl alcohol evaporates from a glass (d)

A hydrogen molecule dissociates to hydrogen atoms (e) A piece of paper is combusted forming CO2(g) and H2O(g)

20.46 Given the following information, determine the boiling point of TiCl4:

∆Hf° TiCl4(l) = −804.2 kJ/mol ∆Hf° TiCl4(g) = −763.2 kJ/mol

S° TiCl4(l) = 221.9 J/mol•K S° TiCl4(g) = 354.9 J/mol•K

20.47 Tungsten melts at 3377°C and has an enthalpy of fusion of 35.23 kJ/mol. Calculate the entropy of fusion in J

mol–1K–1.

20.48 Given the following information, determine the boiling point of iron pentacarbonyl, Fe(CO)5.

∆Hf° Fe(CO)5(l) = −774.0 kJ/mol ∆Hf° Fe(CO)5(g) = −733.9 kJ/mol

S° Fe(CO)5(l) = 338.0 J/mol•K S° Fe(CO)5(g) = 445.2 J/mol•K 374 K

20.49 Given the following information, determine the boiling point of sodium metal (Na).

∆Hf° Na(l) = 2.41 kJ/mol ∆Hf° Na(g) = 107.68 kJ/mol

S° Na(l) = 57.85 J/mol•K S° Na(g) = 153.61 J/mol•K

20.50 Given these reactions and their ∆G° values,

NH3(g) + HCl(g) NH4Cl(s) ∆G° = −21.98 kcal

COCl2(g) + 4 NH3(g) CO(NH2)2(s) + 2 NH4Cl(s) ∆G° = −79.36 kcal

COCl2(g) + H2O(l) CO2(g) + 2 HCl(g) ∆G° = −33.89 kcal

Calculate ∆G° for the reaction

CO(NH2)2(s) + H2O(l) CO2(g) + 2 NH3(g) ∆G° = 1.51 kcal

20.51 Calculate ∆G° for each of the following reactions. (a) Al2O3(s) + 3 H2SO4(l) Al2(SO4)3(s) + 3 H2O(l)



20.52 Calculate ∆G° for each of the following. (a) BaCO3(s) + H2SO4(l) BaSO4(s) + H2O(l) + CO2(g)

(b) C2H5OH(l) + 3 O2(g) 2 CO2(g) + 3 H2O(g) (c) 2 Al(s) + 3 H2SO4(aq) Al2(SO4)3(s) + 3 H2(g)


20.53 (a) What are standard conditions? (b) How do the standard conditions in thermodynamics compare to STP in

Chapter 5?

20.54 What is the equation for converting ∆G° to ∆G?

20.55 A L•atm is a quantity of energy equivalent to 101.311 J. Using this relationship, show how R can equal both

0.08206 L•atm/mol•K and 8.314 J/mol•K.

20.56 Q is the reaction quotient. How does the reaction quotient relate to the reactants and products in a reaction?

20.57 (a) Using data from the Appendix, calculate the value of ∆G° for C(s) + CO2(g) 2 CO(g). (b) Calculate ∆G

at 275 °C if 25.00 g of carbon, 2.50 atm of CO, and 1.50 atm of CO2 are placed in a 5.00 L container.

(a) 120.0 kJ/mol (b) 126.5 kJ/mol

20.58 Determine the temperature ranges over which the following reactions will be spontaneous.

(a) CaCO3(s) + H2SO4(l) CaSO4(s) + H2O(l) + CO2(g) (b) 2 Fe2O3(s)4 Fe(s) + 3 O2(g)

20.59 Estimate the free-energy change for each of the following reactions at 1200°C. Assume the variation of the

entropy values and enthalpy values with temperature are small. (a) PbO(s) + CO(g) Pb(s) + CO2(g)

(b) Si(s) + 2 MnO(s) SiO2(s) + 2 Mn(s) (c) FeO(s) + H2(g) Fe(s) + H2O(g)

20.60 Calculate ∆G at 298 K for the reaction below given the following partial pressures: CH3OH = 0.00180 atm, O2

= 0.200 atm, CO = 0.00187 atm, and H2O = 0.000375 atm. CH3OH(g) + O2(g) CO(g) + 2 H2O(g)

20.61 Calculate ∆G at 298 K for the reaction below given the following partial pressures: C2H4 = 1.00 atm, O2 =

0.200 atm, CO = 0.00125 atm, and H2O = 0.0135 atm. C2H4(g) + 2 O2(g) 2 CO(g) + 2 H2O(g)

20.62 Calculate ∆G at 298 K for the reaction below given the following partial pressures: C2H5OH = 0.00137atm,

O2 = 0.200 atm, CO2 = 0.000175 atm, and H2O = 0.00250 atm. C2H5OH(g) + 3 O2(g) 2 CO2(g) + 3 H2O(g)


20.63 (a) For a spontaneous process, is the value of the equilibrium constant large or small? (b) For a

nonspontaneous process, is the value of the equilibrium constant large or small?

20.64 What is the relationship between the change in the standard free energy and the equilibrium constant?

20.65 By using ∆G°, calculate Kp at for each of the following reactions: (a) 2 CO(g) + O2(g) 2 CO2(g)

(b) C2H4(g) + H2(g) C2H6(g) (c) C2H4(g) + 2 O2(g) 2 CO(g) + 2 H2O(g)

(a) K = 1.330 × 1090 (b) K = 5.60 × 1017 (c) K = 10140

20.66 (a) Calculate ∆G° for the following reaction: MgCO3(s) MgO(s) + CO2(g).

(b) Calculate the equilibrium partial pressure of carbon dioxide at 25°C. (c) Calculate the equilibrium partial

pressure of carbon dioxide at 1000.0°C.

20.67 (a) Using Ka from Appendix X, calculate ∆G° for the dissociation of acetic acid. (b) If the system is in

equilibrium, what is the value of ∆G? (c) Calculate the value of ∆G when pH = 2.522, acetate ion concentration is

2.00 × 10–5 M, and acetic acid concentration is 0.100 M.


20.68 Estimate ∆G° for each of the following:

(a) Ag2CrO4(s) 2 Ag+(aq) + CrO42–(aq) K = 1.9 × 10–16

(b) H2CO3(aq) H+(aq) + HCO3–(aq) K = 4.3 × 10–7

(c) PCl5(g) PCl3(g) + Cl2(g) K = 3.30

(d) N2O4(g) 2 NO2(g) K = 40.0

(e) 4 Al(s) + 3 O2(g) 2 Al2O3(s) K = 1.9 × 10276

20.69 Benzoic acid, HC7H5O2, has Ka = 6.5 × 10–5 for the equilibrium shown below. Determine the value of ∆G

when [H+] = 4.5 × 10–3 M, [C7H5O2–] = 1.5 × 10–5 M, and [HC7H5O2] = 0.15 M.

HC7H5O2(aq) H+(aq) + C7H5O2–(aq) –1.2 × 104 J/mol

20.70 Calculate the value of K at 25°C for the following reaction (a carbonylation reaction), used industrially to

produce acetic acid. CH3OH(l) + CO(g) CH3COOH(l) 3.94 × 1015

20.71 The Ka for chloroacetic acid (HC2H2ClO2) is 1.4 × 10–3 at 298 K. What is the value of ∆G when [H+] = 3.0 ×

10–3M, [C2H2ClO2–] = 2.0 × 10–5M, and [HC2H2ClO2] = 0.10 M? The equilibrium is

HC2H2ClO2(aq) H+(aq) + C2H2ClO2–(aq)


20.72 Why is ammonia such an important industrial chemical?

20.73 Industrially, the Haber process uses relatively extreme conditions. Why is the process run at high pressure

and temperature?

20.74 How do industrial chemists take advantage of Le Châtelier's principle in the production of ammonia? What

other ways could they take advantage of Le Châtelier's principle?

20.75 Determine the value of the equilibrium constant for the Haber process at the following temperatures. Assume

∆G° = −33.4 kJ in each case. (a) 300.0 K (b) 400.0 K (c) 500.0 K (d) 600.0 K (e) 700.0 K

20.76 Industrially, the Haber process operates at high pressure. Determine the value of ∆G at the following

temperatures, given that the partial pressure of nitrogen is 33.3 atm, the partial pressure of hydrogen is 100.0 atm,

and the partial pressure of ammonia is 67.7 atm. Assume ∆G° = −33.4 kJ/mol in each case. (a) 300.0 K (b) 400.0 K

(c) 500.0 K (d) 600.0 K (e) 700.0 K

20.77 Write a balanced chemical equation for the respiration of glucose (C6H12O6).

20.78 The Gibbs free energy change for the photosynthesis of glucose indicates that the process is nonspontaneous.

How can a nonspontaneous process, such as this one, take place?

20.80 Determine the enthalpy and entropy changes for the respiration of glucose and compare these values to those

for the photosynthesis of glucose.

20.81 What is ∆G° for the photosynthesis of glucose, given that ∆G° for the oxidation of glucose (respiration) in

solution is −2872 kJ/mol? C6H12O6(aq) + 6 O2(g)6 CO2(g) + 6 H2O(l)

Summary

20.82 Calculate the work done, in joules (1 J = 101.3 L•atm), when 2.00 mol of steam at 100.0°C and 1.00 atm are

condensed to liquid water at the same temperature and pressure. Assume that the steam behaves as an ideal gas, and

that the volume of the liquid water may be neglected.

20.83 Calculate the work done, in joules (1 J = 101.3 L•atm), when 1.35 g of Na reacts with water to form hydrogen

gas at 0.0°C and 1.00 atm. 2 Na(s) + 2 H2O(l)2 NaOH(aq) + H2(g)


20.84 For each of the following processes, indicate whether the signs of ∆S and ∆H are expected to be positive or

negative. (a) Dry ice sublimes (b) The temperature of a solid is lowered by 55°C (c) Methyl alcohol evaporates

from a glass (d) A hydrogen molecule dissociates to atoms (e) A piece of paper is combusted to form CO2(g) and

H2O(g)

20.85 Predict if the entropy change in each of the following will be positive, negative, or about zero.

(a) H2(g) + Br2(l) 2 HBr(g) (b) H2(g) + Br2(g) 2 HBr(l)

(c) CaO(s) + CO2(g) CaCO3(s) (d) F2(l) F2(g)

(e) Ba(OH)2(s) + 2 NaHSO4(s) 2 H2O(l) + BaSO4(s) + Na2SO4(s)

20.86 The heat of vaporization of isooctane at its boiling point (99.3°C) is 37.7 kJ/mol. What is ∆S in cal/mol•K for

the vaporization of 2.00 mol of isooctane?

20.87 Calculate ∆S° for each of the following reactions:

(a) 3 NO2(g) + H2O(l)2 HNO3(l) + NO(g) ∆G° = 1.98 kcal/mol; ∆H° = −17.14 kcal/mol

(b) OF2(g) + H2O(g)O2(g) + 2 HF(g) ∆G° = −358.4 kJ/mol; ∆H° = −323 kJ/mol

20.88 Elemental sulfur exists in several crystalline forms. At 1 atm, two of the forms are rhombic sulfur and

monoclinic sulfur. Using the following information, determine the temperature for the conversion of monoclinic

sulfur to rhombic sulfur.

∆Hf° (kJ/mol) S° (J/mol•K)

S (rhombic) 0.00 31.88

S (monoclinic) 0.30 32.55

20.89 The enthalpy change for the photosynthesis of glucose is 2802.5 kJ/mol. (a) How much energy is necessary to

form one molecule of glucose? (b) If the energy in part (a) comes from a single photon of light, what must be the

wavelength, in nanometers, of the photon?

20.90 Chloroform (CHCl3) melts at −64°C and boils at 62°C. The enthalpy of fusion of chloroform is 8.798 kJ/mol,

and its enthalpy of vaporization is 31.38 kJ/mol. Estimate the entropies of fusion and vaporization for chloroform.

ΔSfus = 42.1 J/mol K ΔSvap = 93.7 J/mol K

20.91 Isopropyl alcohol (C3H7OH) melts at −90.0°C and boils at 82°C. The enthalpy of fusion of isopropyl alcohol

is 5.36 kJ/mol, and its enthalpy of vaporization is 42.11 kJ/mol. Estimate the entropies of fusion and vaporization

for isopropyl alcohol.

20.92 (a) Determine the Gibbs free energy change for the combustion of ethanol (C2H5OH), shown below. (b) A

bottle of wine is 12.0 % ethanol (alcohol) by mass. The density of the wine is 1.090 g/mL. Determine the total

Gibbs free energy change for the combustion of the ethanol in a 1.00 quart bottle of this wine (1.00 L = 1.057 quart).

C2H5OH(l) + 3 O2(g) 2 CO2(g) + 3 H2O(l)

20.93 At one time manganese metal was produced by the reaction shown below. Determine ∆G° for this reaction

using data from Appendix C. 4 Al(s) + 3 MnO2(s) 2 Al2O3(s) + 3 Mn(s)

20.94 Use VSEPR to predict the geometries of each of the following. (a) N2O (Hint: the atoms are arranged NNO)

(b) SiCl4 (c) SF4 (d) PF5 (e) IF7

20.95 Draw Lewis structures and predict geometries for each of the following.

(a) BrF5 (b) ClF3 (c) KrF2

20.96 Assign oxidation numbers to the indicated elements in each of the following:

(a) Chlorine in NaClO3 (b) Molybdenum in Mo6Cl186– (c) Barium in BaO2

(d) Hydrogen in LiAlH4 (e) Uranium in ZnUO2(C2H3O2)4


20.97 Circle the more covalent compound in each of the following pairs.

(a) Magnesium chloride or beryllium chloride (b) Tin(IV) chloride or Tin(II) chloride

(c) Aluminum chloride or rubidium chloride (d) Sodium chloride or sodium fluoride

(e) Potassium chloride or bromine trifluoride

20.98 What is the expected electron configuration for each of the following ions? (a) Co3+ (b) Cr2+ (c) Ru4+

20.99 The major source of magnesium is seawater. The first step in recovering magnesium from seawater is to

precipitate magnesium hydroxide using calcium oxide (see equation below). How many kilograms of calcium oxide

are required to precipitate 1.0 × 109 g of magnesium hydroxide?

Mg2+(aq) + CaO(s) + H2O(l) Mg(OH)2(s) + Ca2+(aq)

20.100 Name or give the formula for each of the following.

(a) Tin(II) fluoride (f) Tl2O

(b) Barium peroxide (g) Ga2S3

(c) Indium sulfide (h) Be3N2

(d) Lithium hydride (i) LiF

(e) Potassium superoxide (j) Rb2O

20.101 Name or give the formula for each of the following. (a) KO2 (b) Barium acetate (c) AlF3 (d) SnCl2

(e) Mercury(I) iodide

20.102 Write balanced chemical equations for each of the following:

(a) When mercury(II) oxide is heated, it decomposes to form oxygen gas and mercury metal.

(b) When copper(II) nitrate is heated strongly, it decomposes to copper(II) oxide,

nitrogen dioxide, and oxygen.

(c) Lead(II) sulfide reacts with ozone to form lead(II) sulfate and oxygen gas.

(d) Sodium metal reacts with hydrogen gas to form sodium hydride.

(e) Potassium peroxide reacts with carbon dioxide to form potassium carbonate and oxygen gas.

20.103 Write chemical equations for the reaction of oxygen with each of the following.

(a) Cadmium (b) Aluminum (c) Calcium (d) Iron (e) Copper


occurs, write NR. If there is a reaction, write a net ionic equation. (a) NaOH(aq) + Mg(NO3)2(aq) (b) Ca(OH)2(aq)

+ NH4Cl(aq) (c) BaCl2(aq) + (NH4)3PO4(aq) (d) KNO3(aq) + SrCl2(aq) (e) Pb(C2H3O2)2(aq) + HI(aq)

20.105 Write balanced equations for each of the following reactions. (a) Potassium plus water (b) Aluminum nitrate

decomposes to aluminum oxide, nitrogen dioxide, and oxygen gas.

20.106 Write balanced chemical equations for each of the following. Complete the reactions where necessary.

(a) sodium plus water (b) calcium carbonate plus heat (c) magnesium plus oxygen (d) lithium plus hydrogen (e)

aluminum plus aqueous sodium nitrate (basic) produces ammonia plus NaAlO2

20.107 Complete and balance the following equations. (a) Sn(s) + HF(aq)

(b) Tl3+(aq) + OH–(aq) (c) Fe2O3(s) + Al(s)

(d) Pb3O4(s) + C(s) + heat (e) Al(OH)3(s) + NaOH(aq)

20.108 Name or give the formula for each of the following. (a) deuterium (b) sodium borohydride (c) diborane

(d) acetylide ion (e) buckminsterfullerene (f) KO2 (g) Na3BO3 (h) Na2B4O7 (i) SiC (j) H2SiF6 (k) protium (l) boron

oxide (m) hexafluorosilicate ion (n) graphite (o) CsO2

20.109 Name or give the formula for each of the following. (a) rubidium superoxide (b) orthoborate ion

(c) tetraborate ion (d) methane (e) sodium hexafluorosilicate (f) T (g) Ca(BH4)2 (h) BN (i) C(dia) (j) CaC2 (k) C60

(l) barium peroxide (m) D (n) potassium tetraborate (o) potassium orthoborate


20.110 Name or give the formula for each of the following. (a) ammonium chloride (b) ammonium sulfate

(c) barium peroxide (d) calcium tetraborate (e) phosphorous acid (f) T (g) H3PO4 (h) (NH4)2SO4 (i) (C2H5)2NH2+

(j) H2SiF6 (k) chloramine (l) hydroxylamine (m) methylammonium ion (n) N2H3CH3 (o) Na5P3O10

20.111 Name or give the formula for each of the following. (a) calcium orthoborate (b) trimetaphosphoric acid

(c) hydrazine (d) sodium hexafluorosilicate (e) ammonium pyrophosphate (f) Na2O2 (g) NH2Cl (h) NH2OH

(i) CH3NH2 (j) P2O74– (k) CH3NH3Cl (l) NH3 (m) diamond (n) K2SiF6 (o) methylamine

20.112 Name or give the formula for each of the following. (a) pyrosulfuric acid (b) paraperiodic acid

(c) tetraborate ion (d) iron(II) disulfide (e) ammonium sulfate (f) O3 (g) KI3 (h) NaBH4 (i) H6TeO6 (j) P2O74–

(k) XeF6 (l) orthotelluric acid (m) H3BO3 (n) potassium tetraborate (o) metaperiodic acid

20.113 Name or give the formula for each of the following. (a) xenon tetrafluoride (b) ammonium pyrophosphate

(c) selenous acid (d) potassium borohydride (e) calcium triiodide (f) KrF2 (g) H5IO6 (h) CH3NH2 (i) Na2B4O7 (j) S22–

(k) H2SeO4 (l) H2S2O7 (m) ozone

20.114 Both hydrogen (H2) and methane (CH4) are useful fuels. (a) Write balanced chemical equations for the

combustion of each of these gases to produce only gaseous products. (b) Calculate the heat of reaction for each of

your reactions in part (a). (c) Compare the amount of heat produced by the combustion of 10.0 g of hydrogen to the

amount of heat produced by the combustion of 10.0 g of methane. (d) Compare the amount of heat produced by the

combustion of 10.0 L of hydrogen gas at 25°C and 760.0 torr to the amount of heat produced by the combustion of

10.0 L of methane at the same temperature and pressure.

20.115 A magnesium boride reacts with acid to form boric acid and a boron hydride. The boron hydride is 84.3

percent boron and 15.7 percent hydrogen. The density of a sample of this gas is 2.55 g/L at 27°C and 745 torr.

What is the molecular formula for the boron hydride? B5H10


21.1 What is the definition of electrochemistry?

21.2 Define oxidation and reduction.

21.3 Define an oxidizing agent and a reducing agent.

21.4 (a) What happens to the oxidizing agent during a redox reaction? (b) What happens to the reducing agent

during a redox reaction?

21.5 The definitions of oxidation and reduction depend on electrons. (a) What happens to the oxidation number of

an element that undergoes oxidation? (b) What happens to the oxidation number of an element that undergoes

reduction?


21.6 What must the oxidation and reduction half-reactions have in common?

21.7 List, in order, the steps for balancing oxidation-reduction equations.

21.8 (a) How should you balance oxygen atoms in a redox reaction in an acid solution? (b) How should you balance

oxygen atoms in a redox reaction in a basic solution?

21.9 (a) How should you balance hydrogen atoms in a redox reaction in an acid solution? (b) How should you

balance hydrogen atoms in a redox reaction in a basic solution?


21.10 (a) On which side of the reaction arrow are the electrons in an oxidation half-reaction? (b) On which side of

the reaction arrow are the electrons in a reduction half-reaction?

21.11 Complete the following redox couples (half-reactions).

(a) 2 e– + Cl2 _____ (b) 3 e– + CrO42– + 4 H2O _____ + 8 OH–

(c) _____ + FeO42– + 8 H+

Fe3+ + 4 H2O (d) Pb4+ + _____ Pb2+

(e) S22– 2 S + _____

21.12 Add electrons to each of the following half-reactions

(a) Cu(s) Cu2+(aq) (b) NO3– + 4 H+(aq) NO(g) + 2 H2O(l)

(c) 2 I–(aq) I2(s) (d) SO42–(aq) + 4 H+(aq) SO2(g) + 2 H2O (l)

(e) O2(g) 2 O2–(s)

21.13 Add electrons to each of the following half-reactions

(a) 14 H+(aq) + Cr2O72–(aq) 2 Cr3+(aq) + 7 H2O(l) (b) MnO4

–(aq) + 8 H+(aq) Mn2+(aq) + 4 H2O(l)

(c) MnO4–(aq) + 2 H2O(l) MnO2(s) + 4 OH–(aq) (d) C2H5OH(aq) + H2O(l) HC2H3O2(aq) + 4 H+(aq)

(e) N3–(aq) + 6 OH–(aq) 3 NO(g) + 3 H2O(l)

21.14 Assign oxidation numbers to each element in the following equations. Label those that are oxidation-

reduction reactions.

(a) 2 KOH(aq) + K2Cr2O7(aq) 2 K2CrO4(aq) + H2O(l)

(b) 3 Cu(s) + 8 HNO3(aq) 3 Cu(NO3)2(aq) + 2 NO(g) + 4 H2O(l)

(c) AlBr3(s) + 3 H2O(l) Al(OH)3(s) + 3 HBr(aq)

(d) 2 H2SO4(aq) + 2 KI(s) I2(s) + SO2(g) + K2SO4(aq) + 2 H2O(l)

(e) 2 Mg(s) + O2(g) 2 MgO(s)

21.15 Solid copper(II) nitrate decomposes when heated, to produce solid copper(II) oxide, nitrogen dioxide gas, and

oxygen gas. (a) Write a balanced chemical equation for this reaction. (b) Identify which element undergoes

oxidation, which element undergoes reduction, and which elements undergo neither oxidation nor reduction.

21.16 Balance each of the following half-reactions. Identify each half-reaction as an oxidation or as a reduction

half-reaction.

(a) Br2(aq) BrO3–(aq) (acidic solution) (b) Cr2O7

2–(aq) Cr3+(aq) (acidic solution)

(c) MnO4–(aq) Mn2+(aq) (acidic solution) (d) MnO4

–(aq) MnO2(s) (basic solution)

(e) Al(s) Al(OH)4–(s) (basic solution)

21.17 Balance each of the following half-reactions. Identify each half-reaction as an oxidation or as a reduction

half-reaction.

(a) IO3–(aq) I2(s) (acidic solution) (b) H2O2(aq) O2(g) (acidic solution)

(c) NO3–(aq) NH3(aq) (basic solution) (d) CrO4

2–(aq) CrO2–(aq) (basic solution)

(e) NH2OH(aq) N2(g) (basic solution)

21.18 Balance the following redox equations:

(a) (acid) Fe2+(aq) + Cr2O72–(aq) Fe3+(aq) + Cr3+(aq)

(b) (acid) C2O42–(aq) + MnO4

–(aq)CO2(g) + Mn2+(aq)

(c) (base) As2S3(s) + H2O2(aq) AsO43–(aq) + SO4

2–(aq)

(d) (base) Br2(l) BrO3–(aq) + Br–(aq)

(e) (acid) BrO3–(aq) + ClO3

–(aq) ClO4–(aq) + HBrO(aq)

(f) (base) OCN–(aq) + Cl–(aq) CN–(aq) + ClO2(g)

21.19 Balance the following redox equations:

(a) (acid) Br2(l) + BiO+(aq) Bi(s) + HBrO(aq) (b) (base) IO3–(aq) + I–(aq) I2(s)

(c) (acid) IO3–(aq) + Mo(s) MoO4

2–(aq) + I2(s) (d) (base) AsO43–(aq)+Be(s) AsO2

–(aq)+Be2O32–(aq)

(e) (acid) H2N2O2(aq) + W(s) N2(g) + WO42–(aq) (f) (base) H3IO6

2–(aq) + In(s) I–(aq) + In(OH)3(s)


21.20 Write a balanced redox equation for the reaction in each of the following descriptions: (a) An aqueous

potassium iodide is oxidized to I2(s) by aqueous mercury(I) nitrate. The mercury reduces to the metal. (b)

Copper(I) nitrate is oxidized to copper(II) nitrate by nitric acid. Nitrate ion reduces to nitrogen dioxide. (c) A basic

sodium hypochlorite solution oxidizes solid chromium(III) hydroxide to sodium chromate. The sodium

hypochlorite reduces to sodium chloride.

21.21 A redox reaction in which one element is both oxidized and reduced is called a disproportionation reaction.

Balance the following disproportionation reactions:

(a) Cl2(aq) Cl–(aq) + ClO–(aq) (basic solution)

(b) MnO42–(aq) MnO4

–(aq) + MnO2(s) (acidic solution)

(c) H2SO3(aq) S(s) + HSO4–(aq) (acidic solution)

(d) Br2(aq) BrO3–(aq) + Br–(aq) (basic solution)

(a) 2 OH–(aq) + Cl2(aq) Cl–(aq) + ClO–(aq) + H2O(l)

(b) 4 H+(aq) + 3 MnO42–(aq) 2 MnO4

–(aq) + MnO2(s) + 2 H2O(l)

(c) 3 H2SO3(aq) S(s) + 2 HSO4–(aq) + 2 H+(aq) + H2O(l)

(d) 6 OH–(aq) + 3 Br2(aq) BrO3–(aq) + 5 Br–(aq) + 3 H2O(l)

21.22 Complete and balance the following equations:

(a) (acid) NbO3+ + S2O32– Nb3+ + H2SO3 (b) (base) Bi + AsO4

3– AsO2– + Bi2O3

(c) (base) ClO2– Cl– + ClO3

–

21.23 In basic solution, zinc metal oxidizes to zincate (ZnO22–). Write a balanced equation for each of the following

reactions of zinc in basic solution. (a) The reaction with water produces hydrogen. (b) The reaction with nitrate ion

produces ammonia (c) The reaction with SnO32– produces metallic tin.


21.24 (a) What is a galvanic cell? (b) How does a galvanic cell relate to a voltaic cell? (c) Give an example of a

galvanic cell.

21.25 What is an electrolytic cell?

21.26 What is electrolysis?

21.27 (a) At which electrode does oxidation occur? (b) At which electrode does reduction occur?

21.28 Identify the following aspects of the cell shown below. (a) anode (b) cathode (c) salt bridge (d) compartment

where reduction occurs (e) compartment where oxidation occurs (f) direction of electron flow (g) direction of cation

flow from the salt bridge (h) direction of anion flow from the salt bridge

21.29 Sketch a simple galvanic cell and indicate the anode, cathode, salt-bridge, and direction of electron flow.

Show the signs of the electrodes.


21.30 Answer the following questions with respect to the cell shown below. (a) Which electrode is the anode?

(b) Which electrode is the cathode? (c) Which electrode is losing mass? (d) Which electrode is gaining mass?

(e) Why would potassium chloride (KCl) not be an appropriate electrolyte for the salt bridge?

V

Zn Ag

Zn2+ Ag+

21.32 What conditions are required for a standard cell?

21.33 (a) Sketch a standard hydrogen electrode, SHE (b) Show the shorthand notation for an SHE operating as an

anode (c) Show the shorthand notation for an SHE operating as a cathode

21.34 A half-cell has a standard reduction potential of 1.05 V. What is the standard oxidation potential?

21.35 Provide the standard cell potential for the galvanic cell described by the following reactions.

Oxidation Al(s) Al3+(aq) + 3 e– E° = 1.66 V

Reduction 3 (1 e– + Ag+(aq) Ag(s)) E° = 0.80 V

Overall 3 Ag+(aq) + Al(s) Al3+(aq) + 3 Ag(s) E° = ?

21.36 The standard reduction potential does not change when a half-reaction is multiplied by a number. Is the

standard reduction potential an intensive property or an extensive property?

21.37 A voltaic cell is constructed that has one electrode of silver metal in an AgNO3 solution and the other

electrode of cadmium metal in a Cd(NO3)2 solution. The salt bridge contains a KNO3 solution. The cell reaction is

2 Ag+(aq) + Cd(s) 2 Ag(s) + Cd2+(aq)

(a) Sketch this cell. (b) Write balanced half-reactions for each electrode (c) Identify the anode and cathode in your

sketch. (d) Indicate the sign of each electrode (e) Indicate the direction the electrons flow. (f) Indicate the direction

of flow for the anions and cations in the salt bridge

21.38 (a) Write the half-reactions and the overall cell reaction for the following cell: Co|Co2+||Ag+|Ag (b) Sketch the

cell in part (a). (c) Write the half-reactions and the overall cell reaction for the following cell: Ni|Ni2+||H+|H2,Pt

(d) Sketch the cell in part (c). (e) Write the half-reactions and the overall cell reaction for the following cell: Pt,

H2|H+||Fe3+, Fe2+|Pt (f) Sketch the cell in part (e).

21.39 In each case, give a reasonable cell notation for cells employing the following half-reactions:

(a) Cd(s) + Cl2(g) Cd2+(aq)+ 2 Cl–(aq) (b) Cl2(g)+ 2 I–(aq) 2 Cl–(aq)+ I2(s)

(c) Zn(s) + Br2(l) Zn2+(aq)+ 2 Br–(aq) (d) Br2(l)+ 2 I–(aq) 2 Br–(aq)+ I2(s)

21.40 (a) Arrange each of the following in order of increasing oxidizing strength under acidic conditions: Fe2+,

HAsO2, VO2+, Cr2O7

2–, and Nb2O5. (b) Arrange each of the following in order of increasing reducing strength under

acidic conditions: Fe2+, Cr, NO2–, Mn2+, and Pb2+.

21.41 A voltaic cell is constructed with the cell reaction shown below. The standard cell potential is 1.26 V.

(a) Calculate the standard reduction potential for the half-reaction In3+(aq) + 2e– In+(aq). (b) Sketch the cell.

3 In3+(aq) + 2 Al(s) 3 In+(aq) + 2 Al3+(aq) E° = –0.40 V

21.42 Calculate the standard cell potentials for (a) Fe(s) + Cd2+(aq) Fe2+(aq) + Cd(s)

(b) 2 Cr(s) + 3 Cu2+(aq) 2 Cr3+(aq) + 3 Cu(s)



21.43 (a) What is a faraday? (b) How many electrons are in a faraday?

21.44 (a) What is the SI base unit for electric current? (b) What is the relationship between a coulomb and the SI

base unit for electric current?

21.45 Write the mole ratio required to convert from electrons to the underlined substance in each of the following

half-reactions.

(a) Cu(s) Cu2+(aq) + 2 e– (b) NO3– + 4 H+(aq) + 3 e– NO(g) + 2 H2O(l) (c) 2 I–(aq) I2(s) + 2 e–

(d) SO42–(aq) + 4 H+(aq) + 2 e– SO2(g) + 2 H2O(l) (e) 4 e– + O2(g) 2 O2–(s)

21.46 (a) How many grams of fluorine gas can be produced by the electrolysis of a solution of potassium fluoride in

hydrogen fluoride if a current of 15.00 A is passed through the solution for 4.000 h? Assume the cell is 75 percent

efficient. (b) What is the energy requirement for this electrolysis per mole of fluorine formed if the applied standard

cell potential is 4.50 V? (a) 32 g F2 (b) ΔG = 8.68 × 105 J/mol

21.48 A sample containing silver is electrolyzed. A current of 5.00 A produces 5.25 g of silver metal at the cathode

in 31.3 min. What is the oxidation state of silver in the sample?

21.49 The passage of a current of 1.500 A for 50.00 min deposited 1.476 g of copper from a Cu(NO3)2 solution.

From this information, calculate the atomic mass of copper.


21.51 Why will the equation E°cell = (0.0592

𝑛) log K not allow the calculation of K for the following reaction?

2 KOH(aq) + K2Cr2O7(aq) 2 K2CrO4(aq) + H2O(l)

21.52 Complete the following table: ∆G E_____

Spontaneous _______ ______

Nonspontaneous _______ ______

Equilibrium ______ ______

21.53 How does the Q in the Nernst equation relate to the reactants and products in the reaction?

21.54 Calculate the standard cell potentials and the equilibrium constant for each of the following reactions:

(a) 2 VO2+(aq) + 4 H+(aq) + Cu(s) 2 VO2+(aq) + 2 H2O(l) + Cu2+(aq) (a) K = 2.48 × 1022

(b) 3 Ce4+(aq) + As(s) + 2 H2O(l) 3 Ce3+(aq) + HAsO2(aq) + 3 H+(aq) (b) K = 4.2076 × 1060

(c) 2 BrO–(aq) + ClO–(aq) + H2O(l) Br2(aq) + 2 OH–(aq) + ClO2–(aq) (c) K = 8.0 × 10–8

21.55 Calculate G° and K for each of the following using standard electrode potentials:

(a) Fe(s) + Cd2+(aq) Fe2+(aq) + Cd(s)

(b) 2 Cr(s) + 3 Cu2+(aq) 2 Cr3+(aq) + 3 Cu(s)

21.56 (a) What is the minimum amount of energy required to operate the cell shown below (i.e., calculate ∆G° for

the cell)? (b) How much energy is required to produce 0.500 mol of zinc?

Pt|Fe3+(1 M), Fe2+(1 M)||Zn2+(1 M)|Zn

21.57 A half-cell with a silver electrode is attached to an SHE. The silver half-cell contains some silver ion and

solid silver bromide and is 1.00 M in bromide ion. If the cell potential of the cell is 0.072 V, what is the Ksp for

silver bromide? The Ksp equilibrium reaction is AgBr(s) Ag+(aq) +Br–(aq)

21.58 A cell constructed with a standard hydrogen electrode (as the anode) and a silver wire immersed in a saturated

solution of silver(I) oxalate has a potential of 0.589 V. Calculate the Ksp for silver(I) oxalate The equilibrium

reaction is Ag2C2O4(s) 2 Ag+(aq) + C2O42–(aq)


21.59 A voltaic cell consisting of two half-cells contains a zinc electrode in a 1.0 M S2– solution (from buffered

Na2S) saturated with ZnS and a copper electrode in a 1.0 M Cu(NO3)2 solution. The observed voltage is +1.78 V.

Calculate K for the following equilibrium: ZnS(s) Zn2+(aq)+ S2–(aq) Ksp = 1.0 10–23

21.60 Use standard electrode potentials to calculate ∆G° and K for the disproportionation of copper(I) ion:

2 Cu+(aq) Cu(s) + Cu2+(aq) K = 1.66 × 106 G° = –3.55 × 104 J/mol

21.61 Given the following half-reactions

3 e– + AuBr4–(aq) Au(s) + 4 Br–(aq) E° = +0.8700 V

3 e– + Au3+(aq) Au(s) E° = +1.498 V

Calculate the instability constant, K, for the AuBr4– ion. The instability constant equilibrium is

AuBr4–(aq) Au3+(aq) + 4 Br–(aq) K = 1.42 × 10–32


21.62 Write the Nernst equation.

21.63 (a) What value in the Nernst equation will change if the temperature is not 298 K? (b) What value in the

Nernst equation will change if the concentration is non-standard?

21.64 The following reaction occurs in a voltaic cell: Al(s) + Cr3+(aq) Al3+(aq) + Cr(s)

(a) Calculate the standard cell potential. (b) Calculate the cell potential when [Cr3+] = 0.20 M and [Al3+] = 1.5 M.

Assume the temperature is 25°(C)

21.65 The following reaction occurs in a voltaic cell:

6 Fe2+(aq) + Cr2O72–(aq) + 14 H+(aq) 6 Fe3+(aq) + 2 Cr3+(aq) + 7 H2O(l)

(a) Calculate the standard cell potential. (b) Calculate the cell potential when [Fe2+] = 0.20 M, [Cr2O72–] = 0.60 M,

[H+] = 1.0 × 10–4 M, [Cr3+] = 1.00 M, and [Fe3+] = 1.5 M. Assume the temperature is 25°(C).

21.66 Determine the cell potentials for each of the following:

(a) Ag|Ag+(0.0010000 M) || H+(0.2000 M)|Nb2O5, Nb

(b) Tl|Tl+(1.0745 M) || Sn2+(2.000 × 10–7 M)|Sn

(c) Ag|Ag+(0.001200 M) || VO2+(0.02002 M), VO2+(0.01000 M), H+(0.001000 M)|Pt

(d) Au|AuCl4–(0.020 M), Cl–(0.010 M) || Fe3+(0.10 M), Fe2+(0.27)|Pt

(a) E = –3.3742 V (b) E = –0.0001 V (c) E = 0.0000 V (d) E = –0.37 V

21.67 Given the following cell, PbO2(s) + Pb(s) + 4 H+(aq) + 2 SO42–(aq) 2 PbSO4(s) + 2 H2O(l)

(a) Calculate E°, G°, and K for the cell under standard conditions. (b) Calculate E and G if [H+] = 0.010 M,

[SO42–] = 0.0010 M, and T = 25°C. (c) If the sulfate ion concentration were 0.010 M, what would be the hydrogen

ion concentration when E = 0.000 V?

21.68 A cell is constructed using a Zn/Zn2+ and a Cu/Cu2+ half-cell with a sodium nitrate salt-bridge. (a) Sketch the

cell. (b) Identify the anode and the cathode (c) Which electrode gains mass, and which electrode loses mass as the

cell operates? (d) What is the overall cell reaction? (e) Calculate the standard cell potential. (f) Determine the cell

potential at 25° if [Zn2+] = 0.35 M and [Cu2+] = 0.25 M.


21.69 Why is it not possible to produce certain metals from the electrolysis of aqueous solutions of their salts?

21.70 Why does the electrolysis of aqueous sodium chloride give different results than the electrolysis of molten

sodium chloride?

21.71 What half-reactions must we always consider in the electrolysis of any aqueous solution?


21.72 The industrial preparation of magnesium uses seawater as the magnesium source. Why does the isolation of

magnesium not employ the electrolysis of seawater?

21.73 Most of the gold in seawater is present as the AuCl4– ion. The reduction of this ion to gold metal has a

standard reduction potential of +0.99 V. Why is the electrolysis of seawater not a commercially important source of

gold?

21.74 (a) Write the half-reactions and the balanced cell reaction for the electrolysis of molten lithium fluoride

(b) Assuming the cell potentials in molten lithium fluoride are the same as in aqueous solution, what is the potential

required for the electrolysis of lithium fluoride? (c) Write the half-reactions for the electrolysis of aqueous lithium

fluoride. (d) What is the cell potential required for the electrolysis of aqueous lithium fluoride?

21.75 Metallic magnesium can be made by the electrolysis of molten MgCl2. (a) What mass of Mg is formed by

passing a current of 3.50 A through molten MgCl2 for a period of 550 min? (b) How many seconds are required to

plate out 2.00 g Mg from molten MgCl2 using 5.00 A current?

21.76 (a) How many grams of lanthanum (La) metal may be produced by the electrolysis of molten lanthanum

chloride (LaCl3) if a current of 3.00 A is passed through the compound for 24 h? (b) What amperage is required to

produce 10.85 g of La from molten LaCl3 in 12.0 h?

21.77 The electrolysis of aqueous potassium chloride produces hydrogen gas, chlorine gas, and potassium

hydroxide. (a) Using half-reactions, write a balanced chemical equation for the electrolysis of aqueous potassium

chloride (b) How many liters of chlorine gas at 25° and 1.00 atm are produced in the electrolysis of the aqueous

potassium chloride solution by a current of 10.0 A for 2.00 h? (c) How many moles of potassium hydroxide would

be formed in part (b)?(a) 2 H2O(l) + 2 KCl(aq) Cl2(g) + H2(g) + 2 KOH(aq) (b) V = 9.12 L (c) 0.746 mol KOH

21.78 Write balanced chemical equations for the reaction that occurs when each of the following solutions is

electrolyzed: (a) An aqueous solution of lithium iodate (b) An aqueous solution of iron(II) nitrite


21.79 Define the three types of batteries (primary, secondary, and fuel cells).

21.81 How does an alkaline battery overcome one of the problems inherent in using a dry cell battery?

21.82 (a) Using the standard cell potentials in Appendix X, predict what combination of half-reactions would give

the greatest cell potential. (b) Why is this combination not used commercially?

21.83 The reaction shown below takes place in a lead storage battery (automobile battery). The battery, producing

48.2 A, takes 10.0 s to start a car. How many moles of electrons has the battery produced?

2 PbSO4(s) + 2 H2O(l) Pb(s) + PbO2(s) + 4 H+(aq) + 2 SO42–(aq)

21.84 The following half-reactions are important to the operation of a nickel-cadmium (NiCad) rechargeable

"battery." Cd(OH)2(s) + 2e– Cd(s) + 2 OH–(aq) E° = −0.76V

NiO2(s) + 2 H2O(l) + 2e– Ni(OH)2(s) + 2 OH–(aq) E° = + 0.49 V

(a) Write a balanced redox equation for the discharging of a NiCad battery. (b) Write a balanced redox equation for

the recharging of a NiCad battery. (c) Determine the standard cell potential for the discharge of a NiCad battery.

(d) Calculate the G° for discharging a NiCad battery. (e) What are the values of G° and of E when a NiCad

battery entirely discharges?

Corresponds to BLBMWS 20.8

21.85 List the macroscopic observations that characterize the rusting of iron.


21.86 What is the best way of representing the formula of the iron(III) hydroxide in rust?

21.87 Locate iron on the Activity Series given in Chapter 4. Which elements in this Activity Series could replace

zinc as a protective coating on steel?

21.88 (a) Why does aluminum not corrode like iron? (b) How does chromium in stainless steel behave like

aluminum?

21.89 Why is the electrochemical method of removing silver tarnish better than using silver polish?

21.90 If an iron object is nickel-plated, does the nickel protect the iron by cathodic protection? Explain.

Summary

21.91 The manganese in a 0.7500-g sample was converted to permanganate ion. The resultant permanganate

solution was titrated with oxalic acid and yielded manganese(II) ion and carbon dioxide gas. The titration required

28.80 mL of 0.1500 M oxalic acid. Calculate the percent manganese in the sample.

21.92 What volume of a 0.1500 M iron(III) chloride solution is required to react completely with 37.45 mL of a

0.01275 M tin(II) chloride solution. The reaction is

2 Fe3+ + Sn2+ 2 Fe2+ + Sn4+

21.93 How many milliliters of 0.1450 M KMnO4 solution are needed to react completely with 50.00 mL of 0.1235

M SnCl2 solution, given the chemical equation for the reaction?

16 H+ + 2 MnO4– + 5 Sn2+ 5 Sn4+ + 2 Mn2+ + 8 H2O

21.94 If the following reaction is to be run using 2.500 L of 0.9000 M sulfuric acid, how many grams of sulfuric

acid must be weighed out?

K2Cr2O7 + 6 FeSO4 + 7 H2SO4 K2SO4 + Cr2(SO4)3 + 3 Fe2(SO4)3 + 7 H2O

21.95 How many grams of potassium iodate are required to react with 14.25 g of sodium sulfite in the following

reaction? KIO3 + 3 Na2SO3 3 Na2SO4 + KI

21.96 Answer the following questions with respect to the cell shown below. (a) What is the standard cell potential

for this cell? (b) If the zinc ion concentration is 0.500 M and the silver ion concentration is 0.250 M, what is the cell

potential? (c) What is the equilibrium constant for the standard cell? (d) Replacing the sodium nitrate in the salt

bridge with sodium chloride causes a drastic change in the cell potential of the standard cell. Why?

21.97 A voltaic cell consisting of two half-cells contains a copper electrode in a 1.00 M copper(II) nitrate solution

and a silver electrode in a solution saturated with silver chromate. Calculate the voltage of this cell.

21.98 For years, sodium bismuthate (NaBiO3) has been used to test for the presence of Mn2+ in solution. Balance

the following unbalanced equation for the reaction: Mn2+(aq) + NaBiO3(s) MnO4–(aq) + BiO3

3–(aq)


21.99 Draw diagrams for the galvanic cells utilizing the following half-cells. Label all of the parts.

Cathode Anode

(a) Cl2/Cl– I2/I–

(b) Co3+/Co2+ Ni/Ni3+

21.100 Sketch the cells for each of the following overall reactions. Indicate the direction of electron flow and the

direction of flow of ions from the salt bridge, and identify the cathode and anode. Give the overall balanced

equation. Assume all concentrations are 1.0 M.

(a) Cr3+ + Br2 Cr2O72– + Br– (b) Ag+ + Mg Mg2+ + Ag

21.101 Under standard conditions, what is the maximum amount of work, in joules, that a cell employing the cell

reaction shown below can accomplish if 0.453 mol of chromium is consumed? (Remember: The Gibbs free energy

is equal to the maximum amount of work.) 3 Ag+(aq) + Cr(s) 3 Ag(s) + Cr3+(aq)

21.102 How long will it take to reduce all of the aluminum to the metal by electrolysis of 100.0 g of molten

aluminum fluoride using a current of 2.25 A?

21.103 How long would electrolysis of aqueous gold(III) sulfate have to continue to produce 5.00 g of gold at the

cathode at a current of 0.150 A? 4.90 × 104 s

21.104 Why is it not possible to produce elemental fluorine from an aqueous solution of potassium fluoride?

21.105 In the cell described below, the observed voltage is 0.875 V at 25°C. Determine the pH of the solution.

Pt, H2(1.00 atm)|H+(?)||Ag+(0.10 M)|Ag pH = 2.28

21.106 The measured voltage for the following cell is 1.0277 V at 25°C. Calculate the pH of the solution.

Pt,H2(1.00 atm)|H+ ||Ag+(0.100 M)|Ag

21.107 Explain why each of the following would not be a useful electrolyte in a salt bridge. (a) C2H5OH

(b) Ca3(PO4)2 (c) CH3Cl (d) H2S (e) AgCl

21.108 What oxidation states are observed for each of the following?

(a) lithium (b) aluminum (c) tin (d) strontium

21.109 Write chemical equations for each of the following reactions. (a) Sodium hydride reacts with water

(b) Sodium peroxide reacts with water (c) Dinitrogen pentoxide reacts with water

(d) Lithium nitride reacts with water (e) Sodium cyanide reacts with hydrochloric acid

21.110 Determine the number of protons, neutrons, and electrons present in the ions or molecules produced in each

of the following half-reactions. (a) 65Cu(s) 65Cu2+(aq) + 2 e–

(b) 37Cl2(g) + 2 e– 2 37Cl–(aq) (c) 56Fe3+(aq) + 1 e– 56Fe2+(aq)

(d) 210At–(aq) 210At(s) + e– (e) 235U3+(aq) 235U5+(aq) + 2 e–


completion and assume that all reactants are in aqueous solution. (a) Hydrochloric acid plus sodium hydroxide

(b) Ammonia plus nitric acid (c) Oxalic acid plus potassium hydroxide

(d) Ammonium chloride plus lithium hydroxide (e) Phosphoric acid plus strontium hydroxide

21.112 Give the oxidation number of the underlined element in each of the following.

(a) HOF (b) XeO2F2 (c) SF4 (d) ClO2– (e) I3

–


occurs, write “NR.” If there is a reaction, write a net ionic equation. (a) NaOH(aq) + Mg(NO3)2(aq) (b)

Ca(OH)2(aq) + NH4Cl(aq) (c) BaCl2(aq) + (NH4)3PO4(aq) (d) KNO3(aq) + SrCl2(aq) (e) Pb(C2H3O2)2(aq) + HI(aq)


21.114 Draw Lewis structures and predict the molecular geometry for each of the following. (a) I3– (b) BrF4

–

(c) ClO2– (d) H5IO6 (e) XeF4

21.115 Draw Lewis structures, indicating any resonance forms, for each of the following. (a) NO2– (b) CO3

2–

(c) OCN– (the C is central) (d) CNO– (the N is central) (e) ClF3

21.116 Write the electron configuration for the ion produced in each of the following half-reactions.

(a) Ca(s) Ca2+(aq) 2 e– (b) Cl2(g) + 2 e– 2 Cl–(aq) (c) Fe3+(aq) + 1 e– Fe2+(aq)

(d) Po(s) Po2–(aq) (e) Cu(s) Cu2+(aq) 2 e–

21.117 Both hydrogen (H2) and methane (CH4) are useful fuels. Their combustions are redox reactions. (a) Write

balanced chemical equations for the combustion of each of these gases to produce only gaseous products.

(b) Calculate the heat of reaction for each of your reactions in part (a). (c) Compare the amount of heat produced by

the combustion of 100.0 g of hydrogen to the amount of heat produced by the combustion of 100.0 g of methane.

(d) Compare the amount of heat produced by the combustion of 100.0 L of hydrogen gas at 25°C and 760 torr to the

amount of heat produced by the combustion of 100.0 L of methane at the same temperature and pressure.

21.118 The following half-reactions illustrate the reductions of the various halogens. (a) Relate these reactions to

the positions of these elements on the periodic table. (b) Relate the trend of their values to the trend in the electron

affinity of these elements. (c) Predict an approximate value for the reduction of astatine.

F2(g) + 2 e– 2 F–(aq) +2.87 V

Cl2(g) + 2 e– 2 Cl–(aq) +1.359 V

Br2(l) + 2 e– 2 Br–(aq) +1.087 V

I2(s) + 2 e– 2 I–(aq) +0.5355 V

21.119 A cell containing an aqueous solution of an unknown metal chloride (XCl3) was connected in series with a

cell containing aqueous silver(I) nitrate solution. After the solutions were electrolyzed, 1.44 g of Ag was deposited

and 0.120 g of an unknown metal X was deposited. Calculate the atomic mass of X.

21.121 In hot basic solution, aluminum metal oxidizes to aluminate (AlO2–). Write a balanced equation for the

reaction of aluminate in hot basic solution with each of the following.

(a) The reaction with water produces hydrogen.

(b) The reaction with nitrate ion produces ammonia.

(c) The reaction with permanganate ion produces manganese(IV) oxide.


22.1 What percentage of the elements are metals?

22.2 Metallic bonding is an extension of what basic bonding theory?

22.3 List some of the properties of metals that metallic bonding explains.

22.4 Define the terms malleable and ductile.

22.5 Define an alloy and give an example.

22.6 Unlike covalent bonds, which are directional, metallic bonds are best described as what?

22.7 Define the terms alloy and intermetallic compound.

22.8 (a) Where are the rare earth metals located on the periodic table? (b) Throughout the rare earth series, which

atomic orbitals are filling?


22.9 In the formation of a complex, many transition metals behave as Lewis acids. What is a Lewis acid?

22.10 Define the terms rock, mineral, and ore.

22.11 What is a native element?

22.12 If you simply look at a piece of the element silicon, it appears to be a metal. List two further tests that could

be conducted to see if it is really a metal.


22.13 (a) Throughout the transition series, which atomic orbitals are filling? (b) In each of the first three transition

series, identify the element in which these orbitals first fill completely.

22.14 Despite their position on the periodic table, zinc, cadmium, and mercury are not true transition metals. Why?

22.15 (a) Most of the metals in the first transition series form +2 ions. Which electrons are typically lost from these

metals in the formation of a +2 ion? (b) Copper is the only metal in the first transition series that forms a reasonably

stable +1 ion. Using the electron configuration of copper, show why this would be expected.

22.16 Write the electron configuration for each of the following: (a) Mn2+ (b) Sc3+ (c) Cu+ (d) Tc4+ (e) Ti3+

22.17 Write the electron configuration for each of the following: (a) Cr3+ (b) Mn6+ (c) Fe3+ (d) Ta3+ (e) Os2+

22.18 (a) Write the electron configuration of manganese. (b) Which electrons does manganese lose when forming a

manganese(II) ion (Mn2+)? (c) Which electrons does manganese lose when forming Mn7+?

22.19 (a) Write the electron configuration for chromium metal. (b) Chromium forms a number of compounds in

which the oxidation number of chromium ranges from Cr2+ to Cr6+. Write the electron configurations for Cr2+, Cr3+,

and Cr6+. (c) Why are there no known compounds containing Cr7+?

22.20 For each of the following elements, write the formula for the oxide in which the element has the highest

expected oxidation state: (a) W (b) Zn (c) Zr (d) Os (e) Y

22.21 Give the formula for the fluoride containing each of the following in their highest oxidation state: (a) V;

(b) Cu; (c) Ti; (d) Re; (e) Y.

22.22 Compounds containing chromium(VI) are strong oxidizing agents. Tungsten is in the same family as

chromium, yet compounds containing tungsten(VI) are weaker oxidizing agents than are similar chromium

compounds. Why is this true?

22.23 Technetium (Tc) is a human-made element that does not occur naturally on the Earth. The first synthesis of

this element occurred in 1937. (a) Do you expect the properties of technetium to be more like manganese (Mn) or

rhenium (Re)? (b) Explain why you made your choice in part (a).

22.24 (a) In Chapter 8, there was a discussion of the general trend in atomic radii of elements in the same column on

the periodic table. What was this trend, and why does it occur? (b) Do the transition metals exhibit this same trend?

Explain why or why not.

22.25 (a) In Chapter 8, there was a discussion of the general trend in the atomic radii amongst elements in the same

period on the periodic table. What was this trend, and why does it occur? (b) Do the transition metals exhibit this

same trend? Explain why or why not. (c) What is the scandide contraction?

22.26 (a) In Chapter 8, there was a discussion of the general trend in the ionization energy of elements in the same

column on the periodic table. What was this trend, and why does it occur? (b) Do the transition metals exhibit this

same trend? Explain why or why not.


22.27 (a) In Chapter 8, there was a discussion of the general trend in electronegativity of the representative elements.

What was this trend? (b) Which representative element has the highest electronegativity? (c) Do the transition

metals exhibit this same trend? (d) Which transition metal has the highest electronegativity?

22.28 (a) Define the terms paramagnetic and diamagnetic. (b) Why are most of the transition metals and their ions

paramagnetic?

22.29 (a) Write the electron configurations of Ti, Ti3+, and Ti4+. (b) Identify each as either paramagnetic or

diamagnetic.

22.30 Define the terms ferromagnetic and antiferromagnetic.

22.31 Many transition metal compounds are colored. Which transition compounds are likely not to be colored?


22.32 What is the lanthanide contraction?

22.33 Write the electron configurations for (a) samarium, (b) erbium, and (c) americium.

22.34 (a) Based upon its electron configuration, is lanthanum (La) a transition metal or a lanthanide? (b) Based

upon its electron configuration, is actinium (Ac) a transition metal or an actinide?

22.35 (a) Most of the lanthanide elements adopt what oxidation state in the majority of their compounds? (b) Which

lanthanide has the most stable higher oxidation state? What is this oxidation state? (c) Which lanthanides have the

most stable lower oxidation state? What is this state?

22.36 For many years, the first few actinides were thought to be the beginning elements in a fourth transition series.

What led to this misconception?

22.37 Why do the lanthanides normally occur together in nature?

22.38 What technological development led to a drastic decrease in the cost of compounds of the lanthanide metals?

22.39 Name the only two actinides that occur in significant amounts in nature.

22.40 How many of the actinides are radioactive?

22.41 What is the only reason why the elements polonium, astatine, radon, francium, radium, actinium, and

protactinium are present on the Earth?


22.42 Define the terms coordination compound, complex, and ligand.

22.43 Define the terms donor atom and coordination number.

22.44 Define the terms monodentate ligand, polydentate ligand, and chelating ligand.

22.45 Define the terms labile complex and inert complex.

22.46 Define isomers, structural isomers, and stereoisomers.

22.49 Define coordination isomers and linkage isomers.


22.50 Define optical isomers, geometric isomers, and cis and trans isomers.

22.51 Can a monodentate ligand be a chelating ligand? Explain why or why not.

22.52 Give the formula for each of the following: (a) Hexaamminevanadium(III) nitrite

(b) Pentaaquabromoiron(III) phosphate (c) Dichlorobis(ortho-phenanthroline)iron(III) chlorate

(d) Tris(bipyridyl)iron(II) nitrate (e) Potassium diaquatetrabromochromate(III)

22.53 Give the formula for each of the following: (a) Potassium tetrabromo(ethylenediamine)ferrate(III)

(b) Bis(ethylenediamine)platinum(II) tetraiodozincate (c) Hexaamminecobalt(III) tris(oxalato)manganate(III)

(d) Tetraamminecarbonatotitanium(III) sulfate (e) Potassium hexafluorophosphate

22.54 Name each of the following: (a) [Ag(CN)2]– (b) trans-[Pt(NH3)2H(Br)]

(c) K3[Fe(C2O4)3] (d) K3[IrCl3(S2O3)] (e) trans-[Cr(NH3)4(H2O)2]3+

22.55 Name each of the following: (a) [AlCl4]– (b) [Co(en)(NH3)2Br2]Cl (c) [PtCl4(en)]

(d) [Pd(en)2][Cr(NH3)2Br4]2 (e) Cs3[Cr(C2O4)2Cl2] (f) [Zn(NH3)4]2+ (g) [Cr(NH3)4Br2] (h) cis-[Co(en)2(NO2)2]+

22.56 Name each of the following: (a) [Cu(NH3)4]2+ (b) [Mn(NH3)4Br2] (c) [Co(en)3]3+ (d) NaBrF4

22.57 (a) Use Lewis structures to show why each of the following ligands may serve as a Lewis base. Indicate

which of these ligands are chelating. (i) H2O (ii) NH3 (iii) CN– (iv) C2H4(NH2)2 (v) C2O42–. (b) Indicate the

coordination number of the central metal in each complex in the following: [Ag(CN)2]–; trans-[Cr(NH3)4(H2O)2]3+;

[Co(en)(NH3)2Br2]Cl; [PtCl4(en)]; [Pd(en)2][Cr(NH3)2Br4]2

22.58 (a) Use Lewis structures to show why each of the following ligands may serve as a Lewis base. Indicate

which of these ligands are chelating. (i) H2O (ii) NH3 (iii) CN– (iv) C2H4(NH2)2 (v) C2O42–. (b) Indicate the

coordination number of the central metal in each complex in the following: [Ag(CN)2]– ; trans-[Pt(NH3)2H(Br)];

K3[Fe(C2O4)3]; K3[IrCl5(S2O3)]; [AlCl4]–; Cs[Cr(C2O4)2Cl2]; [Zn(NH3)4]2+; [Cr(NH3)4Br2]

22.59 Indicate the coordination number of the central metal in each complex in the following: (a) trans-

[Pt(NH3)2H(Br)] (b) trans-[Cr(NH3)4(H2O)2]3+ (c) [Co(en)(NH3)2Br2]Cl (d) [PtCl4(en)] (e)

[Pd(en)2][Cr(NH3)2Br4]2 (f) cis-[Co(en)2(NO2)2]+ (a) 4 (b) 6 (c) 6 (d) 6 (e) 4 and 6 (f) 6

22.60 Use one or more of the compounds or ions [Cr(NH3)4(Cl)2]Br, [Pt(NH3)2(SCN)2], or [Co(en)2Cl2]+ to

illustrate each of the following types of isomerism: (a) coordination isomerism (simply rewrite the formula);

(b) linkage isomerism (either draw or rewrite the formula); (c) geometrical isomerism (draw the isomers).

22.61 The complex ion [PdBr2(CN)2]2– has a square planar geometry. Draw the isomers of this ion.

22.62 The compound [Pt(NH3)2Cl2] has a square planar geometry. Draw the isomers of this compound.

22.63 The complex of ammonia with copper(II) is intensely blue in color, and the complex of ammonia with

cobalt(III) is dark red in color. The addition of ammonia to a solution containing copper(II) immediately causes the

development of the intense blue color. The addition of ammonia to a solution containing cobalt(III) results in no

immediate change. Explain the different responses of the copper and cobalt solutions to ammonia.


22.64 Who first pioneered work in the understanding of coordination chemistry?

22.65 Define crystal field theory.

22.66 Sketch a diagram illustrating what happens to a set of d orbitals in an octahedral crystal field. Label the

crystal field splitting () on this diagram.


22.67 Sketch a diagram illustrating what happens to a set of d orbitals in a tetrahedral crystal field. Label the crystal

field splitting () on this diagram.

22.68 Sketch a diagram illustrating what happens to a set of d orbitals in a square planar crystal field.

22.69 The hexaaquatitanium(III) ion is a d1 ion that is dark purple in solution. (a) Sketch a crystal field diagram for

this octahedral ion. (b) Sketch a second crystal field diagram illustrating the result of the absorption of a photon of

light of sufficient energy to give the complex its characteristic color.

22.70 Define crystal field splitting using diagrams of the d orbitals in an octahedral and a tetrahedral environment.

22.71 Define spectrochemical series.

22.72 Rank the following complexes in order of increasing crystal field splitting.

(a) [Cr(H2O)6]3+ (b) [Cr(SCN)6]3– (c) [CrF6]3– (d) [Cr(CN)6]3– (e) [Cr(NH3)6]3+

22.73 Rank the following tetrahedral complexes in order of increasing crystal field splitting. (a) [Zn(NH3)4]2+

(b) [Zn(OH)4]2– (c) [Zn(en)2]2+ (d) [ZnI4]2– (e) [Zn(H2O)4]2+

22.74 (a) Define high-spin complex. (b) Define low-spin complex.

22.75 The iron(II) ion occurs in both high-spin and low-spin forms. Draw crystal field diagrams showing the

arrangement of the electrons in (a) a low-spin octahedral environment and (b) a high-spin octahedral environment.

22.76 The iron(III) ion occurs in both high-spin and low-spin forms. Two of the complexes containing this ion are

[FeF6]3– and [Fe(CN)6]3–. Choose which of these two ions is likely to be high-spin and which is likely to be low-

spin. Justify your choices.

22.77 Why are tetrahedral complexes not normally low-spin?

22.78 Explain why simple Lewis structures along with valence bond theory are inadequate as theories for predicting

complex formation.

22.79 Give the number of d electrons associated with the central atom in each of the following complexes.

(a) [Co(NH3)4Cl2]+ (b) [MoCl4O]2– (c) [V(C2O4)3]2– (d) [HgCl4]2– (e) [Cu(NH3)4]2+ (a) 6 (b) 2 (c) 1 (d) 10 (e) 9

22.80 Give the number of d electrons associated with the central atom in each of the following complexes. (a) trans-

[Fe(en)2(NO2)2]+ (b) [IrCl3(S2O3)]3– (c) [Mn(H2O)6]3+ (d) [NiF6]4– (e) [Zn(NH3)4]2+

22.81 What is the expected electron configuration for each of the following ions: (a) Co3+ (b) Cr2+ (c) Ru4+

22.82 Which of the following compounds are expected to have color because of d-d electron transitions? (a) CrO

(b) TiO2 (c) [Cu(NH3)4]SO4 (d) Na2[SiF6] (e) [Zn(NH3)4]Cl2 (f) [VO(H2O)4]SO4.

22.83 Predict the number of d electrons remaining on the transition metal ion in each of the following substances.

(a) [Ti(en)3]Cl3 (b) K2[Pt(CN)4] (c) K3[Cr(C2O4)3] (d) [V(EDTA)]ClO4 (e) K3[OsCl6].

22.84 Construct crystal field energy-level diagrams for each of the following complex ions. (a) [TiCl6]2–

(b) [CoF6]3–(a high-spin complex) (c) [Ir(NH3)6]3+ (a low-spin complex) (d) [NiBr4]2–(tetrahedral)

(e) [Pt(NH3)4]2+(square planar).

Summary

22.85 What is metallurgy?


22.86 Define the terms hydrometallurgy, pyrometallurgy, and electrometallurgy.

22.87 Define the terms leaching and refining.

22.88 Define the terms roasting, smelting, and slag.

22.89 The cyanide process is an oxidation-reduction process. Assign oxidation states to each of the elements in the

following reaction, and indicate which element undergoes oxidation and which element undergoes reduction.

4 Au(s) + O2(g) + 8 CN–(aq) + 2 H2O(l) 4 [Au(CN)2]–(aq) + 4 OH–(aq)

22.90 The recovery of silver in the cyanide process involves the oxidation-reduction reaction displayed below.

(a) Assign oxidation states to each of the elements in the reaction and indicate which element undergoes oxidation

and which element undergoes reduction. (b) How many grams of iron(II) sulfate are necessary to reduce the silver

in 100.0 L of 0.01250 M [Ag(CN)2]–? [Ag(CN)2]–(aq) + Fe2+(aq) + 4 CN–(aq) Ag(s) + [Fe(CN)6]3–(aq)

22.91 Why are lead ores, such as galena (PbS) roasted?

22.92 Tantalum (Ta) has many medical applications, because it provokes only minimal rejection in the body. Thus,

tantalum is useful in objects such as the screws used to hold bones together. The initial stage in the isolation of this

metal gives a product contaminated with niobium (Nb). Why might niobium be a common contaminant in

tantalum?

22.93 Zirconium (Zr) may be isolated from the ore baddeleyite (ZrO2). One of the steps in this process is to heat a

mixture of baddeleyite and carbon in an atmosphere of chlorine (Cl2) gas. This reaction produces gaseous

zirconium(IV) chloride (ZrCl4) and carbon monoxide (CO) gas. (a) Write a balanced equation for this reaction.

(b) How many grams of ZrCl4 might form from a mixture of 4.75 kg ZrO2, 3.22 kg C, and 2.75 kg of Cl2?

22.94 Complete and balance each of the following equations: (a) PbCO3(s) heat

(b) Ni(s) + CO(g) (c) Al(OH)3(s) + OH–(aq)

(d) ZrCl4(g) + Mg(l) (e) CaO(l) + SiO2(l)

22.95 The first step in the processing of a particular type of copper ore involves the air oxidation of the sulfur

present to sulfur dioxide. If a sample of ore weighing 1.5 × 105 kg contains 24% covellite (CuS) and 17% pyrite

(FeS2), how many kilograms of sulfur dioxide gas will form when this ore is processed?

22.96 Many substances may serve as iron ore. Examples are hematite (Fe2O3), magnetite (Fe3O4), goethite

(FeO(OH)), and siderite (FeCO3). Each of these minerals will reduce to iron in a blast furnace. For each of them,

calculate the mass of iron produced from 1.000 kg of ore.

22.97 The production of magnesium metal is accomplished either by the electrolysis of molten magnesium chloride

or by the Pidgeon process. The latter process begins with heating the relatively common mineral dolomite,

MgCa(CO3)2, to 1150° under reduced pressure, so that it decomposes to magnesium oxide solid, calcium oxide

solid, and carbon dioxide gas. Further heating of the metal oxides with ferrosilicon, a mixture of iron and silicon

often represented as Fe/Si, yields magnesium vapor; molten calcium silicate, Ca2SiO4; and solid iron. (a) Write

balanced chemical equations for each step in the Pidgeon process. (b) What is the maximum amount of magnesium

that might form from 10.0 kg of dolomite? (c) What is the percent yield if 1.00 kg of magnesium forms?

(a) MgCa(CO3)2(s) MgO(s) + CaO(s) + 2 CO2(g)

Fe/Si(s) + 2 MgO(s) + 2 CaO(s) 2 Mg(g) + Ca2SiO4(l) + Fe(s) (b) 1.32 kg (c) 75.9 %

22.98 Bauxite includes many impurities. One impurity is the metal gallium, in levels of about 0.02 percent.

Gallium is an unusual metal in that it will melt in your hand. Gallium is also important in the production of LED's.

In the Bayer process, gallium extracts along with the aluminum, because gallium, like aluminum, is amphoteric. It

is then possible to separate the gallium from the solution by simple electrolysis. How many metric tons of bauxite

are necessary to produce 10.0 kg of pure gallium? A metric ton is 1000 kg.


22.99 The Kroll process is useful in the preparation of zirconium metal. In this process (also useful in the

preparation of titanium), magnesium metal reduces zirconium(IV) chloride to zirconium metal and magnesium

chloride. The Kroll process requires a temperature in excess of 1000°C. (a) How many kilograms of magnesium

metal are necessary to produce 25 kg of zirconium metal? (b) Hafnium always contaminates zirconium samples.

Why is hafnium more likely than any other element to be a contaminant?

22.100 It is possible to produce small amounts of very pure zirconium, contaminated only with hafnium, by the van

Arkel-de Boer process. This process begins by heating impure zirconium metal, produced by the Kroll process, with

iodine to release volatile zirconium(IV) iodide vapor. The vapor decomposes to the elements on a white-hot

tungsten filament. How many grams of iodine are necessary to produce 75 g of pure zirconium?

22.101 The Goldschmidt process produces chromium metal from chromium(III) oxide. The reaction is

Cr2O3(s) + 2 Al(s) 2 Cr(s) + Al2O3(s). Determine the enthalpy change for the Goldschmidt process.

22.102 Rhenium, the last non-radioactive element discovered, occurs as an impurity in molybdenum ores. During

the processing of molybdenum ores, the rhenium is converted to solid rhenium(VII) oxide (Re2O7). This oxide

dissolves in aqueous sodium hydroxide solution to form aqueous sodium perrhenate (NaReO4). The addition of

aqueous ammonium chloride to this solution precipitates solid ammonium perrhenate. After drying the precipitate,

heating the solid with hydrogen gas yields solid rhenium metal, water vapor, and ammonia gas. Write balanced

chemical equations for the isolation of rhenium.

22.103 One of the important ores of cobalt is the mineral cobaltite (CoAsS). Roasting this ore in air produces solid

cobalt(II) oxide, sulfur dioxide gas, and tetraarsenic decaoxide vapor. The solid, including impurities, is treated with

aqueous sulfuric acid. This dissolves the cobalt, as cobalt(II) sulfate, and some of the impurities. Treatment of the

cobalt(II) solution with a basic hypochlorite solution precipitates the cobalt as cobalt(III) hydroxide. The

hypochlorite ion reduces to aqueous chloride ion during this step. The dried hydroxide precipitate is converted to

the solid oxide by heating. Reduction of the oxide with hydrogen gas yields cobalt metal and water vapor. (a) Write

balanced chemical reactions for the isolation of cobalt metal from cobaltite. (b) How many kilograms of cobaltite

are necessary to prepare 1.00 kg of cobalt metal? (c) Both sulfur dioxide and tetraarsenic decaoxide are pollutants.

How many kilograms of each of these potent pollutants must the chemical plant prevent from entering the

environment (assuming that none can be allowed to escape) for every kilogram of cobalt isolated?

22.104 The following reaction is one way to prepare zinc metal industrially:

ZnO(s) + CO(g) Zn(s) + CO2(g)

(a) Determine the H°, S°, and G° for this reaction. (b) Calculate K for this reaction at 25°C. (c) It is necessary

to carry out this reaction at elevated temperatures to shift the equilibrium to the right. Why is this so?

(a) H° = –199.8 kJ/mol S° = 47.4 J/mol K G° = –181.1 kJ/mol (b) K = 5.60 × 1031

(c) The S° value becomes more important at higher temperatures and this promotes the reaction.

22.114 Spinel (MgAl2O4) is a very hard and durable mineral. Pure spinel is colorless; however, spinels containing

small amounts of impurities are found in a wide variety of colors and are commonly used in jewelry. The most

valuable color of spinel is red. (a) What is the most likely transition metal ion to impart the red color? (b) Spinel

has a cubic structure with a unit cell edge of 8.103 Å and a density of 3.55 g/cm3. How many MgAl2O4 units are

present in a unit cell? (c) Are the unit cells in spinel likely to be simple cubic, body-centered cubic, or face-centered

cubic?

22.115 Pyrope (Mg3Al2Si3O12), one member of the garnet group of minerals, has a unit cell edge of 11.459 Å, and

there are eight formula units per unit cell. Determine the density of pyrope in g/cm3.

The following two questions refer to BLBMWS Figure 6.4

22.118 Gemologists may use a spectroscope to identify gemstones. Certain natural gems have unique spectra, which

makes identification simple. The presence of certain transition metal ions gives rise to absorption bands, which

produce distinctive colors. Emerald for example, has bands at 6835 Å, 6805 Å, 6620 Å, and 6460 Å. (a) To what


frequencies of electromagnetic radiation do these wavelengths correspond? (b) Consulting Figure 7.2, give the

approximate color absorbed by each of these absorption bands.

22.119 Gemologists may use a spectroscope to identify gemstones. Certain natural gems have unique spectra, which

makes identification simple. The presence of certain transition metal ions gives rise to absorption bands, which

produce distinctive colors. Green tourmaline for example, has bands at 6400 Å and 4975 Å. (a) What is the energy

of each of these bands (in joules)? (b) Consulting Figure 7.2, give the approximate color absorbed by each of these

absorption bands.

Summary

22.120 How many protons, neutrons, and electrons are present in each of the following atoms or ions?

(a) titanium-50 (b) 99Mo (c) 184Re4+

22.121 The primary process for the production of sodium metal is the electrolysis of molten sodium chloride, which

reacts as shown in the first reaction displayed below. Although the other alkali metals can be produced in an

analogous manner, potassium is normally prepared by means of the second reaction displayed below. Why do the

lattice energies of the compounds (KCl and NaCl) in the second reaction favor this reaction? (Hint: Sodium

chloride and potassium chloride adopt the same crystal structure.)

2 NaCl(l) iselectrolys 2 Na(g) + Cl2(g)

Na(g) + KCl(l) C850 NaCl(l) + K(g)

22.122 Calculate the enthalpy change for each of the following reactions.

(a) FeO(s) + CO(g) Fe(s) + CO2(g) (b) Si(s) + 2 PbO(s) SiO2(s) + 2 Pb(s)

(c) MnO(s) + H2(g) Mn(s) + H2O(g) (d) 2 ZnS(s) + 3 O2(g) 2 ZnO(s) + 2 SO2(g)

(e) TiO2(s) + 2 Cl2(g) TiCl4(g) + O2(g)

(a) H° = –11.1 kJ/mol (b) H° = –479.8 kJ/mol (c) H° = 143.4 kJ/mol (d) H° = –661.8 kJ/mol

(e) H° = 181.5 kJ/mol

22.123 Titanium (Ti) may be isolated from the ore rutile (TiO2). One step in the process is to heat a mixture of rutile

and carbon in an atmosphere of chlorine (Cl2) gas. This reaction produces gaseous titanium(IV) chloride (TiCl4) and

carbon monoxide (CO) gas. (a) Write a balanced equation for this reaction. (b) Determine the enthalpy change for

the reaction.

22.124 The nitrite ion (NO2–), the cyanide ion (CN–), and the azide ion (N3

–) may serve as ligands. Draw the Lewis

structure of each of these three ions, showing all resonance forms. (b) For each of your Lewis structures, determine

the formal charge of each atom. (c) Predict the molecular geometry of each ion. (d) Based on the formal charges,

which of the atoms in the ligand is most likely to be the donor atom?

22.125 Molybdenum is often isolated from the ore molybdenite (MoS2). Roasting of molybdenite in air produces

sulfur dioxide (SO2) and molybdenum(VI) oxide (MoO3). The molybdenum(VI) oxide may be reduced to the metal

using aluminum. (a) Write balanced chemical equations for these reactions. (b) Sulfur dioxide is a pollutant. How

many kilograms of sulfur dioxide form when 100.0 kg of molybdenum form?

22.126 A particular lead ore is 35.0% cerussite (PbCO3). (a) Calculate how many kilograms of lead may be

produced from 10.0 metric tons of this ore. (b) What is the percent yield if only 2500 kg of lead form? A metric ton

is 1000 kg.

22.127 The industrial preparation of calcium metal involves the electrolysis of molten calcium chloride. Chlorine

gas is the other product of this reaction. (a) Assuming this process to be 95% efficient, how many grams of calcium

metal could be produced by a current of 9500 A operated for 24 h? (b) What volume of chlorine gas would form at

the same time, assuming the gas is collected at 25°C and a pressure of 755 torr?

(a) 1.6 × 105 g Ca (b) V = 1.0 × 105 L


22.128 A cell has a half-cell containing a silver electrode in a 0.100 M solution of calcium chloride with some solid

silver(I) chloride resting at the bottom, and a half-cell containing a cobalt electrode in a 0.100 M solution of barium

hydroxide with some solid cobalt(III) hydroxide resting at the bottom. A salt bridge containing 5.0 M potassium

nitrate solution connects the two half-cells. The observed cell voltage is –0.768 V. Calculate the Ksp, the

equilibrium constant, for the following equilibrium: Co(OH)3(s) Co3+(aq) + 3 OH–(aq)

You will need the following additional half-reaction information:

Co3+(aq) + 3 e– Co(s) E° = 0.33 V AgCl(s) + 1 e– Ag(s) + Cl–(aq) E° = 0.2223 V

22.129 A voltaic cell has a half-cell containing a silver cathode in a 0.100 M solution of calcium chloride with some

solid silver(I) chloride resting at the bottom, and a half-cell containing a mercury electrode in a solution that is 0.100

M in barium cyanide and 0.0150 M in potassium tetracyanomercurate(II), K2[Hg(CN)4]. A salt bridge containing

6.0 M potassium nitrate connects the two half-cells. The observed cell voltage is 0.605 V. Calculate Kf, the

equilibrium constant, for the following equilibrium: Hg2+(aq) + 4 CN–(aq) [Hg(CN)4]2– Kf = 2 1041

22.130 The Mond process is used in the production of pure nickel. How many liters of carbon monoxide gas at 1.00

atm and 75°C are necessary to process 1 metric ton (1.00 × 103 kg) of nickel that is initially 98.2 % pure?

Ni(s) + 4 CO(g) Ni(CO)4(g) V = 1.91 × 106 L

22.131 A graduate student isolates 17.5486 g of a purplish-black solid during her investigation of some chromium

compounds. She prepared the compound by adding ammonia to a basic chromium solution. The elements in the

sample are chromium, nitrogen, hydrogen, and oxygen. She submits 0.6003 g of this solid for combustion analysis

and learns that combustion with excess oxygen gas produces 0.5533 g of water, a gas, and 0.2223 g of a solid. The

gas is later found to be nitrogen, with a volume of 1.103 L at 27.0 °C and 149.0 torr. The solid is chromium(III)

oxide. Spectral analysis of another portion of the sample indicates that there is no coordination of the chromium by

oxygen. Analysis of a third portion indicates that all the nitrogen is in the form of ammonia and all the oxygen is

present as hydroxide. (a) Determine the formula of the compound. (b) Draw the structure of the most likely

complex ion present in this compound. (c). Name this compound.


23.1 Define radioactivity and radioisotope.

23.2 How many protons, neutrons, and electrons does each of the following isotopes have? (a)4

2He (b)

23

11Na

(c) 56

26Fe (d)

98

43Tc (e)

238

92U

23.3 How many protons, neutrons, and electrons does each of the following isotopes have? (a) 3

2He (b)

40

19K

(c) 59

27Co (d)

188

76Os (e)

238

94Pu

23.4 How many protons, neutrons, and electrons does each of the following isotopes have? (a) carbon-12 (b) C-14

(c) bromine-79 (d) Re-190 (e) gold-197

23.5 How many protons, neutrons, and electrons does each of the following isotopes have? (a) boron-11 (b) N-15

(c) krypton-85 (d) Pb-206 (e) silver-107

23.6 Construct a table listing the mass number, the number of protons, and the number of neutrons in each of the

following nuclei: (a) 55Mn (b) 235U (c) 18F (d) strontium-90 (e) lead-205.



23.7 Indicate the symbol for (a) electron (b) positron (c) neutron. (a) 𝒆−𝟏𝟎 (b) 𝒆+𝟏

𝟎 (c) 𝒏𝟎𝟏

23.8 Indicate the symbol for (a) proton (b) alpha particle (c) beta particle.

23.9 Define each of the following. (a) alpha emission (b) beta emission (c) gamma emission

23.10 Define each of the following. (a) electron capture (b) positron emission (c) spontaneous fission

23.11 There are eight types of radioactive decay mentioned in Section 23.2. Which two types are rare?

23.12 (a) Determine the frequency of a gamma ray having an energy of 2.30 × 10–16 J. (b) Determine the

wavelength in meters of the gamma ray in part (a).

23.13 (a) Determine the frequency of a gamma ray having an energy of 8.25 × 10–16 J. (b) Determine the

wavelength in meters of the gamma ray in part (a).

23.14 (a) Determine the wavelength in meters of a gamma ray having a frequency of 3.82 × 1017 s–1. (b) Determine

the energy in joules of the gamma ray in part (a).

23.15 (a) Determine the wavelength in meters of a gamma ray having a frequency of 7.92 × 1017 s–1. (b) Determine

the energy in joules of the gamma ray in part (a). (c) The energy of gamma rays is usually expressed in MeV

(million electron volts). Determine the energy of the gamma ray in part (b) in MeV. An electron volt equals 96.48

kJ/mol. (a) = 3.79 × 10–10 m (b) E = 5.25 × 10–16 J (c) E = 3.28 × 10–3 MeV


23.16 The following questions refer to even or odd numbers of protons and neutrons in a nucleus. (a) The most

common isotopes have what combination of protons and neutrons? (b) The least common isotopes have what

combination of protons and neutrons?

23.17 (a) Define transmutation. (b) What is another name for transmutation?

23.18 (a) What are magic numbers? (b) List the values of the magic numbers.

23.19 (a) What type of nuclear decay process is likely to occur in nuclei above the band of stability? (b) What type

of nuclear decay is likely in nuclei below the band of stability? (c) Where on the proton-neutron plot does alpha

decay most often occur?

23.20 (a) What is the slope of the field of stability for low atomic mass nuclei in the proton-neutron plot? (b) What

is the maximum slope of the field of stability in the proton-neutron plot?

23.21 List the observed modes of decay for potassium-40.

23.22 (a) Why is a positron considered a form of anti-matter? (b) What happens when a positron encounters its

counterpart?

23.23 What is the balanced nuclear equation for each of these reactions? (a) Aluminum-30 undergoes beta decay.

(b) Einsteinium-252 (Es) undergoes alpha decay. (c) Molybdenum-93 undergoes electron capture. (d) Phosphorus-

28 undergoes positron emission.

23.24 Write balanced nuclear equations for each of the following reactions: (a) Barium-126 undergoes electron

capture. (b) Krypton-94 undergoes -decay. (c) Radon-204 undergoes -decay. (d) Bromine-77 produces a

selenium isotope. (e) Iron-52 undergoes positron decay.


23.25 Predict the nucleus obtained by each of the following processes. (a) 234Pu decays by emission (b) 248Bk

decays by emission (c) 196Pb decays through two successive electron capture processes (d) 226Ra decays through

three successive emissions (e) 69As decays by positron emission

23.26 Balance the following reactions by filling in one of the blanks in each equation with the appropriate species.

It is not necessary to include gamma rays. Note: Only one blank per reaction is to be filled (i.e., some blanks will be

left empty).

(a) 𝐶𝑒58137 + _____ 𝐿𝑎57

137 + _____ (b) 𝐻𝑔80205 + _____ 𝑇𝑙81

205 + _____

(c) 𝐶𝑚96245 + _____ 𝑃𝑢94

241 + _____ (d) 𝐿𝑟103256 + _____ _____ + 𝐻𝑒2

4

(e) 𝑈92238 + 𝑛0

1 _____

23.27 Balance the following by filling in one of the blanks in each part with the appropriate species. Gamma rays

are not needed. Only one blank per part is to be filled (i.e., some will be left empty).

(a) 𝐶𝑙1733 𝑆16

33 + _____ (b) 𝐶𝑎2041 + _____ 𝐾19

41

(c) 𝑆𝑚62147 + _____ 𝑁𝑑60

143 + _____ (d) 𝐴𝑢79199 + _____ 𝐻𝑔80

199 + _____

(e) 𝑀𝑑101257 + _____ 𝐹𝑚100

257 + _____

23.28 Balance the following by filling in one of the blanks in each part with the appropriate species. Gamma rays

are not needed. Only one blank per part is to be filled (i.e., some blanks will be left empty).

(a) 𝐵𝑒47 + _____ 𝐿𝑖3

7 + _____ (b) 𝑀𝑜4298 + 𝐻1

2 𝑛01 + _____

(c) 𝐴𝑢79194 + _____ _____ + 𝑒−1

0 (d) 𝑈92235 + 𝑛0

1 𝑋𝑒54135 + 2 𝑛0

1 + _____

(e) 𝐶𝑓98252 + 𝐵5

10 3 𝑛01 + _____

23.29 For each of the following write balanced equation: (a) 𝑈92238 (n, ) 𝑈92

239 ; (b) 𝑁714 (p, ) 𝐶6

11 ; (c) 𝑂818 (n, ) 𝐹9

19 ; (d)

𝑁714 (, p) 𝑂8

17 ; (e) 𝑈92234 ()__.

23.30 What is the final isotope produced by the thorium decay series? In this series, thorium-232 loses a total of 6

particles and 4 particles in a 10-stage process. 𝑷𝒃𝟖𝟐𝟐𝟎𝟖

23.31 Balance the following reactions by filling in one of the blanks in each equation with the appropriate species.

It is not necessary to include gamma rays. Note: Only one blank per reaction is to be filled in (i.e., some of the

blanks will be left empty).

(a) 𝑇ℎ90232 + 𝐶6

12 + _____ 𝑃𝑢94240 + _____ (b) 𝐵𝑟35

80 + _____ 𝑆𝑒3480 + _____

(c) 𝑃1532 + _____ 𝑆16

32 + _____ (d) _____ + 𝐻𝑒24 𝐵𝑒4

7

(e) 𝑆𝑚62149 + _____ _____ + 𝐻𝑒2

4

23.32 Chose the more stable isotope in each of the following pairs. (a) 16O 17O (b) 23Na 24Na (c) 43Ti 48Ti

(d) 53Mn 58Mn (e) 114Sn 123Sn


23.33 The radioactive decay process is an example of what type of kinetics?

23.34 (a) Rearrange the equation ln [𝐴]0

[𝐴]𝑡 = kt to solve for t. (b) What is the value of t when [A]t = 1/2 [A]0? (c) How

does your result from part b relate to the equation for the half-life of a first-order reaction?

23.35 Tritium (3H) is a radioactive isotope of hydrogen with a half-life of 12.33 years. A sample contains 0.100 g

of tritium. How many micrograms of tritium will remain after 125 years? 8.89 × 10–5 μg

23.36 Sodium-24 has a half-life of 14.96 h. What percent of the sodium-24 will remain after the following numbers

of half-lives have passed? (a) 1 (b) 2 (c) 3 (d) 4 (e) 5

23.37 Potassium-44 has a half-life of 22.1 min. What is the value of k for potassium-44 in min–1? 0.0314 min–1


23.38 (a) How much iron-55 will remain in a 1.000-g sample after 3.750 years, if the half-life of iron-55 is 2.7

years? (b) The half-life of krypton-85 is 10.76 years. What percentage of a sample of krypton-85 will remain after

21.50 years?

23.39 (a) A sample of the radioactive isotope argon-41 produces 555 d/s. After 90.0 min, it produces 314 d/s.

Determine the half-life of argon-41. (b) After 182 min, the amount of arsenic-78 in a particular sample has

decreased from 5.00 mg to 1.25 mg. What is the half-life of arsenic-78?

23.40 (a) The half-life of strontium-90 is 29.1 years. How many years will it take for the radiation from a sample of

strontium-90 to drop to 0.0100 % of its original value? (b) The half-life of iodine-131 is 8.040 days. How many

days will it take for the radiation from a sample of iodine-131 to drop to 0.0100 % of its original value?

(a) 3.87 × 102 y (b) 107 d


23.41 How is the Einstein relationship in conflict with the law of conservation of mass?

23.42 Define the mass defect and the nuclear binding energy.

23.43 What nucleus has the greatest nuclear binding energy per nucleon?

23.44 The energy released when a helium atom forms from its constituent particles is 4.537 × 10–12 J/atom. How

much energy is necessary to separate a helium atom into separate protons, neutrons, and electrons?

23.45 Rank the following in order of increasing nuclear binding energy per nucleon. (a) 96Mo (b) 56Fe (c) 120Sb

(d) 238U (e) 151Sm

23.46 (a) An iron-56 atom weighs 55.9349 amu. How much energy is required to convert this atom to separate

protons, neutrons, and electrons? (b) How much energy is required to break apart 1.000 mol of iron-56 atoms?

(a) 7.887 × 10–11 J (b) 4.750 × 1013 J/mol

23.47 (a) What is the energy change when a mole of barium-130 atoms (atomic mass = 129.9062 amu) is formed

from protons, electrons, and neutrons? (b) Determine the nuclear binding energy (in J/mol) and the binding energy

per nucleon for 209Bi (208.9804 amu). (c) What is the energy change when a mole of bromine-81 atoms (atomic

mass = 80.9163 amu) is formed from protons, electrons, and neutrons?

23.48 One possible reaction of many in a hydrogen bomb explosion is

2H + 3H 4He + 1n

Calculate the energy change for the production of 1.0000 mol of helium if the reactants and products have the

following masses: hydrogen-2 = 2.01410 amu, hydrogen-3 = 3.01605 amu, and helium-4 = 4.00260 amu. (Hint:

mass defect = mass of products − mass of reactants.) –1.697 × 1012 J


23.49 Define nuclear fission and nuclear fusion.

23.50 (a) Define critical mass. (b) Does a nuclear reactor have a supercritical mass?

23.52 The uranium deposit in Okla shows an unusually low concentration of uranium-235. Explain why.

23.53 (a) What type of nuclear weapon relies only on fission? (b) What type of nuclear weapon relies on fusion?

23.54 The fission of a uranium-235 nucleus releases an average of 2.9 × 10–11 J. If the conversion of this energy to

electrical energy is only 42% efficient in a typical nuclear power plant, how many grams of uranium-235 must


undergo fission to operate a 2000-MW nuclear power plant for a year? (Assume the output of the power plant is

2000 MW during the year.) A watt is 1 J/s.


23.56 Define and give the abbreviation for becquerel, curie, gray, and sievert.

23.57 (a) What is the meaning of the acronym rad? (b) What is the meaning of the acronym rem? (c) Which are

more useful in evaluating a radiation hazard, rads or rems?

23.58 The three most common types of radiation were also the first to be discovered: alpha emission, beta emission,

and gamma emission. (a) Rank the three most common types of radiation in order of relative penetrating power.

(b) Rank the three most common types of radiation in order of relative ionizing power.

23.59 Which poses a greater health risk, a sample of a powdered radioactive substance on your hands or a sample of

the same powder in the air you inhale?

23.60 What three natural radioisotopes are in every living organism?

23.61 Your body contains the following naturally occurring radioisotopes. Rank them in order of increasing

potential damage. (a) 14C, a emitter with a half-life of 5730 years (b) 3H, a emitter with a half-life of 12.33 years

(c) 40K, a emitter with a half-life of 1.28 × 109 years (this isotope also emits other types of radiation).

23.62 What are the factors to consider when evaluating a radiation hazard?

23.63 Why is it so difficult to determine the adverse affects of low doses of radiation?

23.64 Describe the inverse square relationship that applies to radiation hazards.


23.65 (a) What is one diagnostic use of radioisotopes? (b) What is one therapeutic use of radioisotopes?

23.66 List three radioisotopes and their medical uses.

23.67 How does a nuclear pacemaker operate?

23.68 (a) What is the meaning of the acronym PET? (b) What radioisotopes are useful for PET scans? (c) What

type(s) of radiation are important in a PET scan?

23.69 List three radioisotopes useful in the treatment of tumors.

23.70 A solution of sodium iodide containing iodine-131 was injected into a patient to test for a malfunctioning

thyroid gland. The half-life of iodine-131 is 8.07 days. If the sample originally contained 5.00 g of iodine-131,

how many micrograms would remain after 7.00 days? 2.74 μg

23.71 Cobalt-60 is necessary for certain types of radiation therapy. This isotope has a half-life of 5.26 years. The

cobalt-60 must be replaced when the amount of cobalt-60 falls to 80.0% of its initial value. If a sample of cobalt-60

was placed in the radiation therapy unit on March 10, 2009, when will it need to be replaced?


23.72 Carbon-14 has a relatively short half-life for a useful radioactive dating tool. Why is this isotope still present

on the Earth?


23.73 What radioisotopes useful for radioactive dating have a half-life over a billion years?

23.74 Which radioisotopes useful for radioactive dating produce a lead isotope?

23.75 Tritium, with a half-life of 12.33 years, is present in all living organisms. Why is tritium not as useful as

carbon-14 for dating ancient artifacts?

23.76 What is the source of carbon-14 in all living organisms?

23.77 The half-life of 14C is 5730 years. The current decay rate of the 14C in a certain sample is 8.900% of the initial

rate. How old is the sample?

23.78 A sample of hair from an Egyptian mummy gives off radiation from carbon-14 at a rate of 60.1% of a present-

day sample. How old is the mummy if the half-life of carbon-14 is 5730 years? 4.21 × 103 y

23.79 An artifact from an Egyptian tomb has a 14C activity of 12.4 counts per minute per gram of carbon, as

measured on a Geiger counter. For living organisms, the 14C activity is 15.2 counts per minute per gram of carbon.

The half-life of 14C decay is 5.73 × 103 years. Calculate the age of the artifact. 1.68 × 103 y

23.80 The process 238U 206Pb has a half-life of 4.5 × 109 years. A particular sample of a mineral contains 55.5 mg

of 238U and 13.5 mg of 206Pb. How old is the mineral sample? 1.6 × 109 y

23.81 A total of 54.2 mg of argon-40 (produced by the radioactive decay of potassium-40) was found in a rock

sample. The rock also contained 62.3 mg of potassium-40. If the half-life of the potassium-40 is 1.28 × 109 years,

how old is the rock?

Summary

23.91 The human body contains about 140 g of potassium. Potassium-40 makes up 0.0117% of this potassium. The

half-life of potassium-40 is 1.27 billion years. How many potassium-40 atoms in the human body decay every

minute?

23.92 The composition of gasoline may be taken to be C8H18. When this gasoline burns in an automobile engine, it

combines with gaseous oxygen to produce carbon dioxide gas and water vapor. (a) Write a balanced chemical

equation for the combustion of gasoline. (b) Using standard heats of formation, determine the energy in joules from

the combustion of 1.00 gal of gasoline. (The density of gasoline is 0.8242 g/mL.) (c) Use the Einstein relationship

to determine the mass equivalent, in grams, of the energy released in part (b).

(a) 2 C8H18(l) + 25 O2(g) 16 CO2(g) + 18 H2O(g) (b) –1.39 × 108 J (c) 1.54 × 10–6 g

23.93 (a) The analysis of a rock sample showed 2.53 × 10–4 mol of uranium-235 and 1.3 × 10–5 mol lead-207. The

half-life of uranium-235 is 0.704 billion years. What was the age of the sample? (b) What percent of the potassium-

40 originally present in the same sample remains? The half-life of potassium-40 is 1.27 billion years.

23.94 Polonium-210 is an alpha emitter with a half-life of 138.4 days. Upon release, the alpha particles will attract

electrons and become helium atoms. How many helium atoms will a sample containing 1.0 ng of 210Po produce in

30.0 days?

23.95 Uranium-235 decays, through a series of steps, to lead-207. During this decay, seven alpha particles are

released. The half-life of uranium-235 is 0.704 billion years. Uranium-238 decays, through a series of steps to lead-

206. During this decay, eight alpha particles are released. The half-life of uranium-238 is 4.47 billion years. Each

of the emitted alpha particles becomes a helium atom. A uranium deposit contains an estimated 1.25 × 106 metric

tons of uranium ore. This ore currently has an average of 2.76 % uranium. The present-day composition of the

uranium is 99.274 % uranium-238 and 0.720 % uranium-235 (there is a small amount of uranium-234, which can be

ignored.) The deposit is 2.50 billion years old. How many liters of helium gas, measured at 25°C and 1.00 atm

have been released from this deposit since its formation?


23.96 Water contains 1.6 × 10–19 g of tritium per gram of water. Tritium is a beta emitter with a half-life of 12.33

years. A 750.0-mL bottle of wine contains 12.5 % alcohol, along with other ingredients, and the remainder is water.

The density of water is 1.00 g/mL and the density of the wine is 1.090 g/mL. (a) How many grams of tritium are

initially present in the bottle? (b) If this wine ages for 25 years, what percentage of the tritium will remain?

(c) How many total grams of tritium will remain? (d) What was the original molarity of tritium in the wine?

(a) 1.1 × 10–16 g T (b) 24.5 % (c) 2.81 × 10–17 g T (d) 5 × 10–17 M T

23.97 When a positron encounters an electron, they annihilate each other. The result is two identical gamma rays.

(a) What is the energy, in joules, of these gamma rays? (b) What is the frequency of these gamma rays? (c) What is

the wavelength, in nanometers, of these gamma rays?

23.98 There is a very small quantity of uranium-234 (0.0055 %) in natural uranium. This isotope is part of the

uranium-238 decay series, with a half-life of 2.46 × 105 years. Why is this isotope not useful in radioactive dating?

23.99 Nuclear fission produces a number of different isotopes. Three of the observed fission reactions for uranium-

235 are

235

92U +

1

0n

144

56Ba +

90

36Kr + 2

1

0n

235

92U +

1

0n

138

53I +

95

39Y + 3

1

0n

235

92U +

1

0n

140

55Cs +

92

37Rb + 4

1

0n

The product isotopes are radioactive and undergo beta emission. The masses are 1

0n = 1.008665 amu/neutron,

235

92U

= 235.043922 amu/atom, 144

56Ba = 143.92294 amu/atom,

90

36Kr = 89.91953 amu/atom,

138

53I = 137.9224 amu/atom,

95

39Y = 94.91279 amu/atom,

140

55Cs = 139.91727 amu/atom, and

92

37Rb = 91.91968 amu/atom. Determine the energy

change in joules for the fission of a uranium-235 atom by each of these processes.

23.100 One of the observed fission reactions for uranium-235 is 235

92U +

1

0n

138

53I +

95

39Y + 3

1

0n

Both of the product nuclei are radioactive. The iodine-138 undergoes three successive beta emissions before

becoming a stable isotope. The half-lives of these steps are 6.5 sec, 14.1 min, and 32.2 min, respectively. The

yttrium-95 also undergoes three successive beta emissions before becoming a stable isotope. The half-lives of these

steps are 10.3 min, 64.02 days, and 34.97 days, respectively. (a) Write balanced nuclear equations for each of the

six beta emissions. (b) Assuming the radiation from these isotopes drops to a relatively "safe" level after ten half-

lives, how long will it take for the radiation from each of these isotopes to reach that level?

23.101 Nuclear fission produces a number of different isotopes. Three of the observed fission reactions for

uranium-235 are: 235

92U +

1

0n

144

56Ba +

90

36Kr + 2

1

0n

235

92U +

1

0n

138

53I +

95

39Y + 3

1

0n

235

92U +

1

0n

140

55Cs +

92

37Rb + 4

1

0n

The product isotopes are radioactive and undergo beta emission. Their half-lives are: 𝐵𝑎56144 = 11.4 s, 𝐾𝑟36

90 = 32.3 s,

𝐼53138 = 6.5 s, 𝑌39

95 = 10.3 m, 𝐶𝑠55140 = 1.06 m, and 𝑅𝑏37

92 = 4.48 s. (a) Write a balanced nuclear equation for the nuclear

decay of each of the isotopes formed in the above reactions. (b) Assuming the radiation from these isotopes drops to

a relatively "safe" level after ten half-lives, how long will it take for the radiation from each of these isotopes to drop

to a "safe" level?

23.102 There are three radioactive decay series currently in operation on the Earth. These series begin with

thorium-232, uranium-238, and uranium-235. It is thought that soon after the formation of the Earth there was a

fourth series. This series began with neptunium-237 and decayed through the following sequence of emissions: ,

, , , , , , , , , and . (There is also an alternate series of steps.) (a) Write balanced nuclear equations for

each step in the neptunium-237 decay series. (b) The longest half-life of any isotope in the series is that of

neptunium-237 (2.14 × 106 years). What is the maximum number of years after the formation of the Earth that

neptunium-237 would be useful as a radioactive dating tool?


23.103 The energy changes in nuclear processes are often reported in electron volts (eV) instead of joules. An

electron volt equals 96.48 kJ/mol. Determine the energy change in the following reaction in MeV (million electron

volts): 2H + 3H 4He + 1n

The masses are 1.008665 amu/neutron, hydrogen-2 = 2.01410 amu/atom, hydrogen-3 = 3.01605 amu/atom, and

helium-4 = 4.00260 amu/atom.

23.104 The major source of radon in homes is the alpha decay of radium-226. Radium-226 is part of the uranium-

238 decay series and has a half-life of 1599 years. How many liters of radon gas will exactly 1.00 mol of radium-

226 produce in 1.00 year? Assume the volume is measured at 25°C and 1.00 atm. 0.01060 L

23.105 During the atmospheric testing of nuclear weapons during the 1950s, strontium-90 was released into the

Earth's atmosphere. Dust containing this radioisotope settled to the surface. Cattle ingested grass with this dust on

it. Due to the chemical similarity between calcium and strontium, the strontium-90 entered the milk produced by

these cattle and was then deposited in the place of calcium in the bones of people who drank the milk. The half-life

of strontium-90 is 28.78 years. Assuming a child absorbed strontium-90 from milk on January 1, 1955, on what date

will the strontium-90 level be reduced to 33 % of the initial value?

23.106 Pierre and Marie Curie discovered radium and polonium through the processing of over a ton of the uranium

ore known as pitchblende. Each ton of pitchblende contained about 0.14 g of radium. The final processing step for

radium involved fractional crystallization, a process in which a less soluble component precipitates from a mixture.

Originally, the Curies used chloride ion to separate radium from the main contaminant barium. However, it is

possible to separate radium from most other metallic elements by precipitation of radium nitrate. Calculate the

grams of radium remaining in 1.00 L of a solution with a nitrate concentration of 0.500 M. The Ksp of radium nitrate

is 1.1 × 10–3.

23.107 Give the name or formula for each of the following compounds of some radioactive elements.

(a) uranium(III) oxide (b) thorium(IV) fluoride (c) radium sulfate (d) polonium(II) oxide (e) francium nitrate

(f) U(C2O4)2 (g) Th3(PO4)4 (h) RaHPO4 (i) Po(NO2)4 (j) FrC2H3O2

23.108 Give the name or formula for each of the following compounds of some radioactive elements. (Note:

americium = Am, curium = Cm, promethium = Pm, technetium = Tc, rutherfordium = Rf.) (a) americium(III) oxide

(b) curium(III) fluoride (c) promethium(III) sulfite (d) technetium(VII) oxide (e) rutherfordium(IV) nitrate

(f) Am(C2O4)2 (g) Cm3(PO4)4 (h) PmPO4 (i) Tc(NO2)4 (j) Rf(C2H3O2)4

23.109 The energies of emitted radiation are typically expressed in MeV (million electron volts). An electron volt

equals 96.48 kJ/mol. For example, some beta particles emitted by cobalt-61 atoms have an energy of 0.158 MeV.

The mass of a beta particle is the same as that of an electron (9.1089 × 10–28 g). (a) Determine the energy of a

cobalt-61 beta particle in joules. (b) Determine the velocity of a cobalt-61 beta particle in meters per second.

(c) Determine the de Broglie wavelength, in meters, of a cobalt-61 beta particle.

23.110 The energies of emitted radiation are typically expressed in MeV (million electron volts). An electron volt

equals 96.48 kJ/mol. For example, some positrons emitted by zinc-64 atoms have an energy of 0.163 MeV. The

mass of a positron is the same as that of an electron (9.1089 × 10–28 g). (a) Determine the energy of a phosphorus-29

positron in joules. (b) Determine the velocity of a phosphorus-29 positron in meters per second. (c) Determine the

de Broglie wavelength, in meters, of a phosphorus-29 positron.

23.112 A 1.23-mg sample of americium (Am) metal was dissolved in hydrochloric acid with the formation of

americium chloride and hydrogen gas. Evaporation of the resultant solution gave 1.77 mg of pink americium

chloride. (a) Determine the formula of the americium chloride formed. (b) Write a balanced chemical equation for

the reaction of americium with hydrochloric acid.


CHAPTER 24 Corresponds to BLBMWS Chapter 24 24.1 Corresponds to BLBMWS 24.1

24.1 Define catenation and give two examples.

24.2 Determine how many σ bonds and how many π bonds are in each of the following. (a) a single bond (b) a

double bond (c) a triple bond (d) two double bonds

24.3 How do π bonds differ from σ bonds?

24.4 How many covalent bonds does carbon have in all organic compounds?

24.5 Label each of the following compounds as organic or inorganic. (a) CH4 (b) CH3CH2OH (c) FeCO3 (d) Al4C3

(e) C6H12O6

24.6 Draw Lewis structures for each of the following and label each carbon atom as sp, sp2 or sp3 hybridized.

(a) methane, CH4 (b) hydrocyanic acid, HCN (c) acetic acid, CH3COOH (d) oxalic acid, H2C2O4 (e) methylamine,

CH3NH2

24.7 Define functional group.


24.8 List the four types of hydrocarbons.

24.9 Define saturated and unsaturated hydrocarbons and give an example of each.

24.10 Use Lewis structures to illustrate resonance in benzene, C6H6.

24.11 The compound pyridine, C5H5N, has a structure similar to benzene. In this structure, a nitrogen atom replaces

one of the CH groups in one corner. (a) Is resonance possible in pyridine? (b) Use Lewis structures to illustrate

why there is or is not resonance in pyridine.

24.12 Convert each of the following condensed structural formulas to structural formulas. (a) CH4 (b) CH3CH3

(c) CH3(CH2)2CH3 (d) (CH3)2CHCH3 (e) C(CH3)4

24.13 Draw the structural isomers of pentane, C5H12. Name each of the isomers.

24.14 Draw the structural isomers of C6H12. Only draw isomers with single bonds.

24.15 Draw structural formulas for all of the structural and geometrical isomers of dibromopropene.

24.16 Note the following combustion reactions for cyclopentane and cyclopropane:

2 C5H10(g) + 15 O2(g) 10 CO2(g) + 10 H2O(l) ∆H = –6634 kJ

2 C3H6(g) + 9 O2(g) 6 CO2(g) + 6 H2O(l) ∆H = –4178 kJ

(a) Draw the structures of the two hydrocarbon reactants. (b) For each compound, calculate the average heat of

combustion for each mole of CH2 present. (c) Explain any observed differences in the two values obtained in part

(b).

24.17 Draw all possible isomers for C3H5Br.

24.18 Draw the structural formulas for all isomers of C3H4F2.

24.19 Naphthalene, like benzene, is an aromatic hydrocarbon. This compound consists of two six-membered rings

sharing an edge. Draw the separate resonance structures for naphthalene.


24.20 Draw structural formulas for each of the following compounds:

(a) CH3CH(CH3)CH2CH3 (b) 2-butyne; (c) CH2=CHCH2CH2CH3

24.21 Draw structural formulas for each of the following: (a) 5-methyl-cis-2-octene (b) 3-bromo-1-butyne (c) meta-

difluorobenzene (d) 2,2,5,6-tetramethylnonane (e) 2-methyl-3-isopropylhexane (f) 2-bromo-6-methyl-3-heptyne

(g) 2,3-dimethyl-2-butene (h) 1,6-octadiene

24.22 Draw the condensed structural formula of each of these compounds: (a) 2,4-dimethylhexane (b) 3,3-

dimethylpentane (c) trans-3-heptene (d) 1,2-dimethylcyclobutane (e) 1,2-dibromoethane (f) 1,3-dichlorobenzene

(g) cis-1,2-dimethylcyclopentane (h) 3-isopropyl-1-hexyne (i) 2-methyl-3-ethyloctane (j) 1-fluoro-4-propylnonane

24.23 Name the following compounds:


24.25 Write a balanced chemical equation for the dehydrogenation of propane, CH3CH2CH3.


24.26 (a) When propane, CH3CH2CH3, undergoes dehydrogenation, why is one of the two hydrogen atoms lost

always from the central carbon atom? (b) During the dehydrogenation of propane, both hydrogen atoms cannot

come from the central carbon. Why?

24.27 (a) The dehydrogenation of butane, CH3CH2CH2CH3, can give three products. Draw structural formulas and

name each of these compounds. (b) The compound methylpropane is an isomer of butane. Dehydrogenation of

methylpropane only gives one product instead of three. Why? (c) Draw the structural formula of the product of the

dehydrogenation of methylpropane.

24.28 The reaction of chlorine, Cl2, with pentane, C5H12, can produce a number of monochlorinated products,

C5H11Cl. Draw structural formulas and name each of the possible products.

24.29 When ethene, C2H4, and hydrogen gas, H2, are mixed there is no reaction. Why?

24.30 Propene reacts with hydrogen gas in the presence of a nickel catalyst. Write a balanced chemical equation,

using structural formulas, to indicate this reaction.

24.31 Write balanced chemical equations, using structural formulas, to indicate the reaction of one molecule of 2-

butyne with one molecule of hydrogen gas and for the reaction of two molecules of hydrogen gas. Assume the

reaction takes place in the presence of a platinum catalyst.

24.32 (a) Using structural formulas write a balanced chemical equation for the addition of bromine to 1-butene.

(b) Using structural formulas, write two balanced chemical equations for the addition of bromine to 1-butyne. In the

first reaction, assume the reactants react in a one to one ratio. In the second reaction, assume the ratio is two

bromines to one butyne.

24.33 Ignoring cis and trans isomers, the reaction of one molecule of CH2=CH-CH2-CH=CH2, with one molecule of

chlorine only gives one product; however, the reaction of one molecule of its isomer, CH2=CH-CH=CH-CH3, can

give two possible products. Explain.


24.34 Identify the functional group in alkenes and the functional group in alkynes.

24.35 (a) List the functional groups that contain oxygen. (b) In which of these is a carbonyl group present?

(c) Which of these functional groups are capable of forming hydrogen bonds?

24.36 (a) List the functional groups that contain nitrogen. (b) In which of these is a carbonyl group present?

(c) Which of these functional groups are capable of forming hydrogen bonds?

24.37 (a) Define primary, secondary, and tertiary alcohol. (b) Draw the isomers of C5H12O that are alcohols.

(c) Name each of the alcohols drawn in part b. (d) Classify each of the alcohols in part b as primary, secondary, or

tertiary. (e) Draw the isomers of C5H12O that are not alcohols.

24.38 Rank the following alcohols in order of increasing solubility in water. (a) CH3CH2OH

(b) CH3CH2CH2CH2OH (c) CH3CH2CH2CH2CH2CH2OH (d) CH3CH2CH2CH2CH2OH

(e) CH3CH2CH2CH2CH2CH2CH2CH2CH2CH2OH (e) < (c) < (d) < (b) < (a)

24.39 Classify the functional groups in Table 24.X as acidic, basic, or neutral.

24.40 (a) How do aldehydes differ from ketones? (b) Draw structural formulas for the smallest possible aldehyde

and the smallest possible ketone. (c) Name each of the compounds in part b.

24.41 The general formula for an alkane is CnH2n+2. What are the general formulas for an alcohol and an ether?

24.42 The general formula for an alkene is CnH2n. What are the general formulas for an aldehyde and a ketone?


24.43 The general formula for an alkane is CnH2n+2. What are the general formulas for a carboxylic acid and an

ester?

24.44 The general formula for an alkane is CnH2n+2. What are the general formulas for a primary amine, a secondary

amine, and a tertiary amine?

24.45 (a) What type of alcohol can be converted to an aldehyde? (b) What type of alcohol can be converted to a

ketone? (c) What type of alcohol can be converted to a carboxylic acid?

24.46 Decanoic acid, C10H22O2, is not soluble in water; however, it is readily soluble in aqueous base. Why?

24.47 What is a condensation reaction?

24.48 (a) Use structural formulas to illustrate the formation of an ester by the reaction of ethanol, CH3CH2OH, with

methanoic acid, HCOOH. (b) Name the ester produced in part (a) (c) The common name for methanoic acid is

formic acid. What is the common name for the ester formed in part a?

(a)

(b) Ethyl methanoate (c) Ethyl formate

24.49 Use structural formulas to illustrate the acid-base reaction of ethylamine, CH3CH2NH2, with methanoic acid,

HCOOH.

24.50 Use structural formulas to illustrate the formation of an amide by the reaction of ethylamine, CH3CH2NH2,

with ethanoic acid, CH3COOH.

24.51 Identify the functional groups in each of the following compounds:

24.52 Identify the functional groups in each of the following compounds:

(a) aldehyde (b) alkyne and amine (c) alcohol and carboxylic acid (d) phenol and ketone

(e) aldehyde and amine (f) alkene (2) and ether

24.53 Identify the carbon atom(s) in the following structure that has each of the following hybridizations: (a) sp3

(b) sp2 (c) sp.


24.54 Using structural formula write a chemical equation for the saponification of ethyl formate.

24.55 Draw the structural formula for each of the following: (a) 2-chloro-2-butanol; (b) l,3-propanediol; (c) methyl

acetate; (d) 3-pentanone; (e) propanal.

24.56 The IUPAC name for a carboxylic acid is based on the name of the hydrocarbon with the same number of

carbon atoms. Name the following carboxylic acids using the IUPAC system.

24.57 Draw structural formulas for the following: (a) butanal; (b) 3-hexanone; (c) 3-ethyl-2-octanone; (d) 2-

isopropylbutanal.

24.58 Draw the complete structural formulas for all isomers of C3H8O.

24.59 Draw the complete structural formula for an example of each of the following: (a) an ether that has the

formula C5H12O; (b) an aldehyde that has the formula C6H12O; (c) a ketone that has the formula C4H8O; (d) a

carboxylic acid that has the formula C10H20O2; (e) an ester that has the formula C2H4O2.


24.60 Define the terms monomer and polymer.

24.61 Define the terms homopolymer and copolymer.

24.62 Define the terms condensation polymer and addition polymer.

24.63 Write a balanced chemical equation, using structural formulas, to illustrate the synthesis of polypropylene

from propylene: CH2=CH-CH3.

24.64 The common name of chloroethene, C2H3Cl, is vinyl chloride. Use structural formulas to illustrate the

formation of polyvinyl chloride, also known as PVC. Begin with three molecules of vinyl chloride.

24.65 Ethylene glycol (HOCH2CH2OH) and oxalic acid may be used to prepare a polyester. Using two molecules of

each of these compounds draw the start of the polymer that will form.


24.66 Ethylene diamine, NH2CH2CH2NH2, and oxalic acid can be used as starting materials in the formation of a

nylon. Draw the structural formula of a portion of the polymer formed containing at least two molecules of each of

the starting materials.

24.67 The amino acid glycine, NH2CH2COOH, is capable of serving as the monomer in the formation of a

condensation polymer. Draw the structural formulas for three molecules of glycine and show how they might

combine to form a condensation polymer.


24.73 (a) Why was there interest in the synthesis of quinine? (b) Identify the functional groups present in quinine.

(c) How many σ bonds and how many π bonds are present.

24.74 A chemical company needs to produce 10.0 metric tons of aniline. The chief chemical engineer suggests the

following reaction scheme:

As with most reactions, the actual yield is lower than the theoretical yield. In this case, the percent yield for the first

step is 82.0 %, and for the second step, the percent yield is 75.0 %. How many metric tons of benzene does the

chemical engineer need to order? A metric ton is 1000 kilograms.

24.75 A chemist wishes to produce a small quantity of cinnamic acid in the laboratory utilizing the following

reaction:

She begins with 125 milliliters of benzaldehyde, and isolates 125 grams of cinnamic acid. The density of

benzaldehyde is 1.043 g/cm3. What was the percent yield of this reaction?

24.76 The first step in the synthesis of indigo involves the condensation reaction of anthranilic acid, HC7H6NO2,

with chloroacetic acid, HC2H2ClO2. This reaction is:

If the percent yield for this step is only 62.25 %, how many kilograms of each of the starting materials are necessary

to produce 100.0 kilograms of product? 11.3 kg HC7H6NO2 7.78 kg HC2H2ClO2


24.77 Many dyes are derivatives of naphthalene, C10H8. In most cases, the first step in dye synthesis is the

sulfonation of naphthalene. In this process, a sulfonic acid group, -SO3H, substitutes for one of the hydrogen atoms

on the naphthalene. Two different products will form. 1-naphthalenesulfonic acid is the major product is the

reaction takes place at about 80-90 °C, and 2-naphthalenesulfonic acid is the major product if the reaction takes

place at 160 °C. The reaction is:

Why are these the only two naphthalenesulfonic acids that can form?

Summary

24.78 Give the IUPAC name for each of the following molecules:

24.79 The bromination of ethane, C2H6, can be repeated until a bromine atom substitutes for each of the hydrogen

atoms. Draw structural formulas and name the products of the bromination of ethane beginning with C2H5Br and

ending with C2Br6.

24.80 Two alcohol molecules may undergo a condensation reaction to form an ether. Use structural formulas to

illustrate the formation of diethyl ether, (C2H5)2O, through the condensation reaction of two alcohol molecules.

24.81 Construct a table similar to Table 24.X for compounds with a molecular weight of about 72 g/mole. Some

categories will have more than one answer; however, only one of each category is necessary for this table.

24.82 Draw the structures for ammonia, methylamine, dimethylamine, and trimethylamine. Which of these

compounds is capable of hydrogen bonding?

24.83 What do all electrically conducting polymers have in common?

24.84 A hydrohalogenation reaction is an addition reaction where a hydrogen halide, such as HCl or HBr, adds to a

π bond. In this reaction, a hydrogen atom goes to one side of the double bond and the halogen atom goes to the

opposite side, instead of a halogen atom on both sides as in a halogenation reaction. Given a choice, the hydrogen

atom will add to the side with the carbon atom already having more hydrogen atoms. (a) Using structural formulas,

illustrate the hydrohalogenation reaction of ethene with hydrogen chloride. (b) Using structural formulas, illustrate

the hydrohalogenation reaction of propene with hydrogen bromide. (c) Using structural formulas, illustrate the

stepwise hydrohalogenation reaction of ethyne with hydrogen bromide. (Add one hydrogen bromide, and then add a

second hydrogen bromide to the product.)

24.85 A hydration reaction is an addition reaction where water adds to a π bond in the presence of an acid catalyst.

In this reaction, a hydrogen atom goes to one side of the bond and a hydroxyl group goes to the opposite side,

instead of a halogen atom on both sides as in a halogenation reaction. Given a choice, the hydrogen atom will add to

the side with the carbon atom already having more hydrogen atoms. (a) Using structural formulas, illustrate the

hydration reaction of ethene with water. (b) Using structural formulas, illustrate the hydration reaction of propene

with water. (c) Using structural formulas, illustrate the stepwise hydration reaction of ethyne with water. (Add one

water, and then add a second water to the product.)


24.86 Draw structural formulas for each of the following: (a) 2-methyl-trans-4-nonene (b) 4-iodo-2-hexyne (c) para-

dihydroxybenzene (d) 2,3,4,5-tetrachlorooctane (e) 3-methyl-3-isopropylheptane (f) 2-fluoro-3-ethyl-4-decyne

(g) cis-2,3-dichloro-2-butene (h) 1,4-heptadiene

24.87 Draw the condensed structural formula of each of these compounds: (a) 3,4-diethyloctane (b) 3,3-

dibromohexane (c) trans-3-nonene (d) trans-1,3-dimethylcyclobutane (e) 1,2-dibromopropane (f) 1,2-

dibromobenzene (g) cis-1,3-dimethylcyclohexane (h) 3-isopropyl-2-heptene (i) 2-methyl-5-butyldecane (j) 1-fluoro-

4-propylbenzene

24.88 Use structural formulas to illustrate the reaction of isopropyl alcohol with propanoic acid to form an ester.

24.89 Use structural formulas to illustrate the formation of an amide by the reaction of dimethylamine with ethanoic

acid.

24.90 Draw the structural formula for each of the following: (a) 1-bromo-3-pentanol; (b) l,2,3-propanetriol; (c) ethyl

propanoate; (d) 2-hexanone; (e) butanal.

24.91 The common name of cyanoethene, C2H3CN, is acrylonitrile. Use structural formulas to illustrate the

formation of polyacrylonitrile, also known as Orlon®. Begin with three molecules of acrylonitrile.



24.94 The following is the structure of indigo. (a) How many σ bonds are present? (b) How many π bonds are

present? (c) What is the hybridization of each of the numbered atoms?

C

C

N

C3

C4

C

C1

N2

C

C

O O

H H

24.95 The monomers chloroethene and 1,1-dichloroethene are used in the formation of a copolymer used in Saran

Wrap®. Assuming that the addition copolymer has alternating chloroethene and 1,1-dichloroethene units, draw the

structural formula of a segment of the polymer containing two of each monomer.


24.96 Aromatic compounds, such as benzene, are not very reactive. The presence of resonance makes the rings

exceptionally stable. When a reaction occurs on the ring of an aromatic compound, it tends to be a substitution

reaction, like alkanes, rather than an addition reaction, like alkenes. In 1877, the French chemist, Charles Friedel,

and the American chemist, James Crafts, discovered than in the presence of a catalyst, such as aluminum chloride,

an alkyl group could undergo a substitution reaction onto an aromatic ring. In honor of the discoverers, this is a

Friedel-Crafts reaction. An example of this reaction is:

The catalyst is a Lewis acid. This reaction takes place in organic solvents where the aluminum chloride behaves as a

nonelectrolyte. Use the Lewis structure of unionized aluminum chloride to show why this catalyst is a Lewis acid.

24.97 Toluene reacts with a mixture of concentrated nitric and sulfuric acids to form nitrotoluene. There are three

different nitrotoluenes formed. The reaction is:

The percentages below the structures indicate how much of each nitrotoluene forms. (a) Name each of the

nitrotoluenes. (b) Why can there be no other nitrotoluenes?

24.98 The reaction of a halogen such as chlorine or bromine with an alkane leads to a substitution of a halogen atom

for one of the hydrogen atoms and the other halogen atom combines with the displaced hydrogen atom to produce a

hydrogen halide. Under the same conditions, an alkene or an alkyne will undergo an addition reaction. On the other

hand, aromatic hydrocarbons do not react. However, in the presence of a catalyst, FeCl3 for Cl2 or FeBr3 for Br2,

aromatic compounds will undergo substitution reactions like alkanes. Use structural formulas to indicate the

substitution reactions that occurs when chlorine reacts with benzene in the presence of iron(III) chloride. Show only

the possible products of the addition of either one or two chlorine atoms.

24.99 In 1900, Victor Grignard discovered how to make an important class of compounds used in organic synthesis.

For his discovery of these compounds, now known as Grignard reagents, he received the 1912 Nobel Prize in

Chemistry. He discovered that magnesium metal would reaction with organic halides in dry ether to insert a

magnesium atom between the organic portion of the molecule, symbolized as R, and the halogen atom, symbolized

as X, in this way RX becomes RMgX. Even though a metal is present, the organomagnesium compound formed has

a covalent bond to the organic portion of the molecule. One use of Grignard reagents is to produce alcohols from

either aldehydes or ketones. The R from the Grignard reagent attaches to the carbonyl carbon and the carbonyl

oxygen becomes an alcohol oxygen. The reaction requires two steps, one where ether is the solvent, and one with an

acid:

(a) Repeat the above reaction using the same Grignard reagent, methylmagnesium chloride, reacting with propanal.

(b) Repeat the above reaction with ethylmagnesium chloride reaction with acetone (propanone).


24.100 The porphyrin ring system, illustrated below, is an important structure present in such biologically important

compounds as hemoglobin, myoglobin, and chlorophyll.

(a) How many σ bonds are present? (b) How many π bonds are present? (c) What is the hybridization of each of the

nitrogen atoms?

CHAPTER 25 Corresponds to BLBMWS Chapter 24 25.1 Corresponds to BLBMWS 24.6

25.1 Define biochemistry.

25.2 List two factors responsible for the usefulness of carbon to biochemistry.

25.3 (a) How does the list of elements present in living organisms compare to the abundances of those elements in

the Earth's crust? (b) Identify one element that is an exception to the comparison in part a.

25.4 Most biologically important materials utilize covalent bonds. Identify two materials where ionic bonding is

important.

25.5 List five elements that are biologically important because they exist as ions in solution.

25.6 (a) Define osmosis. (b) How does the osmotic pressure relate to the presence of ions in solution? (c) What will

happen to a cell if the osmotic pressure within the cell is too low? (d) What will happen to a cell if the osmotic

pressure within the cell is too high?

25.7 (a) List two biologically important oxidation-reduction processes. (b) What type of elements is involved in

most biologically important oxidation-reduction processes? (c) Give one example of an element important to

biologically important oxidation-reduction processes.

25.8 Why is hydrogen biologically important?


25.9 (a) Define a carbohydrate. (b) Define monosaccharide.

25.10 (a) What is an aldose? (b) What is a pentose? (c) What is an aldopentose?

25.11 (a) What is a ketose? (b) What is a hexose? (c) What is a ketohexose?

25.12 What is the significance of a "D" or an "L" in the name of a carbohydrate?

25.13 (a) What is a chiral carbon? (b) Why is the carbon atom in CHFClBr chiral?


25.14 (a) What is a disaccharide? (b) How are the components of a disaccharide linked? (c) What type of reaction

results in the formation of a disaccharide?

25.15 What is a polysaccharide?

25.16 (a) How are the structures of starch and cellulose similar? (b) How do the structures of starch and cellulose

differ?

25.17 Circle each of the chiral carbons in each of the following molecules.

D-arabinose D-galactose D-fructose

O

C

H

C HHO

CH

CH OH

CH2 OH

O

C

H

C OHH

CHO H

CHO H

CH

CH2

OH

OH

OH

CH2

C O

CHO H

CHO H

CH

CH2

OH

OH

OH

25.18 One anomer of mannose has the following structure. The α-anomer, like most monosaccharides, tastes sweet,

whereas the β-anomer tastes bitter.

OH

OH

H

OH

OH

HH

OH

CH2

H

HO

(a) Is this the α or the β anomer? (b) Draw the structure of the other anomer.

25.19 (a) Show how two galactose molecules can react to form a disaccharide. (b) Draw the structure of the

disaccharide. The structure of galactose is:

OOH

H

H

OH

H

OHH

OH

CH2

H

HO


Many of the problems in this section refer to BLBMWS Figure 24.18.

25.20 (a) What is an amino acid? (b) Draw the general formula for an α-amino acid.

25.21 (a) How many of the amino acids in Figure 25.X contain a chiral carbon? (b) How many amino acids in

Figure 25.X contain more than one chiral carbon?

25.22 (a) What is a zwitterion? (b) What is the isoelectric point for an amino acid?

25.23 Draw the amino acid alanine in its zwitterion form.

25.24 (a) What is a peptide bond? (b) Draw a peptide bond.


25.25 (a) What type of reaction is responsible for the formation of a peptide bond? (b) What functional groups react

in the formation of a peptide bond?

25.26 Draw the structure of the amino acid phenylalanine below, at, and above its isoelectric point.

25.27 Draw the structures of the following amino acids: (a) Cysteine (R = -CH2SH)

(b) Glutamine (R = -CH2CH2CONH2) (c) Alanine (R = -CH3) (d) Lysine (R = -CH2CH2CH2CH2NH2)

25.28 Refer to Figure 25.X when answering the following questions. (a) Which of the amino acids have nonpolar

side-chains? (b) Which of the amino acids have polar side-chains? (c) Which of the amino acids have acidic side-

chains? (d) Which of the amino acids have basic side-chains?


25.29 (a) Define protein. (b) What monomers are present in all proteins?

25.30 (a) Is a protein an addition polymer or a condensation polymer? Explain. (b) Is a typical protein a

homopolymer or a copolymer? Explain.

25.31 Define the terms primary structure, secondary structure, tertiary structure, and quaternary structure as these

terms apply to proteins.

25.32 When drawing a typical protein should the leftmost end be an amino group or a carboxylic acid group?

Explain.

25.33 Draw a diagram to illustrate how the secondary structure of a protein arises.

25.34 What are the two most common examples of secondary structures of proteins?

25.35 What type(s) of interactions may be present in forming the tertiary structure of a protein?

25.36 Draw a reaction to illustrate the formation of a disulfide linkage.

25.37 Draw all the dipeptides formed by a combining one molecule of phenylalanaine with one molecule of

cysteine.

25.38 Draw the tripeptide gly-ser-cys.

25.39 Using the three-letter abbreviations, list all possible tripeptides containing one molecule each of

phenylalanine, phe, leucine, leu, and aspartic acid, asp.

25.40 Draw the complete structural formula of the following tripeptides: (a) val-lys-gly; (b) phe-gly-ser;

(c) ser-val-phe; (d) ala-ala-gly; (e) asp-tyr-phe.


25.41 What property do all lipids have in common?

25.42 Define hydrophilic and hydrophobic.

25.43 Sketch the general structure common to all steroids.

25.44 How many chiral carbon atoms are in each of the following molecules? (a) cholesterol (b) testosterone

(c) estradiol


25.45 Define a fat and sketch a general structure for all fats.

25.46 What is the difference between a saturated fat and an unsaturated fat?

25.47 (a) What is margarine? (b) Why does margarine have a higher melting point than that of the oil from which it

was made?

25.49 Glycerol tristearate is a fat where all the fatty acid portions on the molecule come from stearic acid. Stearic

acid is an eighteen carbon saturated fatty acid. (a) Draw the structure of glycerol tristearate. (b) Draw the structure

of the products of the acid hydrolysis of glycerol tristearate. (c) Draw the structure of the products of the

saponification of glycerol tristearate.

25.50 How do detergents differ from soaps?

25.51 Why are soaps less effective in hard water?


25.53 What happens to food molecules, such as polysaccharides, fats, and proteins, during digestion?

25.54 What is an enzyme?

25.55 What do the letters in RNA and DNA represent?

25.56 What are the three components of all nucleic acids?

25.57 (a) What are the names of the nitrogen bases present in nucleic acids? (b) Name the nitrogen bases present in

RNA. (c) Name the nitrogen bases present in DNA. (d) Give the one-letter abbreviation of each.

25.58 (a) What is the structural difference between thymine and uracil? (b) Identify which of the nitrogen bases in

part (a) appears in DNA and which appears in RNA.

25.59 (a) What is the structural difference between ribose and deoxyribose? (b) Identify which of the sugars in part

(a) appears in DNA and which appears in RNA.

25.60 Identify which pairs of nitrogen bases can effectively form hydrogen bonds to each other.

25.61 (a) What are the components present in a nucleoside? (b) What are the components present in a nucleotide?

25.62 What are AMP, ADP, and ATP?

25.63 What does each of the following process produce? (a) replication (b) transcription (c) translation

25.64 (a) What is a codon? (b) How many codons are known? (c) Identify the codons that do not refer to any amino

acid.

25.65 Draw the structure of a nucleotide containing: ribose, uracil, and a phosphate group.

25.66 Write the base sequence for the DNA strand complimentary to a strand with the sequence: AATCGCGTA.

25.67 What are the structural differences between DNA and RNA?


Summary

25.79 The human body maintains a temperature of 37°C. What would happen to the rates of biological processes

should a person's body temperature drop to 27°C? This is a condition known as hypothermia.

25.80 Various processes can disrupt the secondary, tertiary, and quaternary structure of a protein. These processes

result in a primary structure, which does not function properly. A protein with a disrupted structure is said to be

denatured. Heat will denature some proteins as is apparent when egg white cooks. A change in the pH may also

denature a protein. This occurs when an acid, such as vinegar is added to milk. (a) Which amino acids have side

chains that are likely to be susceptible to acids? (b) If the amino acids in part (a) are affected, the most likely

disruption occur in the tertiary structure of the protein. Why?

25.81 Identify the amino acids present in the following protein fragment.

H2N CH C

CH3

N

O

CH C

CH2

N

O

CH2

C

NH2

O

CH C

CH2

N

O

CH2

S

CH3

CH C

CH2

OH

O

OH

H H H

Glycine, Asparagine, Methionine, Tyrosine

25.82 Identify a codon sequence necessary to produce the following protein fragment.

H2N CH C

CH2

N

O

CH C

CH2

N

O

CH2

C

NH2

O

CH C

CH2

N

O

CH2

SH

CH C

CH2

OH

O

H H H

OH

25.83 Partial hydrolysis of a nonaapeptide, a protein fragment containing nine amino acids, produced a mixture of

smaller fragments. Chromatographic separation of the smaller fragments identified them as ala-gly, phe-ser-ala,

phe-phe, gly-phe-ala, ala-cys, and cys-phe. Complete hydrolysis and separation found the nonapeptide contained 2

cys, 1 ser, 2 ala, 3 phe, and 1 gly. What was the original amino acid sequence in the nonapeptide?

25.84 Why can the carbonyl carbon in a monosaccharide never be chiral?

25.85 Identify the chiral carbon atoms in mannose:

OH

OH

H

OH

OH

HH

OH

CH2

H

HO

25.86 Identify the chiral carbon atoms in the following protein fragment:

H2N CH C

CH2

N

O

CH C

CH2

N

O

CH2

C

NH2

O

CH C

CH2

N

O

CH2

SH

CH C

CH2

OH

O

H H H

OH

25.87 A solution of a protein contains 0.382 grams of the material in 1.00 L of solution. At a temperature of 25°C,

this solution has an osmotic pressure of 1.07 torr. Calculate the molar mass of the protein.


25.88 A sample of the hormone vasopressin contains 0.150 grams of this compound in 250. mL of solution. This

solution has an osmotic pressure of 10.4 mmHg at 30.°C. Determine the molar mass of vasopressin.

25.89 A solution contains 1.25 grams of a saturated fatty acid in 75.0 grams of cyclohexane. This solution has a

freezing point that is 1.30°C lower than that of pure cyclohexane. The freezing point depression constant, Kf, for

cyclohexane is 20.0°C/m. Determine the molecular formula of this fatty acid.

25.90 The chemical reduction of pyruvic acid in the laboratory produces equal amounts of D-lactic acid and L-lactic

acid. The biochemical reduction of pyruvic acid in your body yields only L-lactic acid. Why are the results of this

reduction different?

25.91 The compound L-Dopa, shown below, is useful in the treatment of Parkinson's disease. The isomer D-Dopa is

ineffective in the treatment of this disease. Explain.

OH

OH

HO

O

H NH2

25.92 A wax, like a fat, is an ester. Beeswax contains a number of waxes. Hydrolysis of beeswax results in a

variety of long chained alcohols with from 24 to 36 carbon atoms. The fatty acid chains have up to 36 carbon

atoms. The esters are mixed with hydrocarbons containing from 21 to 33 carbon atoms. The alcohols and acids

have even numbers of carbon atoms while the hydrocarbons have odd numbers of carbon atoms. All the carbon

chains are unbranched. Saponification of a sample of beeswax results in a number of products. One of the products,

isolated through a chromatographic procedure, is soluble in cyclohexane. A total of 1.57 grams of this component

dissolves in 65.0 grams of cyclohexane to produce a solution with a melting point 1.22°C lower than pure

cyclohexane. The freezing point depression constant, Kf, for cyclohexane is 20.0°C/m. Chemical analysis of this

compound shows that it does not contain oxygen. Determine the molar mass of this compound and its molecular

formula.

25.93 Some phospholipids have structures similar to fats except that one of the ester groups is a phosphate ester.

The generic structure for many phospholipids is:

CH2

CH

CH2

O

O

O

C

C

P

O

R

O

R'

O

OH

O

R"

Acid hydrolysis of the compound having this generic formula will yield what products?

25.94 Two of the hormones produced by the pituitary gland are oxytocin and vasopressin. Oxytocin causes smooth

muscles to contract, and vasopressin regulates the excretion of water by the kidneys and affects blood pressure. The

structures of these hormones are:

IleGln

TyrAsn

CysCys

SSPro

Leu

Gly

H2N

oxytocin

PheGln

TyrAsn

CysCys

SSPro

Arg

Gly

H2N

vasopressin

The structures are obviously similar with a disulfide linkage giving the tertiary structure. (a) What are the

differences between these hormones? (b) Classify the amino acids that are different as polar, nonpolar, acidic, or

basic. (c) The receptor site for one of these hormones has a negative charge at physiological pH. Which of these

hormones is more likely to bind to this site?


25.95 Vitamin B-12 is unique among vitamins in that it contains a metal ion. The metal is cobalt, which may be in a

+1, +2, or +3 oxidation state. This compound is important in many biological processes. It is an essential vitamin

for all mammals, and it is necessary in their diets because animals cannot synthesize this compound. (a) A solution

containing 0.0870 grams of vitamin B-12 in 175 mL of water has an osmotic pressure of 0.206 torr at 25 °C.

Determine the molecular weight of vitamin B-12. (b) Analysis of a sample of vitamin B-12 found 55.83% C, 6.54%

H, 4.35% Co, 14.47% N, 16.53% O, and 2.29% P. What is the molecular formula of vitamin B-12?

(a) 4.49 × 104 g/mol (b) C63H88CoN14O14P

25.96 Name or give the formula for each of the following. (a) Tritium (b) Potassium borohydride (c) Diborane

(d) Diamond (e) Calcium acetylide (f) RbO2 (g) BO33- (h) B4O7

2- (i) SiC (j) Na2SiF6 (k) Buckminsterfullerene

(l) Tetraboric acid (m) BaO2 (n) BN (o) H2SiF6

25.97 Name or give the formula for each of the following. (a) Cesium superoxide (b) Potassium orthoborate

(c) Potassium tetraborate (d) Methane (e) Potassium hexafluorosilicate (f) 1H (g) BH4– (h) B2O3 (i) C(gr) (j) C2

2–

(k) Deuterium (l) CaB4O7 (m) Ca3(BO3)2 (n) T (o) Sodium peroxide

25.98 Name or give the formula for each of the following. (a) Sodium orthoborate (b) Pyrophosphate ion

(c) Methylammonium ion (d) Barium acetylide (e) Ammonia (f) NH2Cl (g) Na2B4O7 (h) H5P3O10 (i) (NH4)2SO4

(j) NH4Cl (k) Orthophosphoric acid (l) Hydroxylamine (m) N2H2(CH3)2 (n) Sodium borohydride

(o) Diethylammonium ion

25.99 Name or give the formula for each of the following. (a) Sodium tripolyphosphate (b) Ammonium chloride

(c) C(dia) (d) Methylhydrazine (e) Potassium tetraborate (f) N2H4 (g) (C2H5)2NH (h) K3BO3 (i) B2H6 (j) (NH4)4P2O7

(k) Ammonium sulfate (l) H3PO3 (m) (HPO3)3 (n) Hexafluorosilicic acid (o) Chloramine

25.100 Name or give the formula for each of the following. (a) Dioxygen (b) Metaperiodic acid (c) Disulfide ion

(d) Diethylamine (e) Silicon carbide (f) XeF6 (g) H2SeO3 (h) I3– (i) H5P3O10 (j) Na3BO3 (k) Ca(I3)2 (l) Sodium

hexafluorosilicate (m) Dimethylhydrazine (n) Calcium acetylide (o) Orthotelluric acid

25.101 Name or give the formula for each of the following. (a) Ozone (b) Potassium triiodide (c) Selenic acid

(d) Sodium tripolyphosphate (e) Magnesium orthoborate (f) H2S2O7 (g) (C2H5)2NH2+ (h) FeS2 (i) H5IO6 (j) C2

2–

(k) XeF4 (l) H6TeO6 (m) Triiodide ion (n) Trioxygen (o) Selenous acid

25.102 Give the formula for each of the following. (a) dichlorobis(ethylenediamine)platinum(IV) bromide

(b) Potassium diaquatetrabromovanadate(III) (c) Sodium tetrabromo(ethylenediamine)cobaltate(III)

(d) Dichlorobis(ortho-phenanthroline)iron(III) perchlorate (e) pentaaquabromomanganese(III) sulfate

(f) Tris(bipyridyl)ruthenium(II) nitrate (g) Hexaamminenickel(II) tris(oxalato)chromate(III)

(h) Hexaamminechromium(III) nitrate (i) Bis(ethylenediamine)zinc(II) tetraiodomercurate(II)

25.103 Name each of the following: (a) [Pd(en)2][Cr(NH3)2Br4]2 (b) [Zn(NH3)4]2+ (c) [Co(en)(NH3)2Br2]Cl

(d) [Ru(H2O)Cl5]2– (e) [PtCl4(en)] (f) trans-[Pt(NH3)2H(Br)] (g) [Cr(NH3)4Br2] (h) Cs[Cr(C2O4)2Cl2] (i) cis-

[Cr(NH3)4(H2O)2]3+ (j) K3[IrCl5(S2O3)]

chemistry 134 problem set introduction

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