Solution Manual for Chemistry , 10th Edition

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Solution ManualTo AccompanyChemistry 10theditionSteven S. ZumdahlUniversity of IllinoisSusan A. ZumdahlUniversity of IllinoisPrepared by:Donald DeCosteUniversity of Illinois at Urbana-Champaign

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Scheduling and Laboratory3EXPERIMENTAL CHEMISTRYTenthEditionbyJohn G. LittleIntroductionPreface.....................................................................................................................viiLaboratory Glassware and Other Apparatus..............................................................ixSafety in the Chemistry Laboratory..........................................................................xviSafety Quiz...........................................................................................................xxviiExperiments1.The Determination of Mass...................................................................................12.The Use of Volumetric Glassware........................................................................93.Density Determinations......................................................................................214.The Determination of Boiling Point....................................................................295.The Determination ofMelting Point....................................................................396.The Solubility of a Salt.......................................................................................477.Identification of a Substance..............................................................................558.Resolution of Mixtures 1: Filtration and Distillation............................................639.Resolution of Mixtures 2: Chromatography.........................................................7310.Stoichiometry 1: Limiting Reactant...........................................................................................8511.Stoichiometry 2: Stoichiometry of an Iron(III)-Phenol Reaction..................................................9312.Composition 1: Percentage Composition–The Empirical Formula of Magnesium Oxide.........10113.Composition 2: Hydrates and Thermal Decomposition.............................................................10914.Preparation and Properties of Hydrogen and Oxygen Gases...................................................12115.Gas Laws 1: Molar Volume and The Ideal Gas Constant........................................................13116.Gas Laws 2: Graham’s Law....................................................................................................13917.Calorimetry............................................................................................................................14718.Heats of Reaction and Hess’s Law...................................................................15919. Spectroscopy 1: Spectra of Atomic Hydrogen and Nitrogen...........................17120. Spectroscopy 2: Emission Spectra of Metallic Elements.................................18321. Molecular Properties: Molecular Shapes and Structures.................................19522. Properties of Some Representative Elements..................................................20123. Classification of Chemical Reactions...............................................................21724. Colligative Properties 1:Freezing Point Depression and the Determination ofMolar Mass............................................................................................................23525. Colligative Properties 2: Osmosis and Dialysis................................................24526. Rates of Chemical Reactions...........................................................................253

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4Scheduling and Laboratory27. Chemical Equilibrium 1: Titrimetric Determination of an Equilibrium Constant26328. Chemical Equilibrium 2: Spectrophotometric Determination of an EquilibriumConstants...............................................................................................................27329. Chemical Equilibrium 3: Stresses Applied to Equilibrium Systems..................28130. The Solubility Product of Calcium Iodate........................................................29131. Acids, Bases, and Buffered Systems...............................................................30132. Acid–Base Titrations 1: Analysis of an Unknown Acid Sample.......................31733. Acid–Base Titrations 2: Evaluation of Commercial Antacid Tablets.................32734. The Determination of Calcium in Calcium Supplements..................................33535. Determination of Iron by Redox Titration.........................................................34336. Determination of Vitamin C in Fruit Juices.......................................................35137. Electrochemistry 1: Chemical Cells..................................................................36138. Electrochemistry 2: Electrolysis.......................................................................37139. Gravimetric Analysis 1: Determination of Chloride Ion.....................................38140. Gravimetric Analysis 2: Determination of Sulfate Ion.......................................39141. Preparation of a Coordination Complex of Copper(II).....................................40142. Inorganic Preparations 1: Synthesis of Sodium Thiosulfate Pentahydrate.......40943. Inorganic Preparations 2: Preparation of Sodium Hydrogen Carbonate..........41744. Qualitative Analysis of Organic Compounds....................................................42545. Organic Chemical Compounds........................................................................43546. Ester Derivatives of Salicylic Acid....................................................................44747. Preparation of Fragrant Esters.........................................................................45748. Proteins............................................................................................................46549. Enzymes..........................................................................................................47350. Polymeric Substances 1: Amorphous Sulfur....................................................48151. Polymeric Substances 2:Preparation of Nylon................................................48752. Qualitative Analysis of Selected Cations and Anions.......................................493AppendicesA. Plotting Graphs in the Introductory Chemistry Laboratory................................503B. Errors and Error Analysis in General Chemistry Experiments...........................505C. Vapor Pressure of Water at Various Temperatures...........................................509D. Density of Water at Various Temperatures........................................................510E. Solubility Information.........................................................................................511F. Properties of Substances...................................................................................512G. Concentrated Acid Base Reagent Data...........................................................516

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5PART II: CHAPTER DISCUSSIONSCOURSE CONTENT: DESCRIPTIVE CHEMISTRY AND CHEMICAL PRINCIPLESThere has been much discussion and an overall agreement that more descriptivechemistry needs to be included in the general chemistry course. This has beenone ofthe major considerations in the recent curriculum reform movement. How to accomplishthis goal has resulted in much less agreement. General chemistry has to move beyonda rote presentation of the facts and principles of chemistry. We feel that the facts(descriptive chemistry) and principles must be an integrated whole that is aimed atfulfilling the basic goals of the course.TheChemistrytextbook follows what has been for the past several years a fairlytraditional order of topics. The chapters on principles (1-18) are followed by nuclearchemistry (19),a more descriptive chapteronchemistry of the elements (20),a chapteron transition metals and coordination chemistry (21),and organic chemistry andbiochemistry (22). Different instructors will have different approaches to descriptivechemistry and, in fact, will probably give greatly different definitions of descriptivechemistry. We have organized the text this way because we feel this order offersinstructors the greatest flexibility.In the fifties and early sixties, general chemistry texts essentially went through theperiodic table; principles were introduced as needed. Observations were made of theproperties and reactions of matter; then we used chemical principles to try tosystematically organize and understand those observations. This approach may closelyresemble how chemists do chemistry. But is it the best way to learn chemistry? Wedon’t think so. We are all driven by curiosity as to how things work, and an unfortunateconsequence of marching through the periodic table is to dull the students’ curiosityabout why things happen, replacing it with the impression of chemistry as a vast body ofunrelated facts and reactions that must be memorized. This approach tends to makelearning chemistry both frustrating and uninteresting. In the sixties, a principles-dominated outline began to be used in texts, but we overcompensated. The principleswere emphasized. The framework was put in place to deal with chemical facts, but alltoo often the facts were left out. Only the reaction A + BC seems to have beencovered. Learning chemistry was still frustrating, but in a different way from before.We must integrate the two approaches. In truth, the principles of chemistry anddescriptive chemistry are two indispensable parts of the whole. In particular, we mustuse all of our resources in presenting reactions. Can we honestly expect a neophyte tochemistry to appreciate the difference between2H2+ O22H2OandA + BC

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6Chapter Discussionsif they are only written on the blackboard or in a book? We must make use of all of thesenses. Students can appreciate what 2H2+ O22H2O means if the instructor ignitesa mixture of hydrogen and oxygen. They can see, hear, and sometimes feel what thatequation represents.We can’t just talk about reactions. We must show the students reactions or, better yet,let them see the reactions on their own in the laboratory. We should choose moreexperiments for the labs in which students observe reactions or synthesize compoundsand do fewer experiments that involve a measurement that confirms (often poorly) someprinciples discussed in class. The use of lecture demonstrations also allows students tosee chemistry. These two approaches represent powerful tools for presentingreactions.In the text we have tried to emphasize the framework of principles by which chemicalfacts can be organized. There are several important features that facilitate theintegration of fact and principle.1.To assist instructors in providing interesting and relevant classroomdemonstrations of chemical phenomena for their students, theInstructor’sAnnotated Editionindicates via the marginal icon (flask with stirring rod) directreferences to more than 750 demonstrations from several authoritative sources.2.Chapter 4 presents an early, thorough discussion of solution reactions, using theclassifications that chemists typically use (acid-base, precipitation, and oxidation-reduction).3.The illustrations and color photography show the student many reactions. Thismerely supplements lecture demonstrations and labs but right in the text studentscan see the beauty of the reactions and substances composing this body ofknowledge called chemistry.4.Real examples are used to illustrate chemical principles.5.Much of the discussion in the descriptive chapters is organized according to aframework of the early chapters. For example, Chapters 19-21 emphasizeperiodic relationships; the discussion of polymers in Chapter 22 emphasizesstructure-property relationships.6.Many of the end-of-chapter exercisesdeal with reactions or substances that areimportant to our lives. In the problem statement, these relationships arementioned.7.The end-of-chapter exercises in Chapters 19-22 are designed to illustrate therelationships between the descriptive material and the principles chapters.Energy relationships (particularly bond energies), periodicity, structure, and

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Chapter Discussions7equilibrium are emphasized.In the following chapter discussions, we will further point out how descriptive materialfrom the later chapters can be brought into discussion of principles. The organization ofthe text allows each instructor greater flexibility in presenting the material. The uniquefeatures of the text allow for a more thorough integration of principles and facts.Chemistry isn’t just principles; it isn’t just reactions and properties of elements andcompounds. It is an amalgam of both. In the text we have tried tobring both together,while allowing instructors maximum flexibility.CHAPTER ONE: CHEMICAL FOUNDATIONSChapter Learning Goals:Section One:To appreciate the importance of creative problem solving.Section Two:To identify the principal operations andlimitations of thescientific method.Section Three:To describe the SI system of units and prefixes.Section Four:To identify causes of uncertainty in measurement.To show how significant figures are used.To compare precision and accuracy in measurement.Section Five:To show how to determine the number of significant figures in acalculated result.Section Six:To show a general method of solving problems.Section Seven:To show how to convert units between the English and metricsystems.SectionEight:To demonstrate conversions among the Fahrenheit, Celsius,and Kelvin temperature scales.SectionNine:To illustrate calculations involving density.SectionTen:To show how matter can be classified into subgroups.The time spent on the first two chapters may show the greatest differences from class toclass. Much of the first chapter is probably review for students with good high schoolmath and chemistry backgrounds.This chapter lays the foundation for dealing with measured quantities and performingcalculations. Instructors wishing to treat uncertainty in greater detail can discuss thesection Uncertainties in Measurements, from Appendix One. This section of theappendix could also be used in the laboratory.Instructors should be careful to point out that the Sample Exercises are worked in adifferent manner from what students should do. Intermediate answers are rounded offto show the correct number of significant figures at each stage of the calculation. Theserounded values are then used to complete the calculation. Students, when working

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8Chapter Discussionsproblems, should round only at the end. In theSolutions Guide, we have followed thesame convention as the text and have rounded at the points where intermediateanswers are shown. If this would result in excessive round-off error, we have carriedextra digits and noted this in the solution. Most commonly we havecarried extra digitswhen solving two simultaneous equations and in doing equilibrium problems.One of the primary reasons for the discussion of units is to introduce the student to theuse of dimensional analysis as a problem-solving technique. In our classes weemphasize using units as a check. Some students may tend to be sloppy about units incalculations because they are familiar with the quantities being used. The nice featureof dimensional analysis is that it works well even if we don’t have a good intuitive feel forthe quantities and units encountered in a particular problem. On the other hand, it isimportant to remember that correct use of dimensional analysis does not mean that astudent has a conceptual understanding of the chemical concepts in a particularproblem. One of the reasons we have included theActive LearningQuestions is so thestudents have a chance to vocalize their ideas about the concepts covered, and wehave a better chance to try to understand these.Teaching TipsConnecting the real world to the atomic/molecular world is difficult for manystudents. Encourage students to visualize theatomic/molecular world bydrawing pictures. The student responses will give you an idea of their enteringperceptions. Don’t worry about some of the details such as subatomic particleshere. For example, point out the representations of hydrogen, oxygen,and waterin Section 1.1.Consider spending some time discussing Figure 1.4. Some students believe thattheories become laws once a theory has been accepted, but this figure clearlyshows laws and theories as two separate entities. Theories never become laws.Laws tell us what happens, and theories are our attempts to explain why. Interms of a “scientific method”, students should understand that much of scientificthinking is logical or analytical thinking.To illustrate uncertainty in measurements bring to class several pieces ofglassware used to measure volume. Show the students the glassware. Drawthe shape of a representative piece of the glassware on the board marking thedrawing to show the meniscusof a liquid between two of the lines.Since thevolume falls between the lines there is uncertainty involved in the reading. Havethe students write down the result of the measurement. Then compare results.The number recorded for this measurement will vary from student to student onlyin the last digit. Thus there is only one uncertain digit.A relatively easy way to determine the number of significant figures in a numberis to write the number in scientific notation. In this notation there are no leading

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Chapter Discussions9zeros, all zeros in the middle of the number are significant, and trailing zeros areonly recorded if they are significant.To illustrate the difference between a physical and chemical change comparerepresentation (b) and (c) from Figure 1.11 (a physical change from liquid waterto gaseous water) to the chemical change when liquid water is electrolyzed toproduce gases H2and O2. Use the equation in the text to represent theelectrolysis of water chemical change.CHAPTER TWO: ATOMS, MOLECULES, AND IONSChapter Learning Goals:Section One:To give a brief account of early chemical discoveries.Section Two:To describe and illustrate the laws of conservation of mass,definite proportion, and multiple proportions.Section Three:To describe Dalton's theory of atoms and show the significanceof Gay-Lussac's experiments.Section Four:To summarize the experiments that characterized the structureof the atom.Section Five:To describe features of subatomic particles.Section Six:To introduce basic ideas of bonding in molecules.To show various ways of representing molecules.Section Seven:To introduce various features of the periodic table.Section Eight:To demonstrate how to name compounds given their formulasand to write formulas given their names.Chapter 2 is the second of the two introductory chapters. As with Chapter 1, the lengthof time spent on this chapter will depend on the background of the students. It goesover the background in chemical topics that provide the foundation for the rest of thecourse.This chapter takes a historical approach to the development of chemistry. It begins witha discussion of the discoveries leading to Dalton’s atomic theory, continues with theexperiments elucidating the structure of the atom, and ends with chemical nomenclatureand an introduction to the periodic table.The section on nomenclature and formula writing in the laboratory manual by Hall canbe effectively used at this point. Instructors may also wish to discuss nuclear decay inmore detail at this time. Chapter 19, Nuclear Chemistry, can be covered here ifinstructors prefer. However, if the chapter is presented early, the sections on decaykinetics and thermodynamic stability of the nucleus would have to be delayed and couldbe covered with Chemical Kinetics (Chapter 12) or Spontaneity, Entropy, and FreeEnergy (Chapter 17).

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10Chapter DiscussionsTeaching TipsConsider using an analogy between the letters in the alphabet and words and theatoms of the various elements and compounds. In the same way we can makemany words from 26letters, millions of compounds can be made using only the100 or so elements. Putting the same elements together in different ways resultsin different compounds with very different properties. For example, the words“ADD” and “DAD” each consist of the same letters. However, the order of theletters gives them different meanings. This discussion will help later (forexample, in Chapter 3 you can use this analogy to help the students understandwhy we do not change subscripts in a chemical formula).Students often have difficulty understanding what is meant by chemical termssuch as element. It is important to emphasize the various ways we use the termelement in chemistry. When we say element we might mean a single atom, wemight mean a molecule such as N or O or we might mean a large sample suchas a bar of aluminum or the graphite (“lead”) in a pencil. Be sure to help studentsrealize that chemists use terms in many ways and they need to look at thecontext of the word to be sure that they understand what is intended. Chemistsare accustomed to thinking about things at the macroscopic and microscopiclevel simultaneously. Encourage students to think about what they read and toconsider how the terms are being used.Sections 2.2 and 2.3 provide another opportunity to discuss the differencebetween laws and theories. Dalton’s atomic theory is relatively simple in scope,and this is as it should be. John Dalton was trying to explain laws such as thelaw of constant composition. The success of this makes the model successful,but not absolutely correct (a model is always a simplification). For example,Dalton’s theory does not explain questions such as “Why/how do atoms sticktogether to form molecules?”, and “Why/howdo molecules stick together to formliquids and solids?”. But no model answers all questions.The discovery of the proton leads to an excellent example of how and whymodels change. Dalton’s model of the atom did not account for isotopes sincehe assumed that all atoms of an element are exactly alike. As knowledgeexpands models change to accommodate the new informationStudents should learn the names of the common groups on the periodic table(see Figure 2.19). This will simplify your discussion later in the course. Reasonsbehind the structure of the periodic table and trends it shows are not introduceduntil Chapter7. Focus the students’ attention of the location of groups ofelements and the separation between metals and nonmetals.

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Chapter Discussions11Emphasize that ions are formed by the gain and loss of electrons. This is difficultfor students because they tend to think about gain and loss using positivenumbers. It is very important that they understand from the beginning that apositive ion is formed by losing electrons and a negative ion is formed by gainingelectrons. In chemistry the positive number (protons) remain the same.Be sure to point out to students that atoms do not spontaneously gain or loseelectrons. The gain and loss of electrons are always paired. One atom loseselectrons simultaneously with another atom gaining electrons.Stress that learning to name compounds does not consist of memorizing aseemingly endless list of chemicals. There are systematic rules for namingcompounds and, by knowing only a few rules, the students can name most anycompound they will encounter in this course.The students should also understand that we keep the rules as simple aspossible (much the same way we keep our scientific models as simple aspossible). We only add complications (such as prefixes and Roman numerals)when it is required for clarity. Forexample, the name sodium(I) chloride is notnecessarily wrong, merely redundant. Therefore, the name sodium chloridesuffices.The students can determine the charges of all the ions in Table 2.3 from theirpositions on the periodic table, with the exception of Ag+. The students shouldappreciate that the periodic table contains a great deal of information.It is avaluable resource.Make a clear distinction between Type I and Type II cations. This reinforces thereasons for the Roman numerals in the names of the compounds. Ask students,for example “How is ‘iron oxide’ different from ‘magnesium oxide’?”Students often make the mistake that the Roman numeral tells us thatnumber ofions present in the compound. This is because in many cases the Romannumeral is the same as one of the subscripts. Use an example such as iron(II)oxide with the formula FeO to show that the Roman numeral relates only to thecharge.For Type III compounds, consider starting with examples such as “NO” and“NO2”. Ask the students to name these compounds. This will lead naturally to adiscussion of prefixes. Like Roman numerals, prefixes are used only for clarity.For example, there are two possible “carbon oxides” and we must use prefixes todifferentiate between these (carbon monoxide and carbon dioxide). We onlymake the rules more complicated when it is required.

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12Chapter DiscussionsFigure 2.22 and Figure 2.23 are convenient flow charts to help studentssystematically name compounds. Students need to be encouraged to use thistype of device to help them think through a difficult problem.Students should think of polyatomic ions as a “unit”. Many students know that ifNaCl is dissolved in water, Na+ions and Cl-ions are present. However, studentsare often confused about the ions present when NaNO3is dissolved. Instead ofthe NO3–ion students often think individual nitrogen and oxygen ions arepresent.CHAPTER THREE: STOICHIOMETRYChapter Learning Goals:Section One:To describe the modern atomic mass scale and explain howatomic masses are determined experimentally.Section Two:To explain atomic mass and its experimental determination.SectionThree:To explain the importance of themole concept.To show how to convert among moles, mass, and number ofparticles for a given sample.SectionFour:To show how to calculate values for molar mass.To show how to convert among molar mass, moles, and numberof particles in a given sample.Section Five:To describe a conceptual problem solving approach tochemistry.SectionSix:To demonstrate the calculation of the mass percent of a givenelement in a compound.SectionSeven:To demonstrate the calculation of the empirical formula of acompound.To show how to obtain the molecular formula, given theempirical formula and the molar mass.Section Eight:To identify the characteristics of a chemical reaction and theinformation given by a chemical equation.SectionNine:To show how towrite a balanced equation to describe achemical reaction.SectionTen:To show how to calculate the masses of reactants and productsusing the chemical equation.Section Eleven:To show how to recognize the limiting reactant.To demonstrate the use ofthe limiting reactant in stoichiometriccalculations.Chapter 3 deals with the fundamental measurement unit in chemistry, the chemicalmole. The law of conservation of mass is the unifying principle of the chapter. Thechapter deals with compounds firstand then reactions.

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Chapter Discussions13An important point that can be made in lecture is the convenience of the mole as a unit.Chemical reactions, chemical formulas, and structures of molecules focus on numbersof atoms, molecules, or ions. Atoms are so small that we cannot see them to countthem. Even if we could see them, the numbers that we would encounter are so large itwould take an immense period of time to count them. The mole provides a unit thatallows us to connect the number of atoms (what we are interested in knowing) tosomething we can measure (such as mass). This approach makes it easy to introducemolarity in Chapter 4 as a similar type of unit. Molarity enables us to convert betweenthe quantity we need to deal with in reactions (numbers of moles) and the mostconvenientmeasurement we can make for the amount of a solution (volume).Teaching TipsSection 3.5 discussing a conceptual problem solving strategy for the students. Itoffers suggestions on thinking about the problem so that they learn how to solveproblems in general, not simply use the right algorithm for a given problem. It iswell worthwhile to spend some time with the students discussing this.The concept of counting atoms from the mass of a sample can be difficult forstudents to understand.Develop the idea slowly. Many students learn to workthe problems without really understanding this section. This leads to trouble laterwhen they tryto solve more complex problems. Return to the candy analogy toassist in developing this concept. You can also use a hardware store analogy ofbuying nails by weight, or a banking analogy of counting coins into rolls byweighing them.The idea that different samples with identical mass ratios contain the samenumber of objects is adifficult one for the students. Consider using the followingexample.Suppose we have two blocks, a red block and a yellow block. The red blockweighs 1.0 ounce, and the yellow block weighs 4.0 ounces. Now suppose wehave 16 of each color block. What is the mass of each sample? The sample ofred blocks weighs 16.0 ouncesand the sample of yellow blocks weighs 64.0ounces. But note that 16.0 ounces is also 1.0 pound, thus 64.0 ounces is 4.0pounds. The relative masses of the blocks are:one blocksixteen blocksred1.0 ounce1.0 poundyellow4.0 ounces4.0 poundsThe relative masses stay the same (1:4) but the units are changed.You can make an analogy between ounces and pounds in this case and amu’sand grams in the case of the periodic table. In the blocks example, the number16 is analogous to Avogadro's number.

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14Chapter DiscussionsIt is important for students to realize that a mole really describes the number ofobjects present, just as a dozen means 12. This connection between a mole asa number and the mass of a sample is essential because we count atoms byweighing samples containing large numbers of them.Notice that the text defines the mole differently from the SI definition: “the mole isthe amount of substance that contains as many entities as there are in exactly0.012 kg of carbon 12”. The text definition is “the mole is the number equal tothe number of carbon atoms in 12.01 grams of carbon”. We believe the textdefinition is easier for the students to understand at this point as it more stronglyemphasizes both that the mole is fundamentally a number and that the periodictable contains average atomic masses.In Section 3.9 students learn that a chemical reaction involves rearrangement ofthe elements. It is very helpful to use simplemodels of the reactants that youcan take apart and rearrange to show students that atoms are conserved in achemical reaction (toothpicks and gumdrops or clay work well if you do not haveball-and-stick models). Use the molecular level graphics of chemical reactions inSection 3.9 to aid your discussion.The text lists three steps for writing and balancing chemical equations. Step 3tells the students to start “with the most complicated molecule(s).” For thereactionC2H5OH(g) + O2(g)CO2+ H2O(l)consider balancing the equation in the following way:C2H5OH(g) + O2(g)2CO2+ H2O(l)(carbon is balanced)C2H5OH(g) + 2O2(g)2CO2+ H2O(l)(oxygen is balanced)C2H5OH(g) + 2O2(g)2CO2+ 3H2O(l)(hydrogen is balanced, oxygennowunbalanced)C2H5OH(g) + 3O2(g)2CO2+ 3H2O(l)(the equation is balanced)Notice that we eventually get the right answer (so trial and error is a reasonablemethod). However, it would have been better not to have balanced the oxygenas the second step. This is why Step 3 tells the students to start with the mostcomplicated molecule. Encourage the students to leave the elements anddiatomic molecules until last when balancing equations. This simplifies the trialand error process.

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Chapter Discussions15Emphasize that the coefficients in a balanced equation represent the moleratios.An individual coefficient is meaningless (just as an amount of a single ingredientin a recipe is meaningless without the other amounts).It is useful to use a real life example whenintroducing limiting reactants.Students seem to relate easily to kitchen experiences yet have difficulty applyingthe same concepts to chemical equations. By working with real life experiencesfirst and addressing the concepts in this context students have an opportunity tobecome comfortable with the concept before applying it to an unfamiliar setting.Using a concrete analogy (such as making sandwiches) helps the students betterunderstand the calculations. This is especially true for problems involving alimiting reactant. Problems in which a reactant is limiting are conceptually thesame as the cases in which no reactant is limiting. However, students seem tohave much more difficulty dealing with limiting reactants. For example, somestudents calculate the amount of product that could be made from each reactantand then add them. Or, they add themoles of each given reactant for the total.They almost never make these types of mistakes in the analogies, though, so it ishelpful to be very explicit about these pointsUse molecular diagrams like those shown in Section 3.11 of the text to makecounting molecules more concrete (and similar to real life experience).CHAPTER FOUR: TYPES OF CHEMICAL REACTIONSAND SOLUTION STOICHIOMETRYChapter Learning Goals:Section One:To show why the polar nature of water makes it an effectivesolvent.Section Two:To characterize strong electrolytes, weak electrolytes, and non-electrolytes.Section Three:To define molarity and demonstrate calculations involving thecomposition of solutions.Section Four:To introduce several types of solution reactions.Section Five:To show how to predict whether a solid will form in a solutionreaction.Section Six:To describe reactions in solution by molecular,complete ionic,and net ionic equations.Section Seven:To demonstrate stoichiometric calculations involvingprecipitation reactions.

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16Chapter DiscussionsSection Eight:To show how to perform calculations involved in acid-basevolumetric analysis.Section Nine:To characterize oxidation-reduction reactions.To describe how to assign oxidation states.To identify oxidizing and reducing agents.Section Ten:To describe the oxidation statesmethod for balancing oxidationreduction reactions.The placement of this chapter differs from its placement in many general chemistrytexts. We feel there are several good reasons for including these topics at this pointand many other texts have come to agree withus.A thorough discussion of the types of reactions in solution allows for an earlyintroduction of descriptive chemistry. Descriptive chemistry is very important in the first-year course. Although virtually every instructor will give a different definitionofdescriptive chemistry and a different list of topics to be covered, a thorough discussionof chemical reactions in aqueous solution is probably central to all of those definitionsand lists. Hence, the expanded discussion of reactions at an early point.Several years ago texts began introducing molarity and solution stoichiometry in thestoichiometry chapter to give flexibility to the lab program. The expanded discussion inChapter 4 gives a great deal of flexibility to the lab. The table on schedulinglecturesand labs in Part I amply illustrates this point. The stoichiometry for all of the commontypes of solution reactions is discussed. Students can begin to see all types ofreactions in the lab very early in the course. The expanded discussion inthe text willhelp the student focus on what is actually going on in a solution, what species are reallypresent, and how they interact with one another.A chemical equation written on the page of a textbook probably has little meaning to aneophyte. Students must have the opportunity to see reactions. Lecturedemonstrations and the color photographs in the text help, but the most useful place forseeingchemistry is by doing chemistry in the laboratory. Chapter 4 expands the vista ofwhat can be done in the lab.In Chapter 4, reactions are classified as precipitation, acid-base, and oxidation-reduction. The use of terms such as metathesis, combination, displacement, doubledisplacement, and so forth, is avoided because these terms are not used by chemists.Theseterms can be confusing. For example, areAgNO3(aq) + HCl(aq)AgCl(s) + HNO3(aq)andAg+(aq) + Cl-(aq)AgCl(s)the same or different? Is the first a double displacement and the second acombination? We strongly feel reactions should be classified on the basis of what can

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