Solution Manual For Introduction to Chemistry, 4th Edition

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Bauer, Birk, and MarksInstructor AnnotationsChapter 1Section 1.1 Compounds•Teaching TipAsk students to generate a list of examples of matter in the room. In small groups, havethem classify the examples as mixtures or pure substances. Then require them toclassify the pure substances as elements or compounds. Lastly, have them classify themixtures as heterogeneous or homogeneous.Section 1.1 Representations of Matter•Teaching TipProvide opportunities for students to draw pictures of matter on a particulate levelwhenever possible.Section 1.1 Representations of Matter•Teaching TipGiven particulate representations, students often mistake diatomic elements forcompounds. In order for a molecule to be classified as a compound, there must be twoor more atoms of different elements.Section 1.1 States of Matter•MisconceptionMany students believe that there is no molecular motion in solids.Section 1.1 States of Matter•MisconceptionMany students think that condensation on the outside of a glass containing a coldbeverage comes from the liquidinsidethe glass.Section 1.2 Physical and Chemical Changes and Properties of Matter•Teaching TipRather than teaching students to convert from English to SI units, provide them withpractical examples to allow them to think in the SI system. Tell them that a small paperclip has a mass of approximately 1 gram.Section 1.2 Physical and Chemical Changes and Properties of Matter•Teaching TipEmphasize the value of dimensional analysis and that it can be used throughout thecourse to solve problems.Section 1.2 Volume•Teaching TipTell students that a milliliter of water fills a small thimble.Section 1.2 Density•Demonstration

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Materials: 2 index cards, four 250 mL Erlenmeyer flasks, red and blue food dye, hot andcold water.Preparation: Add a few drops of red food coloring to two 250 mL flasks. Add hot wateruntil the flasks are completely filled. Swirl to mix the food coloring. Repeat for the othertwo flasks using blue food coloring and cold water.Presentation: Place the index card over the top of one of the flasks containing the hot(red) water. Carefully invert the flask and place it on top of one of the flask containingcold (blue) water. Remove the card and ask the students to observe. Have them predictwhat would happen if the flask containing the cold (blue) water was placed on the onecontaining the hot (red) water. Repeat the demonstration with the flask with hot water onthe bottom. Ask the students to explain their observations. They should discuss howthe density of water changes with temperature on a macroscopic and particulate level.Safety: Wear safety goggles in case glassware breaks.Section 1.2 Density•DemonstrationsShakhashiri, Bassam Z., “Density and Miscibility of Liquids,”Chemical Demonstrations:A Handbook for Teachers of Chemistry,Vol. 3 (The University of Wisconsin Press,Madison, 1985) pp. 229-233.Summerlin, L., Borgford, C., & Ealy, J., “The Mysterious Sunken Ice Cube,”ChemicalDemonstrations: A Sourcebook for Teachers,Vol. 2 (American Chemical Society, 1988)pp. 15-16.Section 1.2 Temperature•DemonstrationMaterials: Galilean thermometer, large beaker containing ice water.Presentation: Show students the thermometer. Briefly place the bottom of the deviceinto an ice bath until a change is observable. Have students describe the changes andexplain how the thermometer works.Section 1.2 Temperature•MisconceptionsThere are many known misconceptions held by students about heat and temperature:Heat and temperature are the same thing; temperature is transferred, not heat; and atemperature decrease is due to the introduction of “cold.”Section 1.2 Temperature•Teaching TipStudents often do not recognize that pure substances melt and freeze at the sametemperature, even a substance with which they have extensive experience, such aswater.Section 1.2 Physical Changes•MisconceptionsMany students think that a phase change is achemicalchange. A very commonmisconception is that water boiling produces hydrogen and oxygen gases. Alternatively,students who believe that a phase change is a physical change think that such a changeis not evident on a particle level.Section 1.2 Chemical Changes

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•DemonstrationsSummerlin, L., & Ealy, J., “Copper into Gold: The Alchemist’s Dream,”ChemicalDemonstrations: A Sourcebook for Teachers,Vol. 1 (American Chemical Society, 1988)pp. 137-138.Summerlin, L., & Ealy, J., “Catalytic Decomposition of Hydrogen Peroxide: FoamProduction,”Chemical Demonstrations: A Sourcebook for Teachers,Vol. 1 (AmericanChemical Society, 1988) pp. 101-102.Section 1.4 Scientific Inquiry•Teaching TipWhen discussing scientific inquiry, emphasize an iterative process employing differentactivities (observation, asking questions, designing experiments, collecting data,developing models, etc.) rather than a series of sequential steps.Section 1.4 Hypotheses•Teaching TipMany students have trouble distinguishing a prediction from a hypothesis. Use theexample with the pennies to point out a prediction based on Anna and Bill’s hypothesis:If Bill and Anna scratched the pennies more, then the reaction would proceed faster.Math Toolbox 1.1 Scientific Notation•Teaching TipGive the students a list of positive numbers in scientific notation including ones withlargest. Add zero to the list. Many students will place zero between the numbers withnegative and positive exponents. Be sure they understand that the sign on the exponentdoes not relate to the whether or not the number is more or less than zero.Math Toolbox 1.1 Scientific Notation•Teaching TipEncourage students to make sure they can perform computations on their calculatorsusing numbers in scientific notation.Math Toolbox 1.2 Significant Figures•Teaching TipPoint out that in chemistry many of the numbers are actually measurements having amagnitude, unit, and degree of certainty.Chapter 2Section 2.2 Subatomic Particles•Teaching TipA chocolate chip cookie is more familiar to students than plum pudding. Ask thestudents to explain what the chocolate chips and the cookie dough represent if thecookie resembles Thomson’s model.Section 2.4 Atomic Mass•Teaching Support ActivityMaterials: Baggies each containing 12 navy beans, 6 kidney beans, and 8 lima beans(dried) for each group of 3-4 students.

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Directions: Hand out one baggie to each group of students and tell them that the beansrepresent the atoms of a fictional element named beanium. Tell the students the massof each type of atom (bean) and instruct them to count the number of each. Ask them tofind the percentage abundance of each type isotope of beanium and compute theaverage atomic mass (relative atomic mass) of the element.Discussion: Dalton stated in his atomic theory that all atoms of an element are identical.Did Dalton know about isotopes? Why or why not?Section 2.4 Atomic Mass•Teaching TipStudents are likely to already be familiar with weighted averages from the experience ofcomputing their grades in a course in which the various course components (exams,quizzes, labs, homework, etc.) are not weighted equally. Make up a problem withgrades for a hypothetical student and have the class find the average.Section 2.5 Classification of Elements•DemonstrationBring in samples of metals, nonmetals, and metalloids. Show the samples to thestudents and have them describe the properties and classify each element.Chapter 3Section 3.1 Ionic and Molecular Compounds•DemonstrationShakhashiri, Bassam Z., “Conductivity,”Chemical Demonstrations: A Handbook forTeachers of Chemistry,Vol. 3 (The University of Wisconsin Press, Madison, 1985) pp.326-328.Section 3.1 Ionic and Molecular Compounds•Teaching TipEmphasize thatdissolvingdescribes behavior on a macroscopic level and substancesthat dissolve do not necessarilydissociate. Give the formula for sucrose and askstudents if it is ionic or molecular. Ask if sucrose dissolves in water and then if itdissociates in water.Section 3.1 Ionic and Molecular Compounds•Teaching TipPoint out to students that the process of sodium chloride dissolving is a physical change,although we represent it with a chemical equation.Section 3.2 Monatomic Ions•Teaching TipStudents often memorize the charges of many common monatomic ions by theirpositions in the periodic table without thinking about the number of protons and electronsin each particle or its position relative to the noble gases. Emphasizethatmonatomicions in Groups I, II, V, & VII have characteristic chargesthat give each of them thenumber of electrons as the nearest noble gas.Section 3.2 Polyatomic Ions

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•Teaching TipAfter students identify the relationship between the names and formulas for most of theoxoanions, recommend that students memorize the common polyatomic ions with thesuffix–ate, such as sulfate, nitrate, carbonate, phosphate, and chlorate, and simply takeone oxygen away to form the polyatomic ion with the suffix–ite.Section 3.2 Polyatomic Ions•Teaching TipSelect 10-15 common polyatomic ions for students to memorize. Not having to refer totheir notes saves time when writing formulas and naming compounds.Section 3.3 Formulas for Ionic Compounds•Teaching TipPoint out the distinction betweenatom, molecule, ion,andformula unit,and stress theimportance of using these terms accurately.Section 3.3 Formulas for Ionic Compounds•Teaching TipHave the students write the formula for magnesium nitrate with and without theparentheses and explain why they are needed.•DemonstrationBring in vials of FeCl24H2O and FeCl36H2O (or Fe(NH4)2(SO4)26H2O andSection 3.4Naming Ionic CompoundsFeNH4(SO4)212H2O) to show students the difference between iron(II) and iron(III)compounds.Section 3.4 Naming Ionic Compounds•Teaching TipSome students worry about knowing all the possible charges on ions for metals that formmore than one ionic compound. Remind them that the Romannumeral in the name tellsthe charge, and if they are naming, they can deduce the charge of the metal from thecharge on the anion.Section 3.4 Naming Ionic Compounds•Teaching TipRemind students to include the Romannumeral in the name when it is necessary.Section 3.4 Naming Ionic Compounds•Teaching Support ActivityMaterials: Various food, vitamin, nutritional supplement, and beverage labelsDirections: Have the students write chemical formulas for as many ionic compounds asthey can.Discussion Questions: Did most of the names use the common or Stock system? Whichmonatomic ions were common among the compounds? Which polyatomic ions werecommon among the compounds? For what names or categories of names were youunable to write formulas?Section 3.5 Naming and Writing Formulas for Molecular Compounds•Teaching Tip

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To help studentsmemorize the Greek prefixes, remind them of the common ones theyare likely to already know. Bicycle–two wheels; tricycle–three wheels; octopus–eighttentacles; etc.Section 3.6 Acids and Bases•Teaching TipThe treatment of acids and bases here is in the context of nomenclature and is verybrief. Acids and bases are presented in detail in Chapter13.Section 3.6 Acids and Bases•Teaching TipPoint out the three different ways to write the formula for acetic acid (HC2H3O2, CH3-COOH, and CH3CO2H) and for the acetate ion (C2H3O2−,CH3COO−, and CH3CO2−).Section 3.6 Acids and Bases•Teaching TipMention that bases commonly are ionic compounds or ammonia, and formulas formolecular substances ending in–OH are known as alcohols.Section 3.7 Predicting Properties and Naming Compounds•Teaching TipStudents often have a great deal of difficulty recognizing when to apply certainnaming/formula writing rules. Recommend that they practice with the flowchart (Figure3.37), but emphasize that they should eventually be able to correctly write names andformulas without it.Chapter 4Section 4.1 Mole Quantities•DemonstrationTo help students understand the magnitude of 6.022 x 1023, describe the spendingpower of a mole of pennies and the volume of a mole of marshmallows. If a mole ofpennies were evenly distributed to each and every person on Earth and they spent $1million dollars every hour, 24 hours per day, they would have half of their money unspentat death.One mole of marshmallows would cover the USA to a depth of 105,000 km(6500 miles).Next, ask students if they can estimate the volume of one mole of watermolecules. After gathering input, show them a graduated cylinder containing 18 mL ofwater. Ask students what this tells them about the number of things in a mole and thesize of a water molecule.Section 4.1 Mole Quantities•Teaching TipAlthough a mole is analogous to a dozen, the dozen analogy seems to break down forstudents. Twelve is a relatively small number and students do not see the necessity inusing dozens when they can just report a counting number. To help them to understandthe utility of the mole, better alternatives are the ream (as in 500 sheets of paper) andthe gross (144 items).Section 4.1 Mole Quantities

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•Teaching TipStudents can easily see that a CO2molecule is composed of 1 carbon atom and 2oxygen atoms. They get confused however when this is scaled up to moles to say that 1mole of CO2is comprised of 1 mole of carbon atoms and 2 moles of oxygen atoms.They often say that 1 + 2 does not equal 1. Emphasize that the C and O atoms aresubsets of CO2molecules much like individual shoes are subsets of pairs of shoes. Askhow many right shoes they have if they havea dozen pairs, and relate it back to atomsand molecules.Section 4.1 Mole Quantities•Teaching TipInstruct the students not to round the molar masses from the periodic table, but rather touse all of the available digits.Section 4.2 Moles, Masses, and Particles•Teaching TipStudents know how to use percentages to compute their grades on assignments.Remind them that a grade of 88% on a testis the ratio between a part (points earned)and the whole (points possible).Section 4.2 Moles, Masses, and Particles•Teaching TipPoint out that students should evaluate their answers to be sure they make conceptualsense. Even students using dimensional analysis may divide when they should multiplyand vice versa. For example, if there are 63.55 g in 1 mol of Cu, then 1.27 mol (morethan 1 mol) should contain more than 63.55 g. Therefore, if students divide 63.55 g by1.27 mol, they should recognize that this is not correct. Encourage them to estimatebefore they calculate and evaluate after.Section 4.3 Determining Empirical and Molecular Formulas•Teaching TipProvide a list of formulas and have students classify them as empirical or molecular.Ask them to write the empirical formula for each one they identified as a molecularformula.Section 4.3 Empirical and Molecular Formulas•Teaching TipPose these questions to check for understanding: When is an empirical formula and amolecular formula the same? What are examples of compounds with the sameempirical and molecular formulas?Ask why the formulas formostionic compounds areempirical formulas. Can you come up with formulas of ionic compounds that are notempirical formulas?Section 4.3 Empirical and Molecular Formulas•Teaching TipStudents tend to refer to all formulas for compounds as molecular, even ionic ones.Emphasize that the term molecule only refers to particles comprised of nonmetallicatoms and that formula units show the simplest whole number ratio of cations (metallicor NH4+ions) to anions (nonmetallic monatomic or polyatomic ions).

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Section 4.3 Empirical Formulas for Compounds Containing More Than Two Elements•Teaching TipRemind students that when converting from grams to moles for hydrogen, oxygen, orany other diatomic element in an empirical or molecular formula problem,they aredetermining the ratio of atoms, not molecules. This means that the appropriate molarmasses for hydrogen and oxygen are 1.008 g/mole and 16.00 g/mole, respectively. Youcan also remind them that hydrogen, oxygen, nitrogen, fluorine, chlorine, bromine, andiodine are diatomic as free or uncombined element, whereas here they have combinedwith other elements and formed a compound.Chapter 5Section 5.1 What is a Chemical Reaction?•Teaching TipStudents should see as many chemical reactions (demonstrations and animations) aspossible while studying this chapter.Section 5.1 What is a Chemical Reaction?•MisconceptionStudents often misinterpret chemical equations in thinking that the substances on eitherside of the arrowrepresenta different physical entity in a different location. Emphasizethat the reactants and products are separatednot by location,but by time and show thesubstances present before and after the chemical reaction has taken place.Section 5.3 Writing Chemical Equations•Teaching TipEmphasize that coefficients in chemical equations are mathematically analogous tocoefficients in algebra. For example, in 2xythe two coefficient doubles both quantitiesxandy, just as 2MgO in a chemical equation means that there are two magnesium ionsand two oxygen ionspresent in two formula units of magnesium oxide.Section 5.3 Writing Chemical Equations•Teaching TipBefore practicing balancing equations, give students the opportunity to practice countingatoms in chemical formulas, such as in (NH4)3PO4, NH4NO3, and Al2(CO3)3, to ensurethey are able to correctly interpret subscripts and parentheses.Section 5.3 Writing Chemical Equations•Teaching TipEmphasize to studentsthat they must writecorrect formulas for all reactants andproducts before attempting to balance chemical equations.Section 5.3 Writing Chemical Equations•MisconceptionStudents learn quickly to correctly balance chemical equations but often do notunderstand what the equations represent, as evidenced by their inability to draw correctparticulate representations of atoms, molecules, and formula units in a reaction. Be sureto ask them to draw particulate diagrams of chemical reactions from the balancedchemical equations. Since coefficients and subscripts serve the same mathematical

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purpose in counting atoms, students do not distinguish between them in particulatedrawings. For example, students often draw two separate spheres to represent2C(correct), but also draw two separate spheres to representO2(incorrect).Section 5.4 Decomposition Reactions•DemonstrationSummerlin, L., & Ealy, J., “Catalytic Decomposition of Hydrogen Peroxide: foamProduction,”Chemical Demonstrations: A Sourcebook for Teachers,Vol. 1 (AmericanChemical Society, 1988) pp. 101-102.Section 5.4 Decomposition Reactions•DemonstrationDecomposition of SodiumHydrogenCarbonateMaterials: Large test tube, 10-15 g of NaHCO3, Bunsen burner, test tube holder orclamp, wood splint, safety gogglesProcedure: Add the sodiumhydrogencarbonate to the test tube, light the burner, andheat the test tube over the flame. While heating the test tube, place a burning splint intothe test tube. The carbon dioxide produced in the decomposition of sodiumhydrogencarbonate will extinguish the burning splint.Questions:What are the products of the reaction?Why did the flame go out?Section 5.4 Decomposition Reactions•DemonstrationBring in small vials of copper(II) sulfate pentahydrate and anhydrous copper(II) sulfate.Pass them around to show students how hydratesand anhydrous compounds differ inappearance.Section 5.4 Combination Reactions•DemonstrationBurning MagnesiumMaterials: 4-6 cm strip of magnesium ribbon, crucible tongs, watch glass, Bunsenburner, safetygogglesProcedure: Hold magnesium ribbon with crucible tongs. Caution students not to lookdirectly at the flame. Ignite the Mg with the burner. After the reaction is complete, placethe magnesium on the watch glass.Questions:Is there evidence that a chemical reaction occurred? Describe the evidence.What are the formulas for the reactants in this demonstration?What is the formula for the product in this demonstration?What is the balanced chemical equation for this reaction?Section 5.4 Combination Reactions•DemonstrationMaking RustMaterials: Small piece of steel wool, crucible tongs, 9-V battery, safetygogglesProcedure: Hold steel wool with crucible tongs. Touch the terminals of the battery to thesteel wool. Once it ignites, gently blow on the glowing steel wool. Tell the students thatthe product of this reaction is iron(III) oxide.

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Questions:Is there evidence that a chemical reaction occurred? Describe the evidence.What are the formulas for the reactants in this demonstration?What is the balanced chemical equation for this reaction?Section 5.4 Single-Displacement Reactions•DemonstrationMetals in HClMaterials: 100 mL of 1MHCl, small pieces of metals such as magnesium, zinc, copper,and lead, 4 large test tubes, 1 test tube rack, safety gogglesProcedure: Pour about 20 mL of HCl into each test tube. Drop one metal into each testtube and observe.Questions:Which metals reacted with HCl?How did you know?What are the products of the reactions in which the metals replaced hydrogen?Section 5.4 Single-Displacement Reactions•DemonstrationSummerlin, L., & Ealy, J., “The Aluminum Cola Can Rip-Off,”Chemical Demonstrations:A Sourcebook for Teachers,Vol. 1 (American Chemical Society, 1988) pp. 152-153.Section 5.4 Gas Formation Reactions•DemonstrationReaction of CaCO3with HClMaterials: About 5 g of CaCO3, large beaker (1 L), 500 mL of 1MHCl, safety gogglesProcedure: Pour the HCl into the beaker. Add the CaCO3and observe.Questions:What are the products of the double-replacement reaction between calcium carbonateand hydrochloric acid?What evidence did you see that carbonic acid decomposed?Section 5.4 Acid-Base Neutralization Reactions•DemonstrationAcid-Base NeutralizationMaterials: 100 mL of 1MHCl, 100 mL of 1MNaOH, thermometer, 2 400-mL beakers,safety gogglesProcedure: Pour the hydrochloric acid into one beaker and the sodium hydroxide into theother. Measure and report the temperature of each. Pour the solutions back and forthbetween the beakers a few times. Take the temperature of the mixture.Questions:What evidence of a chemical reaction was observed?What are the products of a reaction between hydrochloric acid and aqueous sodiumhydroxide?What class of reaction is this?Section 1.4 Combustion Reactions•DemonstrationEthanol Cannon

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