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Solution Manual For Introduction to Chemistry, 4th Edition

  • Chemistry
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    Bauer, Birk, and Marks Instructor Annotations Chapter 1 Section 1.1 Compounds • Teaching Tip Ask students to generate a list of examples of matter in the room. In small groups, have them classify the examples as mixtures or pure substances. Then require them to classify the pure substances as elements or compounds. Lastly, have them classify the mixtures as heterogeneous or homogeneous. Section 1.1 Representations of Matter • Teaching Tip Provide opportunities for students to draw pictures of matter on a particulate level whenever possible. Section 1.1 Representations of Matter • Teaching Tip Given particulate representations, students often mistake diatomic elements for compounds. In order for a molecule to be classified as a compound, there must be two or more atoms of different elements. Section 1.1 States of Matter • Misconception Many students believe that there is no molecular motion in solids. Section 1.1 States of Matter • Misconception Many students think that condensation on the outside of a glass containing a cold beverage comes from the liquid inside the glass. Section 1.2 Physical and Chemical Changes and Properties of Matter • Teaching Tip Rather than teaching students to convert from English to SI units, provide them with practical examples to allow them to think in the SI system. Tell them that a small paper clip has a mass of approximately 1 gram. Section 1.2 Physical and Chemical Changes and Properties of Matter • Teaching Tip Emphasize the value of dimensional analysis and that it can be used throughout the course to solve problems. Section 1.2 Volume • Teaching Tip Tell 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 and cold water. Preparation: Add a few drops of red food coloring to two 250 mL flasks. Add hot water until the flasks are completely filled. Swirl to mix the food coloring. Repeat for the other two 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 containing cold (blue) water. Remove the card and ask the students to observe. Have them predict what would happen if the flask containing the cold (blue) water was placed on the one containing the hot (red) water. Repeat the demonstration with the flask with hot water on the bottom. Ask the students to explain their observations. They should discuss how the density of water changes with temperature on a macroscopic and particulate level. Safety: Wear safety goggles in case glassware breaks. Section 1.2 Density • Demonstrations Shakhashiri, 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,” Chemical Demonstrations: A Sourcebook for Teachers, Vol. 2 (American Chemical Society, 1988) pp. 15 - 16. Section 1.2 Temperature • Demonstration Materials: Galilean thermometer, large beaker containing ice water. Presentation: Show students the thermometer. Briefly place the bottom of the device into an ice bath until a change is observable. Have students describe the changes and explain how the thermometer works. Section 1.2 Temperature • Misconceptions There are many known misconceptions held by students about heat and temperature: Heat and temperature are the same thing; temperature is transferred, not heat; and a temperature decrease is due to the introduction of “cold.” Section 1.2 Temperature • Teaching Tip Students often do not recognize that pure substances melt and freeze at the same temperature, even a substance with which they have extensive experience, such as water. Section 1.2 Physical Changes • Misconceptions Many students think that a phase change is a chemical change. A very common misconception 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 change is not evident on a particle level. Section 1.2 Chemical Changes

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    • Demonstrations Summerlin, L., & Ealy, J., “Copper into Gold: The Alchemist’s Dream,” Chemical Demonstrations: A Sourcebook for Teachers, Vol. 1 (American Chemical Society, 1988) pp. 137 - 138. Summerlin, L., & Ealy, J., “Catalytic Decomposition of Hydrogen Peroxide: Foam Production,” Chemical Demonstrations: A Sourcebook for Teachers, Vol. 1 (American Chemical Society, 1988) pp. 101 - 102. Section 1.4 Scientific Inquiry • Teaching Tip When discussing scientific inquiry, emphasize an iterative process employing different activities (observation, asking questions, designing experiments, collecting data, developing models, etc.) rather than a series of sequential steps. Section 1.4 Hypotheses • Teaching Tip Many students have trouble distinguishing a prediction from a hypothesis. Use the example 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 Tip Give the students a list of positive numbers in scientific notation including ones with largest. Add zero to the list. Many students will place zero between the numbers with negative and positive exponents. Be sure they understand that the sign on the exponent does not relate to the whether or not the number is more or less than zero. Math Toolbox 1.1 Scientific Notation • Teaching Tip Encourage students to make sure they can perform computations on their calculators using numbers in scientific notation. Math Toolbox 1.2 Significant Figures • Teaching Tip Point out that in chemistry many of the numbers are actually measurements having a magnitude, unit, and degree of certainty. Chapter 2 Section 2.2 Subatomic Particles • Teaching Tip A chocolate chip cookie is more familiar to students than plum pudding. Ask the students to explain what the chocolate chips and the cookie dough represent if the cookie resembles Thomson’s model. Section 2.4 Atomic Mass • Teaching Support Activity Materials: 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 beans represent the atoms of a fictional element named beanium. Tell the students the mass of each type of atom (bean) and instruct them to count the number of each. Ask them to find the percentage abundance of each type isotope of beanium and compute the average 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 Tip Students are likely to already be familiar with weighted averages from the experience of computing their grades in a course in which the various course components (exams, quizzes, labs, homework, etc.) are not weighted equally. Make up a problem with grades for a hypothetical student and have the class find the average. Section 2.5 Classification of Elements • Demonstration Bring in samples of metals, nonmetals, and metalloids. Show the samples to the students and have them describe the properties and classify each element. Chapter 3 Section 3.1 Ionic and Molecular Compounds • Demonstration Shakhashiri, Bassam Z., “Conductivity,” Chemical Demonstrations: A Handbook for Teachers of Chemistry, Vol. 3 (The University of Wisconsin Press, Madison, 1985) pp. 326 - 328. Section 3.1 Ionic and Molecular Compounds • Teaching Tip Emphasize that dissolving describes behavior on a macroscopic level and substances that dissolve do not necessarily dissociate . Give the formula for sucrose and ask students if it is ionic or molecular. Ask if sucrose dissolves in water and then if it dissociates in water. Section 3.1 Ionic and Molecular Compounds • Teaching Tip Point 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 Tip Students often memorize the charges of many common monatomic ions by their positions in the periodic table without thinking about the number of protons and electrons in each particle or its position relative to the noble gases. Emphasize that monatomic ions in Groups I, II, V, & VII have characteristic charges that give each of them the number of electrons as the nearest noble gas. Section 3.2 Polyatomic Ions

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    • Teaching Tip After students identify the relationship between the names and formulas for most of the oxoanions, recommend that students memorize the common polyatomic ions with the suffix – ate , such as sulfate, nitrate, carbonate, phosphate, and chlorate, and simply take one oxygen away to form the polyatomic ion with the suffix – ite . Section 3.2 Polyatomic Ions • Teaching Tip Select 10 - 15 common polyatomic ions for students to memorize. Not having to refer to their notes saves time when writing formulas and naming compounds. Section 3.3 Formulas for Ionic Compounds • Teaching Tip Point out the distinction between atom, molecule, ion, and formula unit, and stress the importance of using these terms accurately. Section 3.3 Formulas for Ionic Compounds • Teaching Tip Have the students write the formula for magnesium nitrate with and without the parentheses and explain why they are needed. • Demonstration Bring in vials of FeCl 2  4H 2 O and FeCl 3  6H 2 O (or Fe(NH 4 ) 2 (SO 4 ) 2  6H 2 O and Section 3.4 Naming Ionic Compounds FeNH 4 (SO 4 ) 2  12H 2 O) to show students the difference between iron(II) and iron(III) compounds. Section 3.4 Naming Ionic Compounds • Teaching Tip Some students worry about knowing all the possible charges on ions for metals that form more than one ionic compound. Remind them that the Roman n umeral in the name tells the charge, and if they are naming, they can deduce the charge of the metal from the charge on the anion. Section 3.4 Naming Ionic Compounds • Teaching Tip Remind students to include the Roman n umeral in the name when it is necessary. Section 3.4 Naming Ionic Compounds • Teaching Support Activity Materials: Various food, vitamin, nutritional supplement, and beverage labels Directions: Have the students write chemical formulas for as many ionic compounds as they can. Discussion Questions: Did most of the names use the common or Stock system? Which monatomic ions were common among the compounds? Which polyatomic ions were common among the compounds? For what names or categories of names were you unable to write formulas? Section 3.5 Naming and Writing Formulas for Molecular Compounds • Teaching Tip

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    To help studen ts memorize the Greek prefixes, remind them of the common ones they are likely to already know. Bicycle – two wheels; tricycle – three wheels; octopus – eight tentacles; etc. Section 3.6 Acids and Bases • Teaching Tip The treatment of acids and bases here is in the context of nomenclature and is very brief. Acids and bases are presented in detail in Chapter 13. Section 3.6 Acids and Bases • Teaching Tip Point out the three different ways to write the formula for acetic acid (HC 2 H 3 O 2 , CH 3 - COOH, and CH 3 CO 2 H) and for the acetate ion (C 2 H 3 O 2 − ,CH 3 COO − , and CH 3 CO 2 − ). Section 3.6 Acids and Bases • Teaching Tip Mention that bases commonly are ionic compounds or ammonia, and formulas for molecular substances ending in – OH are known as alcohols. Section 3.7 Predicting Properties and Naming Compounds • Teaching Tip Students often have a great deal of difficulty recognizing when to apply certain naming/formula writing rules. Recommend that they practice with the flowchart (Figure 3.37), but emphasize that they should eventually be able to correctly write names and formulas without it. Chapter 4 Section 4.1 Mole Quantities • Demonstration To help students understand the magnitude of 6.022 x 10 23 , describe the spending power of a mole of pennies and the volume of a mole of marshmallows. If a mole of pennies were evenly distributed to each and every person on Earth and they spent $1 million dollars every hour, 24 hours per day, they would have half of their money unspent at 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 water molecules. After gathering input, show them a graduated cylinder containing 18 mL of water. Ask students what this tells them about the number of things in a mole and the size of a water molecule. Section 4.1 Mole Quantities • Teaching Tip Although a mole is analogous to a dozen, the dozen analogy seems to break down for students. Twelve is a relatively small number and students do not see the necessity in using dozens when they can just report a counting number. To help them to understand the utility of the mole, better alternatives are the ream (as in 500 sheets of paper) and the gross (144 items). Section 4.1 Mole Quantities

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    • Teaching Tip Students can easily see that a CO 2 molecule is composed of 1 carbon atom and 2 oxygen atoms. They get confused however when this is scaled up to moles to say that 1 mole of CO 2 is 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 are subsets of CO 2 molecules much like individual shoes are subsets of pairs of shoes. Ask how many right shoes they have if they have a dozen pairs , and relate it back to atoms and molecules. Section 4.1 Mole Quantities • Teaching Tip Instruct the students not to round the molar masses from the periodic table, but rather to use all of the available digits. Section 4.2 Moles, Masses, and Particles • Teaching Tip Students know how to use percentages to compute their grades on assignments. Remind them that a grade of 88% on a test is the r atio between a part (points earned) and the whole (points possible). Section 4.2 Moles, Masses, and Particles • Teaching Tip Point out that students should evaluate their answers to be sure they make conceptual sense. Even students using dimensional analysis may divide when they should multiply and vice versa. For example, if there are 63.55 g in 1 mol of Cu, then 1.27 mol (more than 1 mol) should contain more than 63.55 g. Therefore, if students divide 63.55 g by 1.27 mol, they should recognize that this is not correct. Encourage them to estimate before they calculate and evaluate after. Section 4.3 Determining Empirical and Molecular Formulas • Teaching Tip Provide 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 molecular formula. Section 4.3 Empirical and Molecular Formulas • Teaching Tip Pose these questions to check for understanding: When is an empirical formula and a molecular formula the same? What are examples of compounds with the same empirical and molecular formulas? Ask why the formulas for most ionic compounds are empirical formulas. Can you come up with formulas of ionic compounds that are not empirical formulas? Section 4.3 Empirical and Molecular Formulas • Teaching Tip Students tend to refer to all formulas for compounds as molecular, even ionic ones. Emphasize that the term molecule only refers to particles comprised of nonmetallic atoms and that formula units show the simplest whole number ratio of cations (metallic or NH 4 + 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 Tip Remind students that when converting from grams to moles for hydrogen, oxygen, or any other diatomic element in an empirical or molecular formula problem, they ar e determining the ratio of atoms, not molecules. This means that the appropriate molar masses for hydrogen and oxygen are 1.008 g/mole and 16.00 g/mole, respectively. You can also remind them that hydrogen, oxygen, nitrogen, fluorine, chlorine, bromine, and iodine are diatomic as free or uncombined element, whereas here they have combined with other elements and formed a compound. Chapter 5 Section 5.1 What is a Chemical Reaction? • Teaching Tip Students should see as many chemical reactions (demonstrations and animations) as possible while studying this chapter. Section 5.1 What is a Chemical Reaction? • Misconception Students often misinterpret chemical equ ations in thinking that the substances on either side of the arrow represent a different physical entity in a different location. Emphasize that the reactants and products are separated not by location, but by time and show the substances present before and after the chemical reaction has taken place. Section 5.3 Writing Chemical Equations • Teaching Tip Emphasize that coefficients in chemical equations are mathematically analogous to coefficients in algebra. For example, in 2 x y the two coefficient doubles both quantities x and y , just as 2MgO in a chemical equation means that there are two magnes ium ions and two oxygen ions present in two formula units of magnesium oxide. Section 5.3 Writing Chemical Equations • Teaching Tip Before practicing balancing equations, give students the opportunity to practice counting atoms in chemical formulas, such as in (NH 4 ) 3 PO 4 , NH 4 NO 3 , and Al 2 (CO 3 ) 3 , to ensure they are able to correctly interpret subscripts and parentheses. Section 5.3 Writing Chemical Equations • Teaching Tip Emphasize to students that they must write correct formulas for all reactants and products before attempting to balance chemical equations. Section 5.3 Writing Chemical Equations • Misconception Students learn quickly to correctly balance chemical equations but often do not understand what the equations represent, as evidenced by their inability to draw correct particulate representations of atoms, molecules, and formula units in a reaction. Be sure to ask them to draw particulate diagrams of chemical reactions from the balanced chemical equations. Since coefficients and subscripts serve the same mathematical

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    purpose in counting atoms, students do not distinguish between them in particulate drawings. For example, students often draw two separate spheres to represent 2 C (correct), but also draw two separate spheres to represent O 2 (incorrect). Section 5.4 Decomposition Reactions • D emo n stration Summerlin, L., & Ealy, J., “Catalytic Decomposition of Hydrogen Peroxide: foam Production,” Chemical Demonstrations: A Sourcebook for Teachers, Vol. 1 (American Chemical Society, 1988) pp. 101 - 102. Section 5.4 Decomposition Reactions • Demonstration Decomposition of Sodium Hydrogen Carbonate Materials: Large test tube, 10 - 15 g of Na H CO 3 , Bunsen burner, test tube holder or clamp, wood splint, safety goggles Procedure: Add the sodium hydrogen carbonate to the test tube, light the burner, and heat the test tube over the flame. While heating the test tube, place a burning splint into the test tube. The carbon dioxide produced in the decomposition of sodium hydrogen carbonate will extinguish the burning splint. Questions: What are the products of the reaction? Why did the flame go out? Section 5.4 Decomposition Reactions • Demonstration Bring in small vials of copper(II) sulfate pentahydrate and anhydrous copper(II) sulfate. Pass them around to show students how h ydrates and anhydrous compounds differ in appearance. Section 5.4 Combination Reactions • Demonstration Burning Magnesium Materials: 4 - 6 cm strip of magnesium ribbon, crucible tongs, watch glass, Bunsen burner, safety goggles Procedure: Hold magnesium ribbon with crucible tongs. Caution students not to look directly at the flame. Ignite the Mg with the burner. After the reaction is complete, place the 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 • Demonstration Making Rust Materials: Small piece of steel wool, crucible tongs, 9 - V battery, safety goggles Procedure: Hold steel wool with crucible tongs. Touch the terminals of the battery to the steel wool. Once it ignites, gently blow on the glowing steel wool. Tell the students that the 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 • Demonstration Metals in HCl Materials: 100 mL of 1 M HCl, small pieces of metals such as magnesium, zinc, copper, and lead, 4 large test tubes, 1 test tube rack, safety goggles Procedure: Pour about 20 mL of HCl into each test tube. Drop one metal into each test tube 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 • Demonstration Summerlin, 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 • Demonstration Reaction of CaCO 3 with HCl Materials: About 5 g of CaCO 3 , large beaker (1 L), 500 mL of 1 M HCl, safety goggles Procedure: Pour the HCl into the beaker. Add the CaCO 3 and observe. Questions: What are the products of the double - replacement reaction between calcium carbonate and hydrochloric acid? What evidence did you see that carbonic acid decomposed? Section 5.4 Acid - Base Neutralization Reactions • Demonstration Acid - Base Neutralization Materials: 100 mL of 1 M HCl, 100 mL of 1 M NaOH, thermometer, 2 400 - mL beakers, safety goggles Procedure: Pour the hydrochloric acid into one beaker and the sodium hydroxide into the other. Measure and report the temperature of each. Pour the solutions back and forth between 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 sodium hydroxide? What class of reaction is this? Section 1.4 Combustion Reactions • Demonstration Ethanol Cannon

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