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Solution Manual for Chemistry , 10th Edition

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    Solution Manual To Accompany Chemistry 10 th edition Steven S. Zumdahl University of Illinois Susan A. Zumdahl University of Illinois Prepared by: Donald DeCoste University of Illinois at Urbana - Champaign

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    Scheduling and Laboratory 3 EXPERIMENTAL CHEMISTRY Tenth Edition by John G. Little Introduction Preface ................................ ................................ ................................ ..................... vii Laboratory Glassware and Other Apparatus ................................ .............................. ix Safety in the Chemistry Laboratory ................................ ................................ .......... xvi Safety Quiz ................................ ................................ ................................ ........... xxvii Experiments 1. The Determination of Mass ................................ ................................ ................... 1 2. The Use of Volumetric Glassware ................................ ................................ ........ 9 3. Density Determinations ................................ ................................ ...................... 21 4. The Determination of Boiling Point ................................ ................................ .... 29 5. The Determination of Melting Point ................................ ................................ .... 39 6. The Solubility of a Salt ................................ ................................ ....................... 47 7. Identification of a Substance ................................ ................................ .............. 55 8. Resolution of Mixtures 1: Filtration and Distillation ................................ ............ 63 9. Resolution of Mixtures 2: Chromatography ................................ ......................... 73 10. Stoichiometry 1: Limiting Reactant ................................ ................................ ........................... 85 11. Stoichiometry 2: Stoichiometry of an Iron(III) - Phenol Reaction ................................ .................. 93 12. Composition 1: Percentage Composition – The Empirical Formula of Magnesium Oxide ......... 101 13. Composition 2: Hydrates and Thermal Decomposition ................................ ............................. 109 14. Preparation and Properties of Hydrogen and Oxygen Gases ................................ ................... 121 15. Gas Laws 1: Molar Volume and The Ideal Gas Constant ................................ ........................ 131 16. Gas Laws 2: Graham’s Law ................................ ................................ ................................ .... 139 17. Calorimetry ................................ ................................ ................................ ............................ 147 18. Heats of Reaction and Hess’s Law ................................ ................................ ... 159 19. Spectroscopy 1: Spectra of Atomic Hydrogen and Nitrogen ........................... 171 20. Spectroscopy 2: Emission Spectra of Metallic Elements ................................ . 183 21 . Molecular Properties : Molecular Shapes and Structures ................................ . 195 22. Properties of Some Representative Elements ................................ .................. 201 23. Classification of Chemical Reactions ................................ ............................... 217 24. Colligative Properties 1: Freezing Point Depression and the Determination of Molar Mass ................................ ................................ ................................ ............ 235 25. Colligative Properties 2: Osmosis and Dialysis ................................ ................ 245 26. Rates of Chemical Reactions ................................ ................................ ........... 253

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    4 Scheduling and Laboratory 27. Chemical Equilibrium 1: Titrimetric Determination of an Equilibrium Constant 263 28. Chemical Equilibrium 2: Spectrophotometric Determination of an Equilibrium Constants ................................ ................................ ................................ ............... 273 29. Chemical Equilibrium 3: Stresses Applied to Equilibrium Systems .................. 281 30. The Solubility Product of Calcium Iodate ................................ ........................ 291 31. Acids, Bases, and Buffered Systems ................................ ............................... 301 32. Acid – Base Titrations 1: Analysis of an Unknown Acid Sample ....................... 317 33. Acid – Base Titrations 2: Evaluation of Commercial Antacid Tablets ................. 327 34. The Determination of Calcium in Calcium Supplements ................................ .. 335 35. Determination of Iron by Redox Titration ................................ ......................... 343 36. Determination of Vitamin C in Fruit Juices ................................ ....................... 351 37. Electrochemistry 1: Chemical Cells ................................ ................................ .. 361 38. Electrochemistry 2: Electrolysis ................................ ................................ ....... 371 39. Gravimetric Analysis 1: Determination of Chloride Ion ................................ ..... 381 40. Gravimetric Analysis 2: Determination of Sulfate Ion ................................ ....... 391 41. Preparation of a Coordination Complex of Copper(II) ................................ ..... 401 42. Inorganic Preparations 1: Synthesis of Sodium Thiosulfate Pentahydrate ....... 409 43. Inorganic Preparations 2: Preparation of Sodium Hydrogen Carbonate .......... 417 44. Qualitative Analysis of Organic Compounds ................................ .................... 425 45. Organic Chemical Compounds ................................ ................................ ........ 435 46. Ester Derivatives of Salicylic Acid ................................ ................................ .... 447 47. Preparation of Fragrant Esters ................................ ................................ ......... 457 48. Proteins ................................ ................................ ................................ ............ 465 49. Enzymes ................................ ................................ ................................ .......... 473 50. Polymeric Substances 1: Amorphous Sulfur ................................ .................... 481 51. Polymeric Substances 2: Preparation of Nylon ................................ ................ 487 52. Qualitative Analysis of Selected Cations and Anions ................................ ....... 493 Appendices A. Plotting Graphs in the Introductory Chemistry Laboratory ................................ 503 B. Errors and Error Analysis in General Chemistry Experiments ........................... 505 C. Vapor Pressure of Water at Various Temperatures ................................ ........... 509 D. Density of Water at Various Temperatures ................................ ........................ 510 E. Solubility Information ................................ ................................ ......................... 511 F. Properties of Substances ................................ ................................ ................... 512 G. Concentrated Acid Base Reagent Data ................................ ........................... 516

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    5 PART II: CHAPTER DISCUSSIONS COURSE CONTENT: DESCRIPTIVE CHEMISTRY AND CHEMICAL PRINCIPLES There has been much discussion and an overall agreement that more descriptive chemistry needs to be included in the general chemistry course. This has been one of the major considerations in the recent curriculum reform movement. How to accomplish this goal has resulted in much less agreement. General chemistry has to move beyond a rote presentation of the facts and principles of chemistry. We feel that th e facts (descriptive chemistry) and principles must be an integrated whole that is aimed at fulfilling the basic goals of the course. The Chemistry textbook follows what has been for the past several years a fairly traditional order of topics. The chapters o n principles (1 - 18 ) are f ollowed by nuclear chemistry (19 ), a more descriptive chapter on chemistry of the elements (20 ), a chapter on transition metals and coordination chemistry (21), and organic chemistry and biochemistry (22). Different instructors will have different approaches to descriptive chemistry and, in fact, will probably give greatly different definitions of descriptive chemistry. We have organized the text this way because we feel this order offers instructors the greatest flexibility. In the fifties and early sixties, general chemistry texts essentially went through the periodic table; principles were introduced as needed. Observations were made of the properties and reactions of matter; then we used chemical principles to try to syste matically organize and understand those observations. This approach may closely resemble how chemists do chemistry. But is it the best way to learn chemistry? We don’t think so. We are all driven by curiosity as to how things work, and an unfortunate c onsequence of marching through the periodic table is to dull the students’ curiosity about why things happen, replacing it with the impression of chemistry as a vast body of unrelated facts and reactions that must be memorized. This approach tends to make learning chemistry both frustrating and uninteresting. In the sixties, a principles - dominated outline began to be used in texts, but we overcompensated. The principles were emphasized. The framework was put in place to deal with chemical facts, but all too often the facts were left out. Only the reaction A + B  C seems to have been covered. Learning chemistry was still frustrating, but in a different way from before. We must integrate the two approaches. In truth, the principles of chemistry and descriptive chemistry are two indispensable parts of the whole. In particular, we must use all of our resources in presenting reactions. Can we honestly expect a neophyte to chemistry to appreciate the difference between 2H 2 + O 2  2H 2 O and A + B  C

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    6 Chapter Discussions if they are only written on the blackboard or in a book? We must make use of all of the senses. Students can appreciate what 2H 2 + O 2  2H 2 O means if the instructor ignites a mixture of hydrogen and oxygen. They can see, hear, and sometimes feel what that equation 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 more experiments for the labs in which students observe reactions or synthesize compounds and do fewer experiments that involve a measurement that confirms (often poorly) some principles discussed in class. The use of lecture demonstrations also allows students to see chemistry. These two approaches represent powerful tools for presenting rea ctions. In the text we have tried to emphasize the framework of principles by which chemical facts can be organized. There are several important features that facilitate the integration of fact and principle. 1. To assist instructors in providing interesting and relevant classroom demonstrations of chemical phenomena for their students, the Instructor’s Annotated Edition indicates via the marginal icon (flask with stirring rod) direct references to more than 750 demonstrations from several authoritative sources. 2. Chapter 4 presents an early, thorough discussion of solution reactions, using the classifications that chemists typically use (acid - base, precipitation, and oxidation - reduction). 3. The illustrations and color photography show the student many reactions. This merely supplements lecture demonstrations and labs but right in the text students can see the beauty of the reactions and substances composing this body of knowledge called c hemistry. 4. Real examples are used to illustrate chemical principles. 5. Much of the discussion in the descriptive chapters is organized according to a framework of the early chapters. For example, Chapters 19 - 21 emphasize periodic relationships; the discussion of polymers in Chapter 22 emphasizes structure - property relatio nships. 6. Many of the end - of - ch apter exercises deal with reactions or substances that are important to our lives. In the problem statement, these relationships are mentioned. 7. The end - of - chapter exercises in Chapters 19 - 22 are designed to illustrate the relationships between the descriptive material and the principles chapters. Energy relationships (particularly bond energies), periodicity, structure, and

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    Chapter Discussions 7 equilibrium are emphasized. In the following chapter discussions, we will further point out how descriptive material from the later chapters can be brought into discussion of principles. The organization of the text allows each instructor greater flexibility in presenting the materi al. The unique features 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 and compounds. It is an amalgam of both. In the text we have tried to bring both together, while allowing instructors maximum flexibility. CHAPTER ONE: CHEMICAL FOUNDATIONS Chapter Learning Goals : Section One: To appreciate the importance of creative problem solving. Section Two: To identify the principal operations and limitations of the scientific 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 a calculated result. Section Six: To show a general method of solving problems. Section S even : To show how to convert units between the English and metric systems. Section Eight : To demonstrate conversions among the Fahrenheit, Celsius, and Kelvin temperature scales. Section Nine : To illustrate calculations involving density. Section Ten : To show how matter can be classified into subgroups. The time spent on the first two chapters may show the greatest differences from class to class. Much of the first chapter is probably review for students with good high school math and chemistry backgrounds. This chapter lays the foundation for dealing with measured quantities and performing calculations. Instructors wishing to treat uncertainty in greater detail can discuss the section Uncertainties in Measurements, from Appendix One. This section of the ap pendix could also be used in the laboratory. Instructors should be careful to point out that the Sample Exercises are worked in a different manner from what students should do. Intermediate answers are rounded off to show the correct number of significant figures at each stage of the calculation. T hese rounded values are then used to complete the calculation. Students, when working

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    8 Chapter Discussions problems, should round only at the end. In the Solutions Guide , we have followed the same convention as the text and have rounded at the points where intermediate answers are shown. If this would result in excessive round - off error, we have carried extra digits and noted this in the solution. Most commonly we have carried extra digits when 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 the use of dimensional analysis as a problem - solving technique. In our classes we emphasize using units as a check. Some students may tend to be sloppy about units in c alculations because they are familiar with the quantities being used. The nice feature of dimensional analysis is that it works well even if we don’t have a good intuitive feel for the quantities and units encountered in a particular problem. On the othe r hand, it is important to remember that correct use of dimensional analysis does not mean that a student has a conceptual understanding of the chemical concepts in a particular problem. One of the reasons we have included the Active Learning Questions is so the students have a chance to vocalize their ideas about the concepts covered, and we have a better chance to try to understand these. Teaching Tips  Connecting the real world to the atomic/molecular world is difficult for many students. Encourage students to visualize the atomic/molecular world by drawing pictures. The student responses will give you an idea of their entering perceptions. Don’t worry about some of the details such as subatomic particles here. For example, point out the representations of hydrogen, oxygen, and water in Section 1.1.  Consider spending some time discussing Figure 1.4. Some students believe that theories become laws once a theory has been accepted, but this figure clearly shows laws and theories as two separate entities. Theories never become laws. Laws tell us what h appens, and theories are our attempts to explain why. In terms of a “scientific method”, students should understand that much of scientific thinking is logical or analytical thinking.  To illustrate uncertainty in measurements bring to class several pieces of glassware used to measure volume. Show the students the glassware. Draw the shape of a representative piece of the glassware on the board marking the drawing to show the meniscus of a liquid between two of the lines. Since the volume falls between the lines there is uncertainty involved in the reading. Have the students write down the result of the measurement. Then compare results. The number recorded for this measurement wil l vary from student to student only in the last digit. Thus there is only one uncertain digit.  A relatively easy way to determine the number of significant figures in a number is to write the number in scientific notation. In this notation there are no leading

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    Chapter Discussions 9 zeros, all zeros in the middle of the number are significant, and trailing zeros are only recorded if they are significant.  To illustrate the difference between a physical and chemical change compare representation (b) and (c) from Figure 1.11 (a physical change from liquid water to gaseous water) to the chemical change when liquid water is electrolyzed to produce gases H 2 and O 2 . Use the equation in the text to represent the electrolysis of water chemical change. CHAPTER TWO: ATOMS, MOLECULES, AND IONS Chapter 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 significance of Gay - Lussac's experiments. Section Four: To summarize the experiments that characterized the structure of 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 formulas and to write formulas given their names. Chapter 2 is the second of the two introductory chapters. As with Chapter 1, the length of time spent on this chapter will depend on the background of the students. It goes over the background in chemical topics that provide the foundation for the rest o f the course. This chapter takes a historical approach to the development of chemistry. It begins with a discussion of the discoveries leading to Dalton’s atomic theory, continues with the experiments elucidating the structure of the atom, and ends with chemical nomenc lature and an introduction to the periodic table. The section on nomenclature and formula writing in the laboratory manual by Hall can be effectively used at this point. Instructors may also wish to discuss nuclear decay in more detail at this time. Chapt er 19 , Nuclear Chemistry, can be covered here if instructors prefer. However, if the chapter is presented early, the sections on decay kinetics and thermodynamic stability of the nucleus would have to be delayed and could be covered with Chemical Kinetics (Cha pter 12) or Spontaneity, Entr opy, and Free E nergy (Chapter 17 ).

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    10 Chapter Discussions Teaching Tips  Consider using an analogy between the letters in the alphabet and words and the atoms of the various elements and compounds. In the same way we can make many words from 26 letters, millions of compounds can be made using only the 100 or so elements. Putting the same elements together in different ways results in different compounds with very different properties. For example, the words “ADD” and “DAD” each consist of the s ame letters. However, the order of the letters gives them different meanings. This discussion will help later (for example, in Chapter 3 you can use this analogy to help the students understand why we do not change subscripts in a chemical formula).  Students often have difficulty understanding what is meant by chemical terms such as element. It is important to emphasize the various ways we use the term element in chemistry. When we say element we might mean a single atom, we might mean a molecule su ch as N or O or we might mean a large sample such as a bar of aluminum or the graphite (“lead”) in a pencil. Be sure to help students realize that chemists use terms in many ways and they need to look at the context of the word to be sure that they underst and what is intended. Chemists are accustomed to thinking about things at the macroscopic and microscopic level simultaneously. Encourage students to think about what they read and to consider how the terms are being used.  Sections 2.2 and 2.3 provide another opportunity to discuss the difference between 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 the law of constant co mposition. 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 stick together to form molecules?”, and “Why/how do molecules stick together to form liquids and solids?”. But no model answers all questions.  The discovery of the proton leads to an excellent example of how and why models change. Dalton’s model of the atom did not account for isotopes since he assumed that all atoms of an element are exactly alike. As knowledge expands models change to accommo date 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. Reasons behind the structure of the periodic table and trends it shows are not introduced until Chapter 7. Focus the students’ attention of the location of groups of elements and the separation between metals and nonmetals.

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    Chapter Discussions 11  Emphasize that ions are formed by the gain and loss of electrons. This is difficult for students because they tend to think about gain and loss using positive numbers. It is very important that they understand from the beginning that a positive ion is fo rmed by losing electrons and a negative ion is formed by gaining electrons. In chemistry the positive number (protons) remain the same.  Be sure to point out to students that atoms do not spontaneously gain or lose electrons. The gain and loss of electrons are always paired. One atom loses electrons simultaneously with another atom gaining electrons.  Stress that learning to name compounds does not consist of memorizing a seemingly endless list of chemicals. There are systematic rules for naming compounds and, by knowing only a few rules, the students can name most any compound they will encounter in t his course.  The students should also understand that we keep the rules as simple as possible (much the same way we keep our scientific models as simple as possible). We only add complications (such as prefixes and Roman numerals) when it is required for clarity. For example, the name sodium(I) chloride is not necessarily wrong, merely redundant. Therefore, the name sodium chloride suffices.  The students can determine the charges of all the ions in Table 2.3 from their positions on the periodic table, with the exception of Ag + . The students should appreciate that the periodic table contains a great deal of information. It is a valuable resource.  Make a clear distinction between Type I and Type II cations. This reinforces the reasons 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 that number of ions present in the compound. This is because in many cases the Roman numeral 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 the charge.  For Type III compounds, consider starting with examples such as “NO” and “NO 2 ”. Ask the students to name these compounds. This will lead naturally to a discussion of prefixes. Like Roman numerals, prefixes are used only for clarity. For example, there are two possible “carbon oxides” and we must use prefixes to differentiate be tween these (carbon monoxide and carbon dioxide). We only make the rules more complicated when it is required.

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    12 Chapter Discussions  Figure 2.22 and Figure 2.23 are convenient flow charts to help students systematically name compounds. Students need to be encouraged to use this type of device to help them think through a difficult problem.  Students should think of polyatomic ions as a “unit”. Many students know that if NaCl is dissolved in water, Na + ions and Cl - ions are present. However, students are often confused about the ions present when NaNO 3 is dissolved. Instead of the NO 3 – ion students often think individual nitrogen and oxygen ions are present. CHAPTER THREE: STOICHIOMETRY Chapter Learning Goals: Section One: To describe the modern atomic mass scale and explain how atomic masses are determined experimentally. Section Two: To explain atomic mass and its experimental determination. Section Three : To explain the importance of the mole concept. To show how to convert among moles, mass, and number of particles for a given sample. Section Four : To show how to calculate values for molar mass. To show how to convert among molar mass, moles, and number of particles in a given sample. Section Five: To describe a conceptual problem solving approach to chemistry. Section Six : To demonstrate the calculation of the mass percent of a given element in a compound. Section Seven : To demonstrate the calculation of the empirical formula of a compound. To show how to obtain the molecular formula, given the empirical formula and the molar mass. Se ction Eight : To identify the characteristics of a chemical reaction and the information given by a chemical equation. Section Nine : To show how to write a balanced equation to describe a chemical reaction. Section Ten : To show how to calculate the masses of reactants and products using the chemical equation. Section Eleven: To show how to recognize the limiting reactant. To demonstrate the use of the limiting reactant in stoichiometric calculations. Chapter 3 deals with the fundamental measurement unit in chemistry, the chemical mole. The law of conservation of mass is the unifying principle of the chapter. The chapter deals with compounds first a nd then reactions.

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    Chapter Discussions 13 An 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 numbers of atoms, molecules, or ions. Atoms are so small that we cannot see them to count them. Even if we could see them, the numbers that we would encounter are so large it would take an immense period of time to count them. The mole provides a unit that allows us to connect the number of atoms (what we are interested in knowing) to someth ing we can measure (such as mass). This approach makes it easy to introduce molarity in Chapter 4 as a similar type of unit. Molarity enables us to convert between the quantity we need to deal with in reactions (numbers of moles) and the most convenient measurement we can make for the amount of a solution (volume). Teaching Tips  Section 3.5 discussing a conceptual problem solving strategy for the students. It offers suggestions on thinking about the problem so that they learn how to solve problems in general, not simply use the right algorithm for a given problem. It is well wor thwhile to spend some time with the students discussing this.  The concept of counting atoms from the mass of a sample can be difficult for students to understand. Develop the idea slowly. Many students learn to work the problems without really understanding this section. This leads to trouble later when they try to solve more complex problems. Return to the candy analogy to assist in developing this concept. You can also use a hardware store analogy of buying nails by weight, or a banking analogy of counting coins into rolls by weighing them.  The idea that different samples with identical mass ratios contain the same number of objects is a difficult one for the students. Consider using the following example. Suppose we have two blocks, a red block and a yellow block. The red block weighs 1.0 ounce, and the yellow block weighs 4.0 ounces. Now suppose we have 16 of each color block. What is the mass of each sample? The sample of red blocks weighs 16.0 ounces and the sample of yellow blocks weighs 64.0 ounces. But note that 16.0 ounces is also 1.0 pound, thus 64.0 ounces is 4.0 pounds. The relative masses of the blocks are: one block sixteen blocks red 1.0 ounce 1.0 pound yellow 4.0 ounces 4.0 pounds The 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’s and grams in the case of the periodic table. In the blocks example, the number 16 is analogous to Avogadro's number.

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    14 Chapter Discussions  It is important for students to realize that a mole really describes the number of objects present, just as a dozen means 12. This connection between a mole as a number and the mass of a sample is essential because we count atoms by weighing samples conta ining large numbers of them.  Notice that the text defines the mole differently from the SI definition: “the mole is the amount of substance that contains as many entities as there are in exactly 0.012 kg of carbon 12”. The text definition is “the mole is the number equal to the numbe r of carbon atoms in 12.01 grams of carbon”. We believe the text definition is easier for the students to understand at this point as it more strongly emphasizes both that the mole is fundamentally a number and that the periodic table contains average ato mic masses.  In Section 3.9 students learn that a chemical reaction involves rearrangement of the elements. It is very helpful to use simple models of the reactants that you can take apart and rearrange to show students that atoms are conserved in a chemical reaction (toothpicks and gumdrops or clay work well if you do not have ball - and - stick models). Use the molecular level graphics of chemic al reactions in Section 3.9 to aid your discussion.  The text lists three steps for writing and balancing chemical equations. Step 3 tells the students to start “with the most complicated molecule(s).” For the reaction C 2 H 5 OH(g) + O 2 (g)  CO 2 + H 2 O(l) consider balancing the equation in the following way: C 2 H 5 OH(g) + O 2 (g)  2CO 2 + H 2 O(l) (carbon is balanced) C 2 H 5 OH(g) + 2O 2 (g)  2CO 2 + H 2 O(l) (oxygen is balanced) C 2 H 5 OH(g) + 2O 2 (g)  2CO 2 + 3H 2 O(l) (hydrogen is balanced, oxygen now unbalanced) C 2 H 5 OH(g) + 3O 2 (g)  2CO 2 + 3H 2 O(l) (the equation is balanced) Notice that we eventually get the right answer (so trial and error is a reasonable method). However, it would have been better not to have balanced the oxygen as the second step. This is why Step 3 tells the students to start with the most complicated mo lecule. Encourage the students to leave the elements and diatomic molecules until last when balancing equations. This simplifies the trial and error process.

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    Chapter Discussions 15  Emphasize that the coefficients in a balanced equation represent the mole ratios . An individual coefficient is meaningless (just as an amount of a single ingredient in a recipe is meaningless without the other amounts).  It is useful to use a real life example when introducing limiting reactants. Students seem to relate easily to kitchen experiences yet have difficulty applying the same concepts to chemical equations. By working with real life experiences first and addressing the concepts in this context students h ave an opportunity to become comfortable with the concept before applying it to an unfamiliar setting.  Using a concrete analogy (such as making sandwiches) helps the students better understand the calculations. This is especially true for problems involving a limiting reactant. Problems in which a reactant is limiting are conceptually the same as the case s in which no reactant is limiting. However, students seem to have much more difficulty dealing with limiting reactants. For example, some students calculate the amount of product that could be made from each reactant and then add them. Or, they add the moles of each given reactant for the total. They almost never make these types of mistakes in the analogies, though, so it is helpful to be very explicit about these points  Use molecular diagrams like those shown in Section 3.11 of the text to make counting molecules more concrete (and similar to real life experience). CHAPTER FOUR: TYPES OF CHEMICAL REACTIONS AND SOLUTION STOICHIOMETRY Chapter Learning Goals: Section One: To show why the polar nature of water makes it an effective solvent. Section Two: To characterize strong electrolytes, weak electrolytes, and non - electrolytes. Section Three: To define molarity and demonstrate calculations involving the composition 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 solution reaction. Section Six: To describe reactions in solution by molecular, complete ionic, and net ionic equations. Section Seven: To demonstrate stoichiometric calculations involving precipitation reactions.

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    16 Chapter Discussions Section Eight: To show how to perform calculations involved in acid - base volumetric analysis. Section Nine: To characterize oxidation - reduction reactions. To describe how to assign oxidation states. To identify oxidizing and reducing agents. Section Te n: To describe the oxidation states method for balancing oxidation reduction reactions. The placement of this chapter differs from its placement in many general chemistry texts. We feel there are several good reasons for including these topics at this point and many other texts have co me to agree with us . A thorough discussion of the types of reactions in solution allows for an early introduction of descriptive chemistry. Descriptive chemistry is very important in the first - year course. Although virtually every instructor will give a different definition of descriptive chemistry and a different list of topics to be covered, a thorough discussion of chemical reactions in aqueous solution is probably central to all of those definitions and lists. Hence, the expanded discussion of reactions at an early point . Several years ago texts began introducing molarity and solution stoichiometry in the stoichiometry chapter to give flexibility to the lab program. The expanded discussion in Chapter 4 gives a great deal of flexibility to the lab. The table on scheduling lectures and labs in Part I amply illustrates this point. The stoichiometry for all of the common types of solution reactions is discussed. Students can begin to see all types of reactions in the lab very early in the course. The expanded discussion in the text will help the student focus on what is actually going on in a solution, what species are really present, and how they interact with one another. A chemical equation written on the page of a textbook probably has little meaning to a neophyte. Students must have the opportunity to see reactions. Lecture demonstrations and the color photographs in the text help, but the most useful place for seeing chemistry is by doing chemistry in the laboratory. Chapter 4 expands the vista of what 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, double displacement, and so forth, is avoided because these terms are not used by chemists. These terms can be confusing. For example, are AgNO 3 (aq) + HCl(aq)  AgCl(s) + HNO 3 (aq) and Ag + (aq) + Cl - (aq)  AgCl(s) the same or different? Is the first a double displacement and the second a combination? We strongly feel reactions should be classified on the basis of what can

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