Anderson Cater for Oxtoby, Freeman, and Block’s Chemistry · like Oxtoby, Freeman, and Block’s Chemistry: ... Take notes in class on the topics covered by the instructor

ChemistrySCIENCE OF CHANGE

David Cater Steven J. Anderson

ChemistrySCIENCE OF CHANGEfourth edition

SG/Oxtoby,Freeman, and Block, Chemistry Science of Change, Fourth Edition ISBN: 0-03-033231-1 ©2003Text printer: ABC Co. Cover printer: Transcon-Louiseville Binding: Paper Trim: 8" x 9.875" 2-Color: PMS 2995C/PMS 138C

Visit Wadsworth online at www.wadsworth.com

For your learning solutions: www.thomsonlearning.com

ISBN 0-03-033231-1

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PREFACE

Why should students buy a Study Guide to accompany a complete and well-written textlike Oxtoby, Freeman, and Block’s Chemistry: Science of Change, 4th ed.? The answer isthat, confronted by the wealth of material in each chapter, students may need some overview ofwhere the chapter is leading, and some assistance in deciding how to approach each topic.

We have tried to provide this help with Section-by-Section Study Suggestions for eachchapter, pointing out an efficient route to mastering the material. These are followed by a list ofthe Key Terms in the chapter and a section-by-section series of Study Questions and ProblemTypes that cover the main principles and point out the kinds of calculations that might appear onan exam. The Study Questions and Problem Types have answers found directly in the text. Theyare perhaps the most useful part of the Study Guide. Finally, we have provided a Practice Exam,mostly multiple-choice, with worked-out answers, so that students can practice their exam-takingskills.

We hope this approach will indeed be helpful, and we would appreciate your commentsand suggestions. Email us at [email protected].

It is a pleasure to acknowledge the input of Jean Soper Cater to this work. Herorganizational skill, word processing efforts, and encouragement have helped greatly in thecompletion of this project.

Our appreciation is expressed to colleagues at the University of Iowa, some of whoseexam questions appear in the practice exams. A few items have also been adapted from problemsin the text.

TO THE STUDENT:

We have tried to produce a study guide that will help you organize your study ofchemistry so you can both learn more and achieve a better grade. Chemistry, Science ofChange, 4th ed. by Oxtoby, Freeman, and Block is well written and complete. However, thereis such a wealth of material, that you may wonder how to pick out the key points, and how torelate one part of the book to another.

You may wonder (as have generations of students) “Do I have to memorize all this stuff?”The answer is usually “Definitely not!” You do need to memorize certain chemical facts and alimited number of equations. We've tried to point these out in each chapter, and to provide tipsto help you learn these. Then, armed with these as tools, you can reason out much of the rest ofthe material.

It isn't likely that your chemistry course will cover everything in the textbook, unlessyour class is so well prepared from high school that many of the fundamentals can be glossedover. You must rely on your instructors to tell you specifically what your course will cover, andwhat it will omit. The “extra” material is in the book for two reasons: First, it allows yourteacher to tailor the course for your particular student population. Second, the book can serveyou as a reference work on many topics that come up in daily life - interpreting the daily news,explaining the weather, maintaining you car, etc. If you intend to take more advanced chemistry(and other science) courses, keep your textbook. Use it as a review source and to find lessadvanced (and more understandable) explanations for topics covered in higher level courses.

Here are some suggestions on How to Succeed in Chemistry.

(1) Consult your course syllabus and read the assigned topics before class. You'll find thelecture makes much more sense that way.

(2) Take notes in class on the topics covered by the instructor. Observe what the instructorseems to emphasize most. The topics and problems on the exams will most likely be those thatwere emphasized in lecture.

(3) Review your notes after class, perhaps rewriting them in more complete form. Reread thesections of the textbook that cover this material. Use pencil and paper as you read. Chemistryhas details that can only be learned if you write them down and work through them as you goalong. Just reading is usually not enough.

(4) Follow through the material in this Study Guide as you read the textbook. It identifies keypoints and background topics that should help clarify each new topic. Work through theexamples and check to see that you can define the important terms and state important conceptsin your own words.

(5) Work the exercises and problems your instructor has assigned. The purpose of workingproblems is not just to get the right answer! The purpose is to get practice in thinking about aparticular topic, and to test your understanding by doing a task that requires that understanding.Try to solve the problem yourself first. Only then go to the Student Solutions Manual to checkyour answer or seek a hint about how to approach the problem. If you can restate the problemin your own words, you are halfway to the answer.

(6) Go through the LEARNING GOALS in this Study Guide. Write out your own definitionsand explanations of the KEY TERMS. Check the textbook for any you aren't sure about, andrewrite your statement.

(7) Read down the list of STUDY QUESTIONS AND PROBLEM TYPES in the Study Guide.These should help you focus on the principles in each section and the related types of numericalproblems. They are arranged by chapter section, so you can easily find the authors’ explanation.As you come to each, mentally review whether you have mastered it. Go back over the textbookexamples and your homework problems if a problem type doesn't seem familiar.

(8) Then take the PRACTICE EXAM in the Study Guide as though it were a course examination– without notes or an open book. It covers most of the chapter topics, at a level of a typicalcourse exam. The items are mostly multiple choice and short answer types. Always remember,the goal is to satisfy yourself that you have mastered the principles, and not just to “guess” theright answer. Check your answers with those that follow the Practice Exam. For any youmissed, check whether you were just careless, or whether more study of that topic is needed.

(9) If your instructor makes old exams available for practice, by all means practice on them too.The best guess as to what an instructor's exams will be like, is that they will be like the examsfrom previous terms.

(10) The rule of thumb for college courses is that you should study about two hours for everyhour of lecture and also for every lab period. With chemistry this is probably a minimum to startout with. Don't be surprised if you must put in more time than this. Everyone can learnchemistry, but most of us must work at it in order to succeed.

(11) Find a few fellow students (say, one to three) and study with them. This is particularlyhelpful in working problems. Ask each other questions, and discuss the difficult concepts. Eachof you should be an active participant, and none should merely listen and try to learn withoutalso teaching.

And finally, if after study you are satisfied you really understand the chemistry, you will do wellon examinations. If you study only with the aim of passing the exams, you may pass the exams,but you probably will not remember much of the chemistry afterward.

Good luck in your study of chemistry! We hope you find it as fascinating as we have.

CONTENTS

1. The Atomic Theory of Matter . . . . . . . . . . . . . . . . . . . . 1

Key Terms 7Study Questions and Problem Types 8Practice Exam 11

2. Chemical Equations and Reaction Yields . . . . . . . . . . 19


3. Chemical Periodicity . . . . . . . . . . . . . . . . . . . . . . . . . . . 31

Key Terms 39Study Questions and Problem Types 40Pactice Exam 44

4. Types of Chemical Reactions . . . . . . . . . . . . . . . . . . . 51


5. The Gaseous State . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 69


6. Condensed Phases and Phase Transitions . . . . . . . . . 87


7. Chemical Equilibrium . . . . . . . . . . . . . . . . . . . . . . . . . . 115


8. Acid-Base Equilibria . . . . . . . . . . . . . . . . . . . . . . . . . . . 133


9. Dissolution and Precipitation Reactions . . . . . . . . . . . 157


10. Thermochemistry . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 173


11. Spontaneous Change and Equilibrium . . . . . . . . . . . . 189


12. Redox Reactions and Electrochemistry . . . . . . . . . . . 205


13. Electrochemistry and Cell Voltage . . . . . . . . . . . . . . . 221


14. Chemical Kinetics . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 241


15. Fundamental Particles and Nuclear Chemistry . . . . . 261


16. Quantum Mechanics and the Hydrogen Atom . . . . . . 275


17. Many-Rlectron Atoms and Chemical Bonding . . . . . . 291


18. Molecular Orbitals and Spectroscopy . . . . . . . . . . . . . 309


19. Coordination Complexes . . . . . . . . . . . . . . . . . . . . . . . 335


20. Structure and Bonding in Solids. . . . . . . . . . . . . . . . . . 347

Study Questions and Problem Types 347

21. Silicon and Solid-State Materials . . . . . . . . . . . . . . . . . 351


22. Chemical Processes . . . . . . . . . . . . . . . . . . . . . . . . . . . 355


23. Chemistry of the Halogens . . . . . . . . . . . . . . . . . . . . . . 359


24. From Petroleum to Pharmaceuticals . . . . . . . . . . . . . . 363


25. Polymers: Natural and Synthetic . . . . . . . . . . . . . . . . . 367


CHAPTER 1: THE ATOMIC THEORY OF MATTER

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CHAPTER 1 THE ATOMIC THEORY OF MATTER

INTRODUCTION

Chemistry is the study of matter and the changes it undergoes. Chapter 1 presents theclassification of matter into elements, compounds, and mixtures, and explains how these areconstituted from individual atoms. The experiments and reasoning that led nineteenth centurychemists to develop our present system of writing formulas are outlined. Then early and late20th century experiments are described that confirm the existence of atoms, uncover theirinternal structure, and measure their masses. Finally, we learn about moles and cover the basiccalculations by which molecular formulas are found.

If you have already studied chemistry,be careful! You may be tempted to skip over mostof this chapter. Before you do, be sure you understand the principles covered in Chapter 1,because they underlie everything in the rest of the book. As you study, be sure you try to explainnew terms and concepts in your own words, rather than just memorizing definitions from thebook.

SECTION BY SECTION STUDY SUGGESTIONS

1-1 Chemistry: Science of Change

We are introduced to the dynamic subject of chemistry by a description of the violentreaction of sodium and chlorine to produce common salt. Several rather obvious questions areasked, whose answers will be found as we proceed through the book.

1-2 The Composition of Matter

Note the meanings of analysis and synthesis, and read the Chemistry in Your Lifesection. This is the first of many such sections in the text. Each one spotlights a “real world”chemical topic related to the text material. This one points out one example of Nature’s ability tosynthesize a very complex molecule that is of great benefit to us humans, and how scientistshave been able to use semi-synthesis to produce useable amounts of what Nature produces bybiosynthesis in rare quantities only.


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Substances and Mixtures

Figure 1-4 outlines the classification of matter from the macroscopic (large enough to beheld and measured) to the microscopic (atomic) level. You should learn the boldfaced terms thatare used. Note the difference between physical and chemical properties. Remember that theterms substance and chemical substance refer to elements and compounds, and never tomixtures. Also learn the difference between a material and a substance. Mixtures ofsubstances are of two types, homogeneous and heterogeneous. Homogeneous mixtures will bereferred to later in the book as solutions. All mixtures can be separated into their componentsubstances by physical means such as mechanical separation, distillation and crystallization,without chemical reactions, which would create substances different from the components of themixture.

Homogeneous mixtures can be in any phase, either liquid (e.g. gasoline), gaseous (dust-free air) or solid (sterling silver). Substances can also exist in any phase.

Notice the last paragraph of this section. We often speak of substances as though theywere ideally pure, but in the real world nothing is absolutely pure.

Elements and Compounds

Elements are substances all of whose atoms are chemically alike. Compounds aresubstances that contain two or more elements chemically bonded together in a fixed ratio. Thisratio can be expressed in terms of mass or in terms of number of atoms. Names of 28 of the 112known elements are given here, along with the origin of each name. You should beginmemorizing the names and symbols of the more common elements as you encounter them fromnow on.

1-3 The Atomic Theory of Matter

The text follows a historical approach to outline the development of our present way ofwriting and understanding chemical formulas. Notice the difference between a law and ahypothesis. Be prepared to state or describe the following important principles: the law ofconservation of mass and the law of definite proportions (or the law of constantcomposition). Notice how the foundation of chemistry as a science was dependent on carefulobservation of the relative masses of different elements that combine with each other. Follow theauthors’ discussion of how these observations led to the formulation of Dalton’s atomic theoryof matter. The relatively rare class of nonstoichiometric compounds (discussed further inSection 20-4) requires a more sophisticated approach to the nature of compounds than Dalton’sideas. To conclude this section and demonstrate the reality of atoms, we skip 180 years fromDalton to the images produced by the scanning tunneling microscope (STM).


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1-4 Chemical Formulas and Relative Atomic Masses

Chemical formulas like H2O or CO2 tell you the relative numbers of atoms of differentelements in a compound. Do not confuse the concepts compound and molecule. A molecule is afew atoms that are connected together in an identifiable unit. The word “compound” refers to anamount of substance that may contain one molecule or a huge number of molecules. Manycompounds like sodium chloride (common salt) do not actually contain molecules, as we shallsee, but instead have a continuous, regular stacking of alternate kinds of atoms in the crystal. Infact, some elements occur as molecules in which like atoms are combined (for example, oxygen,O2 , and hydrogen, H2).

This section shows how two experimental generalizations and an insightful hypothesisgave chemists in the late 1800’s an accurate picture of how molecules enter into chemicalreactions, and how to write correct chemical formulas. The experimental principles were thelaw of multiple proportions and the law of combining volumes. Follow the authors’discussion of how these principles, combined with Avogadro’s hypothesis, led to recognitionthat the reactive gaseous elements exist as diatomic molecules (O2, N2, H2, Cl2), and led tocorrect formulas of molecules of compounds (H2O, NO, NH3). Other elements exist aspolyatomic molecules, for example As4, C60, S8. Formulas deduced from relative masses ofelements in compounds are empirical formulas. Actual molecular formulas are confirmed byother kinds of data, particularly from mass spectrometry (Section 1-6).

1-5 The Building Blocks of the Atom

The text now describes the experiments, done about a hundred years ago, that revealedthe inner structure of the atom. The results were able to explain the indirectly deduced chemicalprinciples given in the preceding sections. You should learn both how the experiments werecarried out and what each one revealed about the atom. Ask your instructor whether you areexpected also to be able to give the name of the scientist associated with each experiment. Thefiner details and current theory of atomic structure are presented in Chapters 16 and 17.

Thomson’s experiments with the cathode ray tube showed that cathode rays are actuallyelectrons stripped from gaseous atoms. He showed they were negatively charged particles withmass, and measured the charge-to-mass ratio e/me. Millikan’s oil drop experiments measuredthe electron’s charge e. The charge e, when divided by Thomson’s value for e/me, gives theelectron’s mass me.

Rutherford used the scattering of a beam of alpha particles striking a thin foil to probethe interior of atoms, and proposed a new atomic model consistent with his experiments. In thismodel most of the atom’s mass and all its positive charge (value +Ze) are concentrated in a tinynucleus at the center of the atom. A total of Z electrons (total charge –Ze) are distributed in the


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large space surrounding the nucleus. The combination of Z positive charges and Z electronsrenders the atom electrically neutral. With refinements, this remains our accepted model of theatom’s inner structure.

Subsequent research extended this picture. The nucleus contains protons, thefundamental particles of positive charge, and neutrons, elementary particles with about the samemass as protons, but with no electrical charge. Neutrons are necessary to prevent theelectrostatic repulsion between protons from causing the nucleus to fly apart. A nucleus containsZ protons and N neutrons for a total of (Z+N) nucleons. The nucleus contains nearly all themass of the atom, so the value A = Z+N is called the mass number of the atom; Z is called theatomic number of the atom.

1-6 Finding Atomic Masses the Modern Way

By measuring relative masses of elements that reacted with each other, and deducing theformulas of the compounds, as described in Section 1-4, chemists established a scale of relativeatomic masses. The actual masses of individual atoms could not be determined until the early20th century, when the value of Avogadro’s number, N0, was determined.

Mass Spectrometry and Isotopes

The mass spectrometer is the instrument now used to measure the masses of gaseousatoms and molecules, not by direct weighing, but by giving them an electrical charge andmeasuring their paths in electrical and magnetic fields. Atoms and molecules that carry a netelectrical charge are called ions. The operation of the mass spectrometer is described brieflyhere. For now, remember that atoms have a dense, positively charged nucleus, surrounded bynegatively charged electrons of low mass.

Important results from mass spectrometry:∑ Elements have isotopes, atoms of several different mass numbers, but identical chemical

properties. The isotopes’ fractional abundances are measured mass spectrometrically.∑ Mass number A is the sum of the number Z of protons (positively charged) and N of

neutrons (neutral), A = Z + N.∑ 12C has replaced oxygen as the standard of mass since 1961. 12C atoms have a defined

relative mass of exactly 12. In section 1-7 the unified atomic mass unit, symbol u, is definedas 1/12 the mass of a 12C atom.

∑ The “chemical relative atomic mass” of an element is the sum of the masses of all isotopes,each multiplied by its fractional abundance. These values agree with relative atomic massesfrom chemical experiments but are more accurate. Slight variations of isotopic abundancesfrom one natural source to another occur for all elements. This limits the number ofsignificant figures quoted for each element’s relative atomic mass.


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1-7 The Mole Concept: Counting and Weighing Atoms and Molecules

The Mole

The number of 12C atoms in exactly 12 grams defines Avogadro’s number,N0 = 6.0221420 ¥ 1023.

Avogadro’s number, N0 = 6.022 ¥ 1023, is the number of items in a mole. Memorize it.

For most purposes 4 significant figures will be sufficient. Ask your instructor if you needto know N0 to more than 4.

The relative atomic mass of each element is the average mass in grams of N0 atoms ofthat element as found in nature. This amount of material is a mole, and this mass is also calledthe molar mass, (symbol M) of the element. Similarly, the mass of N0 molecules or formulaunits of a compound is the molar mass, M, of the compound.

When talking about moles and molar masses, be sure to specify the formula of interest.A mole of O2 molecules is twice as much material as a mole of O atoms.

∑ The chemical amount of a substance is the number of moles being considered, and theseterms are used interchangeably.

∑ In case you always wondered where chemists got the peculiar term “mole”, it is anabbreviation of a much older term, “gram molecular weight”.

You should get used to adding up the relative atomic masses to get the relativemolecular mass of a compound that exists as molecules (like water, H2O), or the relativeformula mass of a compound that is not molecular in nature (like sodium chloride, NaCl). Thecalculations are identical, so you don’t have to know which kind of compound it is. Notice thatformulas of more complex molecules can be written in different ways (example: C4H8O2 orCH3COOC2H5), but the masses of the atoms add up the same, no matter how they are written.

Your instructors may occasionally use the informal terms molecular weight for both ofthese, and atomic weight instead of relative atomic mass. It’s a habit that’s hard to break.

Follow through the examples, and practice calculating masses of atoms and molecules ingrams, and converting masses in grams to chemical amounts (moles) and chemical amounts(moles) to grams. Again, always be sure to note the formula for which the chemical amount iscalculated.

The Atomic Mass Unit

A new, very small unit of mass is introduced to simplify the discussion of individualatoms, molecules, and subatomic particles. The unified atomic mass unit, symbol u, is defined


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as exactly one twelfth the mass of one atom of 12C. In this text u is the symbol for the unifiedatomic mass unit, but many other texts use a.m.u. Notice that the mole is defined such that themass of one mole of 12C atoms is exactly 12 g. Here u is defined such that the mass of one atomof 12C is exactly 12 u. Also, 1 g = N0 ¥ u = 6.022 ¥ 1023 u.

1-8 Finding Empirical and Molecular Formulas the Modern Way

A chemical formula tells what elements make up a compound and the relative number ofeach kind of atom. A molecular formula expresses the exact number of each kind of atom in asingle molecule, (but not all substances consist of well-defined molecules.) An empiricalformula can be written for any pure substance, and it expresses the smallest whole numbersubscripts that give the correct relative number of each kind of atom. A formula unit is onemolecule of a molecular compound, but also the simplest number of atoms that correctlyexpresses the relative number of atoms in a non-molecular compound like NaCl or Li2O.

Here’s where you need to be able to convert from mass to number of moles (chemicalamount): converting percentages by mass of the elements making up a pure compound into theempirical formula of the compound. Converting moles to mass lets you calculate percentcomposition from the formula. Mass percentages are determined by analytical experiments, forexample by combustion analysis. The percentage mass of the elements gives us the compound'sempirical formula. If we have a molecular (or molar) mass from some other experiment such asmass spectrometry (Sec. 1-6), we can then convert the empirical formula to a molecular formula.

Follow through the examples in this section, making sure you understand what’s beingdone and why, in each one. These calculations of moles and formulas are the key to calculationsin many later chapters, so master them now!

1-9 Volume and Density

Chemists often measure a volume of a liquid or gas and use its density to calculate themass. The definition of density as mass per unit volume gives the formula for this conversion:

densitymass

volume=

The usual units are g cm-3, but other units can be used. Volumes of liquids in the laboratory areusually measured in liters (symbol L) or milliliters (mL). One mL is exactly one cm3.Remember the density of water, 1.00 g cm-3. We frequently use the fact that 1 mL of waterweighs 1 gram.

Examine the range of densities listed in Table 1-1. Most liquids have densities similar towater, while gases have densities far lower. The “heavy” metals like gold and platinum are even


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more dense than lead, and are among the most dense substances on earth. Follow through theexamples to practice converting between mass and volume by use of density.

Follow through the calculation of the average volume occupied by a molecule in liquidand in gaseous benzene. The result for the gas shows that gases are mostly empty space. Thecalculation for the liquid gives an approximate size for the benzene molecule, about 5 ¥ 10-8 cm

across.

SUMMARY

The chapter Summary gives a good thumbnail sketch of the most important topics in thechapter. Your study of each chapter in the text should include reading through the Summary toguide your review of it. Finally, work through the Cumulative Problem at the very end of theProblems at the end of the chapter. It includes calculations on gallium sulfides that illustratemost of the concepts in Chapter 1.

LEARNING GOALS . . . CHAPTER 1

KEY TERMS - These are the terms in bold faced type in the text. You should be able to defineor explain each of them in your own words.

analysissynthesissemi-synthesis, biosynthesisphysical properties, chemical propertiesheterogeneous and homogeneous mixtures, solutionsphasessubstance, chemical substanceelementcompoundbinary, ternary, quaternary compoundsatom, moleculescanning tunneling microscope (STM)chemical formula, chemical equationempirical formula, simplest formulalaw of definite proportionslaw of constant compositionlaw of combining volumeslaw of multiple proportionsdiatomic and polyatomic molecules


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nonstoichiometric compoundsDalton’s atomic theory of matterAvogadro’s hypothesiscathode rayselectron, proton, neutronnucleusatomic number (Z)neutron number (N)mass number (A)relative atomic mass, chemical relative atomic massmass spectrometerion, molecular ionisotopefractional abundanceAvogadro’s number (N0)relative molecular massrelative formula massmole, molar mass (M)chemical amountunified atomic mass unit (u)molecular formula, molecular substanceformula unitelemental analysishydrocarbonsdensity

STUDY QUESTIONS AND PROBLEM TYPES

These questions cover the most important topics in each section of the chapter. They suggest thediscussion topics and types of calculations that might be expected on an exam covering thesection. Read through the questions and ask honestly whether you can answer each one. Reviewthe section in the text again, if you aren’t sure you can answer or work problems of that type.

Section 1-2

1. How do chemical properties differ from physical properties?

2. What is a phase? What is the difference between heterogeneous and homogeneous matter?What is the difference between a homogeneous mixture and a homogeneous substance? Giveexamples of physical processes that can separate the components of homogeneous mixtures.


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3. What is the difference between an element and a compound? Can both be considered puresubstances?

Section 1-3

4. How are laws, hypotheses, and theories related to each other and to the scientific method?

5. What is the law of conservation of mass? How did Lavoisier’s experiment demonstrate thislaw?

6. State the law of definite proportions. Why is this law also called the law of constancy ofcomposition? What exceptions are now known to this law?

7. Give the five principles of Dalton’s atomic theory of matter. Which principle corresponds tothe law of conservation of mass? Which to the law of definite proportions?

8. Can you state why you believe in atoms? What modern device is able to image atoms onsurfaces? How does it work?

9. Why does the atomic theory of matter explain the constancy of composition in compounds,whereas the theory that matter is continuously divisible does not? Do we need to know theabsolute masses of individual atoms to prove the law of definite proportion? Explain.

Section 1-4

10. What is a molecule? What is the difference between a molecule and a compound?

11. What is the law of multiple proportions? Give some examples that illustrate this law. Howis this law different from the law of definite proportions?

12. State Avogadro’s hypothesis. Explain how Avogadro’s hypothesis and Gay-Lussac’s law ofcombining volumes together showed that Dalton’s rule of simplicity was incorrect. How did thisaffect the relative atomic masses of hydrogen, oxygen, carbon, and sulfur?

13. Explain how to use Avogadro’s hypothesis and Gay-Lussac’s law of combining volumes asin Figure 1-14 to determine the chemical formulas of H2O, NO, and NH3.

Section 1-5

14. Describe briefly Thomson’s and Millikan’s contributions to knowledge about the electron.Thomson’s and Millikan’s experiments did not directly measure the mass of the electron,Describe how their experiments established this value.


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15. Describe briefly how Rutherford’s experiment changed the model of the atom. Explain whythe earlier atomic model could not account for the results of Rutherford’s scattering experiment.

16. State the approximate relative masses of the proton¸ neutron, and electron. Give theelectrical charge of each of these particles. Why are they considered fundamental particles?

17. Mass number A, atomic number Z, and neutron number N are related by A = Z + N. Givenany two, find the third and write the isotope’s symbol (see section 1-6).

Section 1-6

18. Why must atoms (or molecules) be given an electrical charge (ionized) to be studied in amass spectrometer? How does the mass spectrometer measure their masses?

19. All the following express the same information in different ways: (a) the mass of one moleof an isotope in grams, (b) the mass of one atom of an isotope in grams, and (c) the mass of oneatom of an isotope in atomic mass units (sec. 1-7). Given any one, calculate the other two.Given any one and a periodic table, describe the isotope’s composition in elementary particles.

20. Which of Dalton’s atomic principles was found to have an exception revealed by massspectrometry? Hint: it has to do with isotopes.

21. How do you determine the fractional abundance of an element’s isotopes? How do youobtain the chemical relative atomic mass from the measured atomic masses of the isotopes?

Section 1-7

22 What is Avogadro’s number? Describe it in words and quote its numerical value. How doesthis number relate the mass in grams of an atom to the relative atomic mass of that element?How is Avogadro’s number related to Avogadro’s hypothesis?

23. What is meant by the relative molecular mass of a molecule? How is it calculated?

24. What is a mole? What is molar mass? What is the relationship between chemical amountand number of moles? What is the relationship between the molar mass of a molecularsubstance and the mass of one molecule in unified atomic mass units?

25. Given a certain mass of an element, calculate how many atoms and how many moles arepresent. Given a certain mass of a compound, calculate how many moles are present.

Section 1-8

26. What is the difference between a molecular formula and an empirical formula? Why is theempirical formula sometimes called the “simplest formula”? For what kinds of compounds is it


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wrong to call the formula a molecular formula? Name some examples of molecular compoundsand compounds that are not made up of identifiable molecules.

27. From either the molecular formula or the empirical formula, calculate a substance’spercentage composition by mass (i.e., the mass percentage of each element in the compound).

28. Determine the empirical formula of a compound, given amounts of the elements present in astated amount of the compound.

29. From the percentage by mass of each element in a compound, calculate the compound’sempirical formula.

30. What additional information do you need to calculate a compound’s molecular formula fromits empirical formula? What experiments provide such information?

Section 1-9

31. Define density. Given the mass of a sample of material and its density, calculate the volumepresent. Given the density and volume, calculate the mass of a material.

32. Given the density of a substance, and given Avogadro’s number, determine the volumeoccupied by a molecule of that substance. Is volume occupied by a molecule the same as themolecule’s volume?

CHAPTER 1: PRACTICE EXAM

Sections 1-1, 1-2, 1-3, 1-4

1. Pick the false statement. (In this question, count sometimes true statements as being true.)

(a) All the atoms of an element are chemically alike.(b) Elements, but not compounds, are considered to be chemical substances.(c) Sterling silver (95% Ag, 5% Cu) is a homogeneous mixture.(d) A glass of iced tea containing undissolved sugar (and ice) contains three phases.

2. Match the phrases on the left and right:

A. a heterogeneous mixture I. source of a compound does not matterB. Avogadro’s hypothesis II. more than one phase presentC. law of definite proportions III. atoms are indestructible in reactionsD. law of conservation of mass IV. same number of molecules in equal volumes


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(a) A-I B-III C-II D-IV(b) A-II B-I C-IV D-III(c) A-II B-IV C-I D-III(d) A-IV B-III C-I D-II

Section 1-4

3. Copper combines with sulfur to form two different copper sulfide compounds. Compound Icontains 20.15 % sulfur by mass; Compound II contains 33.54 % sulfur. Which of thefollowing statements is not true.

(a) The mass of sulfur that combines with 1.000 g of copper in compound I is 0.252 g.(b) The mass of sulfur that combines with 1.000 g of copper in compound II is 0.505 g.(c) There are twice as many atoms of sulfur per atom of copper in I as in compound II.(d) This is an example of the law of multiple proportions.

Section 1-5

4. Thomson's experimental apparatus for measuring the ratio of charge to mass of the electronwas similar to a modern TV picture tube in all of the following ways except one. Which isthe incorrect statement?

(a) Both involve electrons accelerated through a positive voltage.(b) Both involve emission of light from a fluorescent screen.(c) Both involve deflection of beams of electrons in electrical fields.(d) Both involve electron beams scanned repeatedly across a fluorescent screen.

5. Rutherford's alpha particle scattering experiment established which of the followingcharacteristics of the atom?

(a) the charge to mass ratio of protons.(b) the approximate diameter of gold atoms.(c) the concentration of atomic mass in the nucleus.(d) the uniform distribution of electrical charge throughout the atom.

6. The approximate diameters (order of magnitude) of a gold atom and its nucleus are nearest towhich of these?

(a) 10-5 cm and 10-8 cm (c) 10-10 cm and 10-15 cm(b) 10-8 cm and 10-10 cm (d) 10-8 cm and 10-13 cm

7. The atomic mass of uranium is 238.03 u, and the two most abundant isotopes are 92238U and

92235U . Pick the false statement.


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(a) 238U has more protons than 235U.(b) 235U has fewer neutrons than 238U.(c) 235U has the same number of electrons as 238U (neutral atoms).(d) 238U is more abundant than 235U.

8. The isotope 55132Cs has the following number of protons, neutrons, and electrons, in that order:

(a) 132 55 77 (d) 132 77 55(b) 55 77 55 (e) 55 55 77(c) 77 55 77

9. The correct numbers of protons, neutrons, and electrons in the most abundant isotope ofsodium are, respectively,

(a) 23, 11, 23 (c) 11, 12, 23(b) 11, 23, 11 (d) 11, 12, 11

Section 1-6

10. Chemists analyzing samples of lithium carbonate, LiCO3, from minerals in Antarcticmountains, Arabian deserts, and atmospheric dust have found that the ratio of grams of O pergram of Li in the specimens varies slightly among these sources. The most likely explanationfor this variation is:

(a) experimental error in chemical analyses.(b) slight variation in isotopic content of Li and/or O.(c) contamination of samples during handling.(d) slight nonstoichiometry in the carbonate ions.

11. Isotopic analysis of the iron in a particular meteorite showed the following concentrations ofisotopes, with the appropriate relative atomic mass of each isotope:54Fe 53.94 7.00% 56Fe 55.94 91.00% 57Fe 56.94 2.00%Find the relative atomic mass of this sample of iron.

12. The chemical relative atomic mass of natural boron is 10.81. Boron has two isotopes,10B (10.01) and 11B (11.01). What are the natural abundances of 10B and 11B, respectively?

(a) 35.80%, 64.20% (c) 56.20%, 43.80%(b) 20.00%, 80.00% (d) 80.00%, 20.00%

13. Explain briefly why atomic masses listed on the periodic chart are not integer values.


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Section 1-7

14. What is the molar mass of ammonium carbonate, (NH4)2CO3?

Elemental atomic molar masses in g/mol are H 1.0, C 12.0, N 14.0, O 16.0.

(a) 82.0 g/mol (c) 120.0 g/mol(b) 96.0 g/mol (d) none of the above

15. How many molecules are present in 36.5 g of CO2 (molar mass = 44.01 g/mol)?

(a) 8.29¥10–1 molecules (c) 5.00¥1023 molecules

(b) 5.00 molecules (d) 7.84¥1023 molecules

16. Nitrogen exists as molecules of formula N2. All of the following statements are true butone. Which is not true?

(a) 28 g of nitrogen contains 6.02 ¥ 1023 molecules.

(b) 14 g of nitrogen is a mole of nitrogen atoms.(c) One mole of nitrogen molecules weighs 14 g.(d) 6.02 ¥ 1023 nitrogen atoms weigh 14 g.

17. State in your own words the significance of the carbon-12 isotope to the scale of relativeatomic masses.

18. A car travelling at 25 miles per hour emits about 0.400 kg of carbon monoxide gas per mile.How many moles of CO are emitted in going 3.00 miles?.

(a) 0.429 (b) 42.9 (c) 357 (d) 2.67 (e) 26.7

19. Mendelevium, Md, element 101, was made in a nuclear accelerator. It was reported that inone experiment a total of 3.0¥1014 atoms were obtained. How many moles of Md was this?

(a) 5.0 ¥ 10-7 mol (d) 5.0 ¥ 10-9 mol

(b) 5.0 ¥ 10-8 mol (e) 5.0 ¥ 10-10 mol

(c) 1.0 ¥ 10-9 mol

20. Which of the following quantities for a given molecular compound could not have beendetermined before Avogadro's number was known?

(a) relative molar mass(b) mass of a molecule(c) empirical formula(d) relative atomic masses of the elements in the compound


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21. The relative atomic mass of chlorine Cl is 35.45, and that of potassium K is 39.10. Which ofthe following contains more atoms of chlorine?

(a) 5.0 g Cl2 (b) 0.10 mol of KCl (c) 0.20 mol Cl2

Section 1-7

22. A sample of cobalt(II) chloride hexahydrate (CoCl2 ◊ 6 H2O) contains 1.7 ¥ 1024 atoms of

chlorine. How many molecules of water does it contain?

(a) 1.7 ¥ 1024 (d) 6.8 ¥ 1024

(b) 3.4 ¥ 1024 (e) 10.2 ¥ 1024

(c) 5.1 ¥ 1024

23. Nicotine, a component of cigarette smoke, has the molecular formula C10N2H14. Determineits molar mass and empirical formula.

24. The empirical formula of all true carbohydrates is CH2O. If the molar mass of thecarbohydrate inositol is 180.18 g/mol, then its molecular formula is which of the following?

(a) C2H4O2 (c) C6H12O6(b) C3H6O3 (d) C8H16O8(c) C5H10O5

25. An important reducing agent in organic chemistry, sodium borohydride, contains Na, B, andH. The percent composition by mass and atomic masses in g/mol are as follows:

sodium 60.8% boron 28.5% hydrogen 10.5%22.99 g/mol 10.81 g/mol 1.008 g/mol.

The empirical formula of this compound is:

(a) NaBH2 (b) NaB2H (c) NaBH3 (d) NaBH4 (e) Na2BH4

26. A compound that is known to contain carbon, hydrogen, perhaps oxygen, and no otherelements is burned with oxygen. 2.900 g of compound forms 4.400 g CO2 and 0.900 g H2Owhen burned. The simplest formula (empirical formula) of the compound is which of these?

(a) C2H6O (b) CH3 (c) CHO (d) CH3O (e) C2H

Section 1-8

27. What is the mass of a cube of gold that is 1.50 cm on a side? Density of gold is 19.3 g/cm3.

(a) 5.72 g (d) 29.0 g(b) 12.9 g (e) 65.13 g(c) 19.3 g


16

CHAPTER 1 : PRACTICE EXAM ANSWER KEY

1. (b) Both elements and compounds are chemical substances.2. (c) gives the correct matches.

3. (c) This is not true. Compound I contains half as many S atoms per Cu atom as does II.In 100 g of I there are 20.15 % S and 79.85 % Cu. Mass ratio = 20.15/79.85 = 0.252 g S/g CuIn 100 g of II there are 33.54 % S and 66.46% Cu. Mass ratio = 33.54/66.46 = 0.505 g S/g Cu

4. (d) In Thomson’s apparatus the electron beam was steady, and the position where it hit thescreen was measured.

5. (c) Answer (b) is wrong because it was the diameter of the nucleus, not the atom, that wasmeasured.

6. (d)

7. (a) All isotopes of U have the same number of protons, namely the atomic number 92.8. (b)9. (d) Relative atomic mass is very close to 23.0, so most abundant isotope is 11

23Na . Actually,sodium has only this one naturally occurring isotope.

10. (b) Slight variation of isotopic content from place to place around the world limits thenumber of decimals on each element’s atomic mass in the periodic chart.

11. The relative atomic mass is the sum of isotope mass ¥ isotope fraction:

(0.07 ¥ 53.94) + (0.91 ¥ 55.94) + (0.02 ¥ 56.94) = 55.82 g mol-1

12. (b)A = 10.81, A10 = 10.01, A11 = 11.01. Let p10 and p11 be fractions of atoms of each mass.A = A10p10 + A11p11 = A10p10 + A11(1 - p10) = (A10 - A11)p10 + A11p10 = (A - A11) / (A10 - A11) = (10.81 - 11.01)/(10.01 - 11.01) = 0.20p11 = 0.80

13. Individual isotopic masses are close to integer values, but the atomic masses on the chart areaverages of various amounts of the isotopes, so most are not close to integers.

14. (b) 2¥(molar mass of N) + 2¥4¥(molar mass of H) + (molar mass of C)

+ 3¥(molar mass of O) = 96.0 g/mol.

15. (c) (36.5 g CO2 ) (1 mol CO2 / 44.01 g CO2 ) (6.02¥1023 molecules / mol)

= 5.00¥1023 molecules

16. (a) Molar mass of atomic N = 14 g/mol, molar mass of N2 is 28 g/mol..17. The scale is fixed by defining the mass of a 12C atom as exactly 12 u, then giving the relative

masses of all other isotopes relative to this definition. The relative atomic masses in theperiodic chart are averages of the atoms of each element as found in nature. The mole isdefined as that number of items equal to the number of atoms in exactly 12 g of carbon-12.


17

18. (b) Molar mass of CO = (12.01 + 16.00) = 28.01 g / mol.(0.400 kg CO / 1 mi) ¥ (103 g CO/ 1 kg CO) ¥ (1 mol CO/28.0 g CO)= 14.3 mol CO / mi.

Then, 3.00 mi ¥ 14.3 mol CO / 1 mi = 42.9 mol CO.

19. (e) 3.0 ¥ 1014 atoms ¥ (1 mol / 6.02 ¥ 1023 atoms)= 5.0 ¥ 10-10 mol.

20. (b) Relative masses suffice to define (a), (c), and (d).

21. (c) In (c) we have 0.20 mol Cl2 or 0.40 mol Cl. We can eliminate (a) and (b) by finding thechemical amount of Cl in each sampleIn (a) 5.0 g Cl2 ¥ (1 mol Cl2 / 2 ¥ 35.45 g Cl2) ¥ (2 mol Cl / 1 mol Cl2) = 0.14 mol Cl

In (b) 0.10 mol KCl contains 0.10 mol Cl.

22. (c)(1.7 ¥ 1024 atoms Cl) ¥ (6 molecules H2O / 2 atoms Cl)

= 5.1 ¥ 1024 molecules H2O.

23. (10 ¥ 12.01) + (2 ¥ 14.01) + (14 ¥ 1.008) = 162.2 g / mol. Emprical formula: C5NH7.

24. (c)Molar mass of CH2O = 12.01 + 2 (1.01) + 16.00 = 30.02.(Molar mass inositol) / (molar mass CH2O) = 180.18 / 30.02 = 6.00.

25. (d)Assume a 100 g sample of sodium borohydride. This sample then contains60.8 g Na / (22.99 g Na / mol Na) = 2.64 mol Na28.5 g B / (10.81 g B / mol B) = 2.64 mol B10.5 g H. / (1.008 g H / mol H) = 10.42 mol HWe have 1.00 mol Na / mol B and 3.95 mol H / mol B, satisfied by (d).

26. (c)Molar masses: CO2 = 12.01 + 2 ¥ 16.0 = 44.0 g.; H2O = 2 ¥ 1.01 + 16.0 = 18.0 g

4.4 g CO2 / (44.0 g CO2 / mol CO2) = 0.10 mol CO2 Æ 0.10 mol C

0.9 g H2O / (18.0 g H2O / mol H2O) = 0.05 mol H2O Æ 0.10 mol H

Therefore C/H ratio is 1/1. The only formula among our choices that satisfy this ratio isCHO.To check the O, mass of O = [original mass - mass of (.10 mol C + .10 mol H)] = 1.60 g of O.This corresponds to 0.10 mol O, to give overall empirical formula CHO.

27. (e) Mass = density ¥ volume = 19.3 g cm-3¥ (1.5)3 cm3 = 65.1 g


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