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DOCUMENT RESUME

ED 477 269 SE 067 630

AUTHOR DeVoe, Howard; Hearle, Robert

TITLE Communities of Molecules: A Physical Chemistry Module.Teacher's Guide.

ISBN ISBN-06-561226-4PUB DATE 1978-00-00NOTE 84p.; Produced by the Chemistry Association of Maryland. For

student book, see SE 067 629. For other modules in series,see SE 067 618-629.

PUB TYPE Books (010) Guides Classroom Teacher (052)EDRS PRICE EDRS Price MF01/PC04 Plus Postage.DESCRIPTORS Curriculum Design; *Inquiry; Instructional Materials;

*Interdisciplinary Approach; *Physical Chemistry; ScienceInstruction; Secondary Education; Teaching Methods

ABSTRACT

This teacher's guide is designed to provide science teacherswith the necessary guidance and suggestions for teaching physical chemistry.The material in this book can be integrated with the other modules in asequence that helps students see that chemistry is a unified science.Contents include: (1) "Introduction of Physical Chemistry"; (2) "The GaseousState"; (3) "Liquids and Solids: Condensed States"; (4) "Solutions: Soluteand Solvent"; (5) "The Colloidal State"; (6) "Changes in Energy"; (7) "Ratesof Chemical Reactions"; (7) "Chemical Equilibrium"; and (8)"Electrochemistry". (KHR)

Reproductions supplied by EDRS are the best that can be madefrom the original document.

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interdisciplinaryapproachesto chemistry

lacIAC PROJECT TEAMDirectors of IAC:Marjorie Gardner, 1971-73, 1976Henry Heikkinen, 1973-76

Revision Coordinator:Alan DeGennaro

IAC MODULAR CHEMISTRY PROGRAM MODULE AUTHORS

REACTIONS AND REASON:An Introductory Chemistry Module

DIVERSITY AND PERIODICITY:An Inorganic Chemistry Module

FORM AND FUNCTION:An Organic Chemistry Module

MOLECULES IN LIVING SYSTEMS:A Biochemistry Module

THE HEART OF MATTER:A Nuclear Chemistry Module

THE DELICATE BALANCE:An Energy and the Environment Chemistry Module

COMMUNITIES OF MOLECULES:A Physical Chemistry Module

Teacher's Guides(available for each module)

Gordon Atkinson, Henry Heikkinen

James Huheey

Bruce Jarvis, Paul Mazzocchi

David Martin, Joseph Sampugna

Vic Viola

Glen Gordon, William Keifer

Howard DeVoe

Teacher's Guide Coordinators:Robert Hearle, Amado Sandoval

3

Qinterdisciplinaryapproachesto chemistry

lacTEACHER'S GUIDE

COMMUNITIES OF MOLECULES

1817

Harper & Row, PublishersNew York Hagerstown San Francisco London

A PHYSICAL CHEMISTRY MODULE

4

Howard DeVoeRobert Hearle

AUTHORS

TEACHER'S GUIDE

COMMUNITIES OF MOLECULES:A PHYSICAL CHEMISTRY MODULE

Howard DeVoe

Howard DeVoe calls himself a physical chemist, although his researchinterests range all the way from the theory of liquid structure to reactionsof dyes with DNA. This diversity of subjects simply illustrates the factthat the concepts and theories of physical chemistry can be applied toall fields of chemistry. Howard DeVoe received his Ph.D. from HarvardUniversity and did postdoctoral work at the University of California.Before joining the faculty of the University of Maryland, he was a re-search chemist at the National Institutes of Health. He has had manyarticles published in a wide variety of scientific journals.

Robert Hear le

Robert Hear le came to the University of Maryland in 1970 after workingas a research chemist. While at Maryland he taught in the chemistry andeducation departments, assisted in supervising student teachers in sci-ence, and was part of the IAC development team. Bob is now teachingin the Prince George's County, Maryland school system. His specialfields of interest include analytical and organic chemistry and computerassisted instruction, including the development of a computer managedindependent study program.

Copyright © 1978 by Chemistry Associates of Maryland, Inc.All rights reserved.Printed in the United States of America.

No part of this publication may be reproduced in any form or by any means, photographic, electrostatic or mechanical, or by anyinformation storage and retrieval system or other method, for any use, without written permission from the publisher.

STANDARD BOOK NUMBER 06-561226-4

89012MU0987654321

II

Contents

Introducing Communities of Molecules 1

Special Features in the Student Module 2Managing the Laboratory 2Evaluating Student Performance 3Module Concepts 4Module Objectives 6Teaching Communities of Molecules 13

INTRODUCING PHYSICAL CHEMISTRY 13

P-1 Macroscopic and Molecular: Two Worlds 13P-2 The Size of a Molecule/Experiment 13P-3 Journey into a Drop of Milk 15

THE GASEOUS STATE 16

P-4 Measuring Gases 16P-5 Pressure and VolumelMiniexperiment 17P-6 It's a Hit! 19P-7 Temperature and Volume) Experiment 19P-8 There Ought to Be a Law 20

LIQUIDS AND SOLIDS: CONDENSED STATES 23

P-9 The Attraction of Molecules 23P-10 Special Attraction: Polar Molecules 23P-11 The Vaporization of a Liquid 24P-12 The Boiling Process 24P-13 Boiling at Reduced PressurelMiniexperiment 24P-14 Solid Structures 25

SOLUTIONS: SOLUTE AND SOLVENT 27

P-15 The Dissolving Process 27P-16 Supersaturation/Miniexperiment 28P-17 Dissolved Gases: Fizzle and Pop 28P-18 Antifreeze: Cure for the Common Cold 29

THE COLLOIDAL STATE 30

P-19 Some Properties of Colloidal Particles/Experiment 31

P-20 The Importance of Surfaces 32P-21 Small Wonder: Colloidal Dispersions 32P-22 Making Oil and Water MixIExperiment 33P-23 Making a GellMiniexperiment 35

6 iii

CHANGES IN ENERGY 36

P-24 Counting Calories 36

P-25 The Heat of Fusion of Icel Experiment 37

P-26 Endothermic and Exothermic 38

P-27 Heat of Solution/Experiment 39

P-28 Heat of Reaction/Experiment 40

P-29 The Power of Thermal Energy 40

RATES OF CHEMICAL REACTIONS 42

P-30 Time to React' Experiment 42P-31 Rates and Concentrations 45P-32 Rates and Temperature 45P-33 Catalysts: Essential for Life 46

CHEMICAL EQUILIBRIUM 47

P-34 Equilibrium Constants 47P-35 Shifting an Equilibrium/Miniexperiment 51P-36 The Importance of Size 52P-37 Hydrogen Ions: The Acid Test 52P-38 Show Your Colorsl Experiment 52

ELECTROCHEMISTRY 56

P-39 Oxidation-Reduction Reactionsl Miniexperiment 57P-40 Copper-Zinc Celli Experiment 57P-41 Lead-Acid Celli Experiment 59

P-42 Practical Portable Power 59

P-43 Finishing Touchesl Experiment 59

SUMMARY 61

APPENDIX

SafetyMetric UnitsSuggested Readings and FilmsModule Tests

Knowledge-Centered Module TestSkill-Centered Module Test

Materials ListAcknowledgmentsIndexTable of International Relative Atomic MassesPeriodic Table of the Elements

iv

6364656566697274757677

Introducing Communities of MoleculesAs the title Communities of Molecules suggests, thismodule pictures a chemical system as a com-munity of many members which, by their con-stant movements and interactions, determine thecharacteristics of the system as a whole. The firstexperiment permits the students to measure theapproximate size of a typical member of a molecu-lar community. This is followed by a study ofsome of the major organizational patterns thatare found in these communities: gases, liquids,solids, solutions, and colloids.

An explanation of thermal energy and chem-ical energy paves the way for a series of calorim-etry experiments in which the student measuresthe changes in energy and the changes in tem-perature that accompany the reorganization of amolecular community.

An iodine clock reaction is used to demon-strate quantitative ideas of chemical reaction ratesand the effects of concentration, temperature,and a catalyst. The idea of competing reactionrates in a reversible reaction leads naturally to theconcepts of chemical equilibrium and the equilib-rium constant. The student tests the validity ofan equilibrium constant by measuring the ioniza-tion constant of acetic acid for several dilutesolutions of acetic acid.

Electrochemical cells are introduced as exam-ples of chemical systems that can do useful workas they proceed toward equilibrium. The studentprepares two different kinds of cells that areable to light a flashlight bulb. In the final ex-periment, the chemical energy of a dry cell isused to electroplate zinc onto a strip of copper.

There has recently been a heavy emphasis onprinciples of physical chemistry in high-schoolchemistry. The preparation of this book posedthe problem of how to select the most relevanttopics for one single module. All of the topicscovered in this module are commonly found inother high-school chemistry programs, with theexception of the general topic of colloids (sectionsP-19 through P -23), which is included herebecause of the great number of examples of thisstate of matter encountered in everyday life.Among the physical topics that have not beenincluded because they did not seem to be suf-ficiently important by themselves and are not

needed as preparation for other topics are partialpressure, solubility product, and electrical conduc-tivity. Le Chatelier's principle and the concept ofdecreasing enthalpy and increasing entropy as drivingforces have been omitted because we feel thatthese ideas need a more detailed explanation thanseemed appropriate for the module. These areconcepts that the teacher can introduce to classesof sufficiently interested students.

In keeping with the IAC goals of an inquiry-oriented and innovative chemistry program, thismodule includes among its 18 experiments andminiexperiments 7 that introduce features which,as far as is known, are unique to this program.The contents of experiment P-22 Making Oil andWater Mix and miniexperiment P-23 Making a Gelare not original but have not previously been in-cluded in a high-school program. The miniexperi-ment and experiment on gases (P-5 and P-7)utilize an easily constructed new type of capillarytube container. Three of the experiments areoriginal in their materials and instructions. Theseare experiment P-19 Some Properties of ColloidalParticles, experiment P-38 Show Your Colors, andexperiment P-40 Copper-Zinc Cell. In addition,some recommended demonstrations and supple-mentary experiments with original features aredescribed in this teacher's guide.

The pace at which you proceed through themodule will depend on the interest and abilitiesof your students and on your interest in the in-dividual topics. From past experience, teachershave suggested that 7-10 weeks is the averagetime spent in teaching the entire module.

The intent of the module is to explain someof the important principles of physical chemistryand to relate them to practical experience. Theemphasis is on understanding rather than onmemorization of facts, and the suggested evalua-tion items reflect this emphasis. The module is ingeneral more quantitative than the others in theIAC program, but every effort has been made tokeep the mathematics simple. In the experienceof previous IAC classes, this module is no moredifficult than any of the other modules and con-tains concepts and experiments that are fascinat-ing to students.

81

Special Features in theStudent Module

Metric System Le Systeme Internationale (SI)is used throughout the IAC program. As youwork with this module, you may wish to reviewsome points of the metric system as presentedin Reactions and Reason: An Introductory Chem-istry Module (see section A-8 and Appendix II).There is a metric units chart in the appendix ofthe student module that students can easilyrefer to.

Time Machine A feature we call the Time Ma-chine appears in the IAC modules in order toshow chemistry in a broader context. For somestudents, this may provide a handle on particularaspects of chemistry by establishing the social-cultural-political framework in which significantprogress was made in chemistry. Students mayenjoy suggesting other events in chemistryaround which to create Time Machines of theirown. In the IAC program, Time Machines are notto be memorized by students or to be used as abasis for testing.

Cartoons A popular feature of the IAC programis the use of chemistry cartoons. These cartoonsgive students a chance to remember specificpoints of chemistry in another important wayby humor. Suggest that your students createother chemistry cartoons for their classmates toenjoy.

Safety Laboratory safety is a special concern inany chemistry course. In addition to includingsafety discussion and guidelines in the appendixof each student module and teacher's guide, ex-periments have been developed in a way toeliminate potentially dangerous chemicals orprocedures. Moreover, each experiment that

might present a hazardthrough fumes, cor-rosive chemicals, use of a flame, or otherconditionshas been marked with a safety sym-bol to alert students and teacher to use added,reasonable caution. Caution statements, in boldtype, also appear in experiments to specificallyinstruct the student on the care required.

Selected Readings Articles and books that tie inwith the topics discussed in the IAC programhave been listed in the appendix of the studentmodule as well as in the teacher's guide. En-courage your students to use this section. Youmay wish to suggest other material that youyourself have found interesting and enjoyable.

Illustrations and Photographs The module isextensively illustrated to provide relevant andstimulating visual material to enable students torelate chemistry to everyday life, as well as toprovide for provocative discussion. In using someof these illustrations, it is not the intention ofIAC to endorse any particular product or brand,but only to relate chemistry to life outside theclassroom. As you proceed through each section,encourage students to collect, display, and dis-cuss photos and illustrations that provoke furtherdiscussion.

Questions A number of questions have beeninterspersed throughout the student module, inaddition to the questions that are naturally builtinto the narratives and the laboratory experi-ments. You will find some of these in specificallymarked sections in the student module. Thesequestions can be used in a variety of ways asyou see fit. They are not planned as testsremember, the IAC program was designed sothat mastery of the content and skills wouldbe achieved through the repeated reinforcementof ideas and procedures encountered as a studentprogresses through the various modules.

Managing the Laboratory

In the teacher's guide, hints and suggestions aregiven for managing each experiment in the labo-ratory. Share as many of these hints as possible

2

with your students. Allow them to participatefully in successful laboratory management. Makesure that you rotate assignments so that all stu-dents get a chance to experience this type ofparticipation.

Preparations and Supplies Student aides canbe helpful in preparing solutions, labeling andfilling bottles, cleaning glassware, and testingexperiments. (You should still test each experi-ment in the module to determine any revisionneeded to meet the needs of your students.)

Cleaning Up Involve your students in puttingaway equipment, washing up glassware, andstoring material for the next time it is to be used.Taking care of equipment is part of responsibili-ties we seek to foster in the students' outsideenvironment.

Laboratory Reports You may have your ownmethods of student reporting. We will includesome of the suggestions that IAC teachers havefound successful in the past.

It is helpful for students to keep a laboratorynotebook. A quadrille-ruled laboratory notebookwith a sheet of carbon paper allows a studentto produce two data sheets and report summarycopies. One copy of each page can be permanentlyretained in the notebook, while the duplicate copycan be submitted for evaluation or tabulation.

A realistic view of laboratory work suggeststhat, in the most fundamental sense, there areno wrong laboratory results. All students obtainresults consistent with particular experimentalconditions (either correct or incorrect) that theyestablished. Careful work will yield more preciseresults, of course. Encourage each student to takepersonal pride in experimental work. If studentsdisagree on a result, discuss the factors that might

account for the difference. A student who pro-vides a thoughtful analysis of why a particularresult turned out to be "different" (incom-plete drying, a portion of the original samplewas spilled, etc.) deserves credit for suchinterpretation.

Laboratory Safety To use the IAC programsafely, you should become thoroughly familiarwith all student activities in the laboratory. Doall the experiments and carry out all the demon-strations yourself before presenting them to yourclass. We have tested each experiment and havesuggested the use of chemicals that present theleast chance of a problem in laboratory safety.This teacher's guide has many suggestions forhelping you provide your students with safelaboratory experiences. Have the students readAppendix I: Safety. Then conduct a brief reviewof laboratory safety before the students encountertheir first laboratory experience in this module.Review safety procedures when necessary anddiscuss caution and safety each time a safetysymbol appears in the student text.

Materials for IAC In light of increasing costs forequipment and supplies, as well as decreasingschool budgets, we have tried to produce a ma-terials list that reflects only the quantities neededto do the experiments, with minimal surplus.Thus, the laboratory preparation sections containinstructions for only a 10-20 percent surplus ofreagents. Add enough materials for student re-peats and preparation errors.

Evaluating Student Performance

There are many ways of evaluating your stu-dents' performance. One of the most importantforms of evaluation is observing your studentsas they proceed through the IAC program. IAChas developed skill tests and knowledge tests foruse with this module. These test items have beensuggested and tested by IAC classroom chemistry

teachers. You are encouraged to add these to yourown means of student evaluation.

In addition to the questions incorporated inthe student module test and illustration captions,there are suggestions for evaluation at the endof each module section in the teacher's guide.The module tests are at the end of the teacher'sguide. Answers to all of the evaluation items areincluded to help you in your classroom discussionand evaluation.

10 3

Module Concepts

INTRODUCING PHYSICAL CHEMISTRY

Physical chemistry is the branch of chemistrythat deals with measuring the properties of chem-ical systems and with exploring the reasons for

these properties.Macroscopic properties are the result of the in-

teractions among individual molecules, atoms,and ions.

Molecules are extremely small, and a macro-scopic amount of material contains an enormousnumber of molecules.

THE GASEOUS STATE

Four important properties of a gas are mass,volume, temperature, and pressure.

The pressure of a gas is the force per unit areaexerted by the gas against the walls of its

container.Temperature is a measure of the average kinetic

energy of the molecules.The product of the pressure and volume of a

fixed amount of gas at constant temperature is aconstant (Boyle's law).

The volume of a fixed amount of gas at con-stant pressure is a linear function of the tem-perature (Charles' law).

Equal volumes of two gases at the sametemperature and pressure contain equal numbersof molecules (Avogadro's law).

Absolute zero ( 273°C) is the temperature atwhich the extrapolated volume of a gas equalszero.

The kinetic molecular theory explains the prop-erties of a gas and how they are dependent onone another.

The various relations among the volume, tem-perature, pressure, and amount of a gas aresummarized by the ideal gas law.

LIQUIDS AND SOLIDS: CONDENSED STATES

Liquids and solids exist because of attractiveforces between the molecules.

The London force is a weak attractive forcewhich exists between molecules.

Polar attractions can exist between polar mole-cules or ions.

4

An equilibrium is established between a liquidphase and gas phase when the rate of vaporiza-tion is equal to the rate of condensation. Theequilibrium is a dynamic one.

Boiling occurs when the vapor pressure of aliquid is equal to the pressure of the vapor phase.Lowering the pressure causes the boiling tem-perature to be lowered.

Solids may be classified into four crystallinetypes (molecular crystals, ionic solids, metals, andnetwork solids) and amorphous solids.

SOLUTIONS: SOLUTE AND SOLVENT

A solution is a homogeneous mixture of asolvent and solute.

An equilibrium is established between a solidand a solution when the rate of dissolving isequal to the rate of precipitation.

A supersaturated solution is more concentratedthan the equilibrium concentration of a solute.

A solution freezes at a lower temperaturethan does the pure solvent.

THE COLLOIDAL STATE

Colloidal systems contain particles, fibers, orfilms with at least one dimension in the rangeof approximately 1 nm to 1000 nm.

The properties of colloidal systems are largelythe result of their enormous surface areas.

A colloidal dispersion has properties that de-pend on the nature of the dispersed phase andthe continuous phase.

Surfactants are often necessary to stabilize acolloidal dispersion.

A gel is a semirigid material containing col-loidal fibers or crystals and a liquid.

CHANGES IN ENERGY

The thermal energy of a chemical system ismeasured by its temperature.

The chemical energy of a chemical system de-pends on the way in which the molecules areorganized.

The thermal energy of one gram of water in-creases by one calorie when its temperature israised by one degree Celsius.

The change of temperature measured in acalorimeter may be used to calculate changes inthermal energy and chemical energy.

11

An endothermic reaction is one in whichthermal energy is changed into chemical energy.The opposite change occurs in an exothermicreaction.

RATES OF CHEMICAL REACTIONS

The rate of a reaction is increased by increasingthe concentration of a reactant, by increasing thetemperature, or by adding a catalyst.

Reactions often occur in successive elementarysteps.

Colliding molecules must have a minimumkinetic energy, called the activation energy, beforereaction can occur.

A catalyst increases the rate of a reaction bylowering the activation energy of an elementarystep.

A catalyst is not consumed in a reaction.

CHEMICAL EQUILIBRIUM

Chemical equilibrium is established wheneach forward reaction has the same rate as itsreverse reaction.

At chemical equilibrium, the concentrationquotient of a reaction is equal to a constant called

the equilibrium constant. The value of the equi-librium constant is affected by temperature butnot by the presence of a catalyst.

An equilibrium may be shifted by adding orremoving a reactant or a product.

A reaction having a very large equilibrium con-stant goes to completion at equilibrium, whereasa reaction having a very small equilibrium con-stant proceeds to a very small extent.

The ion product of water is equal to 10-14(mole/liter)2.

The hydrogen ion concentration of a solution isconveniently expressed by the pH.

An acid-base indicator is a weak acid or a weakbase with different colors in the ionized andnonionized forms.

ELECTROCHEMISTRY

An electrochemical cell converts chemical en-ergy into electrical energy.

In an oxidation-reduction reaction, one sub-stance is oxidized and another is reduced.

Two half-reactions take place in an electro-chemical cell: an anode reaction (oxidation) anda cathode reaction (reduction).

In electroplating, electrical energy is used todeposit a metallic film on a cathode.

125

Module ObjectivesWe have made an attempt to group objectivesin th'ree broad categories: concept-centered,attitude-centered, and skill-centered. The cate-gories are not mutually exclusive; there isconsiderable overlap. The conditions for eachobjective are not given here, since they caneasily be found in the respective section in thestudent module. Note also that the concepts andskill objectives are more specific than those in theaffective domain. We readily acknowledge thedifficulties in trying to classify objectives in this

way, but we have been encouraged to do so byclassroom teachers who have helped us withthis task.

Evaluation items keyed to these objectives arepresented at the end of each major unit, andtest items covering major concepts discussed inthis module appear at the end of this teacher'sguide. The choice of level for both the objectivesand the evaluation items ultimately belongs toyou. But remember, the level of the testing itemshould be related to that of the correspondingobjectives.

Concept-Centered Objectives Attitude-Centered Objectives Skill-Centered Objectives

INTRODUCING PHYSICAL CHEMISTRY

P-1

State or identify distinguishingfeatures of the macroscopic andmolecular worlds.

P-3List the four states of matter.

Recognize that manymeasurements need not be exactto be useful.

Attempt to comprehend theenormous numbers of moleculespresumed to be present in eventhe smallest visible portions ofmatter.

P-2Express numbers in scientific

notation.Estimate the thickness of a

single-molecule film layer, giventhe volume and radius of thecircular film.

Convert a measurement given incentimeters to a measurementgiven in nanometers.

THE GASEOUS STATE

P-4List four major measurable

properties of a gas and identifycommon units used by chemists tomeasure each.

Explain how a mercurybarometer measures atmosphericpressure.

P-5State the qualitative and

quantitative effects of changingpressure on the volume of a gas.

Identify a correct statement ofBoyle's law.

P-6Explain macroscopic properties

of gases (mass, volume,temperature, pressure) in terms ofthe kinetic molecular theory.

6

Agree that barometric pressuremeasurements are actuallymeasures of length.

Be willing to test and accept amodel in terms of its ability toaccount for observations and topredict new situations.

Accept the kinetic moleculartheory as a valid model fordescribing some physicalproperties of gases.

Relate the discussion of gasbehavior to everyday observations.

Accept the proposition thatmathematical language provides aprecise means for expressingphysical laws.

i3

P-4Determine the pressure

difference in two gas containers,each connected to an arm of amanometer.

Relate one atmosphere ofpressure to a specific height ofmercury in a barometer.

Determine the pressure of a gasin a manometer, given themanometer column heights andthe atmospheric pressure.

P-5Determine the pressure of a gas

in a capillary tube partiallyfilled with mercury.

Calculate the pressure-volumeproduct for a gas, givenappropriate laboratory data.


P-7

Explain how temperature andvolume data for a gas can be usedto determine the value of absolutezero.

Describe qualitatively andquantitatively the relationshipbetween temperature andvolume of a gas.

P-8

Describe the features of anideal gas.

State or recognize Avogadro'slaw.

State the ideal gas law in wordsor symbols.

Describe the qualitativerelationship between the number ofgas molecules and the gaspressure in a container.

P-7

Extrapolate a temperature-volume graph for a gas to findnew temperature-volume values.

Construct and interpret a graphof volume as a function oftemperature for a gas.

Calculate a temperature inkelvins, given the temperature indegrees Celsius.

LIQUIDS AND SOLIDS: CONDENSED STATES

P-9

Express in terms of Coulomb'slaw the relationship betweenattractive force, charge, anddistance.

Relate the strength of a Londonforce to the number of electronsin a molecule.

Explain observed trends in theboiling point of related moleculesin terms of London force.

P-10

Explain the general physicalproperties of polar compounds(boiling points and solubility) interms of their structures.

Identify from a list of compoundsthose that are capable of forminghydrogen bonds.

P-11

Account for the vaporization of aliquid by using a molecularviewpoint.

Explain the term vapor pressure.

Recognize that molecularstructure and shape helps todetermine macroscopic propertiesand behavior.

Seek out and discuss everydayapplications of basic chemicalprinciples.

14

P-10

Classify molecules as polar ornonpolar by examining molecularmodels or their structural formulas.

P-13

Demonstrate the effect oflowering the pressure on a liquid'sboiling point.

P-14

Classify each substance in alist of solids as either a molecular,ionic, metallic, or network crystal.

Build molecular models toillustrate various types of solidstructures.

Classify given solids as eithercrystalline or amorphous.

7


Explain the concept of dynamicequilibrium.

Describe liquid-vapor equilibriumboth from a macroscopic andmolecular point of view.

P-12Describe the conditions needed

for a liquid to boil.Define normal boiling point.

P-13Predict how given changes in

atmospheric pressure will affectthe boiling point of a liquid.

P-14Identify physical properties

characteristic of each of the fourtypes of crystalline solids.

SOLUTIONS: SOLUTE AND SOLVENT

P-15Illustrate or identify correct

use of the terms solute, solvent,miscible, and immiscible.

Explain the process of a soliddissolving in a liquid from amolecular viewpoint.

Describe in macroscopic andmolecular terms the opposingprocesses and equilibrium existingin a saturated solution that isin contact with excess solid solute.

P-17Describe and give an example

of the relationship betweenpressure and solubility of a gas ina liquid.

P-18Describe and explain how the

freezing point of a liquid ischanged by the presence of adissolved solute.

8

Recognize that solubilityinfluences such diverse things ascarbonated beverages, deep seadivers, antifreeze, and fish survival.

15

P-15Prepare a saturated solution,

given adequate quantities ofsolvent and solute.

P-16Produce crystals by seeding a

supersaturated solution.Classify a solution as

unsaturated, saturated, orsupersaturated, by observingwhat happens when solid soluteis added.


THE COLLOIDAL STATE

P-19

Distinguish the properties of atrue solution, a colloidal dispersion,and a macroscopic precipitate.

P-20

Explain the role played bysurfactants in colloid stability.

P-21

Identify the dispersed phase andthe continuous phase for ordinarymaterials.

Identify the nature of thedispersed and continuous phasesfor any of these colloidaldispersions: fog, smoke, emulsion,sol, foam (liquid and solid).

Give examples of different typesof colloidal dispersions.

P-23Describe the characteristic

structure of a gel.Distinguish a sol from a gel on

the basis of structure.

Try to comprehend the immenseinfluence that particle size has onthe total surface area of asubstance.

P-19

Demonstrate the differencebetween a colloidal dispersion anda true solution, given samples ofeach.

Determine whether a liquidcontains a macroscopic precipitateor a colloid, given dialysis tubingand filter paper.

P-22

Prepare an oil-in-water and awater-in-oil emulsion.

Classify emulsions as eitheroil-in-water or water-in-oilemulsions by simple physical tests.

P-23

Prepare a gel, given ethanoland saturated calcium acetatesolution.

CHANGES IN ENERGY

P-24Explain at the molecular level

thermal energy and temperature.Distinguish thermal energy from

chemical energy.

P-26

Describe energy changesaccompanying melting of a solidsubstance.

Distinguish an endothermicprocess from an exothermicprocess.

Sense the importance of theinterplay between macroscopicobservations and molecular-levelexplanations in scientific activity.

16

P-24

Calculate the change in thermalenergy in a mass of water, givenits temperature change.

P-25

Determine the heat of fusion ofice experimentally, using acalorimeter.

Calculate the molar heat offusion of a substance, given thespecific heat of fusion and themolar mass of the substance.

9


P-27Predict whether a solution

process will be endothermic orexothermic, given the relativestrengths of attactive forceswithin the solid solute and betweenthe dissolved solute and solvent.

Calculate the molar heat ofsolution of a known substance,given the temperature changewhen a certain mass is dissolvedin a given volume of water.

P-29Explain in terms of molecular

changes the key events observedwhen gasoline ignites in anengine's cylinder.

Distinguish the dietetic Caloriefrom the scientific (metric) calorie.

P-27Determine experimentally the

molar heat of solution of a knownsolid substance in water.

P-28Measure the temperature

change that takes place during achemical reaction.

Determine experimentally themolar heat of reaction for waterformation in a neutralizationreaction, given the reactants, thebalanced equation, and acalorimeter.

Calculate the molar heat ofreaction for a neutralizationreaction from laboratory data.

RATES OF CHEMICAL REACTIONS

P-30Decide whether a reaction rate

is proportional to a reactantconcentration, given reactiontimes for different concentrationsof the reactant.

P-31Explain the effect of

concentration on reaction rate froma molecular viewpoint.

Explain the relationship betweenthe rate-determining step and theoverall rate of a reaction.

P-32List or identify reasons why

increasing the temperatureincreases the reaction rate.

Explain the significance ofactivation energy in chemicalreactions.

P-33Identify reasons for the effect

catalysts have on reaction rates.

10

Consider everyday tasks andevents in terms of reaction ratesand a rate-determining step.

Recognize the importance ofcatalysts (enzymes) in our bodychemistry.

17

P-30Demonstrate experimentally the

effects of concentration,temperature, and a catalyst on therate of a chemical reaction.


CHEMICAL EQUILIBRIUM

P-34

Identify the characteristics of asystem at chemical equilibrium.

Distinguish the features of anequilibrium system from those of asteady state system.

Express the mathematicalrelationship between rate constantsand equilibrium constant.

P-35

Describe the effects on achemical equilibrium of addingmore reactant or product.

Write an expression for theconcentration quotient for a systemat equilibrium, given the overallchemical equation for the reaction.

P-36Describe the final state of an

equilibrium system in terms ofrelative amounts of reactants andproducts, given the magnitude ofthe equilibrium constant.

P-37

Choose the corrrect expressionfor the ion product of water, givena list of alternatives.

Convert molar concentrations ofhydrogen ions into pH values, andvice versa.

Relate given pH values to therelative acidity or basidity of thesolution.

Explain, from a molecularviewpoint, what happens when astrong acid or strong base isadded to water.

Distinguish strong acids andbases from weak acids and basesat the molecular level.

Seek out nonchemicalillustrations of both steady stateand dynamic equilibrium systemsand share them with the class.

18

P-35

Demonstrate experimentally theeffect of changing concentrationson a system at equilibrium.

P-38

Prepare by dilution a solutionhaving one-tenth the concentrationof the original solution.

Determine the pH or hydrogenion concentration of an acidsolution, using acid-base indicatorsand known-concentration acidsolutions.

Determine experimentally theionization constant for a weak acid.

11


P-38Describe the role of pH in

determining the color of anindicator in solution.

Predict the color of an indicatorin a solution, given the indicator'sequilibrium constant, colors ofacid and base forms, and thehydrogen ion concentration of thesolution.

ELECTROCHEMISTRY

P-39Distinguish between the

processes of oxidation andreduction.

P-40Distinguish between the terms

anode and cathode in regard tothe kinds of reactions that occurat each electrode.

Classify a given half-reaction asrepresenting either an oxidationor a reduction process.

P-41Identify the half-reactions that

occur when a lead storage batteryis discharging or charging.

Explain why a liquid densitymeasurement for a lead storagebattery is a reliable indication ofthe cell's available electricalenergy.

P-42Contrast the operation of an

ordinary dry cell with that of afuel cell.

P-43Cite ordinary examples of

electroplated materials.Describe the process of

electroplating.

12

Consider the everydayimportance of electrochemistry inour own lives.

P-39Decide which of two possible

oxidation-reduction reactionsinvolving two metals and theirrespective ions actually takesplace, given the two metals andsolutions of their metal ions.

P-40Construct an operating

electrochemical cell.Investigate the conditions under

which an electrochemical cell willoperate by varying thearrangement of the components.

P-41Construct a lead-acid cell,

observing and interpreting itsbehavior while it is charging anddischarging.

P-43Construct and operate a simple

electroplating setup.

19

Teaching Communities of MoleculesIntroducing Physical Chemistry

The first three sections give the students a per-spective on the role of chemistry in everydaylife and on the size of molecules in comparisonwith macroscopic chemical systems. First, there isa list of seven questions about applications ofphysical chemistry to automobilesquestionsthat we hope will arouse the students' curiosity.(Each question will be raised again and exploredfully as students move through Communities ofMolecules.) Point out to your students that thesequestions are examples of practical applicationsof physical chemistry, not examples of examina-tion questions.

Ask your students to try to cite other examplesof materials that illustrate the concept of com-munities of molecules such as those in the exam-ple on page 2the melting of ice cubes and themolecule of water.

In this student module discussion, physicalchemistry is called a branch of chemistry. Inter-ested students may wish to find out the name ofother branches of chemistry and what each in-cludes. Help them discuss the area of physicalchemistry by asking them to find out what aphysical chemist does. Look over the table ofcontents for this module to get a brief idea. Theauthor-biography tells what one physical chemistdoes. Some students may wish to find out aboutothers in this field.

There is a discussion of the macroscopic andmolecular worlds, and an experiment in whichstudents measure the size of an oleic acid mole-cule for comparison with familiar macroscopicdimensions. Finally, in order to drive the pointhome that most molecules are extremely smalland to illustrate the high complexity of somechemical systems, the comparative sizes of vari-ous particles in milk are described by an imagi-nary "journey" into a drop of milk.

P-1 MACROSCOPIC AND MOLECULAR:TWO WORLDS

Usually the two extremes of large size and smallsize are referred to by the terms macroscopic andmicroscopic. We have chosen the term molecular

in place of microscopic to eliminate any implicationthat all molecules can be seen in a microscope.

EXPERIMENTP-2 THE SIZE OF A MOLECULE

The purpose of this experiment is to measure the ap-proximate size of an oleic acid molecule, and to gainexperience with expressing numbers in scientificnotation.

Concepts In doing this experiment, a student will en-counter these important ideas:*

When placed in water, oleic acid forms a mono-molecular film on the surface of the water.The thickness of the oleic acid film may be calculatedif the volume and area are known.The length of the oleic acid molecule is approximately3 x 10-7 cm, or 3 nm.Molecular dimensions are conveniently expressed innanometer units.

Objectives After completing this experiment, a studentshould be able to:*

Calculate the thickness of a monomolecular film,given the volume and the radius.Express numbers in scientific notation.Change a measurement in centimeters to a measure-ment in nanometers.

Estimated Time One period

Student Grouping Pairs

Materials** You will need the following for a class of30 students working in pairs:

4-8 dropper bottles, each containing 10 cm3 oleic aciddiluted 1:200 in ethanol

10 g chalk dust or lycopodium powder15 meter sticks15 shallow trays, approximately 60 cm x 60 cm x 5 cm

(e.g., ripple tanks)8 10-cm3 graduated cylinders

This statement appears only with this first experiment, but it applieseach time this section appears in an experiment, unless otherwise noted.

The Materials list for each laboratory experiment in this module isplanned for a class of thirty students working in pairs. You may haveto adjust this to fit the size of your class.

20 13

Advance Preparation Prepare 100 cm3 of the oleicacid solution by placing 5 cm3 oleic acid in a 50-cm3graduated cylinder and adding ethanol to a final volumeof 50 cm3. Mix this solution, put 5 cm3 of the solutioninto a 100-cm3 graduated cylinder, add more ethanolto get a final volume of 100 cm3, and mix. Put theresulting solution into dropper bottles.

Pre lab Discussion The chemical structure of oleicacid is:

H3C(CH2),CH = CH(CH2),CO2H

The carboxyl (CO2H) end of the molecule is polar andis attracted into the water. The rest of the molecule isa nonpolar hydrocarbon chain with little attraction towardwater, and so in the film this portion stands up per-pendicular to the surface of the water and parallel tothe chains of neighboring molecules. The alcohol in thesolution dissolves in the water underneath the film.

Students will need some help in understanding how thevolume and area of the drop can be used to calculatethe thickness of the film. Try taking a 100-cm3 graduatedcylinder full of Rice Krispies, glass beads, or similarmaterials and spreading them out one-grain thick on thelab table in a roughly circular pattern. Discuss whetherthe volume of the Rice Krispies has been changed bythis process. Estimate the diameter of the pattern and getthe students to calculate the area of the pattern. Pointout that the pattern is actually a short, wide cylinder,and that it is the height of the cylinder which interestsus. Help the students to see that

total volume = area of base x height.

If this calculation is actually performed, the studentswill have a fairly accurate estimate of the size of onegrain.

Remind the students that the volume of oleic acid theywill use is only 1/200 of the volume of one drop. Youcan demonstrate this visually by showing them 2.5cm3 of oleic acid in a small graduated cylinder and a500-cm3 volumetric flask representing the final volume.

Have the students read Appendix II in the studentmodule on powers of ten and scientific notation and askthem to express as many of their numbers in this nota-tion as they can. If this is entirely new to them, youmay wish to give some specific examples.

You may also wish to demonstrate that oleic acid andwater are immiscible and that the oleic acid floats onthe surface.

14

Laboratory Safety Remind your students that thesafety rules they followed in their previous laboratoryexperiences in !AC Chemistry also apply to the studyof physical chemistry. Remind them to be on the lookoutfor situations that may need special safety precautionsas they proceed through a discussion of the first severalexperiments.

Laboratory Tips An alternative to the ripple tanks isa disposable version made by taping meter sticks orother sticks to the laboratory table in a .5-meter squarepattern and draping a plastic sheet inside.

The ripple tanks, if used, must be clean.

Do the experiment in advance so you can judge abouthow much chalk dust should be used. If you use toolittle, the circular film will be poorly defined. If you usetoo much, the film will not be able to spread out properly.

Lycopodium powder works better than chalk dust. If

chalk dust is used, the water may be stirred to breakup clumps of chalk dust, but have the students waituntil the water has stopped moving before adding theoleic acid drop.

Range of Results Typical results are:drops in 1 cm3 = 45 dropsvolume of 1 drop = 1 cm3/45 = 2.2 x 10-2 cm3volume of oleic acid in 1 drop = 2.2 x 10-2 cm3/200

= 1.1 x 10-4 cm3average radius = 11 cmarea = (3.14)(11 cm)2 = 3.8 x 102 cm2thickness (length) = 1.1 x 10-4 cm3/3.8 x 102 cm2

= 2.9 x 10-7 cmthickness (length) = 2.9 nm

Postlab Discussion The actual length of the oleicacid molecule is known to be 2.7 nm.

Since the answer from the experiment is unprecise, itmight be worthwhile to consider some of the sourcesof error in this experiment. Return to the Rice Krispiesmodel. Suppose the grains do not spread out all the way.What effect would this have on the results? Supposethe film has some bare spots or holes in it. What effectwill this have on the answer?

If you want your students to have more practice in theuse of scientific notation, there are two additional cal-

culations that can be done in class.

1. Suppose the oleic acid molecules are cubical inshape. The approximate volume is then equal to

'1

(length)3. You can now estimate the number of mole-cules contained in the film by dividing the total volumeof the film by the volume of one molecule. For thetypical data given above

volume of molecule = (2.9 x 10-7 cm)3=2.4 x 10-20 cm3

number of molecules1.1 x 10-4

2.4 x 10-2°cm3cm3

= 4.6 x 1015 molecules

2. Given the fact that the volume of one mole (282 g)of oleic acid is 317 cm3, it is possible to estimateAvogadro's number (the number of molecules in onemole)

3.17 x 102 cm31.3 x 1022

2.4 x 10-20 cm3

The correct value of Avogadro's number is 6 x 1023.

Do not overdo these calculations. If the students areuncomfortable with the mathematics, you could domore damage than good.

P-3 JOURNEY INTO A DROP OF MILK

The students are not to memorize the detailedstructures and dimensions given in this section.The structures and dimensions are given onlyto impress upon the students that a seeminglysmall sample of material contains a very largenumber of very small particles.

You can go over the calculations mentioned atthe end of this section to impress the studentsfurther with this fact. A 1-cm2 area layer of watermolecules, one molecule thick, contains (3 x 107)2= 9 x 10" molecules. A 1-cm3 volume cube con-tains (3 x 107)3 = 3 x 1022 molecules.

Suggested Reading Boeke, Kees. Cosmic View:The Universe in 40 Jumps. New York: The JohnDay Company, 1957. This ingenious book carriesout a scheme similar to the "journey" into a dropof milk. The book consists of 40 drawings withtext. Each drawing illustrates objects of a differentpower of ten in size. The objects range in sizeall the way from the nucleus of a sodium atomto groups of entire galaxies! The author andartist is a Dutch schoolmaster who has lavishedimmense care in rendering the details accurately.Students who read this book will not learn muchchemistry, but they will gain a "cosmic view"

of the universe and its varied structures at dif-ferent scales.

Film Powers of Ten (Charles Eames Associates.Distributed by Pyramid Films, P.O. Box 1048,Santa Monica, California 90406; color, 8 minutes).This film carries out the same scheme as the bookCosmic View, but in an even more dramaticfashion. The depiction of molecules is, unfor-tunately, vague. A new version of this film wasreleased in 1978. For rentals, contact Film Distri-bution Center, Division of Cinema, Universityof Southern California, University Park, LosAngeles, California 90007; or The PennsylvaniaState University Audio-Visual Services, 17 Wil-lard Building, University Park, Pennsylvania16802.

ANSWERS TO QUESTIONS

(Student module page 7)

1. 0.000 54 = 5.4 x 10-48346 = 8.3 x 10341 648 218 = 4.2 x 107

2. (3 x 106) x (4 x 10-6) = 1 x 10-1

18 x 10-26 x 10-6

3 x 104

31 0001.0 x 105

0.31

3. The length of the macromolecules shown in Figure7 is 1 x 10-6 cm, or 10 nm.

volume 1 x 104 cm34. area = 1 x 1010 cm2

thickness 1 x 10-6 cm

length = \/area = x 1010 cm2 = 1 x 105 cm

length = 1 x 105 cm x11cm

m 1 x 103m

1length = 1 x 10-3 m x103

km

m1 km

Thus only 10 liters of oil are capable of forming anoil slick that is 1 km (0.6 mile) wide! The large sizeof the slick is caused by the small size of the mole-cules in the oil, which determines the minimumthickness of the film.

22 15

5.2.99 x 10-23 gram/molecule

= 3.34 x 1022 molecules

1.00 gram

6.

When this answer is rounded off to one significantfigure, it is the same as the answer calculated insection P-3 (3 x 1022 molecules).

3.34 x 1022 molecules4 x 109 people

= 1 x 1013 molecules per person.

This answer illustrates the astronomical differencebetween the population of a molecular communityand the world's population. It would take each person,counting at the rate of one molecule per second,3,000 centuries to count his or her share of themolecules!

EVALUATION ITEMS

These are additional evaluation items that you may wishto use with your students at various times during thepreceding section. The correct answer to each questionis indicated by shading.

1. In a crystal of platinum, the centers of the atomsare 2.8 x 10-6 cm apart. This distance is equal to:

A. 280 nm B. 28 nm C. 2.8 nm D. 0.28 nm

2. When a drop containing 0.000 10 cm3 of oleic acidis allowed to spread out on the surface of a trayof water, a circular film 20 cm in diameter is formed.From these data, the approximate thickness of anoleic acid molecule is:

A. 4.8 nm B. 16.0 nm C. 3.2 nmD. 6.4 nm E. 2.4 nm

The Gaseous State

A gas is the simplest state of matter in the sensethat there are no interactions between the mole-cules. This results in complete randomness ofstructure. The lack of interactions makes gaseseasy to describe quantitatively (by the ideal gaslaw), and sets them aside from the condensedstates that have pressure-volume-temperaturerelations that cannot be described so easily.

Students can deduce Boyle's law and Charles'law from their own measurements, using a capil-lary tube gas container. After each law is deduced,the text explains it on the molecular level by thekinetic molecular theory. At the same time, thetext explains the molecular basis of temperatureand the meaning of the kelvin scale. Finally,Avogadro's law is stated and the three laws(Boyle's, Charles', and Avogadro's) are combinedinto the ideal gas law.

Because this major section is self-contained andthe material in it is not needed in the later sec-tions, you might consider a quick review orskipping the material entirely if you are short of

16

time or if your students have encountered thematerial elsewhere.

P-4 MEASURING GASES

If your students are not yet familiar with themeasuring instruments on page 9, you may wishto demonstrate these in connection with yourdiscussion of the four important properties of agas. Ask a student to bring in a tire-pressuregauge to show to those students who have neverlooked closely at one. On page 11, two types ofbarometers are illustrated. You may wish to obtainboth and illustrate how each is affected in a dif-ferent way by changes in pressure.

Your students might be interested in some moredetails of how the height of the mercury columnin a barometer is used to measure the atmosphericpressure. The fundamental definition of pressureis that pressure equals the force per unit area.The SI-derived unit for force is the newton, andfor area it is the square meter. Thus in the SIsystem pressure is measured in units of newtonsper square meter, or pascals. That is, one pascalis defined as a pressure of one newton per squaremeter.

23

In a mercury barometer, the height of the mer-cury column adjusts itself until the force exertedby the weight of the mercury on a unit area atthe level of the lower surface is equal to theatmospheric pressure. The force in question is thesame as the total downward force of a separatevertical column of the mercury (a downward forceresulting from gravity), having a unit cross sec-tion in a horizontal plane and the same verticalheight as the mercury column in the barometer.From Newton's second law, this force is equalto the mass of such a column of mercury timesthe acceleration resulting from gravity. Simplegeometry shows that the mass of the column ofunit cross section is equal to the density of themercury times the height.

Putting all this together, we get the final result

P = dgh

where P = atmospheric pressure, d = density ofthe mercury, g = acceleration due to gravity,and h = height of the barometer column.

From this equation, it is evident that theob-served height h in the barometer is an accuratemeasure of the atmospheric pressure P only if thequantities d and g have standard, nonvaryingvalues. However, the density d depends on thetemperature of the mercury, and the accelerationg due to gravity at the earth's surface dependson the latitude. For this reason, in precise deter-minations of the atmospheric pressure expressedin cm Hg, corrections are made to the measuredheight of the barometer column to correct it to theheight it would have at a temperature of 0°C(at which temperature the density of mercury hasthe standard value 1.3595 x 104 kg m-3) and ata latitude of 45° (at which latitude the accelera-tion due to gravity has the standard value 9.8067km s-2).

You will not need to worry about these cor-rections for the next miniexperiment. At roomtemperature, the temperature correction to thebarometer reading is less than 0.5 percent of thereading. The latitude correction, even at the polesor the equator, where it would be greatest, is evenless than the temperature correction.

One atmosphere is defined as the pressure thatwill give a barometer reading of exactly 76 cmHg after the corrections for temperature andlatitude have been made. The relations between

various units that can be used to express pressureare given by the following conversions:

1 atmosphere = 76 cm Hg (corrected)= 29.921 inches Hg (corrected)= 1.013 25 x 105 pascals= 1013.25 millibars= 14.7 pounds per square inch.

MINIEXPERIMENTP-5 PRESSURE AND VOLUME

In this experiment the student measures the length ofan air column at three different pressures and verifiesBoyle's law.

Concepts

The difference in height of the two ends of a columnof mercury in a manometer indicates the differencein the pressure of the gases at each end.The product of pressure and volume for a fixed massof gas at constant temperature is a constant.

Objectives

Calculate the pressure of the gas enclosed in amanometer, given the height of the mercury columnand the pressure of the gas at the other end.State the effect of changing the pressure upon thevolume of a gas, both qualitatively and quantitatively.

Estimated Time One-half period


Materials

15 capillary tubes, made from 1.2-meter (4-foot) lengthsof 0.75 mm-1.25 mm bore capillary tubing

epoxy cement (resin and hardner)5 cm3 mercury1 hypodermic syringe15 15-cm metric rulers

Advance Preparation Five tubes about 24 cm inlength can be cut from a standard 1.2-m (4-foot) lengthof Pyrex capillary tubing (0.75 mm-1.25 mm bore).Cut the tubes to length by scoring with a triangular fileand breaking. Roughen the end of the tube to be epoxiedthis will improve the seal. Then place the tubes on ahorizontal surface. Insert the mercury, using a hypo-dermic syringe with a long needle. If necessary, tap thetube until the mercury column is about 8 cm from one

24 17

epoxy cement mercury

8 cm

VA2;07XeAllIZGel.MGC0717/.0=077 =.0=12107,1XXXXIMIXXIXAGOMMXWASWZMaria

end of the tuble. Then seal this end with epoxy cement.Do not disturb the finished tubes for at least 24 hoursso that the cement can harden. Store the tubes in atray to prevent spilling the mercury.

Pre lab Discussion Obtain the atmospheric pressurefrom a mercury or aneroid barometer. If you are closeto sea level (below 0.3 km elevation) you can take thepressure from a weather report; the weather service re-ports barometric pressures that have been corrected tosea level for comparison with other locations, so athigher elevations these corrected values are sub-stantially different from the actual atmospheric pressurein the classroom. Pressure in inches of mercury can beconverted to centimeters of mercury by multiplying by2.54.

Have the students examine the trapped air in the capil-lary tube. When the tube is held horizontally, the pressureof the trapped air is the same as atmospheric pressure.But, when the tube is vertical, the mercury column inthe tube either adds to the atmospheric pressure (openend up) or subtracts from it (open end down). Do notoverdo this if it appears obvious to most students, butmake sure that everyone understands how and why thepressure of the trapped air column is changing.

Ask the students to explain why measuring the length oftrapped air is a good way of keeping track of the volumeof the air. If you feel it would be less confusing, youcan have the students convert the length to the volumein cubic centimeters by the formula: volume 71-r2ewhere r is the radius of the capillary bore and is thelength.

Laboratory Safety Even though mercury is a curiosityand interesting to observe under certain circumstances,caution your students not to wear gold and silver jewelrywhen they work with this liquid metal. Also caution yourstudents not to throw loose mercury into the sink andnot to play with any mercury droplets that may escape.Remind them that liquid mercury and mercury vapor arepoisonous. You might ask one of your interested stu-dents to do some research on mercury poisoning andpresent this information before the class actually does18

24 cm

12cm

the experiment. Solicit from the students a discussionon using materials such as mercury in a positive andcommonsense manner versus the attitude that calls forbanning these items from the laboratory situation justbecause they are labeled poisonous. Ask them to pointout similar situations in their everyday lives.

Laboratory Tips Students should be cautioned tohandle the capillary tubes carefully. Sharp knocks cancause the mercury column to separate. Should thishappen, the capillary must be broken open, the mercuryremoved, and a new capillary constructed. Or, you maybe able to rejoin the column by inserting a clean copperwire and then slowly withdrawing it. It is wise to haveseveral extra capillary tubes in reserve in the eventthat some are broken.

A safe, easy way to clean used mercury for use in thecapillary tubes is as follows: Place the used mercuryin the bottom of a widemouth plastic bottle. Wash themercury with running water while stirring. Decant mostof the water off, leaving a small layer of water on topof the mercury. Put a cover on the bottle and placeit overnight in a freezer to freeze the water and trapany dirt. Remove the ice and dirt. The mercury is thenready to use.

If any mercury is spilled, it should be carefully collectedand placed in a stoppered bottle or absorbed withpowdered sulfur. It may also be picked up with mask-ing tape.

A mercury spill cleanup kit, containing a proprietarychemical formulation to absorb mercury vapor from theatmosphere, is available from J. T. Baker ChemicalCompany, Phillipsburg, New Jersey 08865.

Range of Results A typical set of results is:

atmospheric pressure = 76.2 cm HgHg column length = 12.3 cm

tube horizontalP = 76.2 cm Hglength = 7.9 cmP x length = 6.0 x 102 cm Hg cm

25

tube open end upP= 76.2 cm Hg + 12.3 cm Hg = 88.5 cm Hglength = 6.9 cmP x length = 6.1 x 102 cm Hg cm

tube open end downP= 76.2 cm Hg 12.3 cm Hg = 63.9 cm Hglength = 9.6 cmP x length = 6.1 x 102 cm Hg cm

The product P x length will vary from one capillary tubeto the next, because the tubes contain different amountsof air.

Postlab Discussion Most students' results will verifythat P x V = constant, but make certain they know whatthis means. Ask what will happen to the volume if thepressure is doubled, tripled, or cut in half.

Point out that Pill, = P2V2, where PI and V, are for case1 and P2 and V2 are for case 2. This relation will beuseful for problem solving.

Miniexperiment A Boyle's law experiment that per-mits the use of gases other than air is described in M.Spritzer, "Testing Boyle's Law," Chemistry, Vol. 43(October 1970), p. 29.

P-6 IT'S A HIT!Demonstration The molecular explanation of pres-sure can be demonstrated with a Random MolecularMotion Demonstration Kit, sold by Science Kit, Inc., ,777East Park Drive, Tonawanda, New York 14150 (catalognumber 16625). This consists of plastic foam ballsrepresenting molecules that are placed in a transparentcylinder having a movable piston. The apparatus mustbe placed on top of a Van de Graaff generator, whichcharges the balls and causes random motion. Show thatdecreasing the volume (by moving the piston) increasesthe collision rate on the walls and so increases thepressure. Also show that increasing the speed of theballs (by turning up the generator) increases the collisionrate, and that to return the pressure to its originalvalue the volume must be increased.

You can carry out a similar demonstration with thehomemade KMT apparatus described in sections A-16,A-17, and A-18 of the Teacher's Guide for Reactionsand Reason: An Introductory Chemistry Module.

Miniexperiment You may wish to try the "perfume"experiment that is discussed in the text of the studentmodule. If so, open a bottle of perfume, after-shave, or

cologne in one corner of the classroom and challengethe students to determine how long it takes for the mole-cules to travel to different parts of the room.

Ask several of your students to bring informationon tire pressure for their own cars or a family car.This information can usually be found in theowner's manual. Ask them to discuss why thereare different tire pressures for different situations.Some of your students may wish to compare thisinformation with tire pressure required for vari-ous types of bicycles and bicycling situations.

EXPERIMENTP-7 TEMPERATURE AND VOLUME

This experiment carries the observations on gas prop-erties one step further by varying the temperatureinstead of the pressure.

Concepts

The volume of a gas increases when the temperatureis increased at constant pressure.There is a linear relation between the volume of agas and its temperature.The volume of a gas extrapolates to zero at 273°C.

Objectives

State the effect of changing the temperature upon thevolume of a gas, both qualitatively and quantitatively.Construct and interpret a graph of volume as a func-tion of temperature.Extrapolate a straight line on a graph.



Materials

30-50 ice cubes15 capillary tubes (see experiment P-5 )15 thermometers, 10°C to 110°C1 or 2 400- or 600-cm3 beakers15 ring stands15 ring-stand rings15 wire gauze, asbestos centers15 Bunsen burners15 rulers, centimeter, heavy plastic (see Laboratory

Tips)30 sheets graph paper, inch or centimeter rulings

26 19

Advance Preparation If ice is in short supply, it maybe necessary to purchase a bag of ice cubes on theday the experiment is to be performed. Or you can startsaving them ahead of time in the freezer compartmentof a refrigerator, storing them in a plastic bag until youhave accumulated enough for the experiment.

Pre lab Discussion It will save time if the studentsbegin heating the water in one of the beakers so thatit will be ready when needed, unless sufficiently hot tapwater is available. Point out to the students that as longas the capillary tube is held upright, the pressure onthe gas will be constant. Bright students may wonder ifthe pressure will change with temperature because themercury column changes length. Point out that theweight of the mercury remains constant, so that the totaldownward force of gravity does not change. Also, thecross section area of the capillary bore does not changewith temperature (except for a very slight change result-ing from the expansion or contraction of the glass).Therefore, the force per unit area, or the pressure,remains unchanged.

Laboratory Tips Use hard plastic rulers. The thin,flexible type curls in hot water.

If you wish, you can have the students use the degreerulings on the thermometers to measure the lengths ofthe gas columns. It does not matter what the spacingof these divisions is as long as the spacing is uniformand the same thermometer is used for all of the mea-surements. Be sure the students make both the highand low temperature measurements while the capillarytube is immersed in the water bath. The ruler (or ther-mometer rulings) can be read through the side of thebeaker by bringing the tube close to the side.

Range of Results Typical results are:

room temperaturet = 25°Clength = 6.9 cmT = 25 + 273 = 298 Klength/t = 0.28 cm/°Clength/T = 0.023 cm/K

ice watert = 0°Clength = 6.3 cmT = 0 + 273 = 273 Klength/t = infinitelength/T = 0.023 cm/K

20

hot watert = 60°Clength = 7.7 cmT = 60 + 273 = 333 Klength/t = 0.128 cm/°Clength/T = 0.023 cm/K

The length-versus-t data are plotted on the followinggraph. Instruct the students to label their graphs asshown on the following page.

Post lab Discussion The temperature of the tubewould have to be about 300°C in order for the lengthof the gas column to double. In the example shown inthe graph, the temperature would have to be 122°Cin order for the length to be one half.

There may be considerable variation in the extrapolationof the temperature to the point of zero volume, sincea slight error in any of the measurements will have alarge effect. (This is a good opportunity to point out thatprojections beyond the range of measured values arealways subject to considerable uncertainty.)

The finding that VIT equals a constant for differentvalues of T means that the volume is inversely propor-tional to the temperature on the kelvin scale. Point outthat this relation can be expressed in the equationV, /T, = V2 /T2, where 1 and 2 represent cases 1 and 2.

Miniexperiment An alternate experimental procedurefor studying the volume-temperature behavior of a gascan be found in M. Spritzer and J. Markham, "Charles'Law: Estimating Absolute Zero," Chemistry, Vol. 42(September 1969), pp. 24-25. This article is reprinted inHarold W. Ferguson and Joseph Schmuck ler, eds., LabBench Experiments in Chemistry (Washington, D.C.:American Chemical Society, 1970), pp. 25-26. A moresophisticated version of the traditional Charles' law ex-periment that may be of interest to more advancedstudents can be found in L. Mc Keen, "An Experimen-tal Method to Study Charles' Law," Chemistry, Vol. 44(February 1971), pp. 27-29.

P-8 THERE OUGHT TO BE A LAW

Since 24 464 cm3 of any ideal gas at a tempera-ture of 25°C and a pressure of one atmospherecontains one mole of molecules, the mass of thisamount of gas in grams is equal numerically tothe relative molecular mass: 4.00 g for helium(He), and 28.0 g for nitrogen (N2). Thus a volume

27

tt

NIEMEN

summattill mmmmm

111111111N1

MN

UNE

II

I IIIIII

mumnimairanno

h....SESEmourion

a I meeUPTIFIVIMI

1111111111111:11

300 270 200 100 0 100

of helium has one-seventh the mass of an equalvolume of nitrogen, or only one-seventh the den-sity, which accounts for the buoyancy of helium-filled balloons.

A temperature of 0°C and a pressure of oneatmosphere are sometimes called standard tem-perature and pressure. Under these conditions,the volume of one mole of an ideal gas is22 413 cm3.



The solutions to questions 1, 2, and 3 for this unit aregiven in the margin of the student text. The remaininganswers and solutions follow.

temperature (T)

4. Pill = niRTP2V = n2RT

Therefore

P2V _n2RTPi V niRT

OrP2 _n2

n,

P2 (4.0 1.0) mol 3.0 mol2.4 atm 4.0 mol 4.0 mol

P2 = 2.4 atm x 3.0 mol /4.0 mol = 1.8 atm

PV 1 atm x 4 x 1010 cm35. nRT (82.1 cm3 atm/K mol) x (27 + 273)K

= 2 x 106 mol

2x 106 mole x 4.00 g/mol = 8 x 106 g

Of course, while this was a large mass of helium,an even larger mass of air was displaced by the blimp,

28 21

so that the blimp was "lighter than air" and could

rise.6. V2/V, =1),/P, = 1/27. T2/T, = 600 K/300 K = 2.00

V2/1/1 = T2/T, = 2.00Notice that the ratio of the temperatures (on the kelvinscale) equals the ratio of the volumesanother formof Charles' law. The ratio of the temperatures ex-pressed in degrees Celsius is 327°C/27°C = 12,a value that has no relevance to the ratio of volumes.

8. In order to increase the volume of a gas container,one moves the walls of the container farther apart.Thus, molecules have farther to travel between col-lisions with the walls. For a fixed number of mole-cules at a constant temperature, this means thateach unit of wall area is hit less frequently. Con-sequently, the pressure is lower.

9. nR = PV/TThus if n is constant: P,Vi/T, = P2V2 /T260 cm Hg x 200 cm3/200 K = P2 x 300 cm3/300 KP2 = 60 cm Hg x 200 cm3 x 300 K/300 cm3 x 200 K

= 60 cm Hg

There is no change in the pressure for the particularchanges in volume and temperature; the increasein volume, which tends to decrease the pressure, isexactly balanced by the increase in temperature,which tends to increase the pressure.

EVALUATION ITEMS


1. Which of the following is an acceptable unit forexpressing pressure?

A. cm3 B. g/cm3 C. nm D. cm Hg

2. 0.01 mol of a certain gas occupies a volume of 100cm3 at a pressure of 2.4 atm. If the temperature iskept constant and the pressure is lowered to 0.6 atm,what volume will the gas occupy?

A. 25 cm3 B. 100 cm3 C. 200 cm3 D. 400 cm3

3. Which of the following is a correct statement ofBoyle's law? (In these equations, k stands for aconstant.)

A. V/P = k B. V = kT 'C. PV = k D. PV = 82.1 nT

22

4. The capillary tube shown here contains a sample ofair trapped beneath a 2-cm slug of mercury. Theatmospheric pressure is 76 cm Hg and the tem-perature is 25°C.

(a) What is the total pressure of the trapped air?

A. 78 cm HgC. 2 cm Hg

B. 76 cm HgD. 74 cm Hg

(b) If the tube were placed in a horizontal position,what would the pressure of the gas be?

A. 78 cm B. 76 cm C. 2 cm D. 74 cm

(c) If the tube were placed in the horizontal

position, what would the length of the trapped aircolumn be?

A. 3.9 cm B. 4.0 cm C. 4.1 cm D. 4.2 cm

5. Complete each of the following statements aboutgases with one of these terms: increase, decrease,remain the same, can't be determined from theinformation given.

29

A. If more molecules are added to a containerof fixed volume while the temperature is heldconstant, the pressure in the container will

increase

B. In order for the pressure of a given sample ofgas to increase at a constant temperature, the

volume must decrease

C. If the temperature of a sample of gas in a rigidcontainer is increased, the pressure in the con-

tainer will increase

D. If the temperature of a gas is reduced, the average

velocity of the molecules will decrease

E. If the pressure of a given sample of gas is to re-main the same when the volume is decreased, the

temperature of the gas must decrease.

6. A temperature of 298°C is equivalent to which of thefollowing?

A. 571 K B. 298 K C. 25 K D. 0 K

7. Which of the following best states Avogadro's law?

A. The number of molecules of gas in a fixed volumeis constant.

B. The pressure of a gas is directly proportional tothe number of molecules.

C. At constant temperature and pressure, equal vol-umes of gas contain equal numbers of molecules.

D. One mole of any ideal gas occupies the samevolume at constant pressure.

8. The ideal gas law can be stated as follows: Thepressure-volume product of an ideal gas is directlyproportional to the temperature and the number ofmoles of the gas. Express this statement in symbolicform.

,PV = knT or PV = pRT

Liquids and Solids:Condensed States

The first two sections explain the forces betweenmolecules, which the text calls the London forceand polar attractions. These types of forces arelargely responsible for the condensation of gasesinto liquids and the change of liquids into solids.Following these rather theoretical sections aredescriptions of the dynamic equilibrium betweenthe liquid and vapor phases that are responsiblefor vapor pressure and boiling. Finally, there isa description of the various types of solids.

P-9 THE ATTRACTION OF MOLECULES

The London force is also known as the dispersionforce, or London dispersion force. It is also oftencalled the van der Waals force, although strictlyspeaking a van der Waals force is any force thatcan account for the nonideality of a gas close toits temperature of condensation.

In the mathematical expression of Coulomb'slaw, force = q iti2/cd2, the constant, c, is neededto relate the units of the force and the distance,d, to the units of the charges, q1 and q2. Whenforce is expressed in newtons (N, the SI-derivedunit of force), d is expressed in meters (m, the SIbase unit of length), and q and q2 are expressedin coulombs (C, the SI-derived unit of charge),then the value of c that correctly relates theseunits to one another in Coulomb's law is c =1.113 x 10-10 c2/N m2.

P-10 SPECIAL ATTRACTION:POLAR MOLECULES

The term polar attraction is one that has beenadopted in the student module to denote anyattraction between molecules or ions on accountof partial charges or whole charges. These forcescan be further classified as dipole-dipole forces,charge-dipole forces, and charge-charge forces.They are sometimes collectively called electro-static forces.

30 23

In addition to the polar attractions describedin the module, there is another class of attractiveforces called polarization forces or induced elec-trostatic forces. These forces arise when the elec-tric field produced by a polar molecule or an iondistorts the electron cloud of a neighboring mole-cule in such a way as to make the neighboringmolecule polar, even if it was originally nonpolar.The polarized molecule is then attracted to thefirst molecule or ion.

P-11 THE VAPORIZATION OF A LIQUID

This section introduces a most important concept:dynamic equilibrium. Be sure that your studentsunderstand the difference between the behaviorof water as a whole and the behavior of the in-dividual molecules. At equilibrium, there is no netchange in the amount of water, but individualmolecules move back and forth between the liquidand vapor phases.

P-12 THE BOILING PROCESS

The pressure in a pressure cooker is alwaysslightly greater than the vapor pressure of thewater because of the air trapped inside. For addi-tional discussion of the autoclave and pressurecooker refer students to pages 52-.53 of Moleculesin Living Systems: A Biochemistry Module. Ask stu-dents who are familiar with the pressure cookerto briefly discuss the advantages and disadvan-tages of its use in the home situation. If you havean autoclave in one of your school laboratories,demonstrate its use for your students.

Miniexperiment Some of your students may wish toobserve the phenomena of boiling. Have them heat abeaker of water and watch what happens to the waterfrom the time the thermal energy enters the bottom ofthe beaker until the liquid actually starts to boil. Thereare photographs illustrating this process in the book,Water, by Luna B. Leopold, Kenneth S. Davis, and theEditors of Time-Life Books, pp. 28-29.

MINIEXPERIMENTP-13 BOILING AT REDUCED PRESSURE

This miniexperiment shows that the boiling temperatureof a liquid depends on the pressure.

24

Concept

The boiling temperature of a liquid is less than thenormal boiling point when the pressure is less thanone atmosphere.

Objectives

Define boiling temperature.Describe the effect of reduced pressure upon theboiling temperature of a liquid.



Materials

15 aspirators15 thermometers (-10°C to 110°C)15 500-cm3 filter flasks or smaller, hard glass15 rubber stoppers, 1-hole15 lengths heavy-wall rubber tubing, 9.5 mm (3/8")

inside diameter15 ring stands15 universal clampsboiling chips

Advance Preparation Caution: In order to minimizethe risk to students, the instructor should insert thethermometers into the rubber stoppers in advance.The use of glycerol as a lubricant is suggested. Slittingthe stopper is an even safer method. The thermometerbulb does not have to extend into the water, as dropsof water at the boiling temperature will condense onthe bulb.

Laboratory Tips The students should clamp the flaskto the ring stand to prevent its being tipped over by theweight of the tubing. Students should'take care that,as they heat the water, the rubber tubing is not con-stricted or clamped shut. Otherwise, the stopper may beblown out of the flask.

Range of Results Depending on the aspirator, thewater may boil anywhere between room temperatureand 70°C. Since boiling is an endothermic process, thetemperature drops, thus lowering the vapor pressure ofthe water. Eventually the boiling will cease, unless anexternal source of heat is applied.

31

Post lab Discussion From the temperature at whichthe water boiled in the partially evacuated flask, thestudent can tell what the pressure was. Assuming thatall air was swept out of the flask during the initialboiling, the pressure during subsequent boiling is equalto the vapor pressure at the particular temperature.

Temperature

(°C)Pressure(cm Hg)

Pressure(atm)

Moleculesper cm3

100 76.0 1.00 20 x 101890 52.6 .69 14 "80 35.5 .47 1070 23.4 .31 760 14.9 .20 450 9.3 .12 340 5.5 .07 230 3.2 .04 1.0 "20 1.8 .02 0.5 "

You might ask your students to calculate how manyH2O molecules were present per cm3 volume in the vaporof their flasks. The calculation goes as follows. First, re-arrange the ideal gas law to solve for nIV, the numberof moles of gas per unit volume: n/V = P1RT. Next,multiply the moles per unit volume by the moleculesper mole (6.02 x 1023) to obtain the molecules per unitvolume. The answer in molecules per cm3 is 7.33 x 1021PIT, where P is in atmospheres and T is in kelvins. Somevalues calculated in this way are shown in the tableabove. Notice that even at the low pressure of 0.02atmosphere each cubic centimeter of vapor contains anenormous number of molecules, much greater than thehuman population of the earth!

In discussing the text that follows experimentP-13 in the student module, you may wish to asksome of your students to check food packagesat home to determine if there are any specialinstructions for preparing foods (such as bakingmixes) in areas of high altitudes.

P-14 SOLID STRUCTURES

In your discussion of molecular crystals, you maywish to recommend that some of your studentstry some of the following miniexperiments, whichinvolve crystals and crystal growing.

Demonstration Solidification Behavior: Place twosmall beakers in a large Petri dish on an overhead

projector. Place some phenylacetic acid (molecular crys-tals) in one beaker, and slivers of paraffin wax (amor-phous solid) in the other beaker. Add hot water (80°Cor warmer) to the Petri dish and observe the meltingof the two solids. Remove the Petri dish from thebeaker and observe the different solidification behavior;the phenylacetic acid solidifies sharply at 76°C, while theparaffin wax has a broad range of solidification.

Miniexperiment Crystal Growing: An interestingsource of suggestions for crystal growing experimentsis: Alan Holden and P. Singer, Crystals and CrystalGrowing (New York: Anchor Books, Doubleday, 1960;paperback).

Miniexperiment Rock Candy: One of your studentsmight wish to demonstrate the making of rock candy byusing the following recipe. Ingredients: 2 cups of tablesugar, clean cotton string, pencil, paper clip, clear jaror drinking glass, saucepan. Dissolve the sugar in 1 cupof boiling water, using continuous stirring. Let the solu-tion cool before pouring into a glass. Meanwhile, tie oneend of the string to the pencil and the other to thepaper clip (which keeps the string from floating). Thestring should hang straight in the solution when you laythe pencil across the mouth of the glass. Crystals startforming after a few hours and, if undisturbed, will growinto mouth-sized chunks in a few days.

You may follow a similar recipe to make salt crystals.Take a break as the class observes and tries the "solidstructures." Ask how dissolving and crystallizationare used in other food preparations and for industrialpurposes.

Miniexperiment Building Molecular Models: Youmay wish to have your students examine the structuresof network solids further by building molecular modelsof the diamond and graphite crystal. Lab-Aids Inc.,130 Wilbur Place, Bohemia, New York 11716, producesinexpensive model kits with, atoms that consist of plasticnuclei. Pegs are set at correct bonding angles to produceaccurate three-dimensional models. The kits containseparate procedures, worksheets, and a teacher guideprepared by the company. There is enough materialfor up to 50 students and materials can be usedagain and again. Write the company for price list anddescription.

32 25

Ask one of your interested students to find outabout and discuss with the class the use of thefreeze-dried process for preparing some foodsnow on the market. What actually happens in theprocess? What kinds of food products are pre-pared in this manner? What are the advantages tousing this process in preparing food? What arethe disadvantages? Keep your discussion of"solid structures" in focus as you relate this con-cept to everyday life.

Miniexperiment The discussion on page 33 of thestudent module introduces metal crystal structures. Youmay wish to ask interested students to obtain and bringto class samples of various metals in both sheet andwire form. Discuss the different properties of these typesof metal crystals as you observe each sample. If pos-sible, test each sample by hammering or bending. Ifyou have enough different kinds of metal samples, trygrouping them according to the properties that your stu-dents have determined. Ask students what an engineerwould have to know about a specific metal before recom-mending it to be used in forming products on a stampingpress as illustrated on page 33. For more informationon metal crystal structures, refer your students to Diver-sity and Periodicity: An Inorganic Chemistry Module.

Miniexperiment Obtain samples of quartz, mica, as-bestos, and graphite and use these to illustrate yourdiscussion of the network solid (page 34). In discussingthe caption at the top of page 34, call the students' at-tention to the measurement of the layers in the graphitestructure. Inform them that in this edition of the studentmodule the measurement of the carbon layers of graphiteis misprinted to read 0.34 mm instead of the correctmeasurement of 0.34 nm. You may wish to stop to dis-cuss the difference, using Appendix II of this module;or refer to section A-8 in Reactions and Reason (metricusage).

If some of your students are interested in con-tinuing their discussion of glass as an amorphoussolid, they can write to several of the glass com-panies for consumer information booklets withillustrations of glass production and glass prod-ucts. Or they can check this topic in their libraryresearch center. Information requests should beaddressed to the Public Relations Department ofthe following companies.

26

Corning Glass WorksCorning, New York 14830

Glass Packaging Institute1800 K Street, N.W.Washington, D.C. 20006

Libby-Owens-Ford Company811 Madison AvenueToledo, Ohio 43695

Owens-Illinois, IncorporatedP.O. Box 1035Toledo, Ohio 43666

Some of your students may have seen craftsexhibits where glass objects were formed frommolten glass. Ask them to briefly describe whatthey saw and relate this to the student modulediscussion of an amorphous solid.



1. (a) CH, and H2O molecules both have 10 electrons,so the London forces have similar strengths inboth substances. The higher boiling point of H2Ois due to polar attractions and hydrogen bondsthat are absent in CH4.

(b) CF4, CCI4, and CBr, molecules are all nonpolar.Their boiling points and freezing points increasein this order because of increasing numbers ofelectrons, thus increasingly strong London forces.

(c) HCI and F2 molecules both have 18 electrons.HCI is a polar molecule, F2 is not.

(d) NaCI exists in solution as Na+ ions and Cl- ions.These ions are strongly attracted to the polarmolecules of water, but not to the nonpolarmolecules of carbon tetrachloride.

(e) Diamond is a network solid with strong covalentbonds joining the atoms. In contrast, N2 moleculeshave only weak London forces between them,thus the molecules can readily separate to forma gas.An iodine crystal (12) is a molecular crystal, withweak London forces between the molecules. Nalcrystals are ionic solids, with strong ionic bondsbetween Na+ ions and I- ions.

2. Crystalline sugar contains sucrose molecules in amolecular crystal form. When the crystals are meltedand the syrup is cooled, the molecules tend to

(f)

33

solidify in a disordered, amorphous form that is tech-nically a type of glass.

3. The five items listed can be classified as follows:Plexiglasamorphous solid; aluminum foilmetal;sucrose (sugar)molecular crystal; sodium hydrox-ide (NaOH) ionic solid; and charcoal (graphite)network solid.

EVALUATION ITEMS

These are additional evaluation items that you may wishto use with your students at various times during thepreceding section. The correct answer to each question 4. Given the following structural formulas, circle theis indicated by shading. letter of those that represent polar molecules.

2. Consider molecules A, B, and C. When would theLondon force be greater between molecules A andB than between molecules A and C?

A. B is more polar than C.B. A is closer to C than to 13.C. B has more electrons than C.D. C has more electrons than B.

3. Which of these compounds can form hydrogenbonds?

A. H2O B. CH, C. HF D. NH3

1. In straight-chain hydrocarbons, as the number of car-bons increases, the boiling point increases. A factorthat influences this behavior is:

A. polar attraction between moleculesB. London forces between moleculesC. carbon-hydrogen attractionD. ionic attraction between molecules

A. H CI

0

B. H

CI

C. CIC CI

CI

D. 0 =C --=0

Solutions: Solute and Solvent

Solutions are important in chemistry becausemost chemical reactions are carried out in solu-tions. Here the concept of a dynamic equilibrium,introduced previously for liquid-vapor equilib-rium, is extended to the equilibrium between asolid and its saturated solution.

Demonstration Measure exactly 50 cm3 of ethanolin a 50-cm3 or a 100-cm3 graduated cylinder. Measureexactly 50 cm3 of water in a 100-cm3 graduated cylinder.Pour the ethanol into the water and stir with a long glassrod. The volume of the ethanol solution is only 96 cm3,instead of 50 cm3 + 50 cm3 = 100 cm3. The ethanolmolecules in the solution are closer, on the average, tothe neighboring water molecules than they were to theneighboring ethanol molecules in the pure ethanol.

Film Molecular Motions (CHEM Study Film num-ber 4115; 13 minutes; color). This film shows thetypes of motion of molecules in gases, liquids,

solutions, and solids by means of mechanicalmodels and animation. Before showing the film,explain to your students the distinction betweentranslational motion (motion of a whole moleculein a straight line), rotational motion (spinning inplace), and vibrational motion (vibrations ofbonds).

Miniexperiment Some of your students may wish toexperiment with the dissolving process after observingand discussing the photographs on pages 37, 38, and39. Material such as tea, instant coffee, freeze-driedcoffee, and a beaker of water is all that is necessary.Observing the dissolving process is the main purposeof this miniexperiment. Ask the students to describewhat they observed.

P-15 THE DISSOLVING PROCESS

Although equilibrium constants have not yet beenintroduced in the student module, they providea simple way to interpret the equilibrium betweena solid and its saturated solution. The dissolving

34 27

process for a molecular crystal may be writtenas a chemical equation

A(solid) =----A(solution)

At equilibrium, the concentration quotient[A(solution)]

[A(solid)]

is equal to a constant. Since the concentration ofa solid is a constant that cannot change, [A(solu-tion)] must also equal a constant at equilibrium.This constant is simply the solubility of the solid.

A similar argument may be used for an ionicsolid. For instance, the dissolving process ofAgC1 (a sparingly soluble salt) is

Ag+ + Cl-

and at equilibrium the concentration quotient[Ag+][C1-][AgCI(solid)]

is equal to a constant. Here [Ag+] and [Cl -] arethe concentrations of the ions in the solution.Since [AgCl(solid)] is a constant, the product[Ag+ ][C1-] is also a constant at equilibrium. Thisconstant is known as the solubility product of thesalt.

MINIEXPERIMENTP-16 SUPERSATURATION

This miniexperiment makes use of an easily preparedsupersaturated solution to demonstrate supersaturation,crystal growth, and heat of crystallization.

Concepts

When a crystal of solute is added to a supersaturatedsolution, crystallization results.The precipitation (crystallization) or sodium thiosulfatepentahydrate is exothermic.

Objective

Classify a solution as unsaturated, saturated, orsupersaturated, by observing what happens whensolid solute is added.

Estimated time One-half period


28

Materials

300 g sodium thiosulfate pentahydrate, Na2S203 5H2015 400-cm3 beakers15 18 x 150-mm test tubes15 ring stands and rings15 clamps, universal15 wire gauze, asbestos centers15 Bunsen burners

Laboratory Tips The exact amount of solid placedin each test tube is not important. You can show thestudents approximately how much to put in. Ask thestudents to save one crystal of the sodium thiosulfatefor the last step. Ask the students to clean their testtubes afterwards by reheating them in the hot water,pouring out the hot solution, and rinsing until clean. Or,the test tubes with crystals may be saved for subsequentclasses. (The experiment also works if you use sodiumacetate trihydrate, NaC2H302-3H20.)

Range of Results Crystallization starts immediatelyafter the seed crystal is added. Crystals grow out in alldirections from the seed crystal, eventually filling thewhole test tube. The test tube becomes quite warm asa result of the thermal energy produced by the crystal-

lization process.

Postlab Discussion Reheat one of the test tubes tomelt the crystals. Show the students that no crystalliza-tion occurs when a crystal of NaCI is added to the super-saturated solution. This is because the two crystalsbelong to different crystal groups; NaCl is cubic, andNa2S203 5H20 is monoclinic.

Ask the students to remember that the crystallizationprocess for this particular crystal is exothermic, as thisfact will be referred to later (in experiment P-27 Heatof Solution).

P-17 DISSOLVED GASES: FIZZLE AND POP

Demonstration To demonstrate that air dissolves inwater, shake some cold water with air in an Erlenmeyerflask. Fill a large test tube with the water and let it warmup in a beaker of hot water. After a few minutes, show

students the bubbles of air clinging to the inside wallsof the test tube. The bubbles appear because air is lesssoluble in warm water than in cold water.

35

Miniexperiment If any of your students have not ex-perienced a "flat" soda drink, ask them to pour sodainto a glass and observe the carbon dioxide returningto the gas phase. Have them taste the drink right afterit has been poured and then after there are no morebubbles in the drink.

If any of your students have done underwaterdiving with tanks, let them share their experienceswith the rest of the class. Some students maywish to find out more about the condition called"the bends," which is mentioned in the module,and briefly relate the concept of dissolved gasesto the human body.

P-18 ANTIFREEZE: CURE FOR THECOMMON COLD

Ethylene glycol used as an antifreeze will givefull protection to a car's cooling system down tothe temperature at which crystals first form in thesolution of ethylene glycol and water. This tem-perature depends on the concentration of thesolution. If the temperature drops below the tem-perature of protection, a slush of solid and solu7tion forms that is difficult for the water pumpto handle. Up to a 72-percent-by-volume solutionof ethylene glycol, the more concentrated thesolution, the lower is the temperature to whichit will protect the cooling system. Crystals willnot form in a 72-percent solution at temperatures

FREEZING TEMPERATURE OF SOLUTIONS OF ETHYLENE GLYCOL AND WATER

0

10

a) -20a.a)

CO

-301E

iow -40

50

60

0 20 40 60

volume percent of ethylene glycol

36

80 100

29

down to 63°C. More concentrated solutionsthan this do not protect to as low a temperature.Pure ethylene glycol freezes at 13°C. The graphshows the freezing behavior of solutions of vari-ous concentrations of ethylene glycol, and thekind of crystals that first appear as a solution iscooled to a sufficiently low temperature. Thesecrystals are either ice, ethylene glycol crystals, ora solid hydrate of ethylene glycol.

Miniexperiment Try the old "table trick" of challeng-ing a student to pick up an ice cube from a glass ofice water with a piece of string. If no one knows thesolution to the challenge, proceed by placing the threadon top of the ice cube and shaking on a little salt. Thesalt should start to dissolve the cube, which should thenrefreeze around the thread. When it does, pick up thecube with the attached thread. Relate this to the car-toons and the photograph of "salting the highways" aftera snowfall.



1 Two immiscible liquids that are placed in contactwith one another never remain completely pure. Inthe case of gasoline mixed with water, the gasolinelayer becomes wet and the water layer becomescontaminated with a small concentration of gasoline.

2. A lump of sugar, dropped into a cup of tea that isunsaturated with sugar, will dissolve. If the tea issaturated, the lump of sugar will undergo no netdissolution (its shape may change, but its mass willremain constant). If the tea is supersaturated, thelump of sugar will act as a "seed crystal," allowingnew crystals to form (rock candy).

EVALUATION ITEMS

These are additional evaluation items that you may wishto use with your students at various times during the

preceding section. The correct answer to each questionis indicated by shading.

1. Explain what is happening at the surface of NaCIcrystals in a saturated solution of NaCI in water.

The rates at which sodium and chloride ions aredissolving and precipitating are equal. That is, thenumbers of sodium and chloride ions becoming de-tached from the crystal surface are the same as thenumbers coming out of solution and reattaching tothe crystals.

2. After putting 5 tablespoonfuls of sugar into a cup ofcoffee and stirring vigorously, George notices thatquite a bit of the sugar is sitting in the bottom ofthe cup. Which of the following statements wouldbest explain the situation?

A. The coffee is supersaturated.B. Sugar is not very soluble in coffee.C. Sugar is polar, but water is not.D. Equilibrium has been established.

3. Liquid A is blue and liquid B is clear. They areimmiscible. When they are poured into the same testtube, what would you observe?

Two layers: one blue, one clear.

4. Give a simple explanation for the fact that a bottleof cola fizzes when it is opened.

The pressure drops, and dissolved carbon dioxidecomes out of solution.

5. As the amount of ethylene glycol in a liter of water isincreased,

A. the freezing point is raised.B. the freezing point is lowered.C. the water evaporates more easily.D. the temperature of the water increases.

The Colloidal State

The study of colloids has tended to be a neglectedtopic in high school and college curricula in recentyears. Colloids are, however, important in in-

30

dustry, biology, and everyday life. Furthermore,the colloidal state rounds out our study of thestates of matter, except for a few less-familiarstates, such as liquid helium II, ion plasmas,and liquid crystals. (More information on liquidcrystals is given below.)

BEST COPY AVAILABLE37

Colloids can be classified in the following man-ner. The student module deals mainly with col-loidal dispersions and gels.

1. Colloidal dispersions, also called lyophobic(solvent-hating) colloids. These consist of adispersed phase and a continuous phase.Generally, a surfactant or adsorbed ions arenecessary to stabilize the dispersion againstbreaking down into separate, noncolloidalphases.

2. Micelles, or association colloids. These are stableaggregates of molecules having both polar andnonpolar portions. Examples are micelles ofsoap molecules in water, and the casein par-ticles of milk.

3. Macromolecules, or lyophilic (solvent-loving) col-loids. In a solvent, these colloids differ fromcolloidal dispersions in having no tendency toseparate into two phases. An example is gela-tin in warm water.

4. Solid surfaces with pores of colloidal dimen-sions, as in dehydrated silica gel and activatedcharcoal.

5. Gels, which are often made from lyophiliccolloids.

Suggested Reading Marjorie J. Vold and RobertD. Vold, Colloid Chemistry: The Science of LargeMolecules, Small Particles, and Surfaces (New York:Van Nostrand Reinhold, 1964; paperback).

Egon Matijevic, "Colloidsthe World of Neg-lected Dimensions," Chem Tech (November 1973),p. 656.

Film Common Colloids (Prism Productions, Inc.,Science Close-ups number 310; color; 7 minutes).This film shows examples of many commonhousehold colloids.

Miniexperiment Liquid Crystals: The liquid crystallinestate, or mesomorphic state, is a state of matter havingcharacteristics of both the liquid and crystalline solidstates. In this state, the molecules are partially orderedas in a true crystal but are also mobile as in a liquid.The molecules are long, rod-shaped ones that tend topack together parallel to one another.

Edmund Scientific Company sells several fascinatingvarieties of liquid crystals whose colors are sensitive totemperature or pressure. Write to Edmund Scientific,

Edscorp Building, Barrington, New Jersey 08007 fortheir catalog.

For further information, see the article "Liquid CrystalsDraw Intense Interest," Chemical and EngineeringNews (November 1, 1971), p. 20.

Miniexperiment Some of your students might enjoypreparing the colloidal dessert illustrated on page 44.The recipe follows. Ask them to find and discuss asmany colloids as they can in the recipe when they sharethis with the class. Interested students may wish to sub-stitute another type of colloidal dessert for this one.

Pineapple Cream Cheese Mold1 567-g can crushed pineapple in juice2 85-g packages lemon or lime flavor gelatin1/2 liter boiling water30 cm3 lemon juice2 85-g packages cream cheese, softened75 g chopped walnuts

Drain pineapple, reserving juice. Dilute juice with waterto make 3/4 liter liquid. Dissolve gelatin in boiling water.Add measured liquid and lemon juice. Gradually add1/2 liter of the gelatin to the cream cheese, blendingwell. Chill until thickened. Stir in walnuts and pour intoa 11/2-liter ring mold. Chill until set, but not firm. Chillremaining gelatin until thickened. Fold in pineapple, andspoon into mold. Chill until firm, about three hours.Unmold. Serve with whipped cream, mayonnaise, orsalad dressing (another colloid!). Makes about 12 serv-ings. (Recipe courtesy General Foods Consumer Center,White Plains, New York.)

EXPERIMENTP-19 SOME PROPERTIES OF COLLOIDALPARTICLES

This experiment demonstrates some of the differencesin the behavior of true solutions, colloidal dispersions,and macroscopic precipitate.

Concepts

Particles of a precipitate are retained by filter paper.Particles of a colloidal sol pass through filter paperbut are retained by dialysis tubing.A solution passes through both filter paper and dial-ysis tubing.The particles of a colloid are smaller than those ofa macroscopic precipitate but larger than those of asolution.

3831

Objectives of the tubing should be tied off with overhand knots,or folded over and tied with string.

Demonstrate the proper techniques for filtering anddialyzing.Compare the relative particle sizes in a precipitate,a colloid and a solution.Determine whether a liquid contains a macroscopicprecipitate or a colloid, given filter paper and dialysistubing.



Materials

300 cm3 colored solution (see Advance Preparation)600 cm3 skim milk (see Advance Preparation)75 cm3 20-percent acetic acid (20 cm3 glacial acetic

acid mixed with 80 cm3 water)45 50-cm3 beakers30 150-cm3 beakers45 18 x 150-mm test tubes15 test-tube racks15 10-cm3 graduated cylinders15-45 75-mm diameter funnels15 ring stands15 funnel supports45 sheets 12.5-cm diameter filter paper600 cm 1.6-cm diameter cellulose dialysis tubing (see

Advance Preparation)15 glass stirring rods

Advance Preparation Prepare the skim milk fromnonfat dry milk powder. It keeps for a day withoutrefrigeration. Whole milk is not satisfactory because thefat globules clog the filter paper.

Any colored solution may be used: CuSO4 solution, foodcoloring in water, and so on. The color should be in-tense enough to be visible as it diffuses through thedialysis tubing.

Precut the dialysis tubing into two 20-cm lengths foreach pair of students. Cellophane sheets from a sta-tionery store may also be used; soak them for severalhours in water and form into bags that the studentscan tie off at the top.

Pre lab Discussion Demonstrate the technique offilling the dialysis tubing. It is a good idea to pour theliquid through a funnel inserted into the tubing. The ends

32

Range of Results Both the filter paper and the cel-lulose of the dialysis tubing are porous materials thatallow some substances to pass through and retainothers. The filter paper retains only relatively coarseparticles, whereas the dialysis tubing is able to retainmuch finer particles (colloidal particles).

The filter paper retains only the milk curd. The dialysistubing retains the colloidal casein particles but not thecolored molecules of the true solution.

Post lab Discussion Ask the students to comparethe relative particle size in the three materials, and togive the reasons for the ordering they choose. Pointout that milk sugars, not visible to the eye, are diffusingout from the skim milk through the dialysis tubing.

You may want to leave one of the milk-filled dialysisbags in water for a day to demonstrate that none of thewhite color leaves the bag. The bag swells up becausewater molecules enter itthe phenomenon known asosmotic pressure. After one day at room temperature,the milk in the bag will spoil because of bacterial action.Various components then will pass through the bag andmake the water cloudly. The casein in the precipitatedcurd is used commercially to make white glue. Trysmearing some on paper and pressing another piece ofpaper on top. When the casein dries, it becomes hardand almost clear. It glues the papers together.

P-20 THE IMPORTANCE OF SURFACES

The unique feature of the colloidal state may besaid to be the large fraction of the molecules thatare at a surface. In general, molecules have dif-ferent properties when they are at a surface com-pared with the interior of a homogenous phase.For this reason, solids with exceptionally largesurface areas, such as dehydrated silica gel andactivated charcoal, are included in the list of col-loidal systems.

P-21 SMALL WONDER: COLLOIDALDISPERSIONS

Demonstration To demonstrate a foam, you canplace a marshmallow, shaving cream, whipped cream,

39

or beat-up egg white in a filter flask. Stopper the flask.Apply a vacuum with a heavy-walled rubber tube con-nected to an aspirator. Give the students an opportunityto observe the foam. Then ask, "Why does the foamswell up? In view of what you have learned about thevolume-pressure relation of gases, can you explain theswelling?" (Refer to miniexperiment P-5: Pressure andVolume.) Remove the vacuum. Call attention to how thefoam then collapses to a remarkably small volume.

Demonstration The Tyndall effect: Fill a rectangularbattery jar or a small flat-sided aquarium with afreshly made solution containing 7 g of sodium thiosulfate(Na2S203 5H20) per liter of solution. Darken the room.Place the container in a beam of light from a slide pro-jector so that the class observes the light from the side.You can project a sharp beam by punching a hole inan idex card and inserting the card in the projector slideslot. Note that no light is scattered as the beam passesthrough the Na2S2O3 solution.

projector battery jar or aquarium Iscreen

Add about 1 cm3 of 6 M hydrochloric acid for each100 cm3 of solution in the container. Colloidal sulfur willbegin to form slowly because of the reaction

S2032- + S + SO2 + H2O

Show that from the side the light scattered by the col-loidal particles appears bluish. Then turn the apparatusso the transmitted light falls on a white screen. The trans-mitted light will be yellow at first, and then progressivelyredder. Finally, it will be completely blocked.

Colloidal dispersions scatter blue light to a greaterextent than yellow or red light. The scattering of theblue component of sunlight by the fog and dust in theatmosphere is responsible for the blue daytime color ofthe sky. In a sunset, the yellow and red componentsare transmitted with less scattering and cause the colorchanges seen at the horizon.

The Tyndall beam effect may also be demonstrated byshining the light beam through cigarette smoke, orNH,,CI smoke prepared by allowing gaseous HCI andNH3 to react.

40

Demonstration Cottrell Precipitator: For details, seedemonstrations 23-26 and 23-42s in Alyea and Dutton'sbook Tested Demonstrations in Chemistry.

Suggested Reading G. P. Turner, Introduction toPaint Chemistry (New York: Halsted Press, 1967;paperback).

Harry Baines, The Science of Photography (NewYork: John Wiley & Sons, 1967). Baines presentsthe history and science of photography at a simplelevel.

George T. Eaton, Photographic Chemistry, 2nd ed.(Dobbs Ferry, N.Y.: Morgan and Morgan, Inc.,1965). This book offers a good description of thechemistry involved in modern photography, bothblack-and-white and color, written at an elemen-tary level. The author is head of the PhotographicChemistry Department of the Kodak ResearchLaboratories.

Miniexperiment Suggest that your students collectand bring to class examples of the materials discussedon pages 50 and 51, such as photographic film (beforeand after developing), printing ink, skim milk, milk ofmagnesia, or chocolate milk. Solicit other examplesduring class discussion. Observe some of these house-hold products and relate to the discussion.

EXPERIMENTP-22 MAKING OIL AND WATER MIX

The purpose of this experiment is to demonstrate onemethod of preparing an emulsion and to introduce somesimple tests that may be used to determine emulsiontype.

Concepts

The ease of formation of an emulsion is enhancedby emulsifiers.An emulsion may invert from one type to anotherwhen the relative volumes of the two phases arechanged.The emulsion type may be determined by severalsimple tests.Many familiar household materials are emulsions.

33

Objectives

Prepare an oil-in-water and a water-in-oil emulsion.Determine the type of an emulsion by simple physicaltests.



Materials

60 cm3 Tween 40 (see Advance Preparation)60 g Span 40 (see Advance Preparation)750 cm3 mineral oil (paraffin oil) or vegetable oil8 dropper bottles of blue food coloring8 dropper bottles of Sudan III or Sudan IV dye in mineral

oil (see Advance Preparation)30 pieces of filter paper, any size15 wash bottles, polyethylene15 400-cm3 or 600-cm3 beakers30 150-cm3 beakers30 glass stirring rods15 10-cm3 graduated cylinders15 50-cm3 graduated cylinders15 ring stands15 10-cm ring-stand rings15 wire gauze, asbestos centers15 Bunsen burners15 thermometers, 10°C to 110°C15 porcelain spot plates, or several small watch glasses8 balances

Advance Preparation Tween 40 (polyoxyethylenesorbitan monopalmitate) and Span 40 (sorbitan mono-palmitate) are commercial emulsifiers manufactured bythe Specialty Chemicals Division of ICI Chemical Spe-cialties Company, Wilmington, Delaware 19899. Theyare used in many products, including foods. They areavailable in small quantities from the manufacturer, andin larger quantities from the following distributors: Emul-sion Engineering, 480 Bennett Road, Elk Grove Village,Illinois 60007, telephone (312) 439-7375; Ruger Chem-ical Company, 83 Cordier Street, Irvington, New Jersey07111, telephone (201) 926-0331; and also from Rea-gents, Inc., P.O. Box 3977, Charlotte, North Carolina28203, telephone (704) 523-4064.

The solution of Sudan III or IV dye should be preparedby dissolving enough of the dye in mineral oil to producea bright red color. The dye dissolves slowly; allow acouple of days for the solution process to take place.

34

Ask the students to bring in small amounts of commonhousehold emulsions that may be tested for emulsiontype. These can include such products as face creams,sun-tan lotion, "waterless" hand cleaners, cold creams,hand lotions, hair creams, and cream-type car and furni-ture wax. To ensure that some of the emulsions arewater-in-oil, suggest that the students bring in Pond'sDry Skin Cream and Wildroot Formula 360 hair cream.

Laboratory Tips Two noticeable changes will occuras the water is added to the oil. When the initial emul-sion forms, the mixture will turn milky. Then, as theemulsion inverts from water-in-oil to oil-in-water, therewill be a drop in viscosity. The second transition is muchmore difficult to detect than the first. Tell the studentsto add the water a little at a time and to rely on theirtests to tell them when the emulsion has inverted.

Many students will have trouble with the dye test. Re-mind them that they must find out which of the dyesactually dissolves in the emulsion. The gentle stirring willdisperse both of the dyes, but one will remain suspendedas tiny droplets while the other forms a true solution inthe continuous phase of the emulsion.

The cleanup after the experiment is easier if acetoneor ethanol is used to rinse out the emulsion.

Range of Results The initial emulsion is water-in-oil.After inversion, this becomes oil-in-water, similar to acommercial hand lotion. The students will find that theycan continue to add any amount of water to their oil-in-water emulsion without affecting its homogeneity. The oildroplets simply move farther apart as the water is added.

Postlab Discussion Most commercial emulsions areof the oil-in-water type. The advantages of this type ofemulsion as a hand lotion are that it does not feel greasy,it provides a convenient way to spread oil evenly overthe skin, and it may be rinsed off.

Miniexperiment Emulsifiers: Make concentrated so-lutions of soap and of detergent or dishwashing liquid.Stir equal amounts of mineral oil or vegetable oil intoboth solutions. Compare the soap and detergent asemulsifiers. For better visibility, food coloring may beadded to the soap and detergent solutions or Sudandye may be added to the oil.

41

Miniexperiment Making Mayonnaise: Follow direc-tions in any cook book. Egg yolk or whole egg, mustard,and paprika usually serve as the emulsifiers. Mayonnaisetends to separate during preparation if the oil is added toorapidly. Commercial mayonnaise must be at least 65percent oil by weight to meet government standards.

Suggested Readings Paul Becher, Emulsions:Theory and Practice, 2nd ed. (New York: ReinholdPublishing Corporation, 1965).

William R. Mimrath, 'Van Nostrand's Practical For-mulary (New York: Van Nostrand Company,1957). This book gives recipes for preparing coldcreams, hand cleaner, sun-tan lotions, and otherpreparations. Many use only familiar ingredients.

MINIEXPERIMENTP-23 MAKING A GEL

Here the students prepare an easy-to-make gel that hasthe spectacular property of readily burning.

Concepts

A gel can be formed by the precipitation of colloidalcrystals in a solution.A gel is a soft, semirigid colloidal system.

Objective

Describe the structure and appearance of a gel.



Materials

750 cm3 ethanol75 cm3 saturated calcium acetate solution (see Ad-

vance Preparation)15 150 -cm3 beakers15 10-cm3 graduated cylinders15 50-cm3 graduated cylinders15 ring stands15 ring-stand rings15 wire gauze, asbestos centers, or 15 asbestos squares

Advance Preparation Prepare the saturated calciumacetate solution by placing 40 g of solid in a 150-cm3

beaker and slowly adding 100 cm3 of water, a little ata time, stirring after each addition. (The calcium acetatetends to float on the water if it is added directly to thewater.)

Laboratory Tips The gel burns with a pale blue flamethat is almost invisible. Caution: Remind the studentsto keep their clothing well out of the way. Have themkeep wet paper towels handy in case of fire.

Postlab Discussion Calcium acetate is only slightlysoluble in ethanol. When the ethanol is added to thesaturated solution, a colloidal network of calcium acetatecrystals forms. This network traps water and ethanolin the spaces between the crystals. Although the result-ing gel is fairly solid, the crystals make up only a smallpercentage of the total mass. The commercial Sterno"canned heat" product is a similar gel made from alcoholand a cellulose ester.



This question allows the student to compare the molecu-lar organizations of several "soft" solids of varyingdegrees of structural complexity. Gold is a metal; itsatoms are arranged in an organized pattern and areheld together by valence electrons. Iodine is a molecularcrystal; 12 molecules are held together in an organizedpattern by London forces. Marshmallows are a solidfoam, consisting of air bubbles separated by thin filmsof a solid (actually the solid phase is gelatin). Butteris a water-in-oil emulsion; the oil phase is liquid butterfat.Gelatin is a gel, consisting of collagen fibers held to-gether by hydrogen bonds and permeated by water. Ineach solid, there are only weak attractive forces betweenthe atoms or molecules so that the solid as a wholecan be easily deformed with a little force.

EVALUATION ITEMS


1. Briefly explain the difference between a sol and atrue solution.

A sol contains colloidal particles that are much largerthan the solute molecules in a true solution.

4 2 BEST COPY AVAILABLE35

2. Complete the chart below by using the words solid,liquid, gas.

Substance

fog

smoke

emulsion

sol

solid foam

Dispersed Phase

liquid

Continuous Phase

gas

solid ,gas

liquid

solid

gas

liquid_

,liquid

solid

3. Classify each of the following substances as eitheraerosol, liquid foam, solid foam, sol, or emulsion.

A. fog ,,aerosol

B. whipped cream

C. bread

liquid foam;

solid foam,

D. printing ink

E. hand lotion

sol

emulsion

4. An emulsifier is best classified as:

A. an emulsion breakerB. a colloid

C. a surfactantD. a solvent

5. Explain briefly why soaps are good surfactants.

Soap molecules have a charged end that is attractedto water and a nonpolar end that is attracted to theair or oil phases.

6. Which of the following properties is not characteristicof a gel?

A. definite shape C., easily deformedB. network of fibers D independent crystals

Changes in Energy

In this portion of the module, students are intro-duced to thermochemistry. They perform a seriesof calorimetry measurements to determine a heatof fusion, two heats of solution, and a heat ofreaction.

If time is short, you can let each of the threeexperiments be performed by one-third of theclass. The class can then discuss the results ofall three experiments as a group.

P-24 COUNTING CALORIES

You may be unfamiliar with the terms chemicalenergy and thermal energy. They are convenientclassifications of ways in which energy can bestored by a chemical system. Other possible cate-gories in this classification would be nuclearenergy, the kinetic energy of motion of a bodyas a whole, and various mechanical forms ofpotential energy, such as the energy stored in araised weight.

Chemical energy is a form of potential energy(energy of position) that depends on the bonding

36

arrangements within the molecules and the forcesbetween the molecules. The chemical energy of achemical system can increase or decrease when-ever a physical or chemical process takes place,such as a phase change or a chemical reaction.One of the purposes of calorimetry experimentsis to measure chemical energy changes and thusprovide information on molecular properties suchas relative bond strengths and intermolecularforces.

Thermal energy is the energy that a chemicalsystem possesses because of its temperature. Thethermal energy increases whenever the tempera-ture of the system increases. Most of the increasein thermal energy results from an increase in thekinetic energy (energy of motion) of the mole-cules, moving in random directions. However,there may also be an increase in the potentialenergy because bonds vibrate with greater ampli-tude as the temperature increases.

Heat is a form of energy transfer from one placeto another. It should not be confused with thermalenergy. Heat is a quantity of energy that flowsfrom a warmer to a cooler body. The warmer bodyloses some of the energy originally stored there,and the cooler body gains an equal amount of

43

energy. When this happens, there is a tempta-tion to say that the heat of the warmer body de-creases and the heat of the cooler body increases.While it does little harm to speak in this way ofchanges in the heat, it is actually incorrect becauseheat is not something that a body has or contains.Instead, a body has a certain amount of energy.Heat is not a commodity, but a process. Heatappears only at the outside surface of a body whenthe body's energy is transferred to another bodyas a result of a temperature difference. The se-mantic confusion regarding the use of heat is dis-cussed in an article by T. B. Tripp, "The Definitionof Heat," Journal of Chemical Education, Vol. 53(1976), p. 782.

It is possible to be precise regarding the relationbetween energy and heat by using the followingmathematical form of what is called the first lawof thermodynamicsthermal energy change + chemical energy change

= heat workThe left side of the equation is the total changein energy of a chemical system, positive for anincrease and negative for a decrease. The rightside of the equation is the net amount of energytransfer, positive for a transfer of energy into thesystem and negative for a transfer out. Thus thelaw states that the energy of the system changesby the amount of energy that enters or leaves,and it is basically a statement of the principle ofthe conservation of energy. The work that appearsin the equation consists of all forms of energytransfer other than heat. For historical reasons,work is taken as positive if energy leaves the sys-tem and negative if energy enters. The only typesof work we need to consider here are expansionwork and electrical work.

work = expansion work + electrical work

The expansion work is zero if the volume of thesystem remains constant. In the calorimetry exper-iments performed by the students, the volumeschange so little that expansion work may beneglected. Electrical work occurs only in electro-chemical cells. Thus, work is approximately zeroin these calorimetry experiments.

Calorimetry measurements are used to obtainwhat is customarily called the "heat" of a process,meaning the amount of energy in the form of heatthat enters the system if the process takes place

44

at a constant temperature. Thus we have the heatof fusion (or heat of melting), the heat of solution,and the heat of reaction. It can be seen from theequation for the first law of thermodynamicsgiven previously that if the temperature does notchange (thermal energy change = 0) and the workis zero, then the heat for the process is equalto the chemical energy change. It is for this rea-son that the student module defines the heat offusion as the chemical energy change during themelting process. By this definition, the heat of aprocess is positive for an endothermic processand negative for an exothermic process. (Theopposite-sign convention is sometimes used; infact, the first edition of Communities of Moleculesused the opposite convention.)

The text does not attempt to define the heatof a process as the amount of heat at constanttemperature. Such a definition might be con-fusing to students, because in the calorimetricexperiments there is actually little or no heat. Theexplanation for this apparent paradox is that theexperiments are not carried out at constant tem-perature. However, the chemical energy changefor each process still has the same value it wouldhave if the temperature were kept constant. Inthe experiments, the purpose of the calorimetricis to make heat = 0. Then, assuming work = 0,the first law equation gives

chemical energy change = thermal energy change

By measuring the temperature change withinthe calorimeter, the student calculates the ther-mal energy change; then the preceding equa-tion allows the chemical energy change, to bedetermined.

EXPERIMENTP-25 THE HEAT OF FUSION OF ICE

The purpose of this experiment is to measure thechemical energy change when one mole of ice melts.

Concepts

Thermal energy changes to chemical energy whenice melts.The molar heat of fusion of a substance does notdepend on the amount of substance melted.

37

Objectives

Calculate the change in the thermal energy of a givenmass of water, given the initial and final temperatures.Calculate the molar heat of fusion of a substance,given the specific heat of fusion and the molar massof the substance.

Estimated Time One-half to one period


Materials

30-45 ice cubes15 100-cm3 graduated cylinders15 Styrofoam cups15 thermometers (-10°C to 110°C)

Advance Preparation The ice should be allowed tomelt partially to ensure that it is at 0°C.

Pre lab Discussion Give the students some prac-tice in calculating the change in thermal energy. For in-stance, if the temperature of 100 cm3 of water decreasesby 24.5 degrees Celsius, what is the thermal energychange? (Answer: 2450 calories.)

Laboratory Tips Do not use shaved or cracked ice.Because of its wetness, it will give poor results.

Range of Results Because of the difficulties in de-termining the amount of ice that actually melted, youmay expect considerable variation in the results.

A typical set of data and calculations is:

initial water temperature (t1) = 28.3°Cfinal water temperature (t2) = 0.0°Ct2 ti = 28.3°Cinitial volume of water = 100 cm3final volume of water = 135 cm3volume of water melted = 35 cm3mass of ice melted = 35 gthermal energy change = (1.00 cal/°C g) x (100 g) x

(-28.3°C) = 2.83 x 103 cal= 2.83 x 103 cal

chemical energy change = 2.83 x 103 calspecific heat of fusion = 2.83 x 103 cal /35 g = 81 cal/gamount of ice melted = 35 g/(18.0 g/mol) = 1.9 molmolar heat of fusion

= 2.83 x 103 cal /1.9 mol = 1.5 x 103 cal/mol

38

Postlab Discussion The accepted value for the heatof fusion of ice is 1.44 x 103 cal/mol. This means thatthe chemical energy of one mole of H2O molecules in-creases by 1440 calories during the melting process.This is the energy acquired by the H2O molecules asthey leave the ice crystal against the pull of the attrac-tive forces that hold the crystal together; the energy isstored in the melted water in the form of potential energy.An analogy is the potential energy that a ball acquireswhen it is thrown upward against the pull of gravity.

Different pairs of students will obtain different values ofthe chemical energy change, because different amountsof ice will melt depending on the initial temperature.However, the values of the specific heat of fusion andthe molar heat of fusion should be relatively consistent,since these quantities are independent of the amountof ice that melted.

P-26 ENDOTHERMIC AND EXOTHERMIC

According to the laws of physics, the natural ten-dency of a mechanical system that has no forcesexerted on it from outside the system is to de-crease its potential energy. For instance, a weightreleased in midair will spontaneously fall so as todecrease its gravitational potential energy.

One might expect a chemical system to obeythe same rule. That is, the system ought to changespontaneously in a way that will decrease itschemical energy, which is a form of potentialenergy. According to this idea, all spontaneousprocesses should be exothermic and only exo-thermic processes should take place in nature.

However, we know that this idea is wrong be-cause many spontaneous processes are actuallyendothermicfor instance, the melting of iceabove 0°C. The explanation is that during anendothermic process a second driving force, thetendency toward an increase in the disorder (en-tropy) of the molecules, outweighs the tendencyof the potential energy to decrease. This subjectwill be discussed again in this guide in sectionP-34.

You might have your students carry out thefollowing miniexperiments.

Miniexperiment If anyone in your class can bring inan old-fashioned ice cream freezer from home, you may

45

wish to try making ice cream in the laboratory whilediscussing endothermic and exothermic energy.

Miniexperiment Mix sodium chloride crystals with iceto produce a mixture of solid NaCI, ice, and NaCI solution.The mixture temperature is 21°C, and can be meas-ured on a thermometer that reads this low. Note that thetemperature decreases as a result of the melting ofsome of the ice, an endothermic process. The minimumpossible temperature of 21°C is known as the eutectictemperature of the salt-water-ice system.

EXPERIMENTP-27 HEAT OF SOLUTION

The purpose of this experiment is to measure changesin the chemical energy when compounds are dissolvedin water.

Concepts

A solution process may be either endothermic orexothermic.In an endothermic process, thermal energy changesto chemical energy.In an exothermic process, chemical energy changesto thermal energy.

Objectives

Measure the temperature change that takes placewhen a substance dissolves in a solvent.Decide whether a given solution process is endo-thermic or exothermic, given the initial and finaltemperatures.Calculate the molar heat of solution of a substance,given the initial and final temperatures, the mass ofthe water, the mass of the substance, and the molarmass of the substance.Describe the changes in thermal energy and chemicalenergy for an endothermic process and an exothermicprocess.



Materials

375 g sodium thiosulfate pentahydrate, Na2S203 5H2060 g sodium hydroxide pellets, NaOH

46

15 100-cm3 graduated cylinders15 Styrofoam cups8 balances, 0.1 g to 0.01 g sensitivity

15 spatulas15 thermometers (-10°C to 110°C)15 glass stirring rods

Laboratory Tips You can provide squares of cleanpaper for use in weighing out the sodium thiosulfatepentahydrate crystals. Dry watch glasses would be bet-ter for the sodium hydroxide pellets since they tend toabsorb water from the atmosphere and become sticky.Caution: Instruct the students not to touch the NaOHpellets with their fingers.

The students should weigh their NaOH pellets as rapidlyas possible. The mass does not have to be exactly 4 g.If the weighing is done quickly, the absorbed water doesnot seriously affect the results of this experiment.

Range of Results Typical results are

Na2S203 -5H20 NaOHmass of solid = 24.80 g 4.12 g

amount of solid = 0.100 mol 0.103 molti = 22.2° C 29.7° Ct2 = 13.8°C 39.8°C

t2 t1 = 8.4°C 10.1°Cthermal energy change = 840 cal 1010 cal

chemical energy change = 840 cal 1010 calmolar heat of solution = 8400 cal/mol 9800

cal/mol

The dissolving of sodium thiosulfate pentahydrate is anendothermic process, and the dissolving of sodiumhydroxide is an exothermic process.

Postlab Discussion There are two energy changesthat occur when a solution is formed. The first one oc-curs when the particles of the solid begin to separatein spite of the forces that are attempting to hold thesolid together. This process is analogous to raising aweight against the force of gravity and thus representsan increase in the chemical energy of the system.

The subsequent solvation of the solute particles be-cause of the mutual attraction between these particlesand the molecules of the solvent is analogous to thelowering of a weight. Therefore, this process tends to de-crease the chemical energy of the system. If the chemical

39

energy change for the first step is greater than the chem-ical energy change for the second step, the overallprocess is endothermic. If the opposite is true, the changeis exothermic. Be sure to point out that both of theseenergy changes occur simultaneously as a materialbegins to dissolve. It is only for convenience that weseparate them into two steps.

EXPERIMENTP-28 HEAT OF REACTION

The purpose of this experiment is to measure the changein the chemical energy that accompanies the neutraliza-tion of hydrogen ions by hydroxide ions.

Concept

There may be a change in the thermal energy ac-companying a chemical reaction.

Objectives

Measure the temperature change that takes placeduring a chemical reaction.Calculate the molar heat of a reaction in an aqueoussolution, given the number of moles of productformed, the mass of the water, and the initial andfinal temperatures.



Materials

750 cm3 2 M HCI (164 cm3 concentrated HCI in1000 cm3)

750 cm3 2 M NaOH (80 g solid NaOH in 1000 cm3)15 50-cm3 graduated cylinders15 Styrofoam cups15 thermometers (-10°C to 110°C)15 glass stirring rods

Advance Preparation Prepare the acid and thebase solutions a day ahead of time so that both solu-tions have time to cool to room temperature.

Pre lab Discussion Point out that the sodium ions andthe chloride ions do not participate in the reaction. Havethe students calculate the number of moles of hydrogenand hydroxide ions present in the reaction mixture.

40

Range of Results Typical results are

temperature of HCI solution = 30.1°Ctemperature of NaOH solution = 29.9°Caverage initial temperature (t1) = 30.0°Chighest temperature reached (t2) = 43.9°C

t2 t, = 13.9°Cthermal energy change = 1390 calchemical energy change = 1390 calamount of product formed = 0.10 molmolar heat of reaction = 1390 cal /0.10 mol

= 13 900 cal/mol

The accepted value of the molar heat of the neutraliza-tion reaction is 13 360 cal/mol.

Post lab Discussion Notice that the molar heat ofneutralization has a larger value than any of the othermolar heats measured by the students. In general, thebreaking or forming of a covalent bond causes a rela-tively large change in the chemical energy, often amount-ing to several hundred kilocalories per mole.

Demonstration Endothermic Reaction: In a dry250-cm3 Erlenmeyer flask, put:

32 g barium hydroxide, Ba(OH)2 8H2016 g ammonium thiocyanate, NH4SCN (keep this com-

pound tightly covered before use, as it absorbs

moisture)

Stopper the flask and shake it vigorously until the solidmaterial begins to liquefy. Place the flask on a hardwoodblock that has been moistened with a few drops ofwater. After a few minutes, the flask will be frozento the wood block.

The equation for the reaction is:

Ba(OH)2 8H20 + 2NH3SCNBa(SCN)2 + 2NH., + 10H20

P-29 THE POWER OF THERMAL ENERGY

According to the second law of thermodynamics,thermal energy can never be converted com-pletely to mechanical energy. This is because thethermal energy consists mainly of the kineticenergy of motion of the molecules in randomdirections, whereas mechanical energy is a kineticenergy of motion in one direction.

Energy values of foods (caloric values) aredetermined from the amounts of protein, carbo-

4

hydrate, and fat in the food. For the calculation,the specific heat of combustion of proteins andcarbohydrates is taken as 4.0 kcal/g, and thespecific heat of combustion of fats is taken as9.0 kcal/g. (The kilocalorie, kcal, is the same asthe dietetic Calorie.) These specific heat valuesare average values for various kinds of proteins,carbohydrates, and fats. They refer to combustionreactions in which the products are CO2, H2O,and urea. For more details, see Ned A. Daughertyand Kenneth W. Watkins, "Energy Value ofFoods," Journal of Chemical Education, Vol. 53(January 1976), p. 80.

If the physics laboratory has a demonstrationmodel of the cylinders of an internal combustionengine, borrow this to add to your discussion ofpages 63 and 64 in the student module.



1. (a) thermal energy change= (1.00 cal/°C g) x 400 g x (31 25)°C= 2.4 x 103 cal

chemical energy change = 0 (no change in

molecular organization)(b) thermal energy change

= 0 (boiling at constant temperature)chemical energy change

= (200 - 20) mol x 9.7 x 103 cal/mol= 1.7 x 106 cal

(c) thermal energy change= (1.00 cal/°C g) x 400 g x (-0.32°C)= -1.3 x 102 cal

chemical energy change= -thermal energy change (insulated

calorimeter)= 1.3 x 102 cal

2. amount of NaCI= 5.85 g/(58.5 g/mol) = 0.100 mol

molar heat heat of solution= (1.3 x 102 cal)/0.100 mol= 1.3 x 103 cal/mol

3. OH- is the limiting reactant for neutralization.amount of H2O produced

= 100 cm3 x (1 liter/1000 cm3) x 1.0 mole/liter= 0.10 mole

chemical energy change= 0.10 mol x (-1.336 x 104 cal/mol)= -1.3 x 103 cal

thermal energy change= 1.3 x 103 cal

t2 t1 = 1.3 x 103 cal/(1.00 cal/°C g) x 200 g= 6.5°C

t2 = 25.0°C + 6.5°C = 31.5°C

EVALUATION ITEMS


1. What is the change in thermal energy that occurswhen the temperature of 4.0 g of water goes from25.0°C to 31.5°C?

A. 26.0 calB. 100.0 cal

C 126.0 calD. 226.0 cal

2. What is meant by the term molar heat of fusion?

This is the chemical energy change per mole of solidmelted.

3. Some ice at 0°C was added to 74 cm3 of water at25°C. When the temperature of the mixture reached0°C, the ice was removed. The new volume of thewater was found to be 97 cm3. Assume that 1 cm3of water has a mass of 1 g and that the mass ofone mole of water is 18 g. Calculated from thesedata, approximately what is the molar heat of fusionof water?

A. 1450 cal/molB. 1900 cal/mol

C. 33 300 cal/molD. 43 650 cal/mol

4. Which of the following statements is true?

A. When ice melts, it loses chemical energy.B. 1 g of ice has more chemical energy than 1 g of

water.C. 1 g of water has more chemical energy than 1 g

of ice.D. When ice melts, chemical energy is changed to

thermal energy.

5. In an endothermic process:

A. thermal energy increasesB. molecular attraction increasesC. temperature increasesD. chemical energy increases

8 REST COPY AVAILABLE 41

6. Which of the following is always exothermic?

A. vaporization:B. combustion

C. meltingD. dissolving

7. Suppose you wanted to make a chemical "hotpack" by dissolving a solid in a solvent. Which ofthe following kinds of attractions between the ionsand molecules would give the best results?

A. strong attraction in the solid and weak attrac-tion in the solution

B. strong attraction in the solid and strong attrac-tion in the solution

C. weak attraction in the solid and strong attractionin the solution

D. weak attraction in the solid and weak attractionin the solution

8. When 11.9 g of KBr were dissolved in 100 cm3 ofwater, the temperature of the water fell from 25°Cto 20.6°C. If the molar mass of KBr is taken as119 g/mol, what is the molar heat of solution ofKBr? (Assume that the mass of 1 cm3 of wateris 1 g.)

A. 8800 cal/molB. 4400 cal/mol

C. 2300 cal/molD. 44 cal/mol

9. What is meant by the term molar heat of reaction?

This is the chemical energy change per mole ofproduct produced in a chemical reaction.

10. The reaction for the burning of octane is

2C81-1,8 + 2502 16CO2 + 18H20

Explain in molecular terms the increase of pressureand temperature that occurs in the cylinder of anautomobile engine when this reaction takes place.

The temperature increases because the productshave less chemical energy than the reactants, mak-ing the reaction exothermic. The pressure increasesbecause there is an increase in the number of gasmolecules during the reaction, and also becausethe temperature increases.

Rates of Chemical Reactions

Here students become involved with the elementof time. A chemical reaction under controlled con-ditions produces a certain amount of the productin a given length of time. The ratio of productproduced to time required is the rate of thereaction.

This portion of the module plunges immediatelyinto an experiment to illustrate some of the fac-tors that influence the rate of a reaction. The ratesof most reactions are difficult to measure quantita-tively in a student laboratory, since they requireinstrumental methods of chemical analysis. Theiodine clock reaction in experiment P-30 Timeto React is an exception, since it makes use of asudden change in color to indicate when a cer-tain amount of product has been produced.Students enjoy this experiment because of itsdramatic quality. It has the disadvantage, how-ever, of a somewhat complicated interpretation,since two reactions are involved. An optionalexperiment, Rate of Hydrolysis of Ethyl Acetate, is

42

described following experiment P-30. This op-tional experiment is simpler to interpret, but itrequires more care to perform accurately. Youmay wish to substitute this experiment for experi-ment P-30 if yours is a more advanced class.

EXPERIMENTP-30 TIME TO REACT

The purpose of this experiment is to determine the effectof reactant concentration, temperature, and presence ofa catalyst on the rate of a reaction.

Concepts

The time needed for a certain amount of product tobe produced is a measure of the rate of the reaction.The rate of a reaction increases when the concentra-tion of a reactant is increased or the temperatureis increased or a catalyst is added.The rate of a reaction is often proportional to the con-centration of each reactant.

49

Objectives

Measure the time required for the color to changein a clock reaction.Compare reaction times and reaction rates for dif-ferent reaction conditions.Decide whether a reaction rate is proportional to theconcentration of a reactant, given the reaction timesfor different concentrations of the reactant.

Estimated Time One-half to one period


Materials

1350 cm3 Solution 1 (0.10 M potassium iodide plussoluble starch)

750 cm3 Solution 2 (0.0025 M sodium thiosulfate)1350 cm3 Solution 3 (0.10 M ammonium peroxydisulfate,

sometimes called ammonium persulfate)8 50-cm3 dropper bottles of Catalyst Solution [0.01 M

copper(11) sulfate]500 cm3 distilled waterice15 150-cm3 beakers15-30 50-cm3 or 100-cm3 graduated cylinders15 10-cm3 graduated cylinders15 glass stirring rods15 thermometers (-10°C to 110°C)1 clock with sweep second hand

Advance Preparation Use distilled water in makingup the solutions. The chlorine in a municipal water supplycan ruin the experiment. To make Solution 1 prepare apaste of 0.5 g soluble starch in a little cold water. Slowlyadd the paste to 100 cm3 of boiling water, boil for a fewminutes, and allow to cool. Put this solution in a flask,add 16.6 g K1, and make up to 1000 cm3 total volume.

Solution 2 contains 0.620 g Na2 S203 5H20 in 1000 cm3.

Solution 3 contains 22.8 g (NH4)2S20, in 1000 cm3.

The Catalyst Solution contains 0.25 g CuSO4 5H20 in100 cm3, or 2.50 g CuSO4 5H20 in 1000 cm3.

The solutions may be stored for several days without anappreciable change in the reaction times.

Prelab Discussion You will want to go over thereactions on page 67 in the student module and pointout that, although it is the rate of production of 12 that is

being measured, the first amount of 12 produced is al-lowed to react with thiosulfate ions. The thiosulfate ionis present as a means of timing the first reaction. Whenno more thiosulfate ion is left, only the first reaction con-tinues to take place and a sudden color change is seen.

Laboratory Tips The reaction may not be successfulif allowed to occur above 40°C.

In Part 2, note that students are given a temperatureoption. You may wish to have different pairs of stu-dents conduct the experiment at different temperaturesin order to construct a graph of the reaction time as afunction of temperature.

Have the students repeat a reaction if they are notcertain of the reaction time to within two or three seconds.

Range of Results The standard reaction, using 20cm3 each of Solutions 1 and 3, takes 40 to 45 seconds(depending on the temperature). When the amount ofeither Solution 1 or Solution 3 is halved, the reactiontakes twice as long. Raising or lowering the temperature10 degrees makes the reaction take approximately halfor twice as long. Four drops of Catalyst Solution makesthe reaction take about half as long.

Postlab Discussion This type of reaction is knownas a "clock" reaction because the change observed issudden and reproducible, like a timer. Stress to the stu-dents that the main reaction, producing 12, is going oncontinuously. The sudden color change results from the12 suddenly appearing when the thiosulfate ions are usedup. Point out the small amount of thiosulfate present,compared with the other reactants.

Ask the students what would happen if the reaction werecarried out without the thiosulfate. Show them the reac-tion with Solution 2 omitted, and then add 10.0 cm3 ofSolution 2 and observe. If Solution 2 is added beforeabout 40 seconds, the 12 will be seen to disappear be-cause of the rapid reaction with thiosulfate, only toreappear about 40 to 45 seconds after the start of thereaction.

EXPERIMENT (OPTIONAL)RATE OF HYDROLYSIS OF ETHYL ACETATE

This optional experiment gives excellent results if thestudents are careful to measure all volumes accurately,using burets. The purpose of the experiment is to mea-sure the amount of ethyl acetate, CH3C00C2H5, that

50 43

reacts with water in a given length of time and to deter-mine how the concentration of the ethyl acetate affectsthe rate of reaction.

Concepts

The volume of sodium hydroxide solution needed totitrate the acetic acid is a measure of the extent ofthe reaction.The rate of the reaction is proportional to the con-centration of the ethyl acetate.

Objectives

Titrate an acidic solution to a phenolphthalein endpoint.Compare the relative reaction rates when given therelative amounts of product formed in a given lengthof time.Predict how the rate of the ethyl acetate reactionwill be affected by a given change in the concentra-tion of the ethyl acetate.



Materials

90 cm3 ethyl acetate140 cm3 ethanol, 95%230 cm3 6 M HCI (492 cm3 concentrated HCI in

1000 cm3)1800 cm3 1 M NaOH (40 g NaOH per 1000 cm3)1 dropper bottle phenolphthalein solution (0.1 g phenol-

phthalein in 100 cm3 95% ethanol)30 50-cm3 burets15 buret clamps (double)15 ring stands45 125-cm3 or 250-cm3 Erlenmeyer flasks45 rubber stoppers or corks to fit flasks45 50-cm3 beakers15 wash bottles of distilled water30 10-cm3 graduated cylinders

Advance Preparation Titrate 5.00 cm3 of the HCIsolution with the NaOH solution, using one drop ofphenolphthalein solution as the indicator. Titrate to apermanent pink color and record the volume. Do thisseveral times until the results are consistent, and postthe results for the students to use.

44

Pre lab Discussion The reaction studied is:

0I-13 00002 H5 ÷ H2O -0 CH, COOH + HOC, H5ethyl acetate + water --0 acetic acid + ethanol

HCI is added as a catalyst; the reaction proceeds at anegligible rate unless there is a large concentrationof H+ ion.

Each pair of students is to prepare the following threereaction mixtures in separate Erlenmeyer flasks.

Ethyl acetate Ethanol 6 M HCI

Flask 1 1.0 cm3 4.0 cm' 5.00 cm3Flask 2 2.0 cm3 3.0 cm3 5.00 cm3Flask 3 3.0 cm3 2.0 cm3 5.00 cm3

The ethyl acetate and ethanol are measured with a10-cm3 graduated cylinder. A buret is used to measureaccurately the HCI solution. The purpose of the ethanolis to allow the ethyl acetate to dissolve in the acid, andto make the total volume of each mixture the same.

As soon as each mixture is prepared, the flask is

stoppered and swirled to mix the solution. The startingtime is recorded. While the mixtures are reacting, aburet is used to deliver accurately into each of three50-cm3 beakers the exact amount of NaOH solutionneeded to neutralize 5.00 cm3 of the HCI solution. Whena mixture has reacted for 20 minutes, the NaOH solutionis rinsed from one of the beakers into the flask to stopthe reaction by neutralizing the HCI catalyst.

The contents of each flask are then titrated with the NaOHsolution, using one drop of phenolphthalein solution asthe indicator, in order to determine the amount of aceticacid produced in the 20-minute period.

The reaction rates in the three mixtures are comparedby comparing the volumes of the NaOH solution requiredin the three titrations.

Laboratory Tips Stress the fact that the volumes ofHCI and NaOH solutions must be accurately measuredin order that they may exactly neutralize one another.Also, the NaOH solution must be completely rinsed intothe reaction mixture.

The NaOH solution should be freshly made up, as itslowly reacts with CO, present in the air.

You may want to demonstrate the use of a buret andthe appearance of the end point. The students shouldmake sure there are no bubbles anywhere in the buret,

51

and that they use the bottom of the meniscus to readthe volume.

The end point is easy to distinguish, but it will fade afterseveral minutes because of a slight further amount ofreaction. Tell the students approximately how muchNaOH solution to add to each mixture before theyshould add it drop by drop. The students should period-ically rinse down the walls of the flask with distilled waterduring a titration so that all of the NaOH reaches thereaction mixture.

Range of Results Typical results are:

30.42 cm3 NaOH solution needed to neutralize thecatalyst. Flask 1 required 3.4 cm3 NaOH solution toneutralize the acetic acid. Flask 2 required 6.7 cm3.Flask 3 required 10.4 cm3. Thus the reaction rate(amount of acetic acid produced in 20 minutes) in Flask2 is almost exactly twice as great as in Flask 1; therate in Flask 3 is almost exactly three times as great.

Post lab Discussion The initial concentrations ofethyl acetate in Flasks 2 and 3 were twice and threetimes as great as in Flask 1. The reaction rate is seento be directly proportional to the initial concentration ofethyl acetate. (The initial water concentration was thesame in each mixture.) Stress that the reactions werefar from complete when the NaOH was added. Thusthe reason that three times as much acetic acid wasproduced in Flask 3 compared with Flask 1 is not thatthere was three times as much ethyl acetate to react,but that it reacted three times faster.

P-31 RATES AND CONCENTRATIONS

Make sure that the students understand the sig-nificance of their results in experiment P-30 orin the optional experiment described above. In-creasing the concentration of a reactant increasesthe rate of the reaction, and often there is a directproportionality between concentration and rate.

The rates measured in these reactions were ini-tial rates in solutions having only low concentra-tions of products. Therefore, the reverse reactionshad little influence on the results. Of course, intime the reactant concentrations become smallerand the product concentrations become largeruntil the rates of the forward and reverse reactionsare equal. At this point, chemical equilibrium is

attained. You will move on to this concept startingin section P-34.

The equation relating the initial rate of a reac-tion to the concentrations of the reactants is calleda rate law. Three examples of rate laws are

(a) rate = k, [A] (b) rate = k2 [A] 2 (C) rate = k 3 [A] [B]

The exponent of a concentration in a rate lawspecifies what is called the order. For instance,rate law (a) describes a reaction that is of order1 with respect to reactant A; rate law (b) is for areaction of order 2 with respect to reactant A; andrate law (c) is for a reaction of order 1 with respectto A and order 1 with respect to B.

Although it is not possible to tell from the ratelaw what the elementary steps in a reaction are,it is possible to deduce the rate law if the ele-mentary steps and their rate constants are known.For instance, if reactant A changes to a productin a single elementary step without a collisionneeded (a unimolecular reaction), rate law (a)results. Rate law (b) will result if two moleculesof reactant A collide in a single elementary step(a bimolecular reaction). The iodine clock reactionstudied in experiment P-30 has a rate law of theform of reaction (c); although this rate law wouldresult from a single elementary step in whichmolecules A and B collide, the situation is in factmore complicated, as described in the studentmodule.

P-32 RATES AND TEMPERATURES

Almost all the increase in rate of a reaction thatoccurs when the temperature is raised resultsfrom the activation energy, not from the in-creased rate of collisions.

To illustrate this point, the reaction 21-IIH2 + 12 taking place in the gas phase can be con-sidered. The reaction is known to have a single,bimolecular elementary step in which two HImolecules collide.

Consider two containers, each of 1 cm3 volumeand containing 1.81 x 105 mol (1.09 x 1019 mole-cules) of HI. One container is at a temperatureof 400°C and the other is at 410°C. The fol-lowing can be calculated from experimentalmeasurements

52 45

Container 1 Container 2

Temperature, °C 400 410Pressure, atm 1.00 1.01

Collisions per second 1.035 x 1028 1.043 x 1028Fraction of collisions

with kinetic energygreater thanactivation energy

1.58 x 10-'5 2.60 x 10-"

Collisions per secondresulting in reaction

1.64 x 10'3 2.71 x 10'3

The number of collisions per second that resultin a reaction (bottom line) is the product of thetotal number of collisions per second and the frac-tion of the collisions that have kinetic energygreater than the activation energy. Notice thatraising the temperature by 10 degrees (from400°C to 410°C) makes the rate increase by 65percent, whereas the number of collisions in-creases by less than 2 percent.

Demonstration Dust Explosion: Show that a moundof lycopodium powder doesn't readily burn. Then shakelycopodium powder inside a dry plastic squeeze bottleand squirt it into a Bunsen burner flame; or place thepowder in a rolled-up piece of paper and blow it intothe flame. A "dust explosion" results. Dry flour may alsobe used, but it is not as effective.

P-33 CATALYSTS: ESSENTIAL FOR LIFE

An example of a catalyst is the enzyme calledcarbonic anhydrase, found (among other places)in the red blood cells of animals. The functionof this catalyst in red blood cells is to catalyze thehydration of dissolved carbon dioxide in theblood: CO2(aq) + H2O H+ + HCO3 -. The re-sulting hydrogen carbonate ion, HCO3 -, is carriedin the blood from the tissues to the lungs, wherethe reverse reaction occurs (again catalyzed bycarbonic anhydrase) and the CO2 is expelled in thebreath. Without the presence of this enzyme, theblood could not carry CO2 rapidly enough to thelungs to prevent the animal's smothering.

The actual elementary steps are probably these

46

CO2(aq) + E ECO,E0O2 + H2O --o E HCO3 + H+E HCO3 0 E + HCO3

Sum: CO2(aq) +11,0 H+ + HCO3

E represents the enzyme molecule, and E0O2and E HCO3 represent intermediate forms inwhich the enzyme forms a covalent bond withother atoms. The second step shown above israte-determining, and it occurs at a faster rate thanthe uncatalyzed reaction because the activationenergy is less. Notice that the enzyme is re-generated in the last step, ready to be used again.The net reaction, obtained by adding the equa-tions for the three elementary steps, is the sameas that of the uncatalyzed reaction.


(Sudent module page 72)

1. Three factors that can increase the rate of a reactionare: an increase in the concentration of a reactant;an increase in temperature; the addition of a catalyst.

2. (a) Assuming that the reaction time is 43 secondsin the first measurement in Part 1 of experimentP-30 Time to React, the calculation of the rateconstant is as follows:rate = (2.5 x 10-4 mole/liter)/43 seconds

= 5.8 x 10-6 mole/liter secondk = (5.8 x 10-6 mole/liter second)/

(0.040 mole/liter)(0.040 mole/liter)= 3.6 x 10-3 liter/mole second

(b) Similar values of k should be obtained from theother measurements in Part 1.

(c) If [I- ] = [S2033- ] = 0.020 mole/liter, and assum-ing the values of k shown in part (a) of this ques-tion, the calculation israte = (3.6 x 10-3 liter/mole second) x

(0.020 mole/liter)(0.020 mole/liter)= 1.4 x 10-6 mole/liter second

Halving the concentration of both of the reactantscauses the rate to be reduced to one-quarterof its original value.

53

A. activation energy C. reaction rateEVALUATION ITEMS B. reaction temperature D. reactant concentration


1. Briefly explain the relationship between activationenergy and the rate of a chemical reaction.

The lower the activation energy, the faster the rate ofthe reaction.

2. Which of the following is increased by the presenceof a catalyst?

3. Predict the effect of each of the following changeson reaction rate.

A. a decrease in temperaturedecreases reaction rate

B. an increase in the concentration of a reactantincreases reaction rate:

C. the addition of a catalystincreases reaction rate

Chemical Equilibrium

In the earlier sections on dynamic equilibria andrates of reaction, the module has been buildingtoward the topic of chemical equilibrium. Nowwe can put these ideas together and derive theconcept of the equilibrium constant.

A demonstration called the Water Game is re-commended in section P-34 Equilibrium Constantsof this guide to help your students visualize theresults of opposing reactions. In miniexperimentP-35 Shifting an Equilibrium, the students caneasily see how the addition of a reactant to anequilibrium mixture shifts the equilibrium. Theexperiment P-38 Show Your Colors allows the stu-dent to measure an equilibrium constant and toconfirm that it does not vary with the concentra-tions. As in the previous portion of the studentmodule dealing with rates of chemical reactions,an optional experiment is suggested in this guidethat may be substituted for experiment P-38.This experiment, Measurement of the Ethyl AcetateEquilibrium Constant, ties in nicely with the pre-ceding optional experiment, Rate of Hydrolysis ofEthyl Acetate. However, it requires more carefulpreparation of solutions than in the case ofexperiment P-38.

P-34 EQUILIBRIUM CONSTANTS

Strictly speaking, instead of defining an equilib-rium constant as the value of a concentrationquotient at equilibrium, we should define anequilibrium constant as the value of an activityquotient at equilibrium. The activity of a com-pound is equal to the concentration multipliedby an activity coefficient. In a dilute solution orin a gas, the value of the activity coefficient ofeach compound or species is close to unity; thusthe equilibrium constant in such systems may bedefined in terms of concentrations, as in thestudent module. In concentrated solutions, how-ever, there may be substantial deviations ofactivity coefficients from unity, and the concen-tration quotient may not be exactly a constantfor various equilibrium mixtures.

A chemical reaction taking place at a certaintemperature proceeds in the direction that allowsthe concentration quotient to approach the valueof the equilibrium constant at that temperature.On the molecular level, this can be explained bypointing out that the rates of certain elementarysteps occur faster than the reverse steps.

A general statement may also be made on amore macroscopic level. If the temperature andpressure stay constant, the reaction tends to

54 BEST COPY AVAILABLE 47

proceed in the direction that lowers the potentialenergy (that is, the chemical energy). However,in addition there is a natural tendency for aquantity called the entropy to increase. The en-tropy may loosely be defined as the moleculardisorder (see below). If it is not possible for thechemical system to simultaneously decrease itspotential energy and increase its entropy, thenwhichever happens to be the stronger tendencywill override the other to determine the directionof the reaction and the position of the finalequilibrium.

The two natural tendencies, decrease in po-tential energy and increase in entropy, have beenlikened to the human drives toward security andfreedom, respectively. (For an interesting elabora-tion of this analogy, see Frederick D. Rossini,"Chemical Thermodynamics in the Real World,"Chemical and Engineering News [April 5, 1971],pp. 50-53.) These tendencies are often incom-patible, and a compromise is established in reach-ing the equilibrium. If AH is the change inpotential energy (equal to the enthalpy changewhen the temperature and pressure do notchange), and AS is the entropy change, then thebest compromise is the one that makes AHTAS as negative as possible. (T is the tempera-ture in kelvins.) AH T AS is known as the freeenergy change, or AG. The free energy changewill be more negative the more negative is thepotential energy change and the more positiveis the entropy change, so the natural tendencyof AG is to become as negative as possible.

The entropy change ( AS ) of a chemical systemhas a precisely defined meaning in the disciplineof thermodynamics, but not one that is particu-larly meaningful to students at the high-schoollevel. Teachers who attempt to give students a"feel" for AS usually say that it is a quantitativemeasure of the change in molecular "disorder" or"randomness." If a system becomes more "disor-dered," AS is positive. If the system becomes lessdisordered (or more ordered), AS is negative.Unfortunately, however, these descriptions mayin some cases be confusing or even misleadingwhen one tries to decide on the sign of AS; in anycase, they do not explain how AS is to be mea-sured quantitatively. For this reason, the conceptsof entropy and free energy have been omittedfrom the student module. (For discussions of thedangers in interpreting entropy changes in terms

48

of molecular disorder, see: M. L. McGlashan,"The Use and Misuse of the Laws of Thermo-dynamics," Journal of Chemical Education, Vol. 43[May 1966], pp. 226-232; and P. G. Wright,"Entropy and Disorder," Contemporary Physics,Vol. 11 [1970], pp. 581-588.)

If you do wish to discuss entropy changes inyour class, the following statement gives a rea-sonably accurate picture. One kind of change thatcauses the entropy of a chemical system to in-crease at constant temperature is an increase inthe volume available for molecules to move in.The analogy of AS to an increase in "freedom,"rather than in increase in "disorder," works besthere. Using the picture, we see that AS is positivewhenever a liquid vaporizes at constant tempera-ture, since the volume of the resulting vaporphase is greater than the volume of the originalliquid. Also, AS is expected to be positive whena solid dissolves in a solvent, since the resultingsolution occupies a greater volume than the ori-ginal solid. (Another contribution to the entropychange may, however, also arise from changesin the solvent structure as the solid dissolves.)

Now it is possible to explain an approach toequilibrium with the use of these macroscopic(thermodynamic) principles. We will use as ourexample the dissolving of a molecular crystal inwater, assuming no changes in solvent structureoccur. The graph shows how AH, TAS, andAG depend on the amount of solid dissolved.Each of these quantities has a natural tendencyto become as negative as possible. (The tendencyof AS is to become as positive as possible; thus

TAS, because of the minus sign, tends to be-come as negative as possible.) However, AH andTAS cannot both become negative at the sametime in this example. AG is the sum of AH and

TAS. The solid dissolves until AG is as negativeas possible.

What happens when the solid is first placed incontact with the water? Each molecule that dis-solves causes an increase in entropy (favorable)and also an increase in enthalpy (unfavorable).However, TAS is larger in magnitude than AH,and so the entropy effect overrides the enthalpyeffect. AG becomes more negative as more soliddissolves.

What causes an equilibrium to be eventuallyestablished? Put another way, why doesn't thesolid continue to dissolve to an indefinite extent,

55

3

2

1

0

1

2

3

0 0.5 1.0 1.5

n = moles crystal dissolved

2.0

Values of OH, TAS, and AG for the process of dissolvingn moles of a molecular crystal in 1000 cm3of water at 25°C.The curves were constructed for a hypothetical crystal on theassumptions that the concentration of the saturated solutionis one molar and the entropy change is due entirely to idealmixing of the solute and solvent molecules.

thereby continually increasing the entropy? Basic-ally, the reason is that the entropy does notincrease by an equal amount as each succeedingmolecule dissolves. The entropy increase permolecule depends on the number of solute mole-cules already present in the solution. (The ex-planation for this fact comes from statisticaltheory.) Consequently, TAS is not a linear

function of the amount dissolved. OH is such alinear function. Thus, when a certain concentra-tion of solution has been reached in the dissolvingprocess, the free energy no longer decreaseswhen a molecule dissolves. The entropy effectis exactly balanced by the enthalpy effect. Thisconcentration is the equilibrium concentration, orthe concentration of the saturated solution.

Why does precipitation occur in a supersat-urated solution? Precipitation occurs because atconcentrations greater than the equilibrium con-centration the enthalpy effect overrides the en-tropy effect. By precipitating, the molecules causethe enthalpy to decrease.

An astute student may ask about the case ofa solid that has a dissolving process that is exo-thermic, instead of endothermic as in the preced-ing example. In this case, it would seem as ifequilibrium could never be established becauseboth the enthalpy and entropy effects are favor-able for dissolving. Yet we know that such a solid(NaOH, for example) can in fact form a saturatedsolution. The answer to this paradox is that inaddition to the entropy increase resulting fromthe increased volume available to the soluteparticles after they dissolve, there is an entropydecrease resulting from the ordering of solventmolecules around dissolved solute particles. Suchan entropy decrease, when it occurs, is knownas the "unitary" entropy or "contact" entropy andis a linear function of the amount of solute dis-solved. The graph above would still illustrate thesituation for the approach to equilibrium of such acrystal, except that the curve labeled AH wouldbe relabeled OH T AS (unitary) and the curvelabeled T AS would represent only the volumechange component (or "ideal mixing" compo-nent) of the entropy.

The complexity of this discussion will serve toillustrate why the concept of entropy is bestavoided unless it can be fully explained. It is evenmore difficult to interpret entropy changes in thecase of a chemical reaction, where there is usuallylittle change in the volume and the "disorder"has to do mainly with the bonding arrangementsin the reactants and the products.

Demonstration The Water Game: The purpose ofthis demonstration is to show the attainment of dynamicequilibrium in a simple physical system.

5649

The demonstration is played as a game between twopeople. Let two students at a time carry out the gamein front of the class. The apparatus needed: two 50-cm3burets clamped in a double buret holder on a ringstand, masking tape, and two 50-cm3 beakers. Placea strip of masking tape along each buret to cover thenumbers but not the divisions. Use the tape to renumberthe divisions from 0 cm3 to 50 cm3 from bottom to topinstead of from top to bottom as the permanent numberson the buret are arranged. Also provide at least 100 cm3of water colored with food coloring for good visibility.

Explain to the class that each player will have a buret.The volume of water in the buret represents the con-centration of a compound. Player A's water representscompound A, and player B's water represents com-pound B. The two compounds can react as follows

A B

At each "move" of the game, one-half of compound Awill turn into compound B and one-quarter of compoundB will turn into compound A. In other words, the amountremoved from A will be proportional to the concentra-tion of A, and similarly for B. Point out that this "ratelaw" is similar to the one governing the reaction ofI- ions with S2082- ions. The fractions one-half and one-quarter are like rate constants.

Have the players start with 25 cm3 of water in eachburet. This represents equal initial concentrations ofcompound A and compound B. Each "move" of the gameis played as follows. Player A withdraws one-half of his/her water into an empty beaker. Player B withdrawsone-quarter of his/her water into another beaker.The two players then exchange the beakers and eachplayer pours all of the water from the beaker he/she hasreceived into the buret. This represents the amount ofcompound B which was produced per 1000 cm3 in acertain interval of time by the reaction of compound A,and vice versa. Ask the class to record the new vol-umes of water in each buret after each move; thesevolumes represent the concentration of A and B aftereach equal interval of time.

The players should repeat these moves for a total ofat least five moves, or until the volumes no longer changemuch. At this point, the situation represents a dynamicequilibrium. Do not give anything away. Just stand backand let the class see what happens. Ask the class toexplain why the water volumes no longer change whenthe exchanges of water take place.

50

Now ask two more students to repeat the game, lettingplayer B have all the water initially. Start with no waterin player A's buret and 50 cm3 of water in player B'sburet. Ask the class to explain the results.

Finally, ask two more students to play the game a thirdtime. Start with 10 cm3 of water in A's buret and 50cm3 in B's buret. This time there is a total of 60 cm3 ofwater instead of 50 cm3 as in the first two games.

After the games have been played and the resultsrecorded, calculate the ratio of the final water volumein B's buret to the final water volume in A's buret ineach of the three games. These ratios are like equilib-rium constants.

The table below lists the expected results. The volumes(in cm3) are rounded off from the exact values cal-culated on the assumption of accurate measurementsand no loss of water during the transfers.

Game 1

A B

Game 2

A B

Game 3

A

Start 25.0 25.0 0 50.0 10.0 50.0

Move 1 18.8 31.2 12.5 37.5 17.5 42.5

2 17.2 32.8 15.6 34.4 19.4 40.6

3 16.8 33.2 16.4 33.6 19.8 40.2

4 16.7 33.3 16.6 33.4 20.0 40.0

5 16.7 33.3 16.7 33.3 20.0 40.0

After each game, the ratio of the amount of water inB's buret to that in A's buret is the same as the ratioof the "rate constants"-namely, 2.00.

The total volume of water used in Games 1 and 2 is50 cm3. In Game 1 the net movement of water is fromA to B, and in Game 2 from B to A, but the volumesin each buret are the same after both games. This illus-trates that reactions can proceed in either direction, de-pending on the starting conditions, but will reach thesame equilibrium. Game 3 shows that even thoughthe final volumes in the burets are different than before,the ratio of the volumes is the same.

Ask the class the following questions about Game 1:

Why does the water level in A go down? (The amountof water going out exceeds that coming in.)Why does the water level go down more slowly astime goes on? (The difference between what is goingout and what is coming in is decreasing.)

57

What is the necessary condition that must be metbefore the water levels will stop changing? (Theamount coming in must equal the amount going out.)Ask a similar set of questions for Games 2 and 3.Ask the students to predict the outcome of similarinvestigations; for example, suppose kA = 1/2 andk8 = 1/10.

MINIEXPERIMENTP-35 SHIFTING AN EQUILIBRIUM

The pupose of this experiment is to show what happensto the equilibrium concentration of a product when morereactant is added.

Concepts

The intensity of the red color in a solution containingiron(III) ions and thiocyanate ions is a measure ofthe concentration of the FeSCN2+ ion formed.The addition of either of the reactants shifts theequilibrium to produce more product.Both reactants are equally effective in shifting theequilibrium.

Objective

State the effect of adding a reactant on the equilib-rium concentration of a product.



Materials

225 cm3 0.01 M potassium thiocyanate (0.97 g KSCNin 1000 cm3)

225 cm3 0.01 M iron(III) nitrate, acidified (4.04 gFe(NO3)3. 9H20 and 2 cm3 concentrated nitric acidin 1000 cm3)

30 150-cm3 beakers15 100-cm3 graduated cylinders15 10-cm3 graduated cylinders15 glass stirring rods

Range of Results The initial solution of 100 cm3volume is colored pale red. When the additional amountsof solution containing Fe3+ ions or SCN- ions areadded to equal portions of the initial solution, a deeperred color results. The students should notice that thecolor changes about equally after both additions; in otherwords, the two reactants give equivalent effects.

53

Postlab Discussion Since the initial mixture con-tained 10 cm3 of 0.01 M potassium thiocyanate solu-tion, diluted to a volume of 100 cm3, the totalconcentration of thiocyanate ion (before reaction) was0.01 M x (10 cm3/100 cm3) = 0.001 M. The total con-centration of Fe3+ ion was likewise 0.001 M. Thereaction

Fe' + SCN- FeSCN2+

occurred and equilibrium was attained immediately. Theequilibrium concentrations listed in the text can beverified by noting that (1) they give the correct value of138 liter/mole for the concentration quotient; and (2) thetotal iron concentration (Fe3+ + FeSCN2+) is 0.001 Mand the total thiocyanate concentration (SCN +FeSCN2+) is also 0.001 M.

The light red color is indicative of a complex concen-tration of 1.1 x 10-4 M.

In the last step an additional 0.000 05 mol of a reactant(SCN- or Fe3÷) was added to half of the initial solution.It made no difference which reactant was added, as theyboth play exactly the same role in forming the complex.If the added reactant had not reacted, the concentrationquotient would have been smaller than 138 liter/mole.Therefore, some of it reacted until equilibrium wasreached. The concentration of the added reactant andof the other reactant decreased and the complexconcentration increased to bring the concentrationquotient back to 138 liter/mole. The equilibrium con-centrations were now

added ion: 16.5 x 10-4 Mother ion: 7.4 x 10-4 M[FeSCN2+] = 1.7 x 10-4 M

In general, when a chemical system is at equilibriumand more of one of the reactants is added, some ofthe added reactant reacts to form more of the product(or products). The equilibrium is shifted "to the right."

If product is added to an equilibrium solution, the reversereaction occurs and more reactant is formed. Theequilibrium is shifted "to the left."

If product is removed from solution, the reactants willreact to form more product. This is the basis of somelaboratory procedures and industrial processes for ob-taining large amounts of product by continuously re-moving it from a liquid reaction mixture as a gas, or asanother immiscible liquid, or as a solid.

These shifts in an equilibrium are one kind of exampleof Le Chatelier's principle. Le Chatelier's principle states

51

that if a dynamic equilibrium is present in a chemicalsystem, and the equilibrium is disturbed in some way,the system will change in such a way as to counteractthe disturbance. For instance, when a reactant is addedto an equilibrium system the equilibrium is shifted to theright so as to use up some of the added reactant.

P-36 THE IMPORTANCE OF SIZE

In the student module, examples are given of anextremely large equilibrium constant and an ex-tremely small equilibrium constant. The valuesof these equilibrium constants cannot be deter-mined by measuring the concentrations present atequilibrium, because some of the concentrationsare, for all practical purposes, zero. Instead, thevalues are calculated from measurements of en-thalpy and entropy changes or by other means.

P-37 HYDROGEN IONS: THE ACID TEST

The concentration of H2O in liquid water iscalculated

(1000 gram/liter)/(18.0 gram/mole) = 55 mole/liter.

The ion product of water has the value 1.00 x10-'4 (mole/liter)2 at 25°C. It is a coincidence thatthis is such a simple value. It is different at othertemperatures; for example, the ion product ofwater has the value 2.4 x 10-'4 (mole/liter)2 at37° C.

The ionization constant of a weak acid is usuallywritten K and is sometimes called the acid dis-sociation constant. The ionization constant of aweak base is usually written Kb and is sometimescalled the base dissociation constant. The expres-sion for the ionization constant of ammonia inwater is

[NH4+][0H-1[NH3]

where the concentrations appearing in the con-centration quotient are the equilibrium values.

By a coincidence, the ionization constants ofacetic acid and ammonia at 25°C both have thesame value, 1.8 x 10-5 mole/liter. Note, however,that the ionization reactions are quite different forthe two compounds. One produces hydrogenions and the other produces hydroxide ions. Asolution of acetic acid is acidic, and a solution ofammonia is basic.

52

EXPERIMENTP-38 SHOW YOUR COLORS

The purpose of this experiment is to determine thevalue of the concentration quotient for the ionizationof acetic acid in solutions of different total concen-tration, and to show that the concentration quotient isa constant.

Concepts

The color of an acid-base indicator depends on thehydrogen ion concentration.Different indicators change color in different rangesof hydrogen ion concentration.A solution of acetic acid in water has a lower hydrogenion concentration than a solution of hydrochloric acidat the same concentration.The concentration quotient for the ionization reactionof acetic acid is a constant.

Objectives

Prepare by dilution a solution having one-tenth theconcentration of a previous solution.Determine the hydrogen ion concentration of a solu-tion by the use of acid-base indicators and standards.Calculate the value of the concentration quotient forthe ionization reaction, given the total acetic acid con-centration and the hydrogen ion concentration.



Materials

300 cm3 0.1 M HCI (8.6 cm3 conc. HCI in 1000 cm3)165 cm3 0.1 M acetic acid (5.7 cm3 glacial acetic acid

in 1000 cm3)dropper bottles of solution of either orange IV (trope-

oline 00) or thymol blue sodium salt (see AdvancePreparation)

dropper bottles of solution of one of the following:methyl orange, or methyl orange sodium salt, orbromocresol green sodium salt (see AdvancePreparation)

7000 cm3 distilled water210 18 x 150-mm test tubes15 10-cm3 graduated cylinders15 100-cm3 graduated cylindersmasking tape for labels15 test-tube racks15 125-cm3 or 250-cm3 Erlenmeyer flasks

59

Advance Preparation The HCI and acetic acid stocksolutions should be carefully made up shortly before theexperiment, using distilled water. To save time duringthe experiment, you may wish to provide some or allof the dilute solutions.

High-purity grades of the indicators should be used.Each of the indicator solutions is prepared by dis-solving 0.1 g of the solid form in 100 cm3 distilled water.

Pre lab Discussion It would be a good idea to demon-strate the dilution technique described in the experiment.The purpose of pouring the solution back and forth be-tween two vessels is to ensure good mixing.

Laboratory Tips A convenient way to label the testtubes is with strips of masking tape. The strips may bewritten on with a soft pencil or ball point pen.

Ammonia fumes in the room can ruin the experimentby neutralizing some of the acid in the solutions.

Distilled water should be used for the dilutions. If it isnecessary to conserve the distilled water, the aciddilutions could be carried out in 10-cm3 graduated cylin-ders instead of 100-cm3 graduated cylinders to givesmaller volumes.

Range of Results The standards containing HCI havethe following colors, depending on which indicators areused:

orange IV[H+] or thymol methyl bromocresol

(mole/liter) pH blue orange green

yellowyellowyellowgreenblue

1.0 x 10-' 1.0 purple-red rose-red1.0 x 10-2 2.0 orange-red rose-red1.0 x 10-3 3.0 orange-yellow orange-red1.0 x 10-4 4.0 yellow orange-yellow1.0 x 10-5 5.0 yellow yellow

Orange IV and thymol blue change color in the pH range1-3, whereas methyl orange and bromocresol greenchange color in the pH range 3-5. By using two indica-tors, the students should be able to determine the pHof any solution in the range 1-5 to the nearest pH unit.

The accepted value of the ionization constant of aceticacid is Ka = 1.8 x 10-5 mole/liter. The theoretical valuesof [H+] for the two acetic acid solutions are 1.3 x 10-3mole/liter for the 10-' M solution and 1.2 x 10' mole/liter for the 10-3 M solution.

The students should calculate the following values forKa by assuming that the acetic acid solutions have[H+] equal to the standards giving the closest colormatch

c = 10-1 mole/liter:[H+] = 10-3 mole/liter[H+]2/c = 10-5 mole/liter

c = 10-3 mole/liter:[H +] = 10-4 mole/liter[H+]2/c = 10-5 mole/liter

The calculated value of [H+]2/c is close to the acceptedvalue of Ka, 1.8 x 10-5 mole/liter.

Miniexperiment Have the students drop a small pieceof magnesium ribbon into some 0.1 M HCI and into some0.1 M acetic acid. Hydrogen gas is more rapidly evolvedfrom the HCI. The concentrations of both acids areequal, but [H+] is greater in the HCI because it is astrong acid whereas acetic acid is not.

Miniexperiment Another activity is to have the stu-dents separately titrate equal volumes of 0.1 M HCI and0.1 M acetic acid with 0.1 M NaOH, using an indicatorsuch as phenolphthalein to detect the endpoints. Thesame volume of NaOH is required to reach the end-point in the titration of the HCI as in the titration of theacetic acid. Thus the two acids have the same amountsof potentially available H+, but only part of this H+ isionized in the acetic acid solution.

EXPERIMENT (OPTIONAL)ETHYL ACETATE EQUILIBRIUM

The purpose of this optional experiment is to measurethe equilibrium constant for the reaction in which ethylacetate is hydrolyzed to acetic acid and ethanol

CH3C00C21-15 <-7= C1-13COOFI + HOC, H5

This experiment may be used in place of experimentP-38 for classes of better-than-average students. It re-quires a more careful measurement technique than ex-periment P-38, but is capable of giving better precisionin the calculated value of the equilibrium constant. Youmight want to consider doing this experiment if your classcarried out the optional kinetics experiment (Rate ofHydrolysis of Ethyl Acetate) to replace experimentP-30, since the reaction is the same in both optional

6Q 53

experiments and the students will already have hadpractice in carrying out titrations.

Concepts

The concentration of each substance in the mixturecan be determined from the concentration of the aceticacid.The concentration quotient for the reaction is aconstant.

Objectives

Calculate the number of moles of acetic acid in eachflask, given the volume and concentration of sodiumhydroxide solution used in the titration.Calculate the concentration of each substance in oneof the flasks from the number of moles of acetic acid.Calculate the value of the concentration quotient forthe equilibrium mixture in each flask.

Estimated Time First day: one-half period. Secondday (two or more days after the first day): one period.


Materials

230 cm3 6 M HCI (492 cm3 conc. HCI in 1000 cm3)3000 cm3 1 M NaOH (40 g NaOH in 1000 cm3)40 cm3 ethyl acetate75 cm3 glacial acetic acid120 cm3 ethanol, 95 percent45 125-cm3 or 250-cm3 Erlenmeyer flasks45 rubber stoppers to fit flasks30 50-cm3 burets15 buret clamps (double)15 ring stands15 wash bottles of distilled water1 dropper bottle phenolphthalein solution (0.1 g phenol-

phthalein in 100 cm3 95 percent ethanol)

Advance Preparation Shortly before the second dayof the experiment, determine the volume of NaOHsolution needed to exactly neutralize 5.00 cm3 of theHCI solution.

Pre lab Discussion The students are to make up thefollowing mixtures in Erlenmeyer flasks, dispensing thematerials carefully from burets.

54

Ethyl Acetic6 M HCI Acetate Acid Ethanol

Flask 1 5.00 cm3 2.50 cm3 0 2.50 cm3

Flask 2 5.00 0 3.00 cm3 2.00

Flask 3 5.00 0 2.00 3.00

The number of moles of each material in the mixtures,before reaction begins, is as follows.

Ethyl Acetate Acetic Acid Ethanol

Flask 1 0.025 mole 0 0.041 moleFlask 2 0 0.053 mole 0.033

Flask 3 0 0.035 0.050

Each flask also contains 0.277 mole of H2O and 0.030mole of HCI. (These amounts are calculated from thevolumes used and the known densities and molarmasses of the compounds.)

The flasks are stoppered tightly, swirled to mix the con-tents, and set aside for at least two days.

In two days, chemical equilibrium will have been reachedin each mixture. Tell the students how much NaOH solu-tion to add to each flask to neutralize exactly the HCI(which is present as a catalyst to shorten the time needed

to reach equilibrium).

The students are then to determine the final amount ofacetic acid in each flask by adding one drop of phenol-phthalein solution and titrating the mixture with 1 MNaOH.

Since each cubic centimeter of the 1 M NaOH neutralizes0.001 mole of acetic acid, it is easy to calculate thenumber of moles of acetic acid in each flask. The totalvolume of reaction mixture may be assumed to be 10cm3. The concentration of acetic acid in each reactionmixture is therefore one hundred times the number ofmoles of acetic acid.

Calculating the amounts of water, ethyl acetate, andethanol in each equilibrium mixture is more difficult. It

can be done by realizing that there is a connection be-tween the change in the amount of acetic acid and thechange in the amounts of each of these other com-pounds. If the acetic acid increased by x moles duringthe reaction, then the amount of ethanol increased byx moles and the amounts of water and ethyl acetatedecreased by x moles. If, on the other hand, the acetic

61

acid decreased by x moles, then the amount of ethanoldecreased byx moles and the amounts of water and ethylacetate increased by x moles. From the known initialamounts of each compound, and the measured finalamount of acetic acid, the students are to calculate theamounts of each compound in the final equilibrium mix-tures and the final concentrations.

Finally, the students are to calculate a value for the con-centration quotient

[CH, COOH][HOC2H5l[C1-1,C00C21-15][H20]

in each of the three equilibrium mixtures. They are tosee whether the concentration quotient is a constant.Note that water concentration can vary in this systemand is thus included in the concentration quotient.

Range of Results A typical set of data is as follows

Volume of 1 M NaOH needed to neutralize 5.00 cm36 M HCI: 29.9 cm3.

Titration results

Flask 1Flask 2Flask 3

11.9 cm3 NaOH solution36.020.2

Acetic acid calculations

Flask 1Flask 2Flask 3

Moles acetic acid Change in Moles

0.0120.0360.020

0.012 (increase)0.017 (decrease)0.015 (decrease)

Calculated concentrations in equilibrium mixtures

(CH3C00C2H5) [H20) [CH3COOH] [HOC2H5]

Flask 1 1.3 M 26.5 M 1.2 M 5.3 MFlask 2 1.7 29.4 3.6 1.6Flask 3 1.5 29.2 2.0 3.5

Values of concentration quotient

Flask 1 0.18Flask 2 0.12Flask 3 0.16

Post lab Discussion The exact value of the equilib-rium constant depends slightly upon the concentrationof the acid, since this in turn affects the activity coef-ficients of the various reactants. For the acid concentra-tion used in this experiment, the value of the equilibriumconstant should be about 0.2.

The variation in the data shown above is typical. Pointout that considering the differences among the threestarting mixtures, the surprising thing is not that theyare so different, but that they are so similar.

Emphasize the fact that if we know the equilibrium con-stant for a certain chemical reaction, we can predictwhether or not any subsequent changes in the com-position of the system can be expected. Thus it is pos-sible to predict that the initial mixture in Flask 1 shouldproduce more acetic acid (by a net forward reaction),whereas the mixtures in Flasks 2 and 3 should produceethyl acetate (by the reverse reaction).



1. The solution becomes more orange. Adding H+, areactant, shifts the equilibrium so as to produce moreof the products.

2. At equilibrium, K = [C][D] /[A][B]= (0.010 mole/liter)(0.040 mole/liter)/

(0.010 mole/liter)(0.020 mole/liter) = 2.0

In the new solution, [C][D]/[A][B] = 2.0;(0.010 mole/liter) [D]/(0.010 mole/liter)2= 2.0; [D] = 0.020 M

3. [1-1+1[1n-V[HIn] = KindAt pH = 2: [HIn] /[1n-] = [H +]/Kind

= 10-2 M/6 x 10-6 M= 2 x 103

The concentration of the red form of the indicatoris 2000 times greater than that of the yellow form;the solution is red.At pH = 5: [HIn]/[In-] = 10-5 M/6 x 10-6 M = 2The concentration of the red form is twice that of theyellow form; the solution is red-orange.At pH = 8: [In-] /[HIn] = Kind[H+]

= 6 x 10-6 M/10-8 M= 600

The concentration of the yellow form is 600 timesthat of the red form; the solution is yellow.

62 55

EVALUATION ITEMS


1. A + B C

C ---0 A + 8rate of production of C = ki [A][8]rate of consumption of C = k2[C]

Using the above information, write an expression forthe overall equilibrium constant (K) for the reaction,first in terms of the rate constants and second interms of concentrations of materials.

K=k2

[K [C]

[A][8]

2. Write an expression for the equilibrium constant ofthe reaction: H2S 4=-0 2H+ + S2-

K [H12[S2-][H2S]

3. Given the reaction Fl2PO4- 2H+ + P043-, whatwill happen to the equilibrium if trisodium phosphate,Na31304, a source of P043- ion, is added?

The equilibrium will shift to the left (toward production

of more H2PO4 ).

4. Which of the following is the correct expression for theion product (equilibrium constant) of water?

A. [H+] = 10-7 mole/literB. [H+][0H-] = 10-14 (mole/liter)2C. [H+][0H-] = 10-7 (mole/liter)2

[H[H20] ]D. 10-14 (mole/liter)

5. The molar concentration of hydrogen ions in a solu-tion was found to be 10-8 M. What is the pH of thesolution?

A. 5 B. 9 C. 5 D. 10-8

6. A concentrated solution of a strong base has a pHclosest to:

A. 3 B. 5 C. 8 D. 10

7. When a strong acid is added to water,

A. the hydrogen ion concentration increases.B. the ion product increases.C. the hydroxide ion concentration increases.D. the pH value increases.

8. The two chemical forms of a certain indicator can berepresented as follows:

Hln H+ + In-red yellow

The equilibrium constant (K ;,,d) is 10-8 mole/liter.

A. If a few drops of the indicator are added to asolution containing a strong base, what color willthe solution become?

yellow

B. At what pH value will the indicator exhibit a colorthat is a combination of red and yellow?

pH = 8

9. A test tube containing 2 x 10-2 M formic acid(CHOOH) with two drops of an indicator matchedthe color of a 2 x 10-3 M HCI solution with the sameindicator. The ionization constant of the formic acidaccording to this test is:

A. 8 x 10-8 mole/liter C. 4 x 10-8 mole/liter2 x 10-4 mole/liter D. 1 x 10-1 mole/liter

Electrochemistry

Electrical energy is a form of energy transfer, as

is heat. An electrochemical cell is a device forchanging chemical energy into electrical energy.

56 BEST COPY AVAILABLE

The electrical energy is not stored in the cell, butinstead flows as electricity and is converted intoother forms of energy, such as thermal energy(in an electric heater) or mechanical energy (inan electric motor).

63

MINIEXPERIMENTP-39 OXIDATION-REDUCTION REACTIONS

The purpose of this miniexperiment is to introduce theconcept of an oxidation-reduction reaction, and to pre-pare the student for the construction of an electro-chemical cell.

Concepts

Oxidation and reduction always occur together.Electrons move from the oxidized substance to thereduced substance.

Objectives

Decide which of two possible oxidation-reductionreactions involving two metals and ions of the twometals actually takes place, given pieces of the twometals and solutions containing the metal ions.Write a balanced equation for the reaction that occurs.

Estimated Time One-half period. If possible, experi-ment P-40 should be run in the same period so as tomake use of the same solutions.


Materials

750 cm3 0.5 M copper(II) sulfate (125 g CuSO4 5H20in 1000 cm3)

750 cm3 0.5 M zinc sulfate (144 g ZnSO4- 7H20 in1000 cm3)

15 copper strips, approx. 2 cm x 8 cm15 zinc strips, approx. 2 cm x 8 cm30 150-cm3 beakers

Laboratory Tips Have the students leave the metalstrips in the solutions long enough so that a red coppercolor is evident on the zinc strips. When the coatingis thick enough to appear red, it is easier to rub it offthe zinc strips.

Postlab Discussion The reaction that is observed is

Zn(solid) + Cu2+ ---* Zn2+ + Cu(solid)

Zinc is oxidized, and copper(II) ions are reduced tocopper metal, which is seen as a coating on the zincstrip.

EXPERIMENTP-40 COPPER-ZINC CELL

The purpose of this experiment is to prepare a simpleelectrochemical cell.

Concepts

An electrochemical cell will operate only if the reac-tants of an oxidation-reduction reaction are present.An electrochemical cell will operate only if electronsare able to move through the wire from the oxidizedsubstance to the reduced substance.The cell changes chemical energy into electricalenergy.

Objectives

Construct an electrochemical cell.Determine the conditions under which the cell willoperate.

Write the half-reactions and the overall oxidation-reduction reaction taking place in the cell.



Materials

You will need the following in addition to the materialsprepared for miniexperiment P-39:30 cloth strips, approx. 4 cm x 10 cm (see Advance

Preparation)15 lengths insulated hookup wire, approx. 30 cm long30 small alligator clips15 type PR-2 or PR-4 flashlight bulbs

Advance Preparation The cloth strips may be cutfrom cotton dish toweling. Each strip should be slightlylarger in width and length than the copper and zinc strips.

Attach an alligator clip to each end of each hookup wire.

Prelab Discussion Explain to the students that thepurpose of the cloth strips is to provide a convenientway to allow a solution to contact a metal electrode.

Range of Results The arrangement CuZn2+Zngives no electrical activity, whereas the arrangementCuCu2+Zn lights the flashlight bulb.

64 57

Post lab Discussion The experiment is performedwithout meters, so there is no need to mention eithervolts or electrode potentials. The emphasis should beplaced on the fact that electrical energy is producedwhen electrons flow through the wire as a result of theoxidation-reduction reaction.

The results obtained by the students show that copper(II)ions are necessary for electrical activity in the cell. Pointout that the copper(II) ion is a reactant in the oxidation-reduction reaction observed in miniexperiment P-39,and so is needed for the reaction. The other reactantis provided by the zinc strip. The zinc ion is not a reac-tant for this reaction.

In the arrangement CuCu2+Zn, the reaction is oc-curring in two different ways. Copper is being producedat the copper electrode, as a result of the flow ofelectricity. More copper is being produced at the zincelectrode, just as in the preceding miniexperiment; thisprocess has nothing to do with the electrical activity,but simply causes the cell to run down faster.

Miniexperiment You may wish to have the studentsuse a voltmeter to demonstrate the emf of the cell (about1.1 volts). Two or more of the copper-zinc cells maybe combined in a battery, which will give a brighter light.Simply build up layers of cells, one on top of another.Be sure that the only metal pieces that touch withinthe battery are the copper and zinc strips of adjacentcells. A battery of about nine cells may be used tooperate a transistor radio.

Miniexperiment You can fire a flashbulb with a chem-ical reaction. Prepare two 30-cm wires with alligatorclips on the ends. Clip one end of the leads to eachof the wires at the base of an AG 1 flashbulb. Attacha strip of copper metal (2 cm x 8 cm) and a 50-cm stripof Mg ribbon folded in a bunch to the opposite ends ofthe leads. Hold the setup by the connecting wires anddip the two metals into a dilute H2SO4 solution. The bulbshould flash. In this cell, magnesium is oxidized andhydrogen ions are reduced. The film of hydrogen gas atthe surface of the magnesium ribbon presumably actsas a barrier to the direct transfer of electrons.

An excellent series of experiments with a copper-magnesium cell is described by Joseph E. Davis, Jr.,Carl Berger, and Luke E. Steiner, in Chemistry, Vol. 42[April 1969], pp. 26-28. This article is reprinted in HaroldW. Ferguson and Joseph Schmuck ler, eds, Lab Bench

58

Experiments in Chemistry (Washington, D.C.: AmericanChemical Society, 1970), pp. 129-131.

The accompanying table of electrode potentialscan be used to predict half-reactions and cell po-tentials. Each electrode potential is the voltageof a cell in which the given metal electrode isplaced in contact with a 1 M solution of the givenmetal cation and is combined with a particularreference half-cell (the standard hydrogen elec-trode). The more positive the value of the elec-trode potential, the more readily the cation isreduced to the metal. You will notice that themetals at the top of the table (positive values) aremost readily reduced or least readily oxidized.Arranged in this order, the series of metals iscalled the electromotive series.

When two half-cells are combined into an elec-trochemical cell, the electrode that has the morepositive (or less negative) electrode potential is theone where reduction most readily occurs; thiselectrode is therefore the cathode. The other elec-trode is the anode. If both half-cells contain 1 Msolutions of the respective metal ions, the cellpotential (or voltage) between the electrodes isthe difference between the two electrode po-tentials. If the concentrations are not 1 M, thecell potential may be slightly different. For exam-ple, for the copper-zinc cell the copper electrodeis the cathode because it has the more positiveelectrode potential, and the calculated cell po-tential is 0.34 volt ( -0.76 volt) = 1.10 volt.

Electrode Potentials (standard reduction potentials)

Electrodepotential

Metal Cation (volts)

Au Au3+ 1.50Ag Ag+ 0.80Hg Hg22+ 0.79Cu Cu+ 0.52Cu Cu 2+ 0.34Pb Pb2+ 0.13Sn Sn2+ 0.14Ni Ni2+ 0.25Co Co2+ 0.28Cd Cd 2+ 0.40Fe Fe 2+ 0.44Zn Zn 2+ 0.76Al Al3+ 1.66Mg Mg 2+ 2.37

65

EXPERIMENTP-41 LEAD-ACID CELL

The purpose of this experiment is to construct the typeof electrochemical cell used in lead storage batteries.

Concepts

It is necessary to coat one electrode of a lead-acidcell with lead dioxide before it will operate.In the lead-acid cell, lead is oxidized to lead(II) ionsand lead dioxide is reduced to lead(II) ions.The half-reactions of the lead-acid cell can be re-versed by charging the cell.

Objectives

Construct a lead-acid cell.Identify the changes taking place at each electrode,given the cell reaction.



Materials

30 lead strips, approx. 2 cm x 10 cm1500 cm3 4 M sulfuric acid (see Advance Preparation)15 150-cm3 beakers30 dry cells, No. 6 (1.5 volts)30 lengths insulated hookup wire, approx. 30 cm long,

with alligator clips at each end15 lengths insulated hookup wire without clips15 flashlight bulbs, type PR-2 or PR-4

Advance Preparation To prepare 1000 cm3 of the4 M sulfuric acid, slowly add 222 cm3 concentratedFI2S0, while stirring to 825 cm3 water in a bottle. Keepthe bottle in a sink or a bucket while adding the acidin case of breakage due to the large increase in tem-perature. Allow to cool.

Laboratory Safety Review with your students thesafety precautions necessary when working with acidssuch as sulfuric acid. You many wish to review a fewpoints listed in Appendix Safety in the student module.

Laboratory Tips A d.c. power supply can be used inplace of the dry cells. In this case, keep the voltagebelow 4 volts to prevent splattering of the sulfuric

acid. It is also advisable to cover the top of the beakerwith a watch glass.

Range of Results There is no electrical activity be-fore the lead dioxide is formed.

During the period that the batteries are attached to thecell, a brown coating of lead dioxide forms on the leadstrip that is connected to the positive terminal. No changeis seen in the appearance of the other electrode.

When the test bulb is connected again to the chargedcell, a brief flash of light is seen.

Postlab Discussion The oxidation reaction is

Pb 0 P132+ + 2e-

The reduction reaction is

Pb02 + 4H+ + 2e- PI32+ + 2H20

It is interesting to notice that the lead(II) ion is theproduct of both half-reactions.

You may wish to have one of your studentsdemonstrate the use of a storage battery testerto the rest of the class. If you use an actual storagebattery, caution the students about the corrosiveaction of the acid in the battery.

P-42 PRACTICAL PORTABLE POWER

A useful description of commercial cells and bat-teries, fuel cells, and electroplating may be foundin G. R. Palin, Electrochemistry for Technologists(Oxford: Pergamon Press, 1969; paperback).

An interesting demonstration of a hydrogen-oxygen fuel cell is described on page 89 of thefollowing book: R. M. Lawrence and W. H.Bowman, Space Resources for Teachers: Chemistry(NASA, 1971; available from U.S. GovernmentPrinting Office, Washington, D.C. 20402; orderNASA EP-87).

EXPERIMENTP-43 FINISHING TOUCHES

The purpose of this experiment is to demonstrateelectroplating.

Concepts

In electroplating, electrical energy is used to depositmetal at the cathode.

6.659

Objectives F. Castka, A Sourcebook for the Physical Sciences(New York: Harcourt, Brace and World, 1961).

Construct an electroplating setup.Identify and write the half-reaction taking place ateach electrode.



Materials

15 zinc strips, approx. 2 cm x 8 cm15 copper strips, approx. 2 cm x 8 cm1000 cm3 0.5 M zinc sulfate (144 g ZnSO4. 7H20 in

1000 cm3)15 dry cells, No. 6 (1.5 volts)30 lengths insulated hookup wire, approx. 30 cm long,

with alligator clips at each end15 150-cm3 beakers

Laboratory Tips The zinc coating sticks best to thecopper strip if the strip is first dipped into concentratedhydrochloric acid and rinsed.

The negative (side) terminal of the dry cell must be con-nected to the copper strip. Other metal objects may besubstituted for the copper strip.

Range of Results A light gray coating of zinc formson the copper strip.

Postlab Discussion The half-reactions are

Zn2+ + 2e- --0 Zn (at the copper electrode)Zn--0 Zn2+ + 2e- (at the zinc electrode)

There is no net change in the solution. Zn2+ ions aresupplied from the zinc strip as rapidly as they are de-posited on the cathode. The only changes occur at theelectrodes; the cathode gains a film of Zn and the zincanode loses weight. The process has no tendency tooccur unless the two electrodes are made part of anelectrical circuit in which electrons are forced to moveby an electrochemical cell or other source of electricalenergy.

Miniexperiment Electroplating: A variety of experi-ments in the electroplating of copper, nickel, silver, andchromium are described in Alexander Joseph, Paul F.Brandwein, Evelyn Morholt, Harvey Pollack, and Joseph

60

Some of your students may wish to collect itemsor pictures of items that have been electroplated.Ask them to explain why the electroplating pro-cess is useful in each of these cases. Can theyrecognize items that are plated? In what waysdoes a plated item appear to differ from an un-plated item? What are the disadvantages and/oradvantages of plated materials?



1. (a) Cu- o Cu2+ + 2e-Ag+ + e- --> Ag

(b) The silver electrode is the cathode and the copperelectrode is the anode.

(c) The silver electrode is positive and the copperelectrode is negative.

(d) Cu Cu2+ + 2e-2Ag+ + 2e- --o 2 Ag

sum: Cu + 2Ag+ 0 Cu2+ + 2Ag

(e) Cu in solution of Ag+ ions: the overall chemicalreaction shown in part (d) will proceed; Cu willdissolve and Ag will form.

2. (a)

Ag in solution of Cu2+ ions: no reaction.

Pb --oPb2+ + 2e-Pb02 + 4H+ + 2e --0 Pb2+ + 2H20

sum: Pb + Pb02 + 4H+---o 2Pb2+ + 2H20

Pb + Pb02 + 4H+ --o 2Pb2+ + 2H202Pb2+ + 25042- . 2PbSO4

sum: Pb + Pb02 + 4H+ + 2S0,2---o 2PbSO., + 2H20

(b) 2PbSO4 + 2H20Pb + Pb02 + 4H+ + 2S0,2-

As the battery is charged, water is consumed andsulfuric acid is produced. Since sulfuric acid ismore dense than water, the solution's densityrises.

67

EVALUATION ITEMS


1. Briefly explain the difference between the terms oxida-tion and reduction.

Oxidation is the process whereby electrons are lost.Reduction is the process whereby electrons aregained.

2. The half-reaction Cu2+ + 2e- 0 Cu occurs at

which electrode? cathode

3. Which of the following is an oxidation half-reaction?

A. Zn Zn 2+ + 2e-B. 2H+ + 2e- H2

C. Pb02 + 4H+ + 2e- Pb2+ + 2H20D. Pb2+ + S042- --0 PbSO4

4. Which of the following occurs at the anode?

A. anions are formedB. reductionC. electrons enter the anode from the solutionD. pure metals precipitate

5. Why is the density of sulfuric acid a good indicatorof the electrical energy available from a lead storagebattery?

As the battery is discharged, it consumes H2SO4,and thus the density decreases. The more H2SO4remaining, the greater the amount of electricalenergy available.

6. Which of the following reactions or half-reactionsoccurs when a lead storage battery is charging?

A. Pb2+ + S042- PbSO,B. Pb --0 Pb2+ + 2e-C. Pb02 + 4H+ + 2e- p Pb2+ + 2H20D. Pb2+ + 2H20 ---0 Pb02 + 4H+ + 2e-

7. Which of the following reactions occurs at the cath-ode of a common dry cell?

A. 2MnO2 + Zn2+ + 2e- ---+ ZnMn204B. Mn Mn2+ + 2e-C. 2ZnO2 + Mn2+ + 2e- --0 MnZn2O4D. Zn Zn2+ + 2e-

8. Briefly describe the process of chromium plating.

The object to be plated is connected at the cathodein a chromic acid solution. Reduction at the cathodeproduces a layer of chromium metal on the object.

Summary

Refer to the questions about the automobile at thebeginning of the student module (pages 2-3).Ask the students to review and discuss answersto these questions. Use this opportunity to sumup your discussions of physical chemistrywhatit is, what types of questions it tries to answerand its applications to everyday life. You maydiscuss and list some problems that may be solv-able by physical chemists, and encourage yourstudents to present hypothetical (or practical)solutions to the problems utilizing the knowledgeand skills acquired from this module or othersthat you have studied.

You can encourage this hypothesizing by yourstudents by discussing the photos on pages96-97. A comparison of the photographs canshow the automation of pH measurement andthe titration process that your students may haveexperienced in the laboratory situation. You maywish to discuss other processes that have beenrevolutionized over the years through advancesin chemistry. All of these arose from scientistsmaking "educated guesses" while constantlyseeking to improve upon existing knowledgeand methods. Students commonly think of a re-search chemist as a person surrounded by a mazeof contorted glassware and bubbling concoctions(left photo). But with modern, computerized

68 BEST COPY AVAILABLE61

electronic equipment (right photo), elaboratetests have been adapted for rapid and depend-able performance by machines. The results canbe quickly calculated, printed out, and com-pared to other experimental results for evaluation.But the chemist is still there. It is the chemistwhose mind is behind the machine and whostudies the data and uses them to solve questions.

You may now want to introduce the next mod-ule and relate the ideas of physical chemistry to

62

the new ideas. Certainly the students will beginto notice the obvious overlapping of material atvarious points. As you reinforce these ideasthrough continual referral, you will help moldthe concept of unity in chemistry.

The study of physical chemistry has the reputa-tion of being "difficult." We hope that this moduleshows that physical chemistry can be relevantand fun.

69

AppendixSafety

SAFETY IN THE LABORATORY

Proper conduct in a chemistry laboratory is really an extensionof safety procedures normally followed each day around yourhome and in the outside world. Exercising care in a laboratorydemands the same caution you apply to driving a car, riding amotorbike or bicycle, or participating in a sport. Athletes con-sider safety measures a part of playing the game. For example,football players willingly spend a great deal of time putting onequipment such as helmets, hip pads, and shoulder pads toprotect themselves from potential injury.

Chemists must also be properly dressed. To protect them-selves in the laboratory, they commonly wear a lab apron or acoat and protective glasses. Throughout this course you willuse similar items. Hopefully their use will become secondnature to you, much as it becomes second nature for a base-ball catcher to put on a chest protector and mask beforestepping behind home plate.

As you read through a written experimental procedure, youwill notice that specific hazards and precautions are called toyour attention. Be prepared to discuss these hazards with yourteacher and with your fellow students. Always read the entireexperimental procedure thoroughly before starting anylaboratory work.

A list of general laboratory safety procedures follows. It isnot intended that you memorize these safety procedures butrather that you use them regularly when performing experi-ments. You may notice that this list is by no means complete.Your teacher may wish to add safety guidelines that arerelevant to your specific classroom situation. It would beimpossible to anticipate every hazardous situation that mightarise in the chemistry laboratory. However, if you are familiarwith these general laboratory safety procedures and if youuse common sense, you will be able to handle potentiallyhazardous situations intelligently and safely. Treat all chem-icals with respect, not fear.

GENERAL SAFETY GUIDELINES

1. Work in the laboratory only when the teacher is present orwhen you have been given permission to do so. In case ofaccident, notify your teacher immediately.

2. Before starting any laboratory exercise, be sure that thelaboratory bench is clean.

3. Put on a laboratory coat or apron and protective glassesor goggles before beginning an experiment.

4. Tie back loose hair to prevent the possibility of its con-tacting any Bunsen burner flames.

5. Open sandals or bare feet are not permitted in the labo-ratory. The dangers of broken glass and corrosive liquidspills are always present in a laboratory.

6. Fire is a special hazard in the laboratory because manychemicals are flammable. Learn how to use the fireblanket, fire extinguisher, and shower (if your laboratoryhas one).

7. For minor skin burns, immediately immerse the burnedarea in cold water for several minutes. Then consult yourteacher for further instructions on possible additionaltreatment.

8. In case of a chemical splash on your skin, immediatelyrinse the area with cold water for at least one minute.Consult your teacher for further action.

9. If any liquid material splashes into your eye, wash theeye immediately with water from an eyewash bottle oreyewash fountain.

10. Never look directly down into a test tubeview the con-tents of the tube from the side. (Why?)

11. Never smell a material by placing your nose directly atthe mouth of the tube or flask. Instead, with your hand,"fan" some of the vapor from the container toward yournose. Inhale cautiously.

12. Never taste any material in the laboratory.

13. Never add water to concentrated acid solutions. The heatgenerated may cause spattering. Instead, as you stir, addthe acid slowly to the water or dilute solution.

14. Read the label on a chemical bottle at least twice beforeremoving a sample. H202 is not the same as H2O.

15. Follow your teacher's instructions or laboratory proce-dure when disposing of used chemicals.

This symbol represents three of the common hazards in a chemistrylaboratoryflame, fumes, and explosion. It will appear with certainexperiments in this module to alert you to special precautions in addition tothose discussed in this Appendix.

7Q 63

Metric Units

PHYSICALQUANTITY

SI BASE OR DERIVED UNIT OTHER UNITS

NAMESYMBOL AND

DEFINITIONNAME

SYMBOL ANDDEFINITION

length meter* m kilometercentimeternanometer

1 km = 103 m1 cm = 10-2 m1 nm = 10-9 m = 10-7 cm

area square meter m2 squarecentimeter 1 cm2 = 10-4 m2

volume cubic meter m3 cubiccentimeter

liter1 cm3 = 10-6 m3

1 I = 103 cm3

mass kilogram* kg gram 1 g = 10-3 kg

time

amount ofsubstance

concentration

second*

mole*

moles per

s

mol

cubic meter mol/m3 moles per liter

molarconcentration

1 mol /1 = 103 mol/m3

1 M = mol /1

(molarity)

Celsiustemperature

thermodynamictemperature kelvin* K

degree Celsius °C

pressure pascal Pa = kg/m s2 centimeterof mercury

atmosphere1 cm Hg = 1.333 x 103 Pa

1 atm = 1.013 x 105 Pa1 atm = 76.0 cm Hg

energy joule J = kg m2/s2 calorie 1 cal = 4.184 J

'SI base unit, exactly defined in terms of certain physical measurements.

7164

Suggested Readings and Films

BOOKS

Boeke, Kees. Cosmic View: The Universe in 40 lumps.New York: The John Day Company, 1957. An in-genious book showing the universe at differentscales, with objects ranging in size from galaxies toan atomic nucleus.

Gosselin, R. E.; Hodge, H. C.; Smith, R. P.; andGleason, M. N. Clinical Toxicology of CommercialProducts, 4th ed. Baltimore: Williams & Wilkins, 1976.Lists the ingredients in more than 17 500 commoncommercial products.

Holden, Alan, and Singer, P. Crystals and Crystal Grow-ing. Garden City, N.Y.: Anchor Books, 1960, paper-back. Suggestions for experiments in growingcrystals.

Kirk-Othmer Encyclopedia of Chemical Technology, 2nd ed.New York: John Wiley & Sons, 1963. This encyclo-pedia, in 22 volumes, is a useful source of informa-tion on industrial applications of chemistry at a fairlydetailed but not overly technical level. Some typicalarticles are: "Gasoline and Other Motor Fuels";"Batteries and Electric Cells"; "Emulsions"; "Milk

and Milk Products." The third edition, projected for25 volumes, is being published between 1977 and1983.

Vold, Marjorie J., and Vold, Robert D. Colloid Chem-istry: The Science of Large Molecules, Small Particles,and Surfaces. New York: Van Nostrand Reinhold,1964, paperback. An excellent overall view of thefield of colloid science.

FILMS

Common Colloids. A Science Close-ups Film. PrismProductions, Inc. Color, 7 minutes.

Molecular Motions. A CHEM Study Film. Rochester,New York: Modern Learning Aids, a div. of Ward'sNatural Science Establishment. Color, 13 minutes.Translational, rotational, and vibrational motions ofmolecules are illustrated with mechanical models andanimation.

Powers of Ten. Charles Eames Associates. Color, 8minutes. A mind-boggling view of the universe atdifferent orders of magnitude.

Module Tests

Two module tests follow: one to test knowledge-centered objectives and the other to test skill-centeredobjectives. If you choose to use either or both of thesemodule tests as they are presented here, duplicatecopies for your students. Or, you may wish to selectsome of the questions from these tests that you feelapply to your introductory chemistry course and addadditional questions of your own. Either way, makesure the test that you give reflects your emphasis onthe chemistry you and your students experienced inthis module. The skill-centered test will require thatyou set up several laboratory stations containing

materials for your students to examine or work with.You may wish to add additional test items to roundout the types of skills you and your students haveworked on. (Answers to the test questions in this sec-tion are provided.) If you wish to use a standard-typeanswer sheet for this test, one is provided in the ap-pendix of Reactions and Reason Teacher's Guide.

ANSWERS FOR THE KNOWLEDGE-CENTEREDMODULE TEST

1. A; 2. D; 3. A; 4. B; 5. C; 6a. C; 6b. D; 7. D; 8. D; 9. A;10. B; 11. B; 12. C; 13. A; 14. D; 15. B; 16. B; 17. B;18. C; 19. D; 20a. A; 20b. D; 21. C; 22. A; 23. B; 24a. D;24b. A.

72 65

COMMUNITIES OF MOLECULESKnowledge-Centered Module Test

1. Bubbles rise in a newly opened soft drink bottle.An explanation for this is that carbon dioxide is lesssoluble in water when theA. pressure is decreased.B. temperature is decreased.C. cap is on the bottle.D. pressure is increased.

2. A micelle is a colloid that formsA. a network crystal structure.B. an ionic crystal structure.C. a molecular crystal structure.D. spherical particles with a structure different from

the individual molecules.

3. An example of a chemical equilibrium is given bythe reaction:

2Cr042- + 2H+ Cr2022- + H2O(yellow) (orange)

If more H+ were added to the solution, one wouldexpectA. the solution's color to become more orange.B. the solution to change to a colloid.C. the solution's color to remain unchanged.D. the solution's color to become more yellow.

4. 1.00 mole of A and 1.00 mole of B are placed in aclosed container. A and B react to form C and D,as in the equation: A + B C + D. When equilib-rium is reached, there are 0.800 mole each of Aand B and 0.200 mole each of C and D in the con-tainer. The equilibrium constant for the reaction isA. 0.0040 C. 16.0B. 0.0625 D. 25.0

5. CaCl2 is an example of a (an)A. molecular crystal.B. network crystal.C. ionic crystal.D. metallic crystal.

6. The capillary tube in the following diagram containsa sample of trapped air below an 8.0 cm slug of Hg.The atmospheric pressure is 76.0 cm Hg and thetemperature is 21°C.

66

open end

9.0 cm

8.0 cm

12.0 cm

closed end

(a) The pressure of the gas in the closed end ofthe tube isA. 64.0 cm Hg. C. 84.0 cm Hg.B. 68.0 cm Hg. D. 88.0 cm Hg.

(b) If the capillary tube were placed in an ice watermixture with its open end up, the length of gasin the closed end would beA. approximately 0 cm.B. 12.0 cm.C. greater than 12.0 cm.D. less than 12.0 cm.

7. When 0.100 mol of KI is dissolved in 100 gramsof water, the temperature of the water changes from25°C to 19.9°C. The molar heat of solution of KI isA. 51.0 calories.B. 100 calories.C. 2500 calories.D. 5100 calories.

8. An example of a colloidal dispersion isA. an emulsion.B. a sol.C. smoke.D. all of the above.

9. In the solid state, P4 crystals are held together byLondon forces. One would expectA. KCI to melt at a higher temperature than P4.B. P4 to melt at a higher temperature than KCI.C. KCI and P4 to melt at approximately the same

temperature.D. P4 to be more soluble in water than KCI.

73

10. Three substances were thought to be a solute insolution, a colloid, and a precipitate. SubstanceA retained by filtration; however, substance Bpassed through the filter but was retained bydialysis tubing. Substance C was retained by neitherfiltration nor dialysis tubing. Substance C is mostlikely to be aA. colloid. C. precipitate.B. solute in solution. D. all of these

11. The boiling point of Hg at a pressure of 76 cm Hgis 357°C. If the pressure is raised to 80 cm Hg,the boiling point of Hg would beA. decreased. C. unchanged.B. increased. D. cannot tell from data provided

12. One major difference between fuel cells and otherelectrochemical cells is thatA. fuel cell reactions usually occur at very high

temperatures.B. fuel cells require zinc electrodes.C. fuel cells are continuously supplied with reac-

tants, and they never run down.D. all of the above

13. When the following electrochemical cell is in opera-tion the strip of Mn is slowly consumed.

porous barrier

The overall cell reaction isA. Mn + Fe2+ Mn2+ + FeB. Mn2+ + Fe2+ Mn2+ + FeC. Fe + Mn2+ Fe24- + Mn

D. Fe + Mn Fe2+ + Mn2+

Mn

Mn2+

14. Consider the reaction: A + B C + D. The rateof this reaction could be increased byA. adding a catalyst.B. increasing the concentration of A.C. increasing the concentration of B.D all of the above

15. A large amount of solid KCI is shaken in a test tubewith water until no further change occurs, and theremaining solid settles to the bottom. From thistime onA. only dissolving takes place.B. dissolving and crystallization occur at equal

rates.C. no further changes take place.D. only crystallization takes place.

16. The equilibrium constant for the reaction shown be-low is a very large number.

A+B--0.0

If a reaction mixture of equal amounts of A and Bis allowed to reach equilibrium, which statementconcerning this equilibrium would be true?A. There is no C formed.B. There is much more C formed than either A or

B remaining.C. There are equal amounts of all three materials

remaining.D. There is much more A + B remaining than C

formed.

17. Antifreeze is mixed with water in a car's radiatorbecause itA. is immiscible with water.B. lowers the freezing point of water.C. raises the freezing point of water.D. lowers the boiling point of water.

18. The equilibrium constant for a chemical reaction is

K [A][B]2[AB2]

The reaction that best illustrates this equilibriumequation isA. AB2 = A2 + B2B. AB2 2A + 2BC. AB2,---= A + 2BD. AB2 A + B2

19. When a drop containing 1 x 10-4 cm3 of stearicacid is allowed to spread over the surface of a trayof water, a circular film of 10 cm in radius is formed.Assuming the layer to be one molecule thick, thelength of a stearic acid molecule is aboutA. 1 x 10-8 cm C. 1 x 10-7 cmB. 3 x 10 -8 cm D. 3 x 10 -' cm

7467

20. An electrochemical cell has the following reaction:

Ca + Zn2+ --*Ca2++ Zn

(a) The cell's anode isA. Ca B. Ca2+ C. Zn

(b) The material being reduced isA. Ca B. Zn C. Ca2÷

21. The density of the sulfuric acid solution inbattery is a good indication of whenneeds recharging becauseA. as the battery runs down, the

increases.as the battery runs down, theremains unchanged.as the battery runs down, thedecreases.as the battery runs down, theat the anode.

B.

C.

D.

a

the

D. Zn2+

D. Zn 2+

storagebattery

solution's density

solution's density

solution's density

solution collects

22. In order to double the pressure of 50 cm3 of a gasat constant temperature, you would change thevolume toA. 25 cm3. C. 75 cm3.B. 50 cm3. D. 100 cm3.

68

23. A marshmallow is an example of a (an)A. sol. C. micelle.B. foam. D. aerosol.

24. In the following structural formulas:

H2O

0/ \H H

F2 NH3

N/INFF H H

H

CH4

H

HCH

(a) The polar molecule(s) is (are)A. F2 C. H2O and F2B. H2O and CH, D. H2O and NH3

(b) The molecule(s) capable of forming hydrogenbonds is (are)A. H2O and NH3 C. H2O, NH3, and CH4B. CH4 D. H2O and CH4

Skill-Centered Module Test

If you decide to use these skill-centered test items, youwill need to make certain advance preparations. Thenumerals in the following list indicate the items forwhich you will have to prepare special laboratory sta-tions. Be sure to test each of the lab stations beforeallowing students to determine the answers to the skill-centered items. When students are ready to answerthese questions, they should go to the numbered sta-tion and follow the directions that are given there andin the printed question item. When they finish with thematerials at the station, instruct them to leave the ma-terials in proper order for the next student.

1. Provide a beaker of water, one medicine dropper,and a 10-cm3 graduated cylinder. Supply papertowels for drying the graduated cylinder.

3. Supply a plastic cm ruler and a piece of glass tub-ing. Be sure the tubing isn't longer than the ruler.

4. Provide the following labeled pieces of equipment(a) Erlenmeyer flask (c) evaporating dish(b) suction flask (d) Florence flask

5. Provide a spot plate and a dropper bottle of bluewater-soluble dye and a dropper bottle of red oil-soluble dye.

Methylene blue is a good water-soluble dye, andSudan III or IV is a good oil-soluble dye.

The emulsion can be a sample of hair conditioneror hand cleaner.

Supply paper towels to dry the spot plate aftercleaning. Be sure to test the emulsion yourself.

6. Provide a test tube marked to the level where 2 gNa2S203 5H20 would fill the tube. Also providea 10-cm3 graduated cylinder, -a beaker of water,

76

and a spatula. A -10°C to +110°C thermometershould be available. Provide the Na2S203 51-120 ina beaker.

7. Place a 0.1 M K2Cra4 solution (19.4 g/1000 cm3water) in a test tube marked #1. Half fill the tube.

Place a 0.05 M K2Cr20, solution (14.7 g/1000cm3) in a test tube marked #2. Half fill the tube.

Provide a solution (X) that is a 50/50 mixture of the2 chromate solutions. Put the solution in a beakerlabeled X.

Provide a test tube with a mark at the 10-cm3level, also labeled X.

Place a 4 M NaOH solution (16 g NaOH /100 cm3H2O) in a beaker with a medicine dropper. Labelthe beaker base.

8. Provide small pieces of copper strips or lengths ofcopper wire in a jar. Clean the oxide off the stripsin advance.

Place 50 cm3 of each solution in an appro-priately labeled beaker.

Solution A is 0.1 M CuSO4 51-120 (1.3 g/50 cm3).Solution B is 0.1 M Pb(NO3)2 (1.65 g/50 cm3).Solution C is 0.1 M Zn(NO3)2 (1.5 g/50 cm3).Solution D is 0.1 M AgNO3 (0.8 g/50 cm3).

10. Provide a cm ruler.

ANSWERS FOR THE SKILL-CENTEREDMODULE TEST

1. *; 2. A; 3. *; 4. B; 5. *; 6. B; 7. A; 8. D; 9. C; 10. C

*evaluate according to teacher standards

69

COMMUNITIES OF MOLECULESSkill-Centered Module Test

Several questions in this section require you to makeobservations and perform chemical manipulations. Thestations where you will do these operations will be in-dicated by your teacher. If the station you are going tois being used, continue with the test and go back later.

1. Go to station #1. Using the 10-cm3 graduatedcylinder and dropper provided, determine the vol-ume of 1 drop of water. Record the volume at thebottom of your answer sheet next to #1.

2. The data below were obtained from an experimentusing the air-mercury tube.

Atmospheric Pressure(cm Hg)

Length of Column(cm)

71 2.573 2.375 2.1

77 2.079 1.9

A graph of the data would resemble which of thefollowing?

A. (a) B. (b)

a)

a-

a)

70

Length

Length

a-

C. (c) D. (d)

Length

3. Go to station #3. Using the ruler provided, measurethe length of the object provided. Record youranswer at the bottom of your answer sheet nextto #3.

4. Go to station #4. Examine the laboratory equip-ment provided. Which piece of laboratory equip-ment is normally used in filtration?A. (a) B. (b) C. (c) D. (d)

5. Go to station #5 and place a drop of the emulsionprovided into two different wells on the spot plate.To one well, add a drop of blue, water-soluble dye;to the second well, add a drop of red, oil-solubledye.

The emulsion would be classified asA. water in oil. C. bothB. oil in water. D. neither

CLEAN THE SPOT PLATE BEFORE LEAVING.

6. Go to station #6 and fill the test tube to the markwith the salt provided. Then add 5 cm3 water,stirring with the thermometer.

During the solution process the temperatureA. increased.B. decreased.C. remained constant.D. decreased, then increased.

7. Go to station #7, where test tube #1 contains asolution of Cr042- ions, while test tube #2 containsa solution of Cr2072- ions. Fill a third test tube tothe mark with Solution X; it contains an equal mix-ture of the two ions. Add a few drops of base totube X and .determine to which ionic form the equilib-rium has been shifted. The equilibrium has shifted to:A. Cr042-B. Cr2072-C. has not shiftedD. shifted first to Cr042- and then back to Cr2072-

RINSE OUT TEST TUBE X BEFORE LEAVING.

77

8. Go to station #8. Using the piece of metal pro-vided, place it in each of the lettered solutions,one at a time. Observe in which beakers a reac-tion has occurred.

(Caution: The metal must be cleaned [rinsed withwater] before placing it in a second beaker. if themetal changes, discard it and continue with anunused piece.)

A reaction was visible in:A. Solution A C. Solution CB. Solution B D. Solution D

9. Arrange the steps below in the proper order to de-termine the molar heat of solution of a salt.a. Add salt to water and stir well to dissolve the salt.b. Calculate the molar heat of solution.c. Determine the change in temperature.d. Measure the temperature of water.e. Record the temperature of the solution.f. Weigh the salt to be dissolved and calculate

number of moles of salt.

The correct order is:A. f, d, e, c, a, bB. f, d, e, a, c, b

C. f, d, a, e, c, bD. f, a, d, e, c, b

10. Using a centimeter ruler, determine the approximatearea of the circle sketched below:

The area is approximately:A. 33 cm2 B. 11 cm2 C. 8 cm2 D. 44 cm2

78 71

Materials List

Quantities are for a class of 30 students working inpairs.

*Optional Items. These items depend on teacherchoice. We have listed substitutions in the experimentdiscussion. Consult the specific experiment in theteacher's guide to determine use and quantities.

NONEXPENDABLE MATERIALS

Item Experiment Amount

Alligator clips 40, 41, 43 30

Aspirators 13 15

Balances, 0.01 g sensitivity 8 8

Beakers, 50-cm3 19 45

Beakers, 150-cm3 19, 22, 23, 30,35, 39, 40, 41, 43 30

Beakers, 400- or 600-cm3 7, 16, 22 15

Bunsen burners 7*, 16, 22 15

Capillary tubing, 0.75 mm-1.25 mm bore 5, 7 5 m

Clamps, universal 13, 16 15

Clock, for timing 30 1

Copper strips, 2 cm x 8 cm 39, 40, 43 15 strips

Dropper bottles, 30- to 60-cm3 2, 22, 30, 38 16

Dry cells, No. 6, 1.5 volts 41, 43 30

Erlenmeyer flasks, 125- or 250-cm3 38 15

File, triangular 5, 7 1

Filter flasks, 500-cm3 or smaller 13 15

Flashlight bulbs, PR-2 or PR-4 40, 41 15

Funnels, 75-mm diameter, with supports 15, 19 15-45

Graduated cylinders, 10-cm3 2, 19, 22, 23,30, 35, 38 15

Graduated cylinders, 50-cm3 22, 23, 28, 30 15

Graduated cylinders, 100-cm3 25, 27, 35, 38 15

Lead strips, 2 cm x 8 cm 41 30

Medicine droppers 2, 22, 38 30

Meter sticks 2 15

Ring stands and rings 7, 13, 16, 19, 22, 23 15

Rubber stoppers, 1-hole, to fit filter flask 13 15

Rubber tubing, thick-walled, 9.5-mm (3/8") I.D. 13 15 lengths

Rulers, metric, 15-cm, heavy plastic 5, 7 15

Spatulas 27 15

Spot plates or watch glasses 22 30

Stirring rods, glass 19, 22, 27, 28, 30, 35 30

Syringe, hypodermic and needle, 5-cm3 5 1

Test tubes, 18 x 150 mm 16, 19, 38 210

Test-tube racks, 14-tube capacity 19, 38 15

Thermometers, -10°C to 110°C 7, 13, 22, 25,27, 28, 30 15

72 79

NONEXPENDABLE MATERIALS (cont.)

Item

Trays, 60 x 60 x 5 cm (approx.)Wash bottlesWire, insulatedWire gauze, asbestos centersZinc strips, 2 cm x 8 cm

Experiment

2

22*40, 41, 437*, 16, 22,39, 40, 43

23

Amount

15

15*

45 lengths15

15

EXPENDABLE MATERIALS

Item

Acetic acid, glacialAmmonium peroxydisulfate (ammonium persulfate), (NH4)3S208Boiling chipsCalcium acetateCopper(II) sulfate, pentahydrateCotton cloth (dish toweling)Cups, StyrofoamDialysis tubing, 1.6-cm diameterEpoxy cementEthanol, 95 percent (ethyl alcohol)

Experiment

19, 383013

2330, 39, 4040

25, 27, 2819

5, 72, 23

Filter paper, 12.5-cm diameter 19, 22Food coloring, blue 19, 22Graph paper, linear 7Hydrochloric acid, conc. 28, 38Ice 7, 25Iron(III) nitrate, Fe(NO3)3. 9H20 35Lycopodium powder or chalk dust 2Masking tape 38Mercury, metal 5, 7Methyl orange, methyl orange sodium salt, or bromocresol green

sodium salt (indicators) 38*Milk, skim or powdered 19Mineral oil 22Nitric acid, conc. 35Oleic acid 2Orange IV or thymol blue sodium salt (indicators) 38*Potassium iodide, KI 30Potassium thiocyanate, KSCN 35Sodium hydroxide 27, 28Sodium thiosulf ate, pentahydrate, Na2S203- 5H20 16, 27, 30Span 40 (sorbitan monopalmitate) 22Starch, soluble 30Sudan III or Sudan IV dye 22Sulfuric acid, conc. 41Tween 40 (polyoxyethylene sorbitan monopalmitate) 22Water, distilled 30, 38Zinc sulfate, heptahydrate, ZnSO4. 7H20 39, 40, 43

0

Amount

35 cm350 g15 chips45 g140 g1200 cm245600-cm length1 set tubes1000 cm375 sheets1 bottle30 sheets190 cm360-100 cubes5 g10 g1 roll

5 cm3

0.1 g*600 cm3750 cm32 cm31 cm30.1 g*40 g1 g150 g750 g60 g1 g5 g350 cm360 g14 000 cm3300 g

73

AcknowledgmentsIAC Test Teachers

Linwood Adams, Bowie High School, Prince George'sCounty, MD

Thomas Antonicci, Archbishop Curley High School, Balti-more, MD

Nicholas Baccala, Milford Mill High School, BaltimoreCounty, MD

Rosemary Behrens, Bethesda-Chevy Chase High School,Montgomery County, MD

Virginia Blair, Holton-Arms School, Bethesda, MDEthyl duBois, Crossland and Oxon Hill High Schools, Prince

George's County, MDSally Buckler, High Point High School, Prince George's

County, MDTherese Butler, Bowie High School, Prince George's

County, MDKevin Castner, Bowie High School, Prince George's

County, MDRobert Cooke, Kenwood High School, Baltimore County, MDWilmer Cooksey, Woodrow Wilson High School, Wash-

ington, DCFrank Cox, Parkville High School, Baltimore County, MDRichard Dexter, John F. Kennedy High School, Montgomery

County, MDElizabeth Donaldson, John F. Kennedy High School, Mont-

gomery County, MDClair Douthitt, Chief Sealth High School, Seattle, WALawrence Ferguson, Milford Mill High School, Baltimore

County, MDHarry Gemberling, Du Val and Eleanor Roosevelt High

Schools, Prince George's County, MDAlan Goldstein, Laurel High School, Prince George's

County, MDMarjorie Green, McLean High School, Fairfax County, VAWilliam Guthrie, Parkdale High School, Prince George's

County, MDLaura Hack, Annapolis High School, Annapolis, MDMargaret Henderson, Fort Hunt High School, Fairfax

County, VAMartina Howe, Bethesda-Chevy Chase High School, Mont-

gomery County, MDGlendal Jenkins, Surrattsville High School, Prince George's

County, MDMartin Johnson, Bowie High School, Prince George's

County, MDHarold Koch, Southwest High School, Minneapolis, MNJane Koran, Arundel High School, Anne Arundel County, MDMarilyn Lucas, Euclid High School, Euclid, OHDavid McElroy, Albert Einstein High School, Montgomery

County, MD

IAC 1978 Revision Teacher Consultants

Robert Andrews, Bothell High School, Bothell, Washington;Minard Bakken, The Prairie School, Racine, Wisconsin; ErvinForgy, J. I. Case High School, Racine, Wisconsin; MargaretHenley, Kennedy High School, Granada Hills, California;Bernard Hermanson, Sumner Community Schools, Sumner,Iowa; Merlin Iverson, Mason City High School, Mason City,Iowa; Harold Koch, Southwest High School, Minneapolis,Minnesota; Philippe Lemieux, Lincoln-Sudbury Regional

74

Marilu McGoldrick, Wilde Lake High School, HowardCounty, MD

John Malek, Meade High School, Ft. Meade, MDRobert Mier, Bowie and Eleanor Roosevelt High Schools,

Prince George's County, MDGeorge Milne, Oxon Hill High School, Prince George's

County, MDDavid Myers, Crossland High School, Prince George's

County, MDGeorge Newett, High Point High School, Prince George's

County, MDDaniel Noval, Patapsco High School, Baltimore County, MDM. Gail Nussbaum, Northwestern High School, Prince

George's County, MDElena Pisciotta, Parkdale High School, Prince George's

County, MDAndrew Pogan, Poolesville High School, Montgomery

County, MDCharles Raynor, Dulaney High School, Baltimore County, MDRosemary Reimer Shaw, Montgomery Blair High School,

Montgomery County, MDE. G. Rohde, Academy of the Holy Names, Silver Spring, MDDoris Sandoval, Springbrook High School, Montgomery

County, MDEarl Shaw, Damascus High School, Montgomery County, MDGeorge Smeller, Robert Peary High School, Montgomery

County, MDHoward Smith, Parkville High School, Baltimore County, MDLarry Sobotka, Parkville High School, Baltimore County, MDRoger Tatum, Takoma Academy, Takoma Park, MDYvette Thivierge, Fairmont Heights High School, Prince

George's County, MDBarbara Tracey, Bishop McNamara High School, Forest-

ville, MDRonald Trivane, Pikesville High School, Baltimore

County, MDJeanne Vaughn, Governor Thomas Johnson High School,

Frederick County, MDDrew Wolfe, Randallstown High School, Baltimore

County, MDPauline Wood, Springbrook High School, Montgomery

County, MDJames Woodward, Walt Whitman High School, Montgomery

County, MDClement Zidick, Dimond and Wasilla High Schools, Anchor-

age, AK

High School, Acton, Massachusetts; Robert Sherwood, NewPalestine High School, New Palestine, Indiana; KennethSpengler, Palatine High School, Palatine, Illinois; David Tanis,Holland Christian High School, Holland, Michigan; DaleWolfgram, Grand Blanc High School, Grand Blanc, Michigan;Clement Zidick, Dimond and Wasilla High Schools, Anchor-age, Alaska

81

Index

Acid(s), 52Air, dissolved, 28Antifreeze, 29-30Atmosphere, 17Atmospheric pressure,

17-18Avogadro's law, 16

Barometer, 16-18Base(s), 52Boiling point, 24-25Boyle's law, 16

Calories, countingof, 36

Calorimetry measurements,36-37

Capillary tubes, 18Catalysts, 46Cells

electrochemical, 56-58leadacid, 57

Charles' law, 16, 22Chemical equilibrium,

45-49Colloid(s)

association, 31lyophilic, 31particles of, 31-32

Coulomb's law, 23Crystal(s)

growing, 25liquid, 31metal, 26in solutions, 28

Dissolving process,27-28

Dynamic equilibrium,24, 27-28, 50

Electrochemistry, 56Electroplating, 59-60Electrostatic forces, 23-24Emulsions, 33-35Endothermic process,

38-39

Energychanges in, 36-38chemical, 36-40electrical, 56kinetic, 36mechanical, 36, 40thermal, 36-37

Entropy, 48-49Equilibrium, 27-28Equilibrium constant(s),

47Exothermic process, 38-40

Foam, 32-33Fuel cells, 59

Gas(es), 16-21dissolved, 27-29measurement of, 16

Heat, 36-37

Ice, 37Ideal gas law, 20-21

Le Chatelier's principle,51-52

Liquid(s)boiling of, 24-25crystals, 31vaporization of, 24

London force, 23

Macromolecules. SeeColloid(s), lyophilic

Macroscopic dimensions,13

Mercury, 16-18Micelles. See Colloid(s),

associationMolecular dimensions, 13Molecules

attraction of, 23in colloids, 31-35forces between,

23-24in gases, 16-20

in liquids, 23-25polar, 23-24in solids, 25-26

Moles, of gas, 20-21

Newton's second law,17

Pascal(s), 16Polar

attraction, 23forces, 24

Precipitation, 49Pressure, and volume,

17-19

Rate(s)and concentration, 45law, 45of reaction, 42and temperature, 45

Solid(s)amorphous, 26concentration of, 28ionic, 28

Solution(s), 27-30heat of, 39supersaturated, 28

Temperaturemolecular basis of, 16and reaction rate, 45-46in solution, 29-30and volume, 19-20

Thermal energy. See Energy,thermal

Thermochemistry, 36-37Thermodynamics

first law of, 37second law of, 40Tyndall effect, 33

Van der Waals force,23

Vaporization, 24Vapor phase, 23

3EST COPY AVAILABLE

75

Table of International Relative Atomic Masses*

Element SymbolAtomicNumber

AtomicMass Element Symbol

AtomicNumber

AtomicMass

Actinium Ac 89 227.0 Mercury Hg 80 200.6

Aluminum Al 13 27.0 Molybdenum Mo 42 95.9

Americium Am 95 (243)- Neodymium Nd 60 144.2

Antimony Sb 51 121.8 Neon Ne 10 20.2

Argon Ar 18 39.9 Neptunium Np 93 237.0

Arsenic As 33 74.9 Nickel Ni 28 58.7

Astatine At 85 (210) Niobium Nb 41 92.9

Barium Ba 56 137.3 Nitrogen N 7 14.0

Berkelium Bk 97 (247) Nobelium No 102 (259)

Beryllium Be 4 9.01 Osmium Os 76 190.2

Bismuth Bi 83 209.0 Oxygen 0 8 16.0

Boron B 5 10.8 Palladium Pd 46 106.4

Bromine Br 35 79.9 Phosphorus P 15 31.0

Cadmium Cd 48 112.4 Platinum Pt 78 195.1

Calcium Ca 20 40.1 Plutonium Pu 94 (244)

Californium Cf 98 (251) Polonium Po 84 (209)

Carbon C 6 12.0 Potassium K 19 39.1

Cerium Ce 58 140.1 Praseodymium Pr 59 140.9

Cesium Cs 55 132.9 Promethium Pm 61 (145)

Chlorine CI 17 35.5 Protactinium Pa 91 231.0

Chromium Cr 24 52.0 Radium Ra 88 226.0

Cobalt Co 27 58.9 Radon Rn 86 (222)

Copper Cu 29 63.5 Rhenium Re 75 186.2

Curium Cm 96 (247) Rhodium Rh 45 102.9

Dysprosium Dy 66 162.5 Rubidium Rb 37 85.5

Einsteinium Es 99 (254) Ruthenium Ru 44 101.1

Erbium Er 68 167.3 Samarium Sm 62 150.4

Europium Eu 63 152.0 Scandium Sc 21 45.0

Fermium Fm 100 (257) Selenium Se 34 79.0

Fluorine F 9 19.0 Silicon Si 14 28.1

Francium Fr 87 (223) Silver Ag 47 107.9

Gadolinium Gd 64 157.3 Sodium Na 11 23.0

Gallium Ga 31 69.7 Strontium Sr 38 87.6

Germanium Ge 32 72.6 Sulfur S 16 32.1

Gold Au 79 197.0 Tantalum Ta 73 180.9

Hafnium Hf 72 178.5 Technetium Tc 43 (97)

Helium He 2 4.00 Tellurium Te 52 127.6

Holmium Ho 67 164.9 Terbium Tb 65 158.9

Hydrogen H 1 1.008 Thallium TI 81 204.4

Indium In 49 114.8 Thorium Th 90 232.0

Iodine I 53 126.9 Thulium Tm 69 168.9

Iridium Ir 77 192.2 Tin Sn 50 118.7

Iron Fe 26 55.8 Titanium Ti 22 47.9

Krypton Kr 36 83.8 Tungsten W 74 183.8

Lanthanum La 57 138.9 Uranium U 92 238.0

Lawrencium Lr 103 (260) Vanadium V 23 50.9

Lead Pb 82 207.2 Xenon Xe 54 131.3

Lithium Li 3 6.94 Ytterbium Yb 70 173.0

Lutetium Lu 71 175.0 Yttrium Y 39 88.9

Magnesium Mg 12 24.3 Zinc Zn 30 65.4

Manganese Mn 25 54.9 Zirconium Zr 40 91.2

Mendelevium Md 101 (258)

*Based on International Union of Pure and Applied Chemistry (IUPAC) values (1975).

"Numbers in parentheses give the mass numbers of the most stable isotopes.

76 63

PE

RIO

DIC

TA

BLE

OF

TH

E E

LEM

EN

TS

IAIIA

6.94

LiLi

thiu

m3

9.01

Be

Ber

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m4

23.0

24.3

Na

Mg

Sod

ium

Mag

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12

1.00

8 HH

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VB

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Bor

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.0 Al

Alu

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Car

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16.0

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Oxy

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31.0

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PS

Pho

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19.0 F

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.5 CI

Chl

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20.2 N

eN

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10 39.9 A

rA

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1839

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sium

19

40.1 C

aC

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45.0 S

cS

cand

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21,

47.9 T

iT

itani

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50.9 V

Van

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m23

52.0 C

rC

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24

54.9 M

nM

anga

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25

55.8 F

eIr

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58.9 C

oC

obal

t27

58.7 N

iN

icke

l28

63.5 C

uC

oppe

r29

65.4 Z

nZ

inc

30

69.7 G

aG

alliu

m31

72.6 G

eG

erm

aniu

m32

74.9 A

sA

rsen

ic33

79.0 S

eS

elen

ium

34

79.9 B

rB

rom

ine

35

83.8 K

rK

rypt

on

3685

.5 Rb

Rub

idiu

m37

87.6 S

rS

tron

tium

38

88.9

Yttr

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39

91.2 Z

rZ

ircon

ium

40

92.9

Nb

Nio

bium

41

95.9 M

oM

olyb

d'm

42

(97) T

cT

echn

etiu

m43

101.

1 Ru

Rut

heni

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102.

9 Rh

Rho

dium

45

106.

4 Pd

Pal

ladi

um46

107.

9 Ag

Silv

er47

112.

4 Cd

Cad

miu

m48

114.

8 InIn

dium

49

118.

7

Tin 50

Sn

121.

8 Sb

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imon

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127.

6 Te

Tel

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126.

9

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131.

3 Xe

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2.9 Cs

Ces

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55

137.

3 Ba

Bar

ium

56

138.

9 La*

Lant

hanu

m57

178.

5 Hf

Haf

nium

72

180.

9 Ta

Tan

talu

m73

183.

8

Tun

gste

n74

186.

2 Re

Rhe

nium

75

190.

2 Os

Osm

ium

76

192.

2 IrIr

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m77

195.

1 Pt

Pla

tinum

78

197.

0 Au

Gol

d79

200.

6 Hg

Mer

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80

204.

4 TI

Tha

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81

207.

2 Pb

Lead

82

209.

0 B i

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) Po

Pol

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) At

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88

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Ac

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Act

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104

105

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106

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150.

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157.

315

8.9

162.

516

4.9

167.

316

8.9

173.

017

5.0

Ce

Pr

Nd

Pm

Sm

Eu

Gd

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Dy

Ho

Er

Tm

Yb

LuC

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5960

6162

6364

6566

6768

6970

71

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(251

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(254

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.

I.

U.S. Department of EducationOffice of Educational Research and Improvement (OERI)

National Library of Education (NLE)Educational Resources Information Center (ERIC)

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-'5&06 3 3o

Title: Teacher's Guide for Communities of Molecules: A Physical Chemistry Module

Author(s): Howard DeVoe and Robert Hearle

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Chemistry Associates of Maryland, Inc.

Publication Date:

1978

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