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Chemical Reactions 1:Energy and Chemical Dynamics

CHE-5042-2

Learning Guide

CHEMICAL REACTIONS 1: ENERGY AND CHEMICAL DYNAMICS

CHE-5042-2LEARNING GUIDE

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Chemical Reactions 1: Energy and Chemical Dynamics is the second of three Learning

Guides prepared for the three courses making up the Secondary V Chemistry

program, which comprises the following three courses:

Gases

Chemical Reactions 1: Energy and Chemical Dynamics

Chemical Reactions 2: Equilibrium and Oxidation–Reduction

The three Learning Guides are accompanied by the workbook, Experimental Activities

of Chemistry, which covers the “experimental method” component of the program.


This Guide was produced by the Société de formation à distance des commissions

scolaires du Québec.

Project Coordinator: Jean-Simon Labrecque (SOFAD)

Project Coordinator: Mireille Moisan (first edition)

Coordinator: Céline Tremblay (FormaScience)

Author: André Blondin

Illustrators: Gail Weil Brenner (GWB)

Jean-Philippe Morin (JPM)

Content Revisors: Céline Tremblay (FormaScience) (French Version)

Stéphanie Belhumeur (English Version)

Layout: I. D. Graphique inc. (Daniel Rémy)

Translator: Claudia de Fulviis

Linguistic Revisor: Patricia Fillmore

Proofreader: Gabriel Kabis

First Edition: November 2000

September 2008

© Société de formation à distance des commissions scolaires du Québec

All rights for translation and adaptation, in whole or in part, are reserved for all countries.

Any reproduction by mechanical or electronic means, including microreproduction, is

forbidden without the written permission of a duly authorized representative of the Société

de formation à distance des commissions scolaires du Québec.

Legal Deposit – 2000

Bibliothèque et Archives nationales du Québec

Bibliothèque et Archives Canada

ISBN 978-2-89493-192-9

TABLE OF CONTENTS

GENERAL INTRODUCTION

OVERVIEW ................................................................................................................... 0.10

HOW TO USE THIS LEARNING GUIDE ............................................................................. 0.10

Learning Activities ................................................................................................. 0.11

Exercises .............................................................................................................. 0.11

Self-evaluation Test ............................................................................................... 0.12

Appendices ........................................................................................................... 0.12

Materials .............................................................................................................. 0.12

CERTIFICATION ............................................................................................................. 0.13

INFORMATION FOR DISTANCE EDUCATION STUDENTS .................................................... 0.13

Work Pace ............................................................................................................ 0.13

Your Tutor ............................................................................................................ 0.13

Homework Assignments ....................................................................................... 0.14

CHEMICAL REACTIONS 1: ENERGY AND CHEMICAL DYNAMICS ...................................... 0.15

CHAPTER 1 – HEAT: ENERGY IN MOTION .................................................................... 1.1

1.1 ENERGY ............................................................................................................... 1.3

Forms of Energy .................................................................................................... 1.4

Chemical Energy in Action ............................................................................... 1.6

Potential Energy .............................................................................................. 1.7

Energy Conversions .............................................................................................. 1.8

Conservation of Energy ......................................................................................... 1.13

1.2 HEAT .................................................................................................................... 1.15

Kinetic Molecular Model of Matter ......................................................................... 1.16

Thermometers and Heat Transfers ......................................................................... 1.21

Experimental Activity 1: Heat Transfers ..................................................... 1.22

Mercury Thermometers ................................................................................... 1.23

Thermal Equilibrium ........................................................................................ 1.26

Temperature Scales ........................................................................................ 1.26

The Mechanical Equivalent of Heat ........................................................................ 1.29

Heat and Thermal Energy ...................................................................................... 1.30

1.3 HEAT EXCHANGES ................................................................................................. 1.33

Experimental Activity 2: Final Temperature of a Mixture .............................. 1.33

Specific Heat Capacity ........................................................................................... 1.34

Sea Breezes and Land Breezes ....................................................................... 1.36

Heat Exchange Equation ........................................................................................ 1.37

Calorimeters ................................................................................................... 1.38

Applications ................................................................................................... 1.39

Chemical Reactions 1 - Table of Contents

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Energy in Phase Changes ...................................................................................... 1.44

Melting ........................................................................................................... 1.46

Boiling ............................................................................................................ 1.48

1.4 TECHNICAL APPLICATIONS ..................................................................................... 1.51

The Bread Oven .................................................................................................... 1.51

Re-entering the Atmopshere ................................................................................... 1.52

Geysers ................................................................................................................ 1.52

Key Words in This Chapter ............................................................................................ 1.54

Summary ..................................................................................................................... 1.54

Review Exercises .......................................................................................................... 1.56

CHAPTER 2 – DISSOLUTION: AN ENERGY PHENOMENON ............................................. 2.1

2.1 MIXTURES AND AQUEOUS SOLUTIONS ................................................................... 2.3

Mixtures ............................................................................................................... 2.3

The Water Molecule ............................................................................................. 2.5

Capillary Rise ................................................................................................. 2.9

2.2 DISSOLUTION AND SOLUBILITY .............................................................................. 2.10

Molecular Dissolution ............................................................................................ 2.10

Ionic Dissolution ................................................................................................... 2.12

Hydration ........................................................................................................ 2.14

Salt Waters .................................................................................................... 2.16

Electrolytes ..................................................................................................... 2.16

Solubility and Precipitation ..................................................................................... 2.18

Precipitation ................................................................................................... 2.20

2.3 DISSOLUTION: AN ENERGY PHENOMENON ............................................................. 2.23

The Process of Dissolution .................................................................................... 2.25

Molar Heat of Solution .......................................................................................... 2.28

Experimental Activity 3: Molar Heat of Solution .......................................... 2.30

2.4 SOLUTIONS IN EVERYDAY LIFE ............................................................................... 2.31

Gases in Aqueous Solutions .................................................................................. 2.31

The Senses of Taste and Smell ............................................................................. 2.31

Pharmaceuticals, Beauty Products and Perfumes ................................................... 2.33

Other Solvents ...................................................................................................... 2.34


Summary ..................................................................................................................... 2.35

Review Exercises .......................................................................................................... 2.37

CHAPTER 3 – CHEMICAL REACTIONS AND ENERGY .................................................... 3.1

3.1 HEAT OF REACTION ............................................................................................... 3.3

Definition and Convention ...................................................................................... 3.4

Measuring the Heat of Reaction ............................................................................. 3.7

Change in Enthalpy ............................................................................................... 3.9


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Enthalpy Diagrams ............................................................................................... 3.12

Activation Energy ................................................................................................... 3.16

3.2 COMBUSTION REACTIONS ..................................................................................... 3.18

Rapid Combustion and Slow Combustion ............................................................... 3.18

A Little Bit of History ....................................................................................... 3.21

Fossil Fuels: A Useful Commodity .......................................................................... 3.24

Industrialization and Social Changes ................................................................ 3.27

A Technical Application .................................................................................... 3.29

Fossil Fuels: The Drawbacks .................................................................................. 3.31

Carbon Dioxide (CO2) ....................................................................................... 3.31

Carbon Monoxide (CO) .................................................................................... 3.34

Nitrogen Oxides and Sulphur Dioxide ............................................................... 3.37

A Few Possible Solutions ................................................................................ 3.37

3.3 HESS’S LAW ......................................................................................................... 3.39

Experimental Activity 4: Hess’s Law .......................................................... 3.41

Summation of Heats of Reaction ........................................................................... 3.41

Application of Hess’s Law ..................................................................................... 3.44

Hess’s Law and Chemical Bonds ........................................................................... 3.48


Summary ..................................................................................................................... 3.53

Review Exercises ......................................................................................................... 3.55

CHAPTER 4 – THE RATE OF CHEMICAL REACTIONS ..................................................... 4.1

4.1 PROGRESS OF A CHEMICAL REACTION ................................................................. 4.3

Rate of a Reaction ................................................................................................ 4.4

Using Graphs to Represent Reaction Rates ............................................................ 4.8

Reaction Rate As a Function of Time ...................................................................... 4.10

4.2 DETERMINING FACTORS ........................................................................................ 4.20

Experimental Activity 5: Rate of a Chemical Reaction ................................. 4.21

Nature of Reactants .............................................................................................. 4.21

Concentration of Reactants ................................................................................... 4.23

Pressure and Gaseous Reactants .......................................................................... 4.24

Temperature ......................................................................................................... 4.25

Surface Area ........................................................................................................ 4.27

Catalysts .............................................................................................................. 4.30

An Overview .......................................................................................................... 4.34

Preparation of a Reference Solution ................................................................. 4.34

Effect of Concentration on the Rate of Reaction ............................................... 4.34

Effect of Temperature on the Rate of Reaction ................................................. 4.35


Summary ..................................................................................................................... 4.37



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CHAPTER 5 – ENERGY AND THE RATE OF REACTION ................................................... 5.1

5.1 MOLECULAR ENERGY ............................................................................................ 5.3

Kinetic Energy of Molecules ................................................................................... 5.4

A Nornal Distribution for Large Populations ............................................................. 5.6

Maxwell-Boltzmann Distribution ............................................................................. 5.10

5.2 CONSEQUENCES OF THE MAXWELL-BOLTZMANN DISTRIBUTION .............................. 5.16

Temperature ......................................................................................................... 5.16

Threshold Energy and Activation Energy .................................................................. 5.18

Other Applications ................................................................................................. 5.24

Kinetic Energy in Our Lives .................................................................................... 5.24

5.3 ENERGY AND THE MECHANISM OF REACTION ........................................................ 5.27

Reaction Mechanism ............................................................................................. 5.27

Catalysts .............................................................................................................. 5.32

Spontaneity of Reactions ....................................................................................... 5.36


Summary ..................................................................................................................... 5.40


CONCLUSION

SELF-EVALUATION TEST ............................................................................................... C.5

ANSWER KEY ............................................................................................................... C.13

CHAPTER 1 – Heat: Energy in Motion ..................................................................... C.13

CHAPTER 2 – Dissolution: an Energy Phenomenon .................................................. C.27

CHAPTER 3 – Chemical Reactions and Energy ........................................................ C.36

CHAPTER 4 – The Rate of Chemical Reactions ....................................................... C.49

CHAPTER 5 – Energy and the Rate of Reaction ....................................................... C.61

ANSWER KEY TO THE SELF-EVALUATION TEST ........................................................ C.68

APPENDIX A – The International System of Units (SI) ...................................................... C.73

APPENDIX B – Mathematical Prerequisites ..................................................................... C.75

Ratios and Proportions .......................................................................................... C.75

Formulas .............................................................................................................. C.76

APPENDIX C – Chemical Prerequisites ........................................................................... C.78

Balancing Equations .............................................................................................. C.78

Calculating Molar Mass ......................................................................................... C.81

APPENDIX D – Table of Figures ..................................................................................... C.83

BIBLIOGRAPHY ............................................................................................................. C.87

GLOSSARY ................................................................................................................... C.89

INDEX .......................................................................................................................... C.97


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GENERAL INTRODUCTION

OVERVIEW

Welcome to the course entitled Chemical Reactions 1: Energy and Chemical Dynamics,

which is part of the Secondary V Chemistry program. This program comprises the

following three courses:

CHE-5041-2 Gases

CHE-5042-2 Chemical Reactions 1: Energy and Chemical Dynamics

CHE-5043-2 Chemical Reactions 2: Equilibrium and Oxidation—Reduction

The three main components of the Chemistry program are related content, the

experimental method and the history-technology-society perspective. Whereas the

experimental method is developed in the workbook Experimental Activities of

Chemistry, the related content and the history-technology-society perspective are

covered in the three Learning Guides accompanying the three courses which must

be taken in sequential order.

Chemical Reactions 1: Energy and Chemical Dynamics is the second in the set of three

Learning Guides. It is divided into five chapters corresponding to the five terminal

objectives of the program.1 This Guide is to be used in conjunction with the workbook

Experimental Activities of Chemistry. You will find references to the appropriate sections

of the Workbook throughout the Guide.

The course Chemical Reactions 1: Energy and Chemical Dynamics will help you gain

a better understanding of chemical dynamics and the energy transfers involved in

chemical reactions, together with the related technical applications, social changes

and environmental consequences.

HOW TO USE THIS LEARNING GUIDE

This Guide is the main work tool for this course and has been designed to meet the

needs of adult students enrolled in individualized learning programs, or distance

education courses.

Each chapter covers a certain number of themes, using explanations, tables,

illustrations and exercises designed to help you to master the different program

objectives. A list of key words, a summary and review exercises are included at the

end of each chapter.

Chemical Reactions 1 - General Introduction

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1. The terminal objectives and associated objectives are listed at the beginning of each chapter.

The conclusion contains a summary covering all the courses in the program along

with a self-evaluation test. It also includes an Answer Key for the self-evaluation test,

for the exercises in each chapter and for the review exercises. A glossary with definitions

of the key words, a bibliography, appendices and an index are also provided in the

conclusion. You may wish to consult the books and publications in the bibliography

for further information on the topics covered in this course.

Learning Activities

The Guide contains theoretical sections as well as practical activities in the form of

exercises. The exercises come with an Answer Key.

Start by skimming through each part of the Guide to familiarize yourself with the

content and the main headings. Then read the theory carefully:

– Highlight the important points.

– Make notes in the margins.

– Look up new words in the dictionary.

– Summarize important passages in your own words, in your notebook.

– Study the diagrams carefully.

– Write down questions relating to ideas you don’t understand.

Exercises

The exercises come with an Answer Key, which is located in the coloured section at

the end of the Guide.

• Do all the exercises.

• Read the instructions and questions carefully before writing your answers.

• Do all the exercises to the best of your ability without looking at the Answer Key.

Reread the questions and your answers, and revise your answers, if necessary. Then

check your answers against the Answer Key and try to understand any mistakes

you made.

• Complete each chapter before doing the corresponding review exercises. Doing these

exercises without referring to the Guide is a good way to prepare for the final

examination.


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Self-evaluation Test

The self-evaluation test is a step that prepares you for the final evaluation. You must

complete your study of the course before attempting to do it. Reread your notebook

and the definitions of the key words in the chapters. Make sure you understand how

they relate to the course objectives listed at the beginning of each chapter. Then do

the self-evaluation test without referring to the main body of the Guide or the Answer

Key. Compare your answers with those in the Answer Key and review any areas you

had difficulty with.

Appendices

The appendices contain a review of some concepts you should be familiar with before

beginning the course. The complete list of appendices appears in the table of contents.

Materials

Have all the materials you will need close at hand:

• Learning material: this Guide and a notebook in which you will summarize important

concepts relating to the objectives (listed in the introduction of each chapter). You

will also need to use your periodic table and the workbook Experimental Activities

of Chemistry.

• Reference material: a dictionary.

• Miscellaneous material: a calculator, a pencil for writing your answers and notes

in your Guide, a coloured pen for correcting your answers, a highlighter (or a pale-

coloured felt pen) to highlight important ideas, a ruler, an eraser, etc.


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CERTIFICATION

To earn credits for this course, you must obtain at least 60% on the final examination

which will be held in an adult education centre.

The evaluation for the course Chemical Reactions 1: Energy and Chemical Dynamics

is divided into two separate parts.

Part I consists of a two-hour written examination made up of multiple-choice, short-

answer and essay-type questions. This part is worth 75% of your final mark and deals

with the objectives covered in this Guide. You may use a calculator.

Part II is designed exclusively to evaluate the experimental method. It will be held

in the laboratory during one 90-minute session. It is worth 25% of your final mark

and deals with the course objectives covered in Section B of Experimental Activities

of Chemistry.

INFORMATION FOR DISTANCE EDUCATION STUDENTS

Work Pace

Here are some tips for organizing your work:

• Draw up a study timetable that takes into account your personality and needs, as

well as your family, work and other obligations.

• Try to study a few hours each week. You should break up your study time into several

one- or two-hour sessions.

• Do your best to stick to your study timetable.

Your Tutor

Your tutor is the person who will give you any help you need throughout this course.

He or she will answer your questions and correct and comment on your homework

assignments.

Don’t hesitate to contact your tutor if you are having difficulty with the theory or the

exercises, or if you need some words of encouragement to help you get through this

course. Write down your questions and get in touch with your tutor during his or


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her available hours. The letter included with this Guide or that you will receive shortly

tells you when and how to contact your tutor.

Your tutor will assist you in your work and provide you with the advice, constructive

criticism, and support that will help you succeed in this course.

Homework Assignments

In this course, you will have to do three homework assignments: the first after

completing Chapter 2, the second after completing Chapter 4, and the third after

completing Chapter 5. Each homework assignment also contains questions on the

experimental method you studied in Experimental Activities of Chemistry.

These assignments will show your tutor whether you understand the subject matter

and are ready to go on to the next part of the course. If your tutor feels you are not

ready to move on, he or she will indicate this on your homework assignment, providing

comments and suggestions to help you get back on track. It is important that you

read these corrections and comments carefully.

The homework assignments are similar to the examination. Since the exam will be

supervised and you will not be able to use your course notes, the best way to prepare

for it is to do your homework assignments without referring to the Learning Guide

and to take note of your tutor’s corrections so that you can make any necessary

adjustments.

Remember not to send in the next assignment until you have received the corrections

for the previous one.


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Forest fires, sudden frost, melting snow, respiration, the shuttle launch and archery

all involve changes that have one thing in common—energy. While we cannot observe

energy directly, we perceive it as light, heat, motion and noise, among others things.

Energy takes different forms and is named according to the way it manifests itself.

For instance, we speak of mechanical energy to describe the thrust, firing or rotation

of a mechanical part, of kinetic energy for the movement of molecules or other bodies,

of radiant energy for light, and of potential energy for all forms of stored energy.

Associated with hot and cold sensations, heat is transferred between systems when

they are at different temperatures. We might say that heat is a mode of energy transfer.

Heat and changes in temperature are therefore closely linked. Heat can either be

absorbed or released during a chemical reaction. For instance, energy is released when

we burn natural gas to heat our homes, whereas enormous amounts of energy are

consumed in aluminum production.

The last topic covered in the course entitled Gases was the energy balance of chemical

reactions. The second course, Chemical Reactions 1: Energy and Chemical Dynamics,

examines energy transfers from a broader perspective as well as the rate of reactions

and the factors that affect it. The third course, Chemical Reactions 2: Equilibrium and

Oxidation—Reduction, provides an in-depth study of two categories of chemical

reactions.

Chapter 1 of this Guide analyzes energy transfers between two liquids that are mixed

together. Energy transfers depend on the amounts and types of liquids involved, and

can be observed through changes in temperature. The first chapter also reviews the

heating curve of a substance and gives a more detailed explanation for rises in

temperature and phase changes.

Chapter 2 deals with dissolution reactions, and describes them macroscopically, that

is, according to what can be observed with the naked eye. They are then examined

at the molecular level, using the kinetic molecular model of matter, which is an

expanded version of the model based on the kinetic theory of gases. You will learn

why water is the most common solvent in nature and the solvent of choice in the

laboratory. This chapter focuses on the energy involved in dissolutions.

Chapter 3 deals with the energy associated with chemical reactions. It also reviews

the energy balance involved in the dissociation and formation of bonds. In this course,


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however, the subject of energy balance is examined in greater detail. The energy

diagrams provide more information about the progress of reactions, and the sum of

the energies involved in the steps of a reaction yields the overall energy of the reaction.

This analysis will be based on combustion reactions.

Chapters 4 and 5 examine kinetic chemistry or the rate of chemical reactions. This

rate depends on the nature of the reactants, their concentration, their surface area

and on environmental conditions such as temperature, pressure and acidity. Catalysts

speed up reactions, but do not alter the nature of the products. Speed is part and

parcel of any process involving energy. The fifth and last chapter in this course examines

the relationship between energy, reaction rate and the various factors that affect this

rate. It also introduces the content of the next course.

As in the first two Guides, a table of contents diagram at the beginning of each chapter

shows you how the chapter fits into the course as a whole. The content of the chapter

you are about to begin is in bold type and in larger characters, whereas the content

of completed chapters is in italics. For example, the table of contents diagram for

Chapter 2 is reproduced below. The section for Chapter 2 is in bold type and the content

of Chapter 1 is in italics and smaller type. You will find this diagram a very useful

tool as you go through the course.

Good luck!

1. Heat: Energy in MotionEnergy: forms, conservationHeat: model, thermometerHeat exchangesEnergy in phase changesApplications

2. Dissolution: An Energy PhenomenonMixtures, solutions

Molecular and ionic dissolutionSolubility, precipitation

Heat of solutionSolutions in everyday life

5. Energy and the Rate of ReactionMaxwell-Boltzmann distributionEnergy threshold and activation energyReaction mechanismsSpontaneity of reactions

CHEMICAL REACTIONS 1:

Energy and Chemical Dynamics

4. Rate of Chemical ReactionsRate of reactionGraphsDetermining factors

3. Chemical Reactions and EnergyHeat of reaction

Rapid and slow combustionFossil fuelsHess’s law


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CHAPTER 1

HEAT: ENERGY IN MOTION

GWB

Terminal Objective 1

To analyze the energy transfers that occur in phase changes and mixtures of substances at different temperatures.

Intermediate Objectives

1.1 To recognize the different forms of energy acting in the phenomena observed in their environment.

1.2 To associate a macroscopic phenomenon with corresponding changes occurring at the atomic or molecularlevel.

1.3 To describe a heat transfer in terms of kinetic energy and the variation in temperature.

1.4 To classify physical and chemical phenomena according to whether they represent endothermic orexothermic reactions, on the basis of observations.

1.5 To determine, through experimentation, the factors that influence the final temperature of a mixture.

1.6 To establish relationships between the definition of specific heat capacity and its units of measurement.

1.7 To describe the energy transfers produced during phase changes of a pure substance.

1.8 To describe briefly how Joule established a relationship between heat and mechanical energy.

1.9 To give examples of energy conversions involving heat.

1.10 To solve problems related to energy transfers that occur during phase changes and mixtures of substancesat different temperatures.

1. Heat: Energy in MotionEnergy: forms, conservationHeat: model, thermometerHeat exchangesEnergy in phase changesApplications

2. Dissolution: An Energy PhenomenonMixtures, solutions

Molecular and ionic dissolutionSolubility, precipitation

Heat of solutionSolutions in everyday life

5. Energy and the Rate of ReactionMaxwell-Boltzmann distributionEnergy threshold and activation energyReaction mechanismsSpontaneity of reactions

CHEMICAL REACTIONS 1:

Energy and Chemical Dynamics

4. Rate of Chemical ReactionsRate of reactionGraphsDetermining factors

3. Chemical Reactions and EnergyHeat of reaction

Rapid and slow combustionFossil fuelsHess’s law

Chemical Reactions 1 - Chapter 1: Heat: Energy in Motion

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Have you ever wondered about all the different ways in which we use the word

“energy”? You are no doubt aware that this term has several meanings. For instance,

in the expression “My child is full of energy,” it refers to the ability to move and do

things; however, when we say that “Solar energy powers satellites that orbit the earth,”

it refers to light; and in “Water heaters consume a lot of energy,” it refers to the source—

more often than not electric—which heats the water.

In this chapter, we will review the various forms of energy and pay particular attention

to heat. We will explore the concept of temperature in greater detail, explain heat

transfers at the molecular level and represent them by means of a mathematical

relation. We will then examine the energy associated with phase changes more closely

and end with an overview of a few applications of heat energy.

1.1 ENERGY

Human beings have used what we now call energy since the dawn of time and have

tried to harness it in different ways. At first, humans used energy simply to eat, keep

warm and bask in the Sun’s light. They then went on to “tame” fire and invent the

wheel, levers and other, more complex devices, such as windmills and watermills.

More recently, humans have mastered the use of certain mixtures of explosive

substances: this has contributed to the development of firearms, combustion engines

and dynamite, used to dig mines and construct roads through mountains.

Humans are still actively involved in the quest for technical advances to meet their

ever-changing needs—the construction of completely automated plants and more

effective power-generating stations are a few examples. Others include launching

satellites into orbit around the Earth, developing the Moon’s mineral resources and

exploring the planet Mars, projects which have been undertaken by NASA1 and other

space agencies worldwide.

Although the word is used frequently, “energy” remains a difficult concept to grasp.

Strictly speaking, energy is defined as the capacity for doing work, or the capacity

for producing an effect such as movement or light, among other things.


1.3

1. National Aeronautics and Space Administration: the organization that oversees the entire American space program.In Canada, the Canadian Space Agency fulfills the same function.

Energy cannot be observed directly. Rather, we observe its manifestations, the main

ones being light, heat and motion. Our eyes take in light in a very precise manner

and very small corpuscles2 in the skin detect sensations of warmth or cold. And of

course, our eyes, ears and muscles can detect motion.

Exercise 1.1

Complete the following table by naming sources of energy that are detected as light,

heat or motion. Write two sources for each form of energy indicated.

Form Sources

Light Le soleil, une ampoule électrique allumée, une flamme, etc.nnn

Heat L’élément électrique d’un rond de poêle allumé, d’un grille-pain, d’unebouilloire, d’un radiateur-plinthe ou d’une ampoule électrique ; le composten décomposition, une flamme, etc.

Motion L’énergie d’un trampoline compressé, d’un arc bandé, d’un élan de bâtonde golf ou de baseball ; une voiture, un vélo ou un train en mouvement, unenfant qui court, la rotation des aiguilles d’une horloge, etc.

FORMS OF ENERGY

Forms of energy are usually named after their source or the process of transformation

that produces it. For instance, mechanical energy is produced by the movement of

mechanical parts in machinery and engines which cut, strike or launch objects; solar

energy includes visible light, ultraviolet rays and the heat released by the Sun;

gravitational energy is associated with the force of gravity which causes objects to

fall to the ground; nuclear energy is produced by the fusion or fission of atomic nuclei

and fuels nuclear power stations and aircraft carriers; magnetic energy causes two

magnets to repel each other and aligns the electron stream that produces the image

on a television screen; electrical energy lights our homes and powers engines of all

kinds; wind energy refers to the wind’s capacity to turn the blades of a mill;

hydraulic energy is provided by rivers, waterfalls and tides; muscular energy refers

to a muscle’s capacity to move an arm or a leg; electromagnetic energy includes visible

light, X-rays, microwaves and other types of electromagnetic radiation.


1.4

?

2. Ruffini’s corpuscles detect sensations of warmth and Krause’s corpuscles detect sensations of cold.

More specifically, kinetic energy is the name given to the energy associated with the

straight-line or spinning motion of an object. By contrast, heat, which is also named

calorific energy, involves a transfer of energy between two systems.

In subsequent chapters, we will focus on a form of energy called chemical energy,

so named because it is associated with chemical reactions that, as you may recall,

change the nature of substances. For instance, the combustion of propane gas releases

heat that comes from the chemical energy contained in the propane and oxygen. Thus,

the bond energy3 discussed at the end of the last course is chemical energy.

Exercise 1.2

State the form of energy (wind energy, nuclear energy or hydroelectric energy) that

best corresponds to the descriptions given below.

The electrical energy flowing out of a generator that is driven by arotating turbine powered by a waterfall.

The light, heat and kinetic energy released by the fission of heavy andunstable atomic nuclei such as those of uranium 235.

The mechanical energy generated by the movement of air across theEarth’s surface.


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3. Lalancette, Pauline and M. Lamoureux, Gases (Chemistry, Secondary V), Chapter 1, Learning Guide producedby SOFAD.

?

Benjamin Franklin (1706-1790)4

Born in the Unites States and the fifteenth child in a family of English immigrants,

Benjamin Franklin was a self-educated and very talented man. In 1752, Franklin

performed his famous experiment that involved going out during a thunderstorm and

flying a kite with a metal key attached to it. He wanted to prove that lightning was

similar to electricity. Without knowing it, he was taking an enormous risk. Franklin

is an excellent example of the long line of curious-minded, ingenious people who

were determined to explore nature.

Franklin invented the stove that bears his name, the Franklin stove, the lightning

conductor (a metal rod placed on buildings to protect them against lightning) and bifocal lenses. He also

developed a theory to explain the absorption of heat. A master of many trades, Franklin was elected to

Pennsylvania’s legislative assembly and helped draft the American Declaration of Independence, which he

signed at the age of 70!

Chemical Energy in Action

Chemical energy is behind the powerful thrust needed to propel space rockets into

orbit. In American space shuttles, the central external tank fuels three engines placed

in a triangular arrangement at the back of the shuttle. Chemical energy is released

by the intense reaction between hydrogen and oxygen.5 In order to save space, the

two gases are first liquefied and stored separately in refrigerated tanks. The formation

of water vapor occurs, and the acceleration produced by the chemical energy that is

released causes it to be forcefully expelled from the rocket. Guided by the nozzles,

the expelled vapour propels the rocket. The shuttle’s three main engines provide a

thrust of more than two million newtons,6 or, in other words, the thrust needed to

propel the equivalent of 165 cars into orbit! At the outset, the shuttle weighs about

2 500 tonnes. The white trail behind the rocket as it rises into the sky results from

the condensation of water vapour, which cools upon contact with air.

The general combustion reaction for hydrogen is:

2 H2 + O2 → 2 H2O + energy


1.6

4. A light bulb indicates additional information: this information is not part of the course as such and will not becovered on the final examination.

5. “Propellant” is the term used to designate one or more substances that react chemically to produce the energyrequired to propel rockets into space. Hydrogen and oxygen constitute the main propellants for space shuttles.

6. The newton is the unit of force in the International System of Units (SI).

GWB

a) Hydrogen and oxygen form the space shuttle’s liquid propellant which is contained incompartments within the central external tank. The side rockets contain the solid propellant.

b) The central rocket engine is composed of tanks and fuel pumps. H2 and O2 combustion takes place in the three engines attached to the back of the shuttle.

Potential Energy

Regardless of its form, energy may be stored in such a way that it can be recovered

at a later time. Any form of stored energy that is ready for use is called potentialenergy. We will now look at three examples that will help us better understand this

concept: the propellant in a rocket, a match and an electric battery.

Let’s consider the two liquids (oxygen and hydrogen) stored in the refrigerated tanks

of a rocket that is about to be launched. Until the rocket has been launched, the liquids

remain in their respective chambers, the combustion reaction does not occur and

no energy is released. Upon ignition, however, the two liquids combine in the

combustion chamber, the reaction takes place and the energy released propels the

rocket into space. Prior to the reaction between the two liquids, the potential chemical

energy was stored in the substances, ready to be released.


1.7

External tank

Oxygen tankO2(l)

H2(l)

H2 + O2

H2 O+

Energy

Hydrogen tank

Combustionchamber

Nozzles

Solidpropellantrocket

Figure 1.1 - Liquid propellant rocket

Fuel pump

GWB

a) b)

}

A match in a matchbox will not light up by itself. By striking the match, we release

the energy contained in the substances that cover the match head. More precisely,

the potential chemical energy which is stored in the substances that make up the match

(sulphur, potassium chlorate, etc.) and in the oxygen in the air, is activated as soon

as the match is rubbed against a rough, specially prepared surface. We then perceive

heat and light.

When an electric battery is connected to a circuit, the substances in the battery will

react chemically to produce the energy needed to light a lamp or drive an engine.

The substances cannot react if the battery is not connected. The battery therefore

contains potential chemical energy.

All forms of energy can be described as potential, provided the energy is not activated,

that is, as long as it cannot be perceived.

Exercise 1.3

Briefly explain in your own words why we can say:

a) that a stone on the edge of a ravine has potential energy.

b) that a drawn arrow has potential energy.

ENERGY CONVERSIONS

The potential energy stored in a battery can be released in different forms depending

on whether the battery is used to power a flashlight, the motor of a toy car or any

other device. We speak of a conversion when energy changes form, that is, when it

is converted or changed from one form to another.


1.8

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Let’s consider an archer who shoots an arrow into the sky at a 45° angle. Between

the moment the archer takes aim and releases the arrow and the moment the arrow

hits the target, energy is converted from one form to another several times.

Figure 1.2 - An arrow in flight

The archer uses his muscles to release the arrow at a 45° angle above the horizontal plane.Muscular energy is transferred from the archer to the arrow. The arrow curves downwards

and strikes a target about 30 m away. Energy is then transferred from the arrow to the target; a part of this energy serves to drive the arrow into the wood and the rest is converted into heat.

Let’s look at this example more closely. Read on and try to imagine what happens

when the arrow is released. You will see that throughout the arrow’s flight, energy,

whether stored or active, is converted from one form to another.

By inhaling air and digesting his food, the archer provides his cells with the sugar

and oxygen they need to release the chemical energy that will allow him to contract

his muscles (muscular energy) and draw the bow (the bow stores potential mechanical

energy). The slow combustion of sugar in the cells converts the potential energy stored

in the sugar and oxygen into muscular energy that can be used to draw the bow.

When the arrow is released, the bow straightens out (mechanical energy of rotation)

and pulls on the bowstring that then propels the arrow upwards at a 45° angle (kinetic

energy). As the arrow rises into the air, the energy in the arrow serves to counteract

the gravitational force that is pulling it towards the ground. As it gains height, however,

the arrow stores potential gravitational energy. At the highest point along its


1.9

GWB

trajectory,7 the arrow changes direction and gains speed as it starts to fall (it acquires

kinetic energy). In short, the energy stored in the bow has been converted into the

kinetic energy of the arrow, which in turn is converted into potential gravitational

energy as the arrow rises and then converted again into kinetic energy as the arrow

falls.

The arrow eventually hits and becomes embedded in its target (mechanical energy).

Immediately after impact, the arrow’s metallic tip will feel slightly warm (heat). In

this case, the kinetic energy of the arrow has been converted once again, but this time

into heat and work, since the arrow has become lodged in the target.

You no doubt noticed in the description you have just read that energy takes a different

form with each conversion. For discussion purposes, imagine that in each case a

quantity of energy is transferred from a source to a receptor, which in turn becomes

the source for the next energy conversion. In this way, we have a continuous chain

of energy conversions.

Let’s now go back to our example of an arrow in flight. This time, however, we will

use the concepts of source and receptor. This exercise will allow you to become familiar

with these two terms used often in this chapter. Figure 1.3 illustrates the description

of the sequence of energy conversions.

Food is the primary source of energy. The potential energy that it contains is transferred

to the first energy receptor, namely, the archer’s muscles that contract. These then

become the source of power needed to draw the bow, which is the new receptor. The

drawn bow then becomes the source of the motion of the bowstring, which in turn

becomes the new receptor. The bowstring then propels the arrow, which receives the

bowstring’s energy. The arrow uses up this energy by gaining altitude. At the same

time, however, it stores potential gravitational energy. This energy is recovered when

the arrow falls towards the ground. Upon impact, the target and the materials that

make up the arrow absorb the arrow’s energy. The partial perforation of the target

and the heating up of the arrow’s tip are proof that they have become the new receptors.

In short, at all moments during the arrow’s path, we have been able to identify a source

of energy and at least one receptor which absorbs the source’s energy.


1.10

7. An arrow’s path always follows a more or less elongated parabola, depending on the arrow’s initial velocity andangle. The vertex is the highest point on the parabola.

Figure 1.3 - An unbroken chain of energy conversions

In the second step in the energy chain, the archer’s (source) muscular energy is transferred to thebow (receptor) in the form of potential mechanical energy. In turn, the bow becomes a source,

and its energy is transferred to the arrow (receptor). Thus, each person or object participating inthe action is both a receptor and a source of energy. The target is the last receptor in the sequence.

Let’s consider a second example that is similar to the arrow in flight. The figure below

illustrates a shotgun being fired.

Figure 1.4 - Trigger mechanism of a shotgun

Detail of a shotgun’s trigger mechanism

When the gunman presses the trigger, the shotgun’s firing pin strikes the detonator (cap) of thecartridge. The heat causes the gunpowder to explode. The shots are forced into the barrel and then

through the air, before reaching the target.


1.11

Chemicalenergy(foods)

Muscularenergy

Potentialmechanical

energy(bow)

Mechanicalenergy ofrotation

(bowstring)

Kineticenergy(arrowrises)

Potentialgravitational

energy (arrowon top)

Kineticenergy(arrowfalls)

Mechanicalenergy +

heat (arrow,target)

→ → → → → → →

GWB

Firing pin Shots

Detonator

Powder

Gun barrel

Firing chamberSpring releasemechanism

Firing pin spring

IDG GWB

Exercise 1.4

The following six steps describe the firing of shots with a shotgun. The steps are not

in chronological order.

• Released gases force the shots to travel down the barrel of the shotgun.

• A finger presses the trigger.

• The firing pin hits the detonator of the cartridge.

• The shots travel through the air and hit the target (a bottle).

• The bottle shatters and the pieces are projected into different directions.

• The heat released by the impact initiates combustion of the gunpowder in the

cartridge case.

Complete the following table by answering questions a) and b). The first line in the

table has been completed for you.

a) Place the steps in the order in which they occur.

b) For each step, describe the energy conversion that takes place and identify the energy

source and receptor.


1.12

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CONSERVATION OF ENERGY

Just as it is difficult to believe that a rabbit actually disappears inside a magician’s

hat (just because we cannot see it does not mean that it no longer exists!), it is equally

difficult to believe that energy can be destroyed. Rather, we observe, as have the leading

scientists of our time, that energy is converted into work or other forms without losing

its force. This principle, which was implicit in the examples we have looked at so far,


1.13

No. Step

1. A finger presses thetrigger.

2. Le percuteur dufusil frappe ledétonateur de laballe.

3. La chaleur dégagéepar l’impact amorcela combustion de lapoudre à canondans la douille.

4. Les gaz libéréspoussent lesplombs dans lecanon.

5. Les plombsvoyagent dans l’airet frappent la cible.

6. La bouteille éclateet les morceauxsont projetés dansdifférentesdirections.

Conversion

Muscular energy in thefinger → mechanical energyin the release mechanism

Énergie potentielle duressort → énergiemécanique + chaleur

Énergie potentielle chimiquedes substances → chaleur+ énergie cinétique desmolécules de gaz

Énergie cinétique desmolécules → énergiecinétique des plombs dechasse

Énergie cinétique desplombs → travail d’impactsur la bouteille + chaleur

Énergie cinétique desmorceaux de bouteille →impact sur les surfacesenvironnantes + chaleur

Source

Muscles in thefinger

Mécanisme dedétente (ressort)

Détonateur

Gaz libérés

Plombs

Morceaux debouteille

Receptor

Releasemechanism(spring)

Détonateur de laballe

Poudre

Plombs

Bouteille

Surfacesenvironnantes

is called the law of conservation of energy. It states that in a closed system, the

total amount of energy remains constant, regardless of the conversions it undergoes.

In other words, the total amount of energy is the same (it is conserved) before and

after the conversion.

Law of conservation of energy:

Total energy before conversion = Total energy after conversion

Remember that a system is a collection of objects that form a whole. In the example

of the bow and arrow, the archer, his bow, the arrow, the ground, the ambient air

and the target are all part of a system. The description implied that when the archer

contracted his muscles, all the chemical energy released by the cells was transferred

to his muscles, which in turn transferred all of this energy to the bow, arrow and

target. The same assumption applies to each energy conversion in the chain. We

therefore intuitively applied the law of conservation of energy. As we mentioned above,

this law states that in a closed system the total amount of energy remains constant.

However, because we wanted to keep the description simple, we omitted certain details.

For example, when the arrow falls, friction with the air produces a small amount of

heat that is transferred to the air. The rest of the arrow’s potential energy is converted

into kinetic energy. The law of conservation of energy still applies, but if we wanted

to be more precise, we would write:

Potential gravitational energy = Kinetic energy (arrow) + Heat (air)

Energy is conserved, since the two receptors, namely, the arrow and the air, absorb

all the potential energy. In fact, in real life situations, several receptors often share

the energy, even though only one of the receptors is relevant to the action being studied.

Exercise 1.5

Jehane Benoît describes the preparation of hard-boiled eggs as follows: “Place the

eggs in a saucepan and cover them with cold water. Place a lid on the pan and bring

the water to a boil over medium heat. As soon as the water starts to boil, remove the

saucepan from the source of heat and wait from 3 to 10 minutes before removing

the eggs from the water, depending on whether you like your eggs soft-boiled or hard-

boiled.”8


1.14

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8. Benoît, Jehane, La nouvelle encyclopédie de la cuisine (Montréal: Les messageries du Saint-Laurent Ltée, 1971),p. 283. (translation)

Using this recipe, illustrate the principle of conservation of energy by either writing

a description or drawing a diagram. Keep in mind all the receptors, including those

that are not directly involved in cooking the eggs.

1.2 HEAT

Heat, which is also called calorific energy, is one of the most common manifestations

of energy. Examples of heat are abundant: the Sun heats beaches, the electromagnetic

energy of microwaves heats and cooks our food, and so on. You may have already

noticed that the head of a nail is warm after it has been struck vigorously with a

hammer. If not, try it—but watch your fingers!

Now that you have felt this heat, how would you explain it? In other words, how does

the hammer’s kinetic energy heat the nail? Do you have any idea? Write it down in

the next exercise. We will come back to this exercise later and you will have a chance

to complete your answer then, if necessary.


1.15

Exercise 1.6

Try touching the head of a nail that you have just hit forcefully with a hammer. How

can you explain the heat emanating from the nail? In other words, what happened

inside the nail? Give an explanation based on what you currently know about energy.

In order to provide an adequate answer to the preceding exercise, we have to try to

imagine what happens inside the nail. But first we will review the model of matter

that we developed in the previous course.9

KINETIC MOLECULAR MODEL OF MATTER

First, let’s look again at the kinetic theory of gases which, as you may recall, describes

the model of an ideal gas. We will then apply this theory to all matter, whether solid,

liquid or gaseous. This model will help us form a mental picture of the fundamental

structure of matter, given that it cannot be observed directly.

At the molecular level, the structure of matter is invisible and cannot be observed

with an ordinary microscope. The purpose of the model is therefore to provide a

microscopic representation of matter. The model suggests images and types of

behaviours that provide an explanation for those properties of matter we can

perceive with our senses and measure with instruments. These observable properties

are said to be macroscopic. The term “microscopic” refers to invisible phenomena

and the behaviour that the model attempts to describe and explain. By contrast, the

term “macroscopic”refers to properties and phenomena that are normally visible to

the naked eye. This will be better understood if we consider the macroscopic and

microscopic views of a gas confined in a cylinder.


1.16

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9. Lalancette, Pauline and M. Lamoureux. Gases (Chemistry, Secondary V), Chapter 1. Learning Guide producedby SOFAD.

Figure 1.5 - Gas confined in a cylinder

a) The gas appears uniform and translucent.

b) The gas is composed of very small particles (atoms or molecules) that are separated from each other by great distances, that move freely in all directions and that do not attract

or repel one another. Note that this drawing is not to scale. A single millilitre of gas contains billions upon billions of particles.

Figure 1.5 shows a gas confined in a cylinder with an immobile piston in the top. At

the molecular level, the model provides the following explanation for what we see:

the gas particles are in motion and collide with the walls of the container and the

piston. The numerous collisions between the particles and the underside of the piston

prevent it from falling, even though its weight pulls it down.10

The model of an ideal gas, as it was described in the previous course in connection

with the kinetic theory of gases, can be summed up by the following hypotheses.

• All gases are composed of very small particles, either atoms or molecules, separated

by a vacuum. The distance between the particles is large compared to their size.

• The particles of a gas are independent, that is, they do not attract or repel one another.


1.17

a) Macroscopic view b) Microscopic view

10.Bernoulli (1728) held that the pressure exerted by a gas on the walls of its container is due to the billions of collisionsbetween the molecules of the confined gas and the walls (kinetic theory of gases). This view, which confirms Boyle’slaw on the compressibility of gases, also accounts for the fact that the temperature of a gas is related to the motionof its particles and therefore to its kinetic energy. From this point onwards, the foundation for the microscopicinterpretation of heat had been laid.

• The particles of a gas are in constant motion (translation, rotation and vibration).

They collide regularly with one another or with the walls of the container in which

they are confined.

• The average kinetic energy of the particles is a function of the temperature of a

gas. An increase in temperature will cause the molecules of the gas to become more

agitated. Conversely, a decrease in temperature will cause them to become less

agitated.

Matter exists not only in the gaseous state, but also in the liquid and solid states. As

a result, the hypotheses outlined above do not provide a satisfactory explanation for

the three states of matter at the microscopic level. If the ideal gas model states that

gas particles neither attract nor repel one another, how can we explain the solidity

of a candy or a stone?

It may be useful to look at an example of a solid to help us answer this question. Try

to imagine the structure of a penny at the microscopic level.

Figure 1.6 - A penny

a) Macroscopic view b) Microscopic view (to be completed)


1.18

Exercise 1.7

Take a penny and observe it carefully. It is made of copper. Complete Figure 1.6b by

drawing a microscopic representation of the penny as you imagine it.

The coin is a solid. It keeps its shape and it is difficult to cut. In your opinion, can

the model of the kinetic theory of gases explain these properties? Can your drawing

explain them? Let’s take a closer look.

Since the coin keeps its shape, we can conclude that the particles—the atoms of

copper—stick together and occupy fixed positions in relation to one another.

According to this description, although the atoms are agitated, their motion is limited

to vibration movements. Furthermore, it is difficult to cut the coin, which, on a

molecular level, means that it is difficult to separate the atoms. We can conclude from

this that an attractive force (cohesive force) is keeping the atoms together, acting

somewhat like a glue. This force explains why the atoms remain together and why

it is difficult to separate them. We can use the same explanation for other solids since

it accounts for the solidity of a stone just as well as that of a grain of sugar.

The model of a solid that we developed in the preceding paragraph is therefore not

consistent with all the hypotheses outlined earlier, which let us recall, are valid for

gases. The second hypothesis is particularly inexact. However, by reformulating these

hypotheses, we can obtain a more general model that describes the behaviour of gases,

liquids and solids. Because this model is based on the kinetic theory of gases, we will

refer to it as the “kinetic molecular model of matter.” It can be summed up by the

following hypotheses.

Kinetic Molecular Model of Matter

• All matter is composed of very small particles (atoms or molecules) separated by

a vacuum. In a gas, the particles are far apart from one another, whereas in a liquid

and in a solid, they are close to one another.

• There are forces of attraction between the particles of a substance. These forces

are negligible in gases but strong in liquids and solids.

• All the particles are in constant motion, be it translation, rotation, vibration, or a

combination of the three.


1.19

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• The average kinetic energy of the particles of a substance is a function of its

temperature. An increase in temperature will cause the particles to become more

agitated. Inversely, a decrease in temperature will cause them to become less agitated.

Before you continue, you may find it useful to compare these hypotheses with those

given for gases earlier. Most scientists today use this model of matter to explain many

common phenomena. You may want to refer to it often to help you understand what

cannot be observed by the naked eye. The table below compares the macroscopic

properties of gases, liquids and solids and the corresponding microscopic view provided

by the model. Take the time to study it carefully.

Figure 1.7 - The three states of matter: properties and model

MACROSCOPICGAS LIQUID SOLIDPROPERTIES

Shape Indefinite Indefinite Definite

Volume Indefinite Definite Definite

Compressibility High Negligible Negligible

MODELGAS LIQUID SOLID(microscopic view)

Diagram

Distance between molecules Large Very small Very small

Principal types of movements Vibration, rotation Vibration Vibrationand translation and rotation

Attractive forcesbetween the molecules

No Yes Yes

Order No No Yes


1.20

Exercise 1.8

Consider again the nail that heats up when it is struck with a hammer (Exercise 1.6).

A nail is made up of atoms of iron.

a) Use the kinetic molecular model of matter to explain the presence of heat in the

head of the nail.

b) Compare your answer with the answer you gave in Exercise 1.6.

THERMOMETERS AND HEAT TRANSFERS

The phenomena that involve a change in temperature are numerous. The nail that

is hit with a hammer is just one example. The nail heats up (its temperature rises)

because the hammer transferred a part of its energy to the nail. Consider a second

example: a drop of alcohol on your skin produces a sensation of coolness as it

evaporates. How can we explain this? The temperature of the skin’s surface decreases

because the skin provides the energy necessary to evaporate the alcohol. A transfer

of energy has occurred between the skin and the alcohol. In the following experimental

activity, you will prepare a series of mixtures and, for each one, you will observe whether

a change in temperature occurs. You will also determine the direction of the energy

transfer.


1.21

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Experimental Activity 1: Heat Transfers

In this first activity, you will observe a series of physical and

chemical changes. For each one, you will determine whether

or not a heat transfer has occurred. In most cases, you will

use a thermometer to detect the changes in temperature.

This is your first “hands-on” experience with the scientific

method in this second chemistry course. You will explore

different types of mixtures, compare the results obtained and interpret them while

remembering to take into account the thermometer’s degrees of accuracy. Allow

approximately 30 minutes to carry out all the steps in the experiment. Although the

steps are relatively simple, you must be meticulous in order to obtain meaningful results.

All of the information you need to carry out this activity is given in Section B of the

workbook Experimental Activities of Chemistry. Enjoy your work!

When a physical or chemical change causes the temperature of its surroundings to

increase, it is said to be exothermic. For instance, the dissolution of NaOH, which

you observed in the experimental activity, is exothermic. By releasing heat, it acted

as a source of energy and transferred heat to the solution (receptor).

Inversely, when a physical or chemical change produces a decrease in energy, energy

is absorbed from the surroundings or from an external source and the change is said

to be endothermic. For instance, ice melting is an endothermic change because energy

is absorbed in the process. In the experiment you conducted, the surrounding water

(source) provided the energy and the ice was the receptor. The temperature of the

water decreased because it gave up some of its energy to melt the ice.

In most of the changes studied in the activity, the thermometer was used to determine

whether a heat transfer had occurred and, if so, the direction of the energy flow. Have

you ever wondered how exactly a thermometer works? How does it measure temperature?

A thermometer is a common instrument used to determine whether a system is hotter

or colder than another by means of the temperature displayed on a scale. Remember

that temperature is associated with the kinetic energy of the molecules of a substance.

If the temperature rises, the kinetic energy of the molecules rises and, inversely, if the

temperature decreases, the kinetic energy also decreases. To better understand how a

thermometer works, let’s take a look at what happens at the microscopic level.


1.22

Mercury Thermometers

A mercury thermometer consists of a glass bulb filled with mercury attached to a

graduated tube whose inside diameter is as fine as a hair (capillary tube). Widely

used in the chemistry laboratory, thermometers are usually graduated in degrees

Celsius. According to this scale, water boils at 100°C, ice forms at 0°C and the

temperature of the human body is 37°C.

Figure 1.8b shows a microscopic view of a thermometer. To keep the diagram simple,

the solid walls of the bulb and of the capillary tube are represented as a single layer

of tightly packed glass particles firmly held together. The bulb contains liquid mercury

(Hg) whose atoms are less tightly packed than the glass molecules. The crowded

mercury atoms are in continuous motion, colliding with each other and with the

molecules that make up the glass wall. However, because the glass molecules are held

in place more firmly, they do not separate but vibrate in fixed positions, thus retaining

the atoms of Hg.

Figure 1.8 - Mercury thermometer

a) b)

a) Macroscopic view: the base of the thermometer consists of a bulb attached to a long graduated capillary tube.

b) Schematic diagram of a thermometer at the molecular level. To keep the diagram simple, only a few particles have been represented. In fact, a thermometer is made

up of billions upon billions of particles.


1.23

Capillary tube

Wall ofbulb

Mercuryatoms

Bulb

Immerse the bulb of a thermometer in a hot gas, such as the vapour escaping from

a saucepan of steaming vegetables, for instance. Be careful! The temperature of the

vapour can go up to 130°C. You may want to use a candy thermometer to conduct

this experiment.

Figure 1.9 - Thermometer immersed in a hot gas

a) b)

a) The bulb of the thermometer is immersed in the vapour escaping from the saucepan.

b) Bombarded by vapour molecules, the bulb of the thermometer heats up. The atoms of mercury (Hg) that it contains become more agitated and tend to occupy more space.

The liquid mercury expands and rises in the capillary tube attached to the bulb.

What happens at the microscopic level? The vapour molecules, which are very agitated,

bombard the walls of the bulb, transmitting a part of their kinetic energy to the glass.

The collisions cause the glass molecules to become agitated. In turn, the glass molecules

transmit energy to the Hg atoms. As they become more agitated, the distance between

the Hg atoms increases and, as a result, the mercury will try to occupy a greater volume.

We say that the mercury expands. The Hg atoms then move into the only available

space, the opening of the capillary tube, and we see the mercury rise. The height reached

by the mercury depends on the energy that has been transmitted to it.

Exercise 1.9

Now place the thermometer in the freezer. The thermometer’s bulb is now immersed

in cold air, whose molecules have a lower average kinetic energy than that of the

mercury atoms in the bulb.


1.24

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Describe what happens:

a) at the macroscopic level.

b) at the microscopic level.

What happens when the molecules in the thermometer reach the same level of agitation

as those in the medium in which the thermometer is immersed? This is a specific

case with very important implications. If we immerse a thermometer in a gas whose

molecules are moving at the same speed as the molecules in the thermometer’s bulb,

then the molecules in the thermometer and those in the ambient gas cannot

accelerate each other’s movements. In this case, there is no transfer of energy between

the gas and the bulb. The mercury level does not change and the thermometer displays

a constant temperature.

The manner in which the mercury moves inside the thermometer is of special

significance. When it rises, the bulb absorbs energy; when it drops, the bulb loses

energy. When it is stable, there is no transfer of energy from the thermometer’s bulb

to the outside. The temperature of the thermometer then equals the temperature of

its surroundings.


1.25

Exercise 1.10

A thermometer is placed in a bowl of water and, after a few minutes, the level of the

mercury stabilizes. Then it suddenly starts to climb. How would you interpret this?

Thermal Equilibrium

A thermometer that is immersed in a liquid for any length of time will display a constant

temperature. The temperature will remain constant provided the surrounding

conditions remain the same, that is, the water does not cool down or heat up. We

say that the thermometer is in thermal equilibrium with the liquid. In other words,

the thermometer and the liquid are at the same temperature. At the microscopic level,

the degree of molecular motion is, on average, the same in both the thermometer

and the liquid.

We can therefore use a thermometer to detect heat transfers, or transfers of kinetic

energy between molecules. If a system comes in contact with an energy source, the

mercury in the thermometer rises. Inversely, if the system gives up energy to a receptor,

then the temperature drops and the mercury contracts.

Temperature can be expressed according to various scales. The most commonly used

ones are the Celsius, Kelvin and Fahrenheit scales. Let’s review here some of the

highlights in the history of thermometers.

Temperature Scales

As we have seen, our skin detects sensations of hot and cold. These sensations are

therefore very familiar to us. The need to define these sensations and quantify them

more objectively led scientists such as Galileo (circa 1592), Torricelli (circa 1672) and

especially Fahrenheit (1714) and Celsius (1742) to determine fixed reference points

between which a numerical scale could be established.


1.26

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Delancé, in 1688, and Newton, in 1701, came up with the first scales for measuring

degrees of heat. For instance, Newton had arbitrarily set 0° as the point at which snow

melts, 12° as the temperature of the human body and 34° as the point at which water

boils vigorously.

Fahrenheit chose the coldest temperature obtained with a mixture of snow and

ammonia salt for the zero on his thermometer, and the boiling point of mercury, or

600°, for the highest point. He then divided the interval between these two points

into 600 equal divisions. On this scale, water freezes at 32°, water boils at 212° and

the temperature of the human body is 98.6°.

Today, Celsius’ reference points have been widely adopted and integrated into the

International System of Units (SI). The zero on the Celsius scale was obtained by

immersing the thermometer in ice water and the 100° mark was obtained by immersing

the thermometer in boiling water at standard atmospheric pressure (101.3 kPa).

Figure 1.10 - The most commonly used temperature scales

The reference points on the Celsius scale are the freezing and boiling points of water, set at 0°Cand 100°C respectively. On the Fahrenheit and Kelvin scales, the freezing point of water is 32°F

and 273 K respectively. Note that the Kelvin scale has only positive values.


1.27

373

273

0

212

32

−459

100

˚C ˚F K

0

−273Absolute 0All molecular

movement ceases.

Exercise 1.11

On my oven, I set the baking temperature for a raisin pie at 450°F. The red pilot light

came on right away and went off about five minutes later. I then opened the oven

door just long enough to put the pie in the oven. The red pilot light came on again

briefly and then went off. Ten minutes later, following Julia Child’s instructions, I

lowered the temperature to 350°F and continued to bake the pie for another

30 minutes. The pilot light did not come on again. The pie was delicious. As it was

piping hot, I had to wait 20 minutes before I could top it with a scoop of vanilla ice

cream and start eating it.

a) What is the purpose of the red pilot light?

b) Had we been able to read the oven thermometer, what would the mercury have

done during the procedure outlined in the problem above? Answer by completing

the following table.

PhasesPilot

EnergyMovement of mercury

light thermometer

Oven turned on, On Elements heating up Mercury rises.T = 450°F

After 5 minutes Éteint Les éléments cessent Le mercure est stablede chauffer.

Door opened Rouge Il y a perte de chaleur ; Le mercure monte.les éléments chauffent.

After 5 minutes Éteint Les éléments cessent Le mercure est stable.de chauffer.

T reduced to 350°F Éteint Les éléments Le mercure descend.ne chauffent pas

After 10 minutes Éteint Les éléments Le mercure descend ne chauffent pas. ou reste stable.


1.28

?

Caloric Fluid

Early scientists had formulated a theory to explain the propagation of heat. These men, including Antoine-

Laurent de Lavoisier (1789) and Sadi Carnot (1824), had imagined that heat was an invisible fluid that could

flow from one body to another. They named this fluid “caloric.” This is how they explained the heat and light

released during combustion.11 To measure heat, this elusive caloric was “broken down”into small units defined

as the quantity of caloric needed to raise the temperature of a given quantity of water by one degree. This

is how the unit of heat called the “calorie” originated.

THE MECHANICAL EQUIVALENT OF HEAT

While it seems natural to us today to consider heat as a form of energy, this relationship

was not so obvious in the early 19th century, when heat and energy were considered

to be two apparently unrelated phenomena. If you stop to think about it, how much

heat is needed to provide enough energy to raise a block of concrete onto a wall ten

metres high? Not so simple, is it?

James Prescott Joule (1818-1889) was the first to determine the number of units of

mechanical energy equivalent to one unit of heat (1 calorie). This relationship is called

the mechanical equivalent of heat. Joule established this correlation by comparing

the mechanical energy required to rotate paddles in a container of water with the

quantity of heat resulting from this action.

He defined his experiment after Count Rumford (Benjamin Thomson) had shown,

in 1798, that the drills used to bore cannons produced frictional heat that raised the

temperature of both the tube and the metal shavings. This is how Joule got the idea

of producing frictional heat by rotating paddles in water and measuring the resulting

rise in temperature. By definition, the calorie is the quantity of heat needed to raise

the temperature of one gram (1 g) of pure water by one degree (1°C) at standard

atmospheric pressure (101.3 kPa). Today, his name denotes one of the units used to

measure energy. For instance, we say that 4.18 joules of work are equivalent to one

calorie of heat.


1.29

11.Brock, William H. The Fontana History of Chemistry, New York: W. W. Norton (1993), p. 119.

Joule published the results of his work for the first time in 1843, but he had to wait

four years before his ideas were noticed. This occurred in Oxford, England, in August

1847. He was 28 years old. He had now obtained more precise measurements with

his paddle-wheel apparatus and wanted to make his results known to the British

Association. All the leading British scientists of the time belonged to this organization

and they periodically held meetings to discuss and exchange their ideas. The

president of the organization had given Joule a very short amount of time in which

to present his findings.

His work would have passed unnoticed if a young scientist named William Thomson,

who later came to be known as Lord Kelvin, had not agreed with Joule and started

a lively discussion. The ideas presented by Joule contradicted the caloric theory of

heat which was upheld by scientists at that time12 (see section entitled “Caloric Fluid”

above). From this point on, the scientific community started to take an interest in

Joule’s ideas, which are at the origin of the law of conservation of energy, considered

to be one of the cornerstones of contemporary science. The unit of work and energy

is named in his honour.

HEAT AND THERMAL ENERGY

Thermal energy and calorific energy (better known as heat) are examples of the two

views of matter we have discussed extensively so far. Heat is a macroscopic

phenomenon because it can be perceived by touch, while thermal energy is the sum

of the kinetic energies of all the particles in a system. It is therefore a microscopic

phenomenon. But what exactly is the link between heat and thermal energy? Can we

define heat in terms of thermal energy? Let’s find out.

When two systems at different temperatures come into contact, the warmer system

cools down and the cooler one heats up, until they both reach the same temperature.

In the process, a part of the thermal energy of the warmer system is transferred to

the cooler system. Today, science tells us that heat is the portion of thermal energy

that is exchanged between the two systems. The sensation of heat that we feel when

we touch a hot object is therefore produced by the thermal energy transferred to us

by the object. Inversely, when we touch an object that is colder than our body

temperature, the sensation of cold that we feel is due to the thermal energy we lose

to that object.


1.30

12.Boorse H., Motz L. Weaver. The Atomic Scientists, a Biographical History, New York: John Wiley & Sons, 1989,p. 61.

An analogy with the operation of a lock may help us distinguish between thermal

energy and heat. The transfer of heat between two systems at different temperatures

can be compared to the changes in water level in a lock (Figure 1.11). In this analogy,

the systems come into contact with each other when the sluice is opened, the water

that is exchanged represents the heat, and the total amount of water represents the

thermal energy. The different levels represent different temperatures.

Figure 1.11 - A boat making its way through a lock

A lock is a structure designed to move vessels from one elevation to another, either upstream ordownstream. A lock consists mainly of gates equipped with sluices that retain or let water outdepending on the direction in which the vessel is being moved. The lock chamber is the central

portion of the lock, and is located between the two gates.

a) A vessel is about to move downstream, that is, from the higher level to the lower level. The twogates and the sluices are closed. The upstream sluice is opened and gravity causes the water to

flow from the higher to the lower level. The lock chamber fills with water.

b) When the water in the lock chamber reaches the level of the water upstream, the gate is openedand the vessel enters the lock chamber. The gate and the upstream sluice are then closed and the

downstream sluice is opened. Gravity causes the water to flow from the lock chamber into the tail bay. The level of water in the lock chamber descends.

c) When the water in the lock chamber reaches the level downstream, the downstream gate isopened and the vessel moves into the tail bay. The vessel has then overcome the changes in

elevation and can continue on its course.


1.31

Lock chamber

a)

b)

c)

Tail bay

Head bay

Closed gate

Closed gate

Closed gate

Closed gate and sluice

Closed gateand sluice

Open gate and sluice

Open sluice

Closed sluice

Open sluice

As in the analogy, when two systems at different temperatures come into contact (open

sluice), a part of the thermal energy of the warmer system is transmitted to the molecules

of the cooler system until the levels of molecular motion are the same in both systems.

The kinetic energy that is transferred takes the form of heat. After the transfer, the two

systems are at the same temperature and have attained thermal equilibrium.

Kinetic energy, which results from the collisions between the molecules, is always

transferred from a warmer system to a cooler system. This idea can be expressed in

three different ways. For instance, we can say that heat flows from the warmer system

to the cooler system; from the system with the higher temperature to that with the

lower temperature; or from the system with the higher average kinetic energy to that

with the lower average kinetic energy.

The sensation of heat that we perceive results from the kinetic energy associated with

molecular motion and which is transferred from molecule to molecule and from atom

to atom as these particles collide. In short, a change in temperature indicates a transfer

of heat at the macroscopic level. At the microscopic level, however, this change indicates

a transfer of kinetic energy between the molecules or the atoms of the two systems

in question.

Exercise 1.12

How can you explain the fact that heat flows from a warm system to a cold system

and not the reverse?


1.32

?

1.3 HEAT EXCHANGES

When two systems at different temperatures come into contact, heat flows from the warmer

system to the cooler one until both reach thermal equilibrium, that is, the same temperature.

This is what happens when we mix a hot and a cold liquid together. After a while, the

temperature of the resulting mixture is uniform and falls somewhere between the initial

temperatures of the two liquids. Because heat is a form of energy, we can measure it and

assign numerical values to it. In the following experimental activity, you will study the

final temperatures of mixtures in order to derive a mathematical relationships for

determining the quantity of heat that is transferred from one liquid to another.

Experimental Activity 2: Final Temperature of a Mixture

In this activity, you will determine whether your predictions

about the final temperature of a mixture of hot and cold water

are true. You will then analyze the factors that affect the final

temperature of a mixture of liquids.

Allow approximately 50 minutes to conduct the experiment.

In order to obtain meaningful results, be meticulous in

carrying out the procedure. In this activity, you will also learn

more about experimental errors and uncertainty in measurements and become familiar

with writing instructions for the experimental procedure.

All of the information you require in order to carry out this activity is given in Section B

of the workbook Experimental Activities of Chemistry. Enjoy your work!

In the experimental activity, you derived an equation that explains heat transfers on

a macroscopic level. For the mixture of hot and cold water, you found that m1 × ΔT1

= –m2 × ΔT2. This equation confirms the law of conservation of energy. In fact, all

things being equal, the energy lost by the mass of hot water (m1) has been gained by

the mass of cold water (m1). In other words, the quantity of heat released by the source

(hot water) is equal to the quantity of heat gained by the receptor (cold water). If we

generalize the results of the experiment, we obtain the well-known relation that governs

heat exchanges and that we write as:

Qlost = Qgained or Qs = Qr,

where Q stands for heat, and the subscripts “s”and “r” for the source and receptor

respectively.


1.33

SPECIFIC HEAT CAPACITY

In the second part of Experimental Activity 2, you observed that windshield washer

fluid cools more rapidly than water when placed in the refrigerator. The nature of

the liquids involved is therefore a determining factor in heat transfers. More

generally, every substance may be characterized by its capacity to absorb or give up

heat. This property is called specific heat capacity13 and is designated by a

lowercase c. Specific heat capacity expresses the ratio between the heat Q supplied

to a system and the product m × ΔT. Mathematically, this is expressed as:

Qc = ––––––

m × ΔT

where Q is the quantity of heat gained or lost,

m, the mass of the substance,

ΔT, the change in temperature,

c, the heat capacity.

The table in Figure 1.12 gives the specific heat capacity of selected substances. Specific

heat capacity is expressed in terms of cal/g•°C or J/g•°C. The values are based on the

calculations done by engineering firms and are published in reference books which

can be consulted in science libraries.

You will note from the table that water has by far one of the highest specific heat

capacities on the planet. Only two gases found in relatively small amounts in the

atmosphere, hydrogen and helium, have a greater specific heat capacity than water.


1.34

13. Also called “heat capacity”

Figure 1.12 - Specific heat capacity of selected substances14

Substances Average specific heat capacity (c) Temperaturein the temperature interval interval

(ΔT) (ΔT)

(J/g•°C) (cal/g•°C) (°C)or

(kcal/kg•°C*)

Solids

Paraffin 2.97 0.710 0-20

Lumber, hard 2.93 0.700 10-60

Ice (water) 2.03 0.485 (−20)-(0)

Leather 1.50 0.360 0-100

Paper 1.30 0.310 20-60

Lumber, soft 1.25 0.300 10-60

Chalk 0.92 0.220 0-200

Clay, dry 0.92 0.220 20-100

Sandstone 0.90 0.215 0-100

Aluminum 0.89 0.214 0-700

Sand, dry 0.82 0.195 0-100

Glass 0.75 0.180 0-100

Nickel ($0.05) 0.46 0.111 0-100

Iron 0.45 0.107 20-100

Copper 0.39 0.094 20-1100

Tin 0.25 0.060 0-100

Tungsten 0.13 0.031 25

Liquids

Water (reference substance for the calorie) 4.19 1.000 0-100

Seawater 3.89 0.930 0-80

Methanol (also called “wood spirit”; highly poisonous) 2.51 0.600 20-25

Acetic acid 2.27 0.542 20-90

Mercury 0.14 0.033 0-100

Gases (constant pressure)

Hydrogen 14.18 3.392 20

Helium 5.23 1.250 20

Vapour (water) 1.90 0.455 20

Nitrogen 1.04 0.248 20

Air 0.99 0.238 20

* The units cal/g•°C and kcal/kg•°C are equivalent.


1.35

14. Based on Tuma, Jan J. Handbook of Physical Calculations. McGraw-Hill Book Company, 1976, pp. 301-303.

Exercise 1.13

a) Compare the specific heat capacity of ice, water and water vapour.

b) What requires more energy: heating a block of ice from –20°C to –10°C or heating

the same mass of water from 10°C to 20°C? Explain your answer.

Sea Breezes and Land Breezes

Along a shore, a wind can usually be felt blowing in from the sea in the daytime.

This is called a “sea breeze.” When the sun sets, the wind dies down and gradually

changes direction, blowing out to sea. This is called a “land breeze.” This phenomenon

is due to the considerable difference between the specific heat capacity of seawater

and that of the substances that make up the land mass (see the specific heat capacity

for sand and sandstone in the table).

While 0.82 J of solar heat are sufficient to raise the temperature of one gram of sand

by one degree, five times as much heat is needed to raise the temperature of one gram

of seawater by one degree. During the day, with the same amount of solar heat, seawater

heats up more slowly than rocks and sand. Land is therefore generally warmer than

the surface of the water. For this reason, the air over the shore is warmer than the

air over the water. Consequently, a current of warm air rises over the shore and a

current of cool air descends over the water. The combined effect of these two currents

creates a cool wind along the shore blowing in from the water. This is called a sea

breeze.


1.36

?

Figure 1.13 - Land breezes and sea breezes

a) Sea breeze b) Land breeze

In daytime, under the influence of the Sun’s rays, the air over the shore warms up more quicklythan the air over the water, producing a sea breeze. At night, the opposite phenomenon occurs. As

the seawater has stored more heat than the ground, a land breeze develops.

At night, the Sun disappears below the horizon. At this point, both the water and the

land begin to cool, but the water loses heat more slowly than the shore. The air over

the water therefore becomes warmer than the air over the land. Consequently, a

descending air current forms over the shore and a rising air current develops over

the sea. The combined effect of these currents produces a wind moving from the shore

out to the sea, or a land breeze.

HEAT EXCHANGE EQUATION

Recall the following equation that defines specific heat capacity:

Qc = ––––––

m × ΔT

If we rearrange this equation, we get the more familiar heat exchange equation shown

below. This equation means that the quantity of heat (Q) that flows from one system

to another is a function of the mass of the substance (m), its nature (c) and the change

in temperature observed (ΔT). We write:

Q = mcΔT = mc(Tf – Ti)

where Tf and Ti stand for the final and initial temperatures.


1.37

GWB

We can interpret this equation in three different ways.

• The quantity of heat (number of calories or joules) needed to raise the temperature

of a substance is proportional to the mass of the substance (Q ∝ m).

• The quantity of heat needed to heat a given mass of a substance is proportional

to the desired increase in temperature (Q ∝ ΔT).

• The ratio of the heat Q to the product (m × ΔT) is characteristic of the substance

and corresponds to c, or the specific heat capacity.

Combined with the law of conservation of energy, the heat exchange equation allows

us to calculate the amount of heat exchanged between two quantities of matter that

come into contact with one another. The operation of the calorimeter is based on

this equation. After taking a look at this instrument, we will consider other

applications of this equation.

Calorimeters

The transfer of heat between two substances is generally measured using a calorimeter.

This instrument is used mainly to determine the amount of energy released or absorbed

by a chemical reaction. The Styrofoam cup used in Experimental Activity 2 is a

simplified version of this instrument and is generally adequate for simple experiments

that do not require a great deal of precision.

Figure 1.14 - Calorimeter

A calorimeter consists of an insulated container at the centre of which is a reaction chamber. Thestirrer keeps the temperature of the water in the container uniform. The quantity of heat

transferred to the water (or absorbed from the water) during the reaction can bedetermined by measuring the temperature of the water before and after the reaction.


1.38

Stirrer

Thermometer

Water

Reactionchamber

A calorimeter consists essentially of an insulated container with a lid through which

a thermometer and a stirrer can be inserted. The container holds a known quantity

of water. The heat absorbed or lost by a reaction is measured by having the reaction

occur in a chamber immersed in the water in the calorimeter. The temperature of

the water is measured before the reaction and after, when the water and the reaction

chamber are in thermal equilibrium. In general, we assume that the thermal energy

lost or gained by the reaction remains in the calorimeter and we do not take into

account any heat lost through the openings or the spaces around the lid. Under these

conditions, all the heat released by the source is transferred to the water that serves

as the receptor. Of course, in special cases, another liquid can be substituted for water.

From now on and in all the following examples, we will assume that heat exchanges

occur within a perfect calorimeter. The systems will therefore be perfectly insulated

and losses will be considered negligible. Mathematically, this situation corresponds

to the equation we saw earlier on in this chapter: the quantity of heat released by

the source (Qs) is equal to that absorbed by the receptor (Qr).

Qs = Qr

If we replace each term in this equation with the terms in the heat exchange equation

and identify each symbol with the subscripts “s”and “r,” we get:

mscs ΔTs = mrcr ΔTr

If the two substances that come into contact with each other are of the same nature,

as in the case of hot and cold water, the specific heat capacity is the same on both

sides (cs = cr) and the equation is simplified. The equation is then the same as that

obtained in Experimental Activity 2:

ms ΔTs = mrΔTr which is equivalent to m1ΔT1 = –m2ΔT2.

Applications

This section includes several examples and a number of exercises. Take the time to

analyze the examples carefully, and try to understand and visualize the concepts which

underlie the equations. Equations represent a phenomenon or situation that you can

observe or that is described in a problem. Identify the source and the receptor, and

assign the appropriate subscripts to the quantities given. Check the values assigned

to each term in the chosen equation. Lastly, solve it by carrying out the necessary

algebraic operations and calculations.


1.39

Be careful to make the distinction between two concepts that are sometimes

confused with each other: heat and temperature. While heat refers to the amount of

energy that flows from one substance to another, a change in temperature indicates

the direction of the heat transfer. In short, remember this basic principle: heat is always

transferred from a warmer system (higher temperature) to a cooler system (lower

temperature). As the change in heat occurs, the temperature of the source drops while

that of the receptor rises.

Example

a) How much heat is needed to raise the temperature of a 100-g strip of copperfrom 20°C to 100°C? Give your answer in calories.

The heat source here is a given, and we assume that it provides all the heat required.The receptor is the strip of copper, whose specific heat capacity is given in thetable in Figure 1.12. The heat absorbed by the copper is proportional to its mass,and the rise in temperature depends on the specific heat capacity of copper. Theheat exchange equation can be used to describe what happens.

The mass of the strip of copper is 100 g, its initial temperature is 20°C and thefinal temperature is 100°C. The table gives c = 0.094 cal/g•°C for copper. Weare looking for Q. Let’s substitute the variables in the equation with the valuesin our problem. We get:

Q = mcΔTcal

Q = 100 g × 0.094 –––– × (100°C − 20°C)g•°C

Q = 752 cal

The heat source will therefore have to supply 752 calories to raise the temperatureof the strip of copper to 100°C.

b) The copper in a) is now replaced by a strip of aluminum. The aluminum has thesame mass, its initial temperature is the same and it absorbs the same amountof heat. Which of the two strips of metal is hotter? Answer first without doing anycalculations.

The strip of copper will be hotter. In fact, the specific heat capacity of aluminumis greater than the specific heat capacity of copper (see the table in Figure 1.12).


1.40

If the two strips of metal have the same mass, more heat is needed to raise thetemperature of the aluminum by one degree. The aluminum will therefore have alower final temperature.

c) Calculate the final temperature of the strip of aluminum.

The quantity of heat absorbed by the strip of aluminum (receptor) is the same asthat absorbed by the strip of copper, that is, 752 cal. Again, we can use the heatexchange equation. We are looking for the final temperature of the aluminum (Tf).We have m = 100 g, Ti = 20°C, c = 0.214 cal/g•°C and Q = 752 cal. By applyingthe heat exchange equation, we get:

Q = mcΔTQ

ΔT = ––––mc 752 cal × g•°C

ΔT = ––––––––––––––––100 g × 0.214 cal

ΔT = 35.1°C

ΔT = Tf − Ti

Tf = ΔT + Ti = 35.1°C + 20°C = 55.1°C

The temperature of the strip of aluminum will therefore be 55.1°C.

Exercise 1.14

How much heat is needed to raise the temperature of a 100-g strip of iron from 20°C

to 90°C? Give your answer in joules.


1.41

?

Example

Peter is about to wash the dishes and is preparing a basin of water. He pours 10 litresof hot water into the basin (T = 40°C) and then adds 1 litre of cold water (T = 12°C).What is the final temperature of the mixture? Remember that the mass of 1 litre ofwater is 1 000 g, since the density of water is 1 g/mL, or 1 kg/L.

The hot water is the source of heat and the cold water is the receptor. This situationis similar to the one you encountered in Experimental Activity 2, where you mixedknown quantities of hot and cold water. According to the law of conservation of energy:

Qs = Qr

mscsΔTs = mrcrΔTr

Since in this case water is both the source and the receptor, the value of c is thesame on both sides of the equation. In other words, cs = cr. By simplifying, we get:

msΔTs = mrΔTr (equivalent to m1ΔT1 = −m2ΔT2 )

The temperature of the hot water (source) will go from 40°C (Ti) to a certain finaltemperature (Tf). The mass is 10 kg since we have 10 L of hot water. The initialtemperature of the cold water (receptor) is 12°C and it will rise to the final temperature(Tf), which will be the same as that of the hot water since we are mixing them together.The mass of 1 L of cold water is 1 kg. We get:

msΔTs = mrΔTr (ΔT = Thigher – Tlower)10 kg × (40°C − Tf) = 1 kg × (Tf − 12°C)

By simplifying the kg, we get:10(40°C − Tf) = (Tf − 12°C)400°C − 10 Tf = Tf − 12°C

−11 Tf = −412°C Tf = 37°C

The temperature of the water prepared by Peter is therefore 37°C.


1.42

Example

How long will it take a 1 500-watt electric kettle to bring one litre of water to a boil?The water’s initial temperature is room temperature (20°C). Assume that 30% of theheat produced by the heating element is lost to the air.

First, determine the energy needed to bring the water to a boil. Remember that 1 Lof water has a mass of 1 kg (1 000 g). The heat exchange equation gives us:

Q = mcΔTJ

Q = 1 000 g × 4.18 –––––– × (100°C – 20°C)g•°C

Q = 3.344 × 105 J = 334.4 kJ

Therefore, 334.4 kJ of energy are needed to bring the water to a boil. However, only70% of the energy provided to the kettle goes towards heating the water. More energymust therefore be provided to the element. If we call the energy that goes towardsheating the water efficient energy, Eeff, and the energy that is dissipated by the element,Eel, we can write Eef f = 0.70 Eel since only 70% of the element’s energy serves toheat the water. We have:

Eef f = 0.70 Eel

Eef f 334.4 kJEel = –––––– = –––––––– = 477.7 kJ

0.70 0.70

A total of 477.7 kJ must be supplied to the kettle in order to bring the water to aboil.

Now, we will determine how long it will take to bring the water to a boil. Electric powerrepresents the amount of energy consumed by the kettle each second. One watt isequal to one joule per second (1 W = 1 J/s). The 1 500-W kettle therefore consumes1 500 J per second. We are now trying to calculate how much time it will take toprovide the kettle with the 477.7 kJ needed to bring the water to a boil.

Mathematically, we have:

EP = –––, where P stands for power, E for energy and t for time.

t


1.43

In our case, P = 1 500 W or 1 500 J/s, E = Eel = 477.7 kJ and t, the time we wantto calculate. We have:

EP = ––––

tEelP = ––––tEel 477.7 kJ 477.7 × 103 J × s

t = –––– = –––––––– = –––––––––––––––––– = 318 sP 1 500 W 1 500 J

mint = 318 s × ––––

60 st = 5.3 min

A little more than 5 minutes are required to boil one litre of water.

Exercise 1.15

An 800-g piece of glass has a temperature of 25°C. What will its temperature be after

it has absorbed 1 500 joules?

ENERGY IN PHASE CHANGES

So far, we have studied transfers of thermal energy from one system to another. Forexample, when a strip of hot copper immersed in a bowl of cold water causes thetemperature of the water to rise, it is an indication that a certain amount of energyhas flowed from the copper (source) to the water (receptor). We can calculate thisamount of energy by applying the law of conservation of energy and the heat exchangeequation.

But do these principles apply in all situations? What happens with regard to energywhen the water being heated starts to boil? Or when ice melts?

Let’s now consider what happens when a block of ice is heated continuously. Thissituation is illustrated by the graph of the temperature as a function of time, knownas the heating curve. Most of the curve can be obtained with a setup such as the oneshown in Figure 1.15.


1.44

?

Figure 1.15 - Heating a block of ice

After immersing a thermometer in a container of water, the entire apparatus is placed in thefreezer for some time. When it is removed, the water is frozen and the thermometer reads –20°C.

The container is then placed on a heating plate, and the temperature on the thermometer is recorded every minute.

Figure 1.16 - Heating curve for ice

(a) In the interval between –20°C and 0°C, the temperature rises and the ice absorbs heat from the source.

(b) Starting at 0°C, liquid water appears and the ice starts to melt. The temperature remains stable at 0°C until the ice is completely melted.

(c) If we continue to supply heat, the temperature will rise and the water will boil at 100°C.

(d) The temperature remains stable as long as some liquid water remains.

(e) If we recover the vapour and continue to supply heat, the temperature will start to rise again, as indicated by the last interval on the graph.


1.45

−20˚C

Block of ice

120

100

80

60

40

20

0

−20

Tem

pera

ture

(˚C

)

Time

Water vapour

Water and water vapour(vapourization)

Water

Ice andwater(melting)

Ice

(e)

(a)

(b)

(c)

(d)

The heating curve for water was covered in the first guide.15 This curve includes five

intervals and features two plateaus corresponding to the two phase changes. Let’s

see what happens in these plateaus at the microscopic level. First, we will analyze

the three increases in temperature, that is, intervals (a), (c) and (e) on the curve.

In interval (a), the temperature of the ice rises. The kinetic energy of the molecules

increases, causing their vibrations to become more intense.

In interval (c), the liquid water heats up. The kinetic energy of the molecules increases

and they become more agitated, causing the mercury in the thermometer to rise.

In interval (e), the temperature of the vapour increases and, consequently, so does

the kinetic energy of the molecules. Their movements become increasingly agitated.

Note that intervals (a), (c) and (e) have different slopes. In fact, the slope of the intervals

is a reflection of the capacity of the ice, water or vapour to absorb heat. Since liquid

water has the highest specific heat capacity, more energy is required to heat it and

its temperature rises more slowly. On the graph, the slope of the interval representing

water as a liquid is clearly less steep than the slopes obtained for ice and vapour.

Exercise 1.16

Why are burns caused by water vapour often more serious than those caused by boiling

water?

Melting

Now let’s examine the first plateau on the curve, or interval (b), in more detail. The

temperature remains stable as the ice gradually melts. The average kinetic energy of

the molecules remains the same (constant temperature) even though the heat

supplied is being absorbed. What happens exactly? If the block of ice is to remain a

solid, in other words, if it is to keep its shape, the forces of attraction between the

molecules must be sufficiently strong enough to keep them together, in fixed

positions in relation to each other. The case of liquid water is different because it


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15. Lalancette, Pauline and M. Lamoureux. Gases (Chemistry, Secondary V), Chapter 1. Learning Guide producedby SOFAD.

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takes the shape of its container. We can conclude from this that the heat provided

during the time the temperature remains constant serves to weaken the attractive

forces until the molecules no longer adhere firmly together in fixed positions. Since

the molecules no longer adhere together as strongly, they start to slide over one another

on the sides of the block of ice and only the walls of the container prevent them from

flowing further. The molecules now form a liquid.

It’s as though the energy from the outside source were “eaten up” by the ice. In more

scientific language, we say that the energy absorbed by the solid is transformed into

potential energy as the attractive forces between the molecules of the solid weaken.

If the water were cooled and once again solidified, the ice would release an amount

of energy equal to the energy that went into melting it, meaning that potential energy

would once again be converted to heat.

The heat of fusion of a substance is the energy required to turn it from a solid into

a liquid at its melting point. It is characteristic of the substance and is measured in

calories per gram (cal/g) or in joules per gram (J/g) of melted solid. This value is

determined experimentally and the heat of fusion of a large number of solids can be

found in specialized reference books and in some commercial periodic tables. The

heat of fusion of ice is 334 J/g or 80 cal/g under standard conditions (0°C and 101.3 kPa).

Exercise 1.17

What is the role of the energy absorbed by a melting solid?


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Boiling

Have you ever noticed that boiling water doesn’t get hotter even if it is heated longer?

The second plateau on the heating curve confirms this. The temperature is constant

in interval (d).

This phenomenon is similar to what happens when ice melts. When the water’s

temperature reaches 100°C, bubbles form in the water and the temperature remains

stable as long as there is liquid water remaining. The energy absorbed by the molecules

of water serves to separate them, making them independent of each other. They then

form a gas.

In the example of the heating plate, the heat source is in direct contact with the bottom

of the container, causing the water at the bottom of the container to heat up faster.

It is not surprising then that the first molecules to succeed in breaking away from

the liquid are those at the bottom of the container. These molecules form bubbles

which quickly rise to the surface, in the same way that a submerged cork or a ping-

pong ball floats to the surface when released. The bubbles burst when they reach the

surface of the water and release the “independent” molecules into the air. These then

form water vapour.

As soon as they are in air, the molecules collide with the molecules of nitrogen and

oxygen, which together make up 98% of air. They lose a part of their kinetic energy

to the air, and this slows them down. Some of the molecules adhere together, creating

minute drops of liquid and producing a white cloud. This is why the vapour that forms

above boiling water is white.

The heat of vapourization of a substance is the heat needed to vapourize a given

mass of the substance at its boiling point. It is specific to each substance and is

expressed in cal/g or in J/g. The heat of vapourization for water is 540 cal/g or 2 255 J/g.

Remember that we use the term “boiling” to refer to the rapid vapourization of a liquid

into gas (e.g. boiling of water, of liquid hydrogen) and “evaporation” to refer to the

slow vapourization of liquid into a gas (e.g. the evaporation of water, of alcohol).

The following table shows the typical heats of fusion and of vapourization of selected

substances. Note that in the case of water, the heat of vapourization is much higher

than the heat of fusion. This is why the vapourization plateau is longer than the melting

plateau on the heating curve for water.


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Figure 1.17 - Heats of fusion and of vapourization of selected substances

NameTfusion ΔΔHfusion* Tvapourization ΔΔHvapourization*

(°C) (J/g) (°C) (J/g)

Alcohol (methanol) 64.5 502

Aluminum 658 396 2 970 10 759

Copper 1 083 206 2 567 4 728

Water 0 333 100 2 255

Tungsten 3 410 192 5 660 2 296

*ΔHfusion stands for heat of fusion and ΔHvapourization stands for heat of vapourization.

Example

A 400-g strip of aluminum whose temperature is 25°C is heated to its melting point.Heat is then applied to it until it is completely melted. How much energy did thealuminum absorb?

Using the specific heat capacity for aluminum, we first calculate the energy neededto heat the aluminum and then the energy it absorbs as it melts. The total energyis obtained by adding these two values together.

Heating the solid

We have m = 400 g; c = 0.89 J/g•°C (table in Figure 1.12); Ti = 25°C and Tf = 658°C (table in Figure 1.17). We are looking for Q. Using the heat exchangeequation, we get:

Q = mcΔTJ

Q = 400 g × 0.89 ––––– × (658°C – 25°C)g•°C

Q = 225 348 J ou 225.3 kJ

225.3 kJ of energy are required to heat the aluminum to the melting point.


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Melting

We know, from the table, that the heat of fusion is 396 J/g. Using the property ofproportions, we have:

400 g × 396 J1 g → 396 J –––––––––––––– = 158 400 J400 g → ? J 1 g

158.4 kJ are required to melt the aluminum.

Total energy

Qtotal = Qheating + Qfusion

Qtotal = 225.3 kJ + 158.4 kJQtotal = 383.7 kJ

The aluminum absorbed a total of 383.7 kJ.

Exercise 1.18

A 200-g strip of copper whose temperature is 22°C is heated to its melting point. Then

it is heated again until it is completely melted. How much energy was supplied?


1.50

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1.4 TECHNICAL APPLICATIONS

THE BREAD OVEN

Readily perceived by the senses, heat is closely associated with human activity. Bread

ovens, whose origins go back to biblical times, are proof of this. Our ancestors built

their ovens with stones capable of storing large amounts of heat. A fire was allowed

to burn in the oven for a certain amount of time and, once the stones were hot, the

fire was put out and the wood was replaced with bread dough. The heat that had

accumulated in the stones then baked the bread. The same principle is used today

in restaurants that serve pizza cooked in a wood oven.

Figure 1.18 - Bread oven

Refractory stones have great thermal capacity. Once heated, they slowly release the stored heat

that serves to bake the bread.

The example of the bread oven is a good illustration of the usefulness of heat. Other

useful applications of heat include the teakettle and heating systems. However, in many

human activities, heat can be a useless, and even undesirable, by-product.

Exercise 1.19

a) List two or three examples where heat is useful for human activity.

b) List two or three examples where heat is undesirable and even useless.


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GWB

To conclude, we will focus on two spectacular examples that involve kinetic energy

and heat simultaneously.

RE-ENTERING THE ATMOPSHERE

In orbit, a spacecraft has an enormous amount of kinetic energy. Just think that it

circles the Earth in just 90 minutes and that it travels at a speed of over 25 000 km/h!

To land, the spacecraft must reduce its speed to about 250 km/hr in approximately

10 minutes. As its speed decreases, the spacecraft loses kinetic energy. However, this

energy is transformed into heat, which is the result of friction between the spacecraft

and the atmosphere. This heat must be re-channelled to prevent it from burning up

the passengers and disintegrating the vehicle. For this purpose, special insulating tiles

are used to line the bottom of the spacecraft. These tiles serve to re-channel part of

the heat but they also absorb most of it. In fact, the tiles contain a layer that melts,

thereby absorbing a great deal of heat without an increase in temperature. They have

a very low density and keep their properties up to temperatures of around 3 000°C.

They therefore form a very effective thermal shield.

Figure 1.19 - A space shuttle losing kinetic energy

When the space shuttle slows down, the lost kinetic energy is transformed into heat. The special tiles that line its surface form a heat shield that protects the crew

as well as the spacecraft and the equipment.

GEYSERS

Geysers are natural “machines” that periodically spout columns of water and vapour

up to 100 metres high. They draw energy from the Earth’s core where melting rocks

reach temperatures of several hundreds of degrees Celsius. The water that drains into

the ground settles in cavities hundreds of metres below the surface of the Earth and

in the channels that lead from these cavities to the surface (Figure 1.20). The water


1.52

that accumulates in the channel compresses the water below it. Because of the pressure,

the water boils at much higher temperatures than normal. After a certain time, the

pressure of the vapours formed at the bottom of the cavities becomes so high that it

succeeds in explosively ejecting a large part of the accumulated water and vapour. This

is how a geyser is formed. Geothermic energy, which comes from the heat released

by magma,16 is transformed into kinetic energy that projects the water outside its natural

well. Runoff water drains into the ground and the cycle begins again.

Figure 1.20 - A geyser

Hot water periodically erupts from the geyser. The water, heated by the surrounding rocks, isforced out with great intensity at more or less regular intervals. The water that shoots up often

contains sulphurous matter and mineral deposits.

In this chapter, we examined the topic of energy in greater detail. The kinetic

molecular model of matter enhanced our understanding of heat transfers and

the examples studied, all involved the concept of energy, along with the kinetic

molecular model of matter helped enhance our understanding of heat transfers.

This concept seems to be at the root of all explanations and of all physical and

chemical changes. Did we not define it as “the capacity of an object to produce

an effect”? Its importance will become even more apparent in the rest of this

course. In the next chapter, we will explore the phenomenon of dissolution by

concentrating on the types of energy involved.


1.53

Superheated

water T > 200°C

16. A very hot viscous liquid formed by the melting of the Earth’s crust or mantle and which, after being ejectedthrough fissures (or volcanoes), forms volcanic rock.

GWB

Calorie CalorimeterCohesive force Conservation of energy

Endothermic EnergyExothermic

Heat (calorific energy) Heat of fusionHeat of vapourization

Joule

Kinetic energy

Macroscopic Melting pointMicroscopic

Potential energy

Specific heat capacity (c)

Thermal energy Thermal equilibrium

Work


1.54

KEY WORDS IN THIS CHAPTER

Energy exists in many different forms, for example, electrical, nuclear, mechanical,

and so on. It can be stored just as it can be converted from one form to another. All

types of energy have one thing in common: the ability to produce an effect that can

be perceived, such as light, movement or heat. Heat is a form of energy associated

with numerous phenomena. Some, like combustion, produce heat and are said to be

exothermic. Others, such as ice melting and water evapourating, absorb heat and

are said to be endothermic.

All substances can store an amount of heat that is determined by their specific heatcapacity (c). The sum of the kinetic energies of the molecules in a system makes

up the thermal energy of the system. Temperature is an indication of the level of

this energy. Two systems that come into contact with each other tend to reach thermal

SUMMARY

equilibrium. Heat transfer always occurs in the same direction, that is, from the system

with the higher temperature (source) to the system with the lower temperature

(receptor). Heat is the sensation we perceive during the transfer of thermal energy.

Changes in heat are governed by the following equations:

Qs = Qr law of conservation of energy, and

Q = mcΔT heat exchange equation

The first equation means that the heat released by the source is absorbed by the

receptor. The second equation determines the amount of heat absorbed (Q) by a

substance with mass m and heat capacity c when its temperature is raised by a number

of degrees ΔT.

A calorimeter serves to measure the amount of energy released or absorbed during

a heat transfer or a chemical reaction. The Styrofoam cup used in the experimental

activities is a simplified version of this instrument.

The heating curve for a substance generally features two constant temperature plateaus

that reflect phase changes. The energy absorbed during melting and vapourization

serves to separate the molecules (or atoms) and it is stored in the form of potentialenergy. The same amount of energy is then released in the opposite phase changes

(melting and solidification). Between the plateaus on the curve, the slope of the intervals

indicates the heat capacity of the different phases.


1.55

Exercise 1.20

Each of the following descriptions corresponds to an instrument that detects,

converts or measures energy. From the list, choose the instrument that best

corresponds to the descriptions given in the table.

Television antenna, radio set, sonar, telescope, television set,

thermometer, voltmeter, wattmeter

Description Instrument Form of energy

A bulb containing a liquid and attached to a capillary tube which detects heat loss or heat gain in the surroundings

An instrument used to measure the magnitude of electric potential difference between two terminals electrique

A device similar to a radar, which uses sound waves to detect and locate underwater objects

énergie acoustique

A device used to capture electromagnetic waves emitted by a remote source telévision electromagnétique

A device used to decode and translate the waves from a television antenna into images electromagnétique

An optical instrument used to observe distant objectsTelescope Énergie lumineuse

A device used to decode and translate the waves from an antenna into sound

Poste de radio Énergie electrique

An instrument for measuring the magnitude of the electrical power in a circuit

Wattmètre Énergie electrique


1.56

REVIEW EXERCISES

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Exercise 1.21

a) Describe the energy conversions that take place when a spring is compressed and

then immediately released. Use the example of the bow and arrow in the “Energy

Conversions” section of this chapter to describe the sequence of conversions that

take place.

b) Discuss how the law of conservation of energy applies in this situation.

Exercise 1.22

Give a microscopic explanation for each of the following phenomena:

a) A wooden stair is gradually worn down.


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b) A small puddle of water evaporates in the Sun.

c) Two small drops of water brought close together form a single drop.

d) Water boiling in a Pyrex container forms small bubbles at the bottom of the

container.

e) Surface runoff water that seeps into a geyser’s underground channel boils at very

high temperatures, sometimes reaching 250°C.


1.58

Exercise 1.23

Indicate whether the following phenomena are endothermic or exothermic. For each

equation, include the heat term on the appropriate side.

a) The combustion of wax: wax(s) + O2(g) → CO2(g) + H2O(g)

b) The formation of ice: water(l) → water(s)

c) Identification test for hydrogen: 2 H2(g) + O2(g) → 2 H2O(g)

d) Lead melting: Pb(s) → Pb(l)

e) Heating a quantity of water: water at 37°C → water at 88°C

Exercise 1.24

Calculate the quantity of heat absorbed by 180 g of water if its temperature is raised

by 20°C. Give your answer in kJ.


1.59

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Exercise 1.25

A 500-g block of ice (T = 0°C) melts in a dish overnight. In the morning, the temperature

of the water is 14°C. Calculate the energy absorbed by the water during the night.

Exercise 1.26

Thirty litres of hot water at 40°C are mixed with 20 litres of cold water at 3°C. What

is the final temperature of the mixture?


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Exercise 1.27

A drop of methanol (wood spirit) is placed on your hand. It evaporates immediately,

producing a sensation of coolness.

a) How much energy did your skin supply to the methanol? The mass of the methanol

is 0.1 g and its initial temperature is 22°C.

b) Does most of the energy go into heating the methanol or in evaporating it?

Exercise 1.28

Aluminum’s molar heat of fusion is the amount of energy needed to melt one mole

of aluminum atoms once the aluminum has reached the melting point. It is expressed

in kJ/mol. Calculate the molar heat of fusion of aluminum.


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Exercise 1.29

In your own words, explain why the comparison of the molar heat of fusion of various

solids may be used to compare the cohesive forces between the molecules (or atoms)

of these solids.

Exercise 1.30

A nail is struck with a hammer. The nail becomes embedded in the plank and

immediately after being struck the head of the nail feels warm. How was the kinetic

energy transferred from the hammer to the nail?


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Exercise 1.31

Explain how a thermostat functions. Refer to a user’s manual or the dictionary.


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