CHE-5041-2 Learning Guide - SOFAD Reactions 1: Energy and Chemical Dynamics Chemical Reactions 2: Equilibrium and Oxidation—Reduction The three learning guides are complemented by

Gases CHE-5041-2

Learning Guide

GASES

CHE-5041-2LEARNING GUIDE

Gases is the first of the three learning guides for 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 complemented by the workbook entitled Experimental Activities

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

GASES

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

scolaires du Québec.

Production Coordinator: Jean-Simon Labrecque (SOFAD)

Production Coordinator: Mireille Moisan (First Edition)

Coordinator: Céline Tremblay (FormaScience)

Authors: Pauline Lalancette (Chapters 1 to 6)

Martin Lamoureux (Chapter 7 and self-

evaluation test)

Illustrators: Gail Weil Brenner (GWB)

Jean-Philippe Morin (JPM)

Content Revisors: Céline Tremblay (FormaScience)

(French Version)

Martin Lamoureux (French Version)

Hélène Leung (English Version)

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

Translator: Claudia de Fulviis

Linguistic Revisor: Kay Flanagan

Translation and Linguistic Revision Direction de la production en langue anglaise

(Chapters 1 to 4): Services à la communauté anglophone

Ministère del’Éducation

Proofreader: Gabriel Kabis

First Edition: October 2000

September 2005

© Société de formation à distance des commissions scolaires du Québec

All rights to 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 978-2-89493-191-2

TABLE OF CONTENTS

GENERAL INTRODUCTION

OVERVIEW ................................................................................................................... 0.12

HOW TO USE THIS LEARNING GUIDE ............................................................................. 0.12

Learning Activities ................................................................................................. 0.13

Exercises .............................................................................................................. 0.13

Self-evaluation Test ............................................................................................... 0.14

Appendices ........................................................................................................... 0.14

Materials .............................................................................................................. 0.14

CERTIFICATION ............................................................................................................. 0.15

INFORMATION FOR DISTANCE EDUCATION STUDENTS .................................................... 0.15

Work Pace ............................................................................................................ 0.15

Your Tutor ............................................................................................................. 0.15

Homework Assignments ........................................................................................ 0.16

GASES ......................................................................................................................... 0.17

CHAPITER 1 – MATTER IN ALL ITS FORMS .................................................................. 1.1

1.1 THE THREE STATES OF MATTER ............................................................................. 1.3

Definitions ............................................................................................................ 1.3

Atoms and Molecules (Review) .............................................................................. 1.5

The Atomic Model ........................................................................................... 1.5

Elements, Molecules and Chemical Formulas ................................................... 1.7

A Model for the Three States of Matter .................................................................. 1.9

Experimental Activity 1: Gases, Liquids and Solids .......................................... 1.9

Solids ............................................................................................................ 1.10

Liquids ........................................................................................................... 1.11

Gases ............................................................................................................ 1.13

The Kinetic Theory of Gases ................................................................................. 1.16

Ideal Gases? .................................................................................................. 1.17

1.2 HISTORY AND GASES: THE BIRTH OF MODERN CHEMISTRY ................................... 1.19

The Four Main Elements ....................................................................................... 1.19

The Late 18th Century: Chemistry Becomes an Experimental Science .................... 1.20

The Early 19th Century: the Atom and the Molecule ............................................... 1.23

1.3 DIFFUSION AND BROWNIAN MOTION ...................................................................... 1.25

Diffusion of Gases ................................................................................................ 1.25

Smells ........................................................................................................... 1.26

Odours and Toxicity ......................................................................................... 1.28

Speed of Diffusion .......................................................................................... 1.28

Diffusion in Liquids and Solids .............................................................................. 1.30

Brownian Motion ................................................................................................... 1.32

Gases - Table of Contents

0.5

1.4 PHASE CHANGES .................................................................................................. 1.33

Melting and Solidification ...................................................................................... 1.35

Vaporization and Liquefaction ................................................................................ 1.36

Sublimation and Crystallization ............................................................................. 1.38

The Heating Curve for Water ................................................................................. 1.39

Double Boilers ................................................................................................ 1.39

The Melting Point and the Boiling Point ............................................................ 1.40

Expansion and Contraction .............................................................................. 1.44

Thermometers ................................................................................................ 1.45

Pressure and Boiling Temperature ......................................................................... 1.46

1.5 TECHNICAL APPLICATIONS ..................................................................................... 1.48

Refrigerators, Freezers and Air Conditioners ........................................................... 1.48

CFC’s: The Two Sides of the Coin .......................................................................... 1.51

1.6 OTHER STATES OF MATTER .................................................................................... 1.52

Plasma ................................................................................................................. 1.52

Amorphous Solids and Liquid Crystal ..................................................................... 1.52

KEY WORDS IN THIS CHAPTER ..................................................................................... 1.55

SUMMARY .................................................................................................................... 1.55

REVIEW EXERCISES ...................................................................................................... 1.58

CHAPTER 2 – THE MANY USES OF GASES ................................................................... 2.1

2.1 THE PROTECTIVE AND LIFE-SUPPORTING PROPERTIES OF AIR ................................. 2.3

The Atmosphere ................................................................................................... 2.3

From the Stratosphere to the Ionosphere ......................................................... 2.4

Composition of the Atmosphere ....................................................................... 2.7

Respiration ........................................................................................................... 2.10

The Exchange of Gases ................................................................................... 2.10

Oxygen or Carbon Monoxide? ........................................................................... 2.13

Air Quality and Health ........................................................................................... 2.15

The Air Quality Index ....................................................................................... 2.16

Pollution and Toxic Effects ............................................................................... 2.16

2.2 NATURAL CYCLES ................................................................................................. 2.19

The Oxygen Cycle .................................................................................................. 2.20

Photosynthesis ............................................................................................... 2.21

The Carbon Cycle .................................................................................................. 2.23

The Water Cycle .................................................................................................... 2.28

Ozone: A Special Case .......................................................................................... 2.29

An Umbrella at 25 Km of Altitude .................................................................... 2.30

Ozone: A Pollutant at Low Altitudes .................................................................. 2.33


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2.3 THE ROLE OF AIR IN TECHNOLOGY ........................................................................ 2.34

Air Transportation ................................................................................................. 2.34

Hot-air Balloons .............................................................................................. 2.34

The Dirigible ................................................................................................... 2.37

The Airplane ................................................................................................... 2.38

Underwater Diving ................................................................................................ 2.40

Preservation in Museums ..................................................................................... 2.43

2.4 OTHER USEFUL GASES ......................................................................................... 2.44

Anesthetizing Gases ............................................................................................. 2.44

Oxygen ........................................................................................................... 2.47

Nitrous Oxide .................................................................................................. 2.47

Energy-producing Gases ........................................................................................ 2.48

Natural Gas .................................................................................................... 2.48

Hydrogen ........................................................................................................ 2.51

2.5 GASEOUS POLLUTANTS ......................................................................................... 2.52

Carbon Monoxide .................................................................................................. 2.54

Hydrocarbons ....................................................................................................... 2.55

Sulphur Dioxide .................................................................................................... 2.55

Nitrogen Oxides .................................................................................................... 2.56


SUMMARY .................................................................................................................... 2.58


CHAPTER 3 – PRESSURE AND VOLUME ...................................................................... 3.1

3.1 VARIATIONS IN VOLUME AND PRESSURE ................................................................ 3.3

Volume and Its Variations ...................................................................................... 3.3

Amount of Gas (Number of Moles) ................................................................... 3.5

Temperature ................................................................................................... 3.6

Pressure ........................................................................................................ 3.6

Pressure and Its Variations ................................................................................... 3.8

Definition and Units of Pressure ...................................................................... 3.8

Temperature ................................................................................................... 3.14

Amount of Gas (Number of Moles) ................................................................... 3.14

Volume ........................................................................................................... 3.15

Pressure and Volume in Competition? ................................................................... 3.17

Movable Piston ............................................................................................... 3.17

Immobile Piston .............................................................................................. 3.18

Other Situations .............................................................................................. 3.19


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3.2 APPLICATIONS OF PRESSURE IN EVERYDAY LIFE ...................................................... 3.21

Pressure in Some Common Products ..................................................................... 3.21

Beer and Soft-drink Delivery Systems ............................................................... 3.21

Aerosol Cans .................................................................................................. 3.23

Pressure in the Air ................................................................................................ 3.25

Atmospheric Pressure ..................................................................................... 3.25

Pressure and Respiration ................................................................................ 3.27

Meteorology .................................................................................................... 3.30

3.3 MEASURING PRESSURE ........................................................................................ 3.35

Barometer ............................................................................................................ 3.35

Manometer ........................................................................................................... 3.40

3.4 BOYLE’S LAW ....................................................................................................... 3.44

Experimental Activity 2: Boyle’s Law ................................................................ 3.44


SUMMARY .................................................................................................................... 3.51


CHAPTER 4 –– VOLUME AND TEMPERATURE ................................................................ 4.1

4.1 CHARLES’ LAW ..................................................................................................... 4.3

Relationship Between Volume and Temperature (°C) ............................................... 4.4

Experimental Activity 3: Charles’ Law ............................................................. 4.5

Absolute Zero and the Kelvin Scale ....................................................................... 4.7

Statement and Applications of the Law .................................................................. 4.13

4.2 TEMPERATURE ...................................................................................................... 4.17

Temperature and Energy ....................................................................................... 4.17

Temperature and Pressure .................................................................................... 4.20

Temperature Scales .............................................................................................. 4.24

Thermometers: Yesterday and Today ...................................................................... 4.25


SUMMARY .................................................................................................................... 4.31


CHAPTER 5 –– VOLUME AND NUMBER OF MOLES ........................................................ 5.1

5.1 RELATIONSHIP BETWEEN THE VOLUME AND THE NUMBER OF MOLES OF A GAS .... 5.3

Experimental Analysis ........................................................................................... 5.5

Experimental Activity 4: Number of Moles and Volume ..................................... 5.5

Electrolysis ........................................................................................................... 5.6

Experimental Activity 5: Electrolysis of Water .................................................. 5.9

The Law and Its Applications ................................................................................. 5.9


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5.2 AVOGADRO’S LAW ................................................................................................. 5.16

The Story of Avogadro’s Law ................................................................................. 5.18

Molar Volume ....................................................................................................... 5.23

Density ................................................................................................................. 5.27

Carbon Dioxide Extinguishers ........................................................................... 5.29

Dirigibles ........................................................................................................ 5.30

Relative Density .............................................................................................. 5.32


SUMMARY .................................................................................................................... 5.33


CHAPTER 6 – GENERAL BEHAVIOUR OF GASES ........................................................... 6.1

6.1 IDEAL GAS LAW .................................................................................................... 6.3

The Law ............................................................................................................... 6.3

Applications .......................................................................................................... 6.9

Three Laws in One ................................................................................................ 6.13

Real or Ideal Gas? ................................................................................................ 6.16

6.2 OTHER APPLICATIONS OF THE IDEAL GAS LAW ....................................................... 6.18

Identifying a Gas ................................................................................................... 6.18

Problems of a Technical Nature ............................................................................. 6.21

6.3 CHEMISTRY THROUGH THE AGES .......................................................................... 6.23

From Antiquity to the Middle Ages ......................................................................... 6.23

Transition from Alchemy to Chemistry .................................................................... 6.25

Chemistry, a Modern Science ................................................................................ 6.27

6.4 DALTON’S LAW ...................................................................................................... 6.32

Experimental Activity 6: Law of Partial Pressures ............................................ 6.32


SUMMARY .................................................................................................................... 6.36


CHAPTER 77 –– REACTIONS INVOLVING GASES .............................................................. 7.1

7.1 FROM THE ATOM TO THE MOLECULE (Review) ........................................................ 7.3

The Structure of Matter ......................................................................................... 7.4

Simple Diatomic Gases ......................................................................................... 7.11

Halogens ........................................................................................................ 7.11

The Oxygen Molecule (O2) ................................................................................ 7.12

The Nitrogen Molecule (N2) .............................................................................. 7.12

Compound Gases ................................................................................................. 7.13

Halogen Compounds ....................................................................................... 7.13


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Oxygen Compounds ........................................................................................ 7.15

Nitrogen Compounds ....................................................................................... 7.16

Carbon Compounds ........................................................................................ 7.18

Activity 7.1: Carbon Compounds ....................................................................... 7.18

7.2 ENERGY BALANCE OF CHEMICAL REACTIONS ......................................................... 7.22

Bond Energy ......................................................................................................... 7.25

Calculating the Energy Balance ............................................................................. 7.35


SUMMARY .................................................................................................................... 7.40


CONCLUSION .............................................................................................................. C.1

SELF-EVALUATION TEST ................................................................................................ C.4

ANSWER KEY

CHAPTER 1 ........................................................................................................... C.19

CHAPTER 2 ........................................................................................................... C.28

CHAPTER 3 ........................................................................................................... C.36

CHAPTER 4 ........................................................................................................... C.47

CHAPTER 5 ........................................................................................................... C.59

CHAPTER 6 ........................................................................................................... C.74

CHAPTER 7 ........................................................................................................... C.93

ANSWER KEY TO THE SELF-EVALUATION TEST ............................................................... C.109

APPENDICES

APPENDIX A: THE INTERNATIONAL SYSTEM OF UNITS (SI) ...................................... C.117

Symbols of Quantity and Their Units ................................................................. C.117

Multiples and Submultiples of SI Units ............................................................. C.117

APPENDIX B: MATHEMATICAL PREREQUISITES ....................................................... C.119

Ratios and Proportions ................................................................................... C.119

Formulas ........................................................................................................ C.120

APPENDIX C: CHEMICAL PREREQUISITES .............................................................. C.122

Balancing Equations ........................................................................................ C.122

Calculating Molar Mass ................................................................................... C.125

APPENDIX D: List of Figures .................................................................................. C.127

BIBLIOGRAPHY ............................................................................................................. C.131

GLOSSARY ................................................................................................................... C.135

INDEX .......................................................................................................................... C.147


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GENERAL INTRODUCTION

OVERVIEW

Welcome to the course entitled Gases, 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 areas of focus of the Chemistry program are related content, the

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

method is developed in the workbook entitled Experimental Activities of Chemistry,

whereas the related content and the history-technology-society perspective are

covered in the three Learning Guides that complete the three courses which must

be taken in sequential order.

The Gases Learning Guide is the first in the set of three. It is divided into seven chapters,

which correspond to seven terminal objectives in the program.1 This Guide is to be

used together with Experimental Activities of Chemistry. You will find references to

the latter at appropriate times throughout this Guide.

The purpose of this course is to help you expand your knowledge of gases and establish

links between this knowledge and technical aspects, social changes and environmental

consequences of gases and their uses.

HOW TO USE THIS LEARNING GUIDE

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

specific needs of adult students in individualized learning programs, or who are enrolled

in 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. Each chapter ends with a list of key words, a summary and review exercises.

Gases - General Introduction

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

The conclusion of the Guide summarizes all the courses in the program and contains

a self-evaluation test. The conclusion also includes an Answer Key for the self-evaluation

test, for the exercises in each chapter and for the review exercises. A glossary containing

definitions of the key words, a bibliography, appendices and an index are also found

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

This Guide includes theoretical sections as well as practical activities in the form of

exercises. These exercises come with an Answer Key.

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

content and 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 found in the coloured section at the end of

this 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 a chapter before doing the corresponding review exercises. Doing these

exercises without referring to the lesson you have just completed is a better way

of preparing 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 this 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, a notebook where 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 entitled Experimental

Activities of Chemistry.

• Reference material: a dictionary.

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

notes in your Guide, a coloured pencil 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.

Evaluation for the Gases course 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 80% of your final mark and deals

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

Part II deals exclusively with evaluation of the experimental method. This second part

consists of a 90-minute written examination and does not require your presence in a

laboratory. It is worth 20% of your final mark and deals with the course objectives

covered in Section A of the experiment kit entitled Experimental Activities of

Chemistry.

INFORMATION FOR DISTANCE EDUCATION STUDENTS

Work Pace

Here are some tips that will help you in 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


0.15

her available hours. If necessary, write to him or her. 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 5, and the third after

completing Chapter 7. 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 the right 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 your 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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GASES

Flowers, smoke, water, rock—nature’s diversity is truly astonishing. From seed to stem,

to flower and to fruit, or from table leftovers to compost, matter undergoes

continuous transformation. Curious about this endless cycle of change, human beings

have observed and wondered about the nature of this varied and ever-changing mystery.

Why does our breath, which is invisible in the summer, have a whitish appearance

in the winter? How does an apple seed produce a tree? Where do the materials that

make up a tree’s trunk, branches and leaves come from? Why do cars rust? Why does

cooking change the colour and flavour of foods? How do we make plastic from

petroleum? What knowledge is hidden behind the magic of fireworks? Chemistry

attempts to answer these questions and many more. It is the study of the properties,

composition and transformation of matter.

Such variety calls on us to classify the many transformations observed in nature. Water

becomes ice or vapor, while its composition remains the same, and this change from

one form of matter to another is reversible. In this case, we speak of phase changes

from one state to another, the three states being gas, liquid and solid. Wood turns to

ash and smoke when burned; this change is profound and irreversible. In this case,

we speak of a chemical reaction.

Changes in matter involve energy. For instance, heat is needed to melt ice; a wood

fire produces heat and light. Movement, light and heat are the main manifestations

of energy. Not all chemical reactions occur at the same rate. Fire takes very little time

to destroy a tree that nature took so many years to build!

Substances are divided into categories. Depending on their properties, such as smell,

colour or reaction with the air or with a metal, they are said to be an acid, a base or

a salt, or classified as a mineral or organic compound or categorized according to

other systems of classification. Chemical reactions involving acids and bases produce

neutral solutions, another type of reaction is called oxidation-reduction, and still others

have different names. When substances and reactions are classified, it helps us learn

more about the organization of matter.

Through their keen curiosity and dedicated and meticulous work, chemists have

discovered some of the secrets of matter. The classification of substances and reactions

has revealed similarities and reduced the complexity of matter to its fundamental

building block, the atom. Consisting of a nucleus surrounded by a cloud of electrons,


0.17

the atom of an element is distinguished from those of other elements by the number

of protons in its nucleus. The diversity we see around us is the result of combinations

of just 112 elements, which constitute the “alphabet” of matter and combine to form

all the different substances we observe, in the same way that letters are combined to

form the words of a language. Molecules are the “words” of chemistry and its language

is that of chemical formulas.

Chemical reactions will be the main focus of study in the second and third courses,

respectively entitled Chemical Reactions 1: Energy and Chemical Dynamics and Chemical

Reactions 2: Equilibrium and Oxidation-Reduction. This course, entitled Gases, is the

first in the series. It examines gases in a broad context and uses them to introduce

chemical reactions.

The first chapter defines gases in terms of their properties, by comparing them with

liquids and solids. A model is used to describe the three states of matter at the molecular

level. The phase change from one physical state to another is reversible and does not

alter the nature of the substances. In the case of water, the molecules are the same

whether they are in the form of ice, water or water vapor. In a phase change, the

molecules move either further away or closer to one other and arrange themselves

differently.

In the second chapter, we will concentrate on gases, which are everywhere around

us. For instance, the atmosphere forms a protective gaseous layer around the planet.

The composition of the atmosphere is maintained by the cycles of generous and

abundant nature. Humans have learned to use gases productively in such applications

as anesthesia, medical treatment, heating systems, pneumatic machinery and space

rockets, among others. The list of the uses of gases, whether natural or synthetic, is

long and impressive. By contrast, industrial processes and the use of internal-

combustion engines release gaseous by-products that pollute the atmosphere.

The four subsequent chapters (Chapters 3 to 6) form the core of the course. They examine

the physical properties of gas samples. Whether air is compressed or hot or cold its

characteristics can be quantified in terms of pressure, volume, temperature or mass,

and by determining the number of moles. The way these factors act on each other

is governed by an equation that is commonly called the “ideal gas law.” The study of

this law and its application call for experimentation and tools such as mathematical

equations, units of measure and graphs.


0.18

Compressing a gas or cooling it does not change the nature of its molecules; the

molecules of a gas change, however, and new substances are produced when it is

involved in a chemical reaction. Some molecules are destroyed and new ones are

formed; atoms are rearranged and the chemical formulas change.

The seventh and last chapter of this course deals with the chemical properties of gases.

The formation of gas molecules, their chemical composition and the energy involved

in chemical reactions are all examined. The practical situations analyzed in this chapter

call into play all of the subject matter covered. The content of the last chapter also

links the study of gases, which is the main focus of this course, with chemical reactions,

which form the central theme of the next course.

A table of contents diagram at the beginning of each chapter shows you where 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 larger bold type and the content of Chapter 1 is in

italics. You will find that this diagram is a very useful tool as you go through the course.

Good luck!

1. Matter in All Its FormsThe Three States of MatterHistory and Gases: the Birth of

Modern ChemistryDiffusion and Brownian MotionPhase ChangesTechnical Applications

2. The Many Uses of GasesThe Protective and Life-supporting

Properties of AirThe Natural CyclesAir and TechnologyUseful GasesGaseous Pollutants

7. Reactions Involving GasesFrom the Atom to the MoleculeBond EnergyEnergy Balance

6. General Behaviour of GasesIdeal Gas LawApplicationsHistory of ChemistryDalton’s Law

GASES

5. Volume and Number of MolesElectrolysis of WaterHistory of Avogadro’s LawMolar VolumeDensity

4. Volume and TemperatureCharles’ LawTemperature and EnergyTemperature and PressureMeasurement of Temperature

3. Pressure and VolumeFactors Affecting VolumeFactors Affecting PressurePhenomena Involving PressureMeasurement of PressureBoyle’s Law


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CHAPTER 1

MATTER IN ALL ITS FORMS

GWB

Terminal Objective 1

At the end of this chapter, you will be able to explain the properties of the three main states of matter andthe general behaviour of substances undergoing phase changes.

Intermediate Objectives

1.1 To compare the observable properties of the solid, liquid and gaseous states of matter.

1.2 To explain the properties of the three main states of matter, using a model.

1.3 To describe molecular motion in the three states of matter.

1.4 To explain the phenomena of diffusion and Brownian motion, using a model.

1.5 To compare the rate of diffusion of a substance in a liquid and in a gas, as well as in two differentgases.

1.6 To describe phase changes, using examples and a model describing the three states of matter.

1.7 To define “melting point” and “boiling point.”

1.8 To compare the melting and boiling points of various substances in relation to their states and to agiven temperature.

1.9 To describe a technical process that relies on a phase change, using examples.

1.10 To give an example of a state of matter other than solid, liquid or gas.

1. Matter in All Its FormsThe Three States of MatterHistory and Gases: the Birth of

Modern ChemistryDiffusion and Brownian MotionPhase ChangesTechnical Applications

2. The Many Uses of GasesThe Protective and Life-supporting

Properties of AirThe Natural CyclesAir and TechnologyUseful GasesGaseous Pollutants

7. Reactions Involving GasesFrom the Atom to the MoleculeBond EnergyEnergy Balance

6. General Behaviour of GasesIdeal Gas LawApplicationsHistory of ChemistryDalton’s Law

GASES

5. Volume and Number of MolesElectrolysis of WaterHistory of Avogadro’s LawMolar VolumeDensity

4. Volume and TemperatureCharles’ LawTemperature and EnergyTemperature and PressureMeasurement of Temperature

3. Pressure and VolumeFactors Affecting VolumeFactors Affecting PressurePhenomena Involving PressureMeasurement of PressureBoyle’s Law

Gases - Chapter 1: Matter in All Its Forms

1.2

The diversity of the matter found in the world around us is truly astonishing. For

instance, the wool used in a sweater, the bricks in the walls of a building, the water

we drink and the air we breathe, seem to have nothing in common. The same substance

can change form depending on the temperature and amount of pressure to which it

is subjected. For example, water freezes when the thermometer drops below zero and

air changes into a liquid when it is highly compressed. Endlessly varied and ever-

changing, matter is indeed a rich and complex subject of study.

In this chapter, we will first examine the three principal states1 of matter (solid, liquidand gas) and, using a model, we will analyze how these states differ at the molecular

level. We will then look at phase changes and at a few technical applications that make

use of these changes.

1.1 THE THREE STATES OF MATTER

What do morning dew, steam rising from a boiling kettle, humidity in your washroom

after you take a shower, and ice cubes in a drink all have in common? The answer

is simple. They are all made of water. Indeed, water can take many different forms

and each has a different name—for example, ice, hail, frost, dew, snow, humidity and

steam, to name but a few.

DEFINITIONS

Like water, almost all of the substances found in the environment can exist in three

distinct states: solid, liquid and gas. However, most substances are most familiar to

us in the state they assume at room temperature. For example, at 25°C, lighter fluid

is a liquid, steel is a solid and helium is a gas.

Exercise 1.1

You are familiar with objects or substances that exist in each of the three states of

matter. Complete the following exercise by giving two additional examples for each

state.

Solid state:

Steel,


1.3

1. Words appearing in boldface in the text are defined in the glossary at the back of the guide.

?

Liquid state:

Lighter fluid,

Gaseous state:

Helium,

We can usually tell at a glance whether a substance is a solid, a liquid or a gas. However,

it is not as simple to define the criteria for differentiating among these three states.

For example, what is the basis for saying that air is a gas? Which criteria allow us

to say for sure that the water in a lake is a liquid? Of course, we can tell by the

appearance of these substances, but let’s try to answer these questions more precisely

by defining the three states of matter.

Exercise 1.2

Define the following terms in your own words.

A solid:

A liquid:

A gas:

In defining these terms, you may have mentioned the hardness of a solid, the lightness

of a gas and the fact that a liquid flows. The dictionary defines a solid as having a

definite shape and volume; a liquid as having a definite volume, but no definite shape;

and a gas as having no definite shape and as being compressible and expandable,

meaning that its volume decreases as pressure is applied to it and that it increases

if the available space increases.

While your definitions may not be the same as the actual dictionary definitions, you

have followed the same process by describing the different states according to their

specific properties (e.g. hardness, shape, compressibility, expansibility).


1.4

?

Technology makes frequent use of these properties. For example, solids are used in

construction because they resist compression, bridges are generally made of metal

and houses of brick and wood. By contrast, your BBQ runs on propane, a gas kept

under pressure in a metal tank. But why can a gas such as propane be compressed

whereas solids resist compression? What distinguishes these two forms of matter?

The answer to this question requires a more in-depth study of matter, since the

behaviour of matter is the result of what happens at the molecular level.

ATOMS AND MOLECULES (REVIEW)

You have defined the states of matter and noted that the shape of a substance varies

from one state to the next. For example, ice has a definite shape whereas liquid water

takes the shape of its container. Yet, regardless of the state, the water molecules remain

the same. In order to understand what distinguishes the different states, we must

examine how molecules are organized in a gas, in a liquid and in a solid. We will use

a model to do this.

Scientists often use models to represent what cannot be seen with the naked eye. Models

are three-dimensional constructions, images or diagrams that are used to simplify

the description of a concept. Models change over time, in the light of new discoveries.

The atomic model, with which you are already familiar, is one such example. It

describes the structure of atoms, which form the basis of all matter, whether solid,

liquid or gaseous. Before we take a closer look at the differences between the three

states of matter, let us consider what they have in common. We will start by reviewing

how the atomic model has changed through the ages and recalling some useful

concepts.

The Atomic Model

In ancient times, several centuries before our era, certain Greek philosophers

suggested that matter was made up of small indestructible particles, or atoms. This

theory was soon discarded, to be revived only two thousand years later, in the late

15th century.

A few centuries later, around 1800, John Dalton maintained that the atoms of a given

element were identical. At the end of the 19th century, Sir Joseph Thomson linked

the concept of electrical charge with matter and described the atom as consisting of

both positive and negative charges (Figure 1.1). About ten years later, Ernest

Rutherford established the nuclear model of the atom and formulated the existence


1.5

of a positive nucleus made up of protons around which negative electrons travel,

similar to the way the planets revolve around the Sun. He also concluded that the

numbers of protons and electrons were equal. Other experiments soon revealed that

electrons revolve around the atom’s nucleus in well-defined levels, and that the nucleus

contains neutrons. The atomic model continued to evolve over the years and has

become very complex.

In summary, the current atomic model, while still evolving, is the result of research

done over a period of 2 400 years; however, the atom has yet to give up all of its secrets

and continues to be the subject of intense study.

Figure 1.1 - History of the atomic model

Date Inventor Concept Innovative Modelof the model of matter aspect

of the model

1808 Dalton Matter is Atomcomposed ofindivisible particlescalled atoms.

1902 Thomson The atom is Electron divisible. Matter (negativeconducts electricity. charge)Electrons are negatively chargedparticles.

1911 Rutherford First nuclear model Nucleusof the atom: a dense (protons)nucleus is composedof protons; the electronstravel around the nucleus;apart from the nucleus, most of the atom consistsof empty space.

1913 Bohr Electrons are Energy levelsarranged in specific positions of electronson the shells (energy levels) around the nucleus.

1932 Chadwick In addition to protons, Neutronsthe nucleus containsneutrons.

+ ++ +-

-

-

-

+ ++ +

-- -

-

-

-

-

-+

+

++

-

-

-

-

+++ +


1.6

par ticles atoms

Electrons nucleus

protons energy

levels neutrons

The complex and precise atomic model used by scientists today provides a good

description of how atoms behave. However, for “simple” scientific work, scientists

use a simplified model. This is the model we will be using (Figure 1.2). For any given

element, this model shows the composition of the nucleus, that is, the number of

protons (p+) and neutrons (n) it contains. It also situates the electrons on energy levels

that are represented by concentric circles. The electrons in the last level, or the outer

shell, are called peripheral or valence electrons. They are the ones that travel and

make it possible for atoms to bond together in order to form new molecules.

Figure 1.2 - Simplified atomic model

At the centre of the atom is the nucleus which contains protons (positive particles) and neutrons(neutral particles). Electrons (negative particles) are found in energy levels around the nucleus.

Elements, Molecules and Chemical Formulas

The basic building blocks of all substances are the elements. There are 112 known

elements, that is, 112 types of atoms, each defined according to the number of protons

and electrons it contains. These 112 elements are arranged in a precise order in the

periodic table, and constitute the “alphabet” of chemistry. They are combined to

produce millions of different substances, in the same way that English words are

formed from an alphabet of only 26 letters. Thus, most substances result from the

combination of small groups of atoms called “molecules.” All the molecules of a given

substance are identical. For example, water is made up of molecules, each containing

two hydrogen atoms (H) and one oxygen atom (O). For this reason, its chemical formula

is written as H2O. Figure 1.3 shows the chemical formula of a few substances and

illustrates how the atoms are arranged in the molecules.

a) Three-dimensionalrepresentation

of a carbon atom

b) Two-dimensionalrepresentation

c) Simplified representation

6p+

6n

Electron Proton Neutron− +


1.7

Figure 1.3 - The chemical formula and molecular structure of a few substances

In the molecular structure, each dash represents a bond between two atoms, formed by two valence electrons.

Note that molecules and atoms are extremely small. In fact, they are so small that

18 g of water (about 18 mL), or the equivalent of a film of water at the bottom of a

glass, contains one mole of molecules, or more than 600 000 billion billion (6 × 1023)

molecules. Whether this quantity of water (18 g) is in the form of a liquid, ice or steam,

it always contains the same number of molecules.

The chemical formula of a substance tells us its composition. A subscript in

parentheses placed after the chemical formula indicates the state of the substance,

that is, whether it is a liquid (l), a solid (s) or a gas (g). The subscript (aq) is used to indicate

that a substance is dissolved in water or that it is in an aqueous solution. Below are

some examples of how these subscripts are used.

H2O(s) ice (solid water)

H2O(l) liquid water

H2O(g) water vapour

NaCl(aq) salt-water solution

H

NH H

O C O

OH H

H

H C H

H

O O

Name Chemical formula Structure

Water H2O

Ammonia NH3

Methane CH4

Oxygen O2

Carbon dioxide CO2


1.8

Carbon

dioxide

Exercise 1.3

Consider sulphuric acid, represented by the formula H2SO4.

a) How many different elements does it contain?

b) How many atoms are there in one molecule of sulphuric acid?

c) The label on a bottle indicates that it contains H2SO4(aq). What does the bottle

contain?

A MODEL FOR THE THREE STATES OF MATTER

Whether in the solid H2O(s), liquid H2O(l) or gaseous H2O(ag) state, water is composed

of identical molecules. Yet, ice, water and vapour are very different in appearance.

How can identical molecules produce such varied forms? Is it the way they are arranged

or the distance between them? What distinguishes the three states? We need to develop

a model in order to answer these questions. We will determine the initial characteristics

of this model in Experimental Activity 1. Then we will use the results of this experiment

to develop the model further.

Experimental Activity 1: Gases, Liquids and Solids

In this activity we will examine the properties of the three

states of matter and establish the basis for a model which will

help us explain these properties.

In this introduction to the scientific method, you will discover

the importance of observation and of the conclusions it allows

us to draw.

All the information you will need to carry out this activity is

given in Section A of Experimental Activities of Chemistry.

Have fun!


1.9

GWB

?

In Experimental Activity 1 you compared solids, liquids and gases and studied the

properties of shape and volume. On the basis of your observations, and guided by

questions, you began to deduce the differences between the states of matter at the

molecular level. You looked at the forces of attraction between molecules and at the

distances that separate them. These two variables, among others, form the basis of

the model we will develop to illustrate the three states of matter. We will now continue

our study by analyzing each state in detail.

Solids

A solid is characterized by having a definite shape and volume and little or no

compressibility. You may be wondering how these properties can be explained at the

molecular level. Let’s proceed by analogy. When one constructs a building several stories

high, the building materials are arranged in a very specific order, otherwise the entire

structure might collapse. The ordered structure of this arrangement gives the

building its solidity. The same is true of the molecules in a solid. They are aligned in

a precise order. In addition, the attractive forces between them allow them to remain

in fixed positions close to one another. This explains why solids have a definite shape.

Even the molecules found close to the surface of a solid are held firmly in place, thus

giving the solid its well-defined outline (Figure 1.4a).

Solids resist compression. From this, we can deduce that the molecules in a solid

have a regular geometric pattern that cannot be packed more closely together by

applying pressure to them. This reasoning stands only if the distance between the

molecules is very small.

The attraction between molecules and the short distance that separate them also

considerably limit their movements. Thus, the molecules in a solid vibrate but remain

in place—they cannot move from one point to another within the solid (Figure 1.4b).


1.10

Figure 1.4 - Solid state model

a) The molecules are represented by dots. The structure is ordered and compact and the attractionbetween the molecules gives the solid a definite shape. Note that because molecules

are extremely small, the number shown above is much smaller than the actual number.

b) The molecules of a solid vibrate but cannot move away from each other.

Exercise 1.4

State the molecular properties that explain why solids have a definite shape.

Liquids

A liquid cannot be compressed. It assumes the shape of its container and occupies

a definite volume, since its volume remains the same regardless of the space

available. This is what you concluded from the experiment.

Since a liquid cannot be compressed, we may conclude that its molecules are so tightly

packed that they cannot be compressed any further. As in a solid, the distance between

the molecules in a liquid is very small.

a)

Vibration

b)


1.11

?

Figure 1.5 - Liquid state model

a) A liquid takes the shape of its container.

b) The particles of a liquid are close together; however, they are not arranged in a specific order and they have the ability to move past one another

within the shape of the container.

c) The molecules of a liquid turn on their axes. This movement is called “rotation.”

The molecules of a liquid are held together by attractive forces. The force of this

attraction allows the liquid to maintain a constant volume, but is not sufficiently strong

to keep the molecules in a fixed position. The molecules are therefore not ordered

as they are in a solid, and the liquid does not have a definite shape (Figure 1.5b).

How do the molecules of a liquid adapt to the shape of different containers? The

molecules turn on their axes and move past one another, like the balls in a ball pit

where children’s feet cause the balls to slide over one another. We may conclude that

in addition to vibrating, liquid molecules turn on their axes. This movement is called

rotation (Figure 1.5c).

Exercise 1.5

State the molecular properties that explain why a liquid resists compression.

10 ml 10 ml

b)a)

Rotation

c)


1.12

?

Gases

A gas is compressible and expandable and takes the shape of its container. This means

that gas molecules are not ordered and that the attractive forces between them are

not strong enough to give the gas a definite shape.

Since a gas is compressible, we may conclude that when pressure is applied to it, its

molecules move closer together. Therefore, unlike the molecules of solids and

liquids, the molecules of a gas are not compact and the distances between them are

large. Gas is also expandable, which means that it tends to occupy all the available

space. This property can be explained by the fact that the molecules move away from

each other as the available space increases. This indicates that there is little or no

attraction between the molecules of a gas; otherwise, they would remain in fixed

position to one another. The attractive forces in a gas are very weak, and the molecules

are independent of one another.

Figure 1.6 - Gaseous state model

a) The molecules in a gas move freely in their container and occupy all the space available.

b) The molecules in a gas vibrate, turn on their axes and move freely. This last movement is called “translation.”

Independent of one another and widely spaced, the molecules of a gas move freely

in their container and occupy all of its space (Figure 1.6a). All movements are possible

for the molecules of a gas: they vibrate, turn on their axes and move freely from one

point to another. This last movement, called translation, is the most important one

in gases. The phenomenon of translation allows the molecules of a gas to occupy all

the space available to them.

Translation

a) b)


1.13

Exercise 1.6

State the molecular properties that explain why a gas is expandable.

Exercise 1.7

Complete the following sentences.

The most structured state is the __________________________________ state. In the

_______________________ state, the molecules are very close together, but they are not

arranged in any particular order. In the __________________ state, the molecules move

freely from one point to another.

Exercise 1.8

What are the main molecular movements that characterize each state?

a) Solid state:

b) Liquid state:

c) Gaseous state:

Now summarize what you have learned so far by completing the table in Exercise 1.9.

The top part of the table shows the observable properties of matter, that is, those that

can be seen with the naked eye, whereas the bottom part of the table indicates the

model for each state and a description of the invisible properties. The table gives the

molecular characteristics of all three states.


1.14

?

?

?

Exercise 1.9

Complete the following synoptic table.

The Three States of Matter: Properties and Model

The above table summarizes the concepts covered so far in this chapter and will be useful

when you review the material. Be sure to make the appropriate connections between

the observable properties and the model that represents them.

The following are important points that bear repeating:

• The rigorous structure and very strong attractive forces between the molecules is

responsible for the cohesion of a solid, its rigid structure and its definite shape.

The molecules of a solid are tightly packed and they vibrate.

PROPERTIES GAS LIQUID SOLID

Shape(Definite or not)

Indefinite

Volume(Definite or not)

Definite

Compressibility(Negligible or high)

Negligible

MODEL GAS LIQUID SOLID

Diagram

Distance between the molecules Large(Large or very small)

Main movements(vibration, rotation, Vibration and rotationtranslation)

Attractive forces between the molecules (Yes or no)

No

Order(Yes or no)

No


1.15

Compressibilityvibration rotation

translation

?

• The attractive forces between the molecules of a liquid are strong enough to maintaina constant volume, but are not strong enough to give the liquid a definite shape.A liquid takes the shape of its container. The molecules of a liquid are held closeto one another, vibrate and turn on their axes (rotation). When a liquid flows, itsmolecules slide over one another.

• The attraction between the molecules of a gas is negligible and is often consideredto be non-existent. The molecules are independent of one another and move aroundfreely in the available space. The distance between the molecules is large. Themolecules vibrate, turn on their axes and move from one point to another(translation).

THE KINETIC THEORY OF GASES

We have compared the properties of the three states of matter and the models thatrepresent the three states at the molecular level. In the rest of this course, we willdeal mainly with gases. To describe the behaviour of gases, scientists use the kinetictheory of gases. This theory can be summarized in a few sentences and complementsthe model we have developed. According to the kinetic model, gases consist of particlesthat have the following properties:

• The particles are very small and they are mainly molecules;

• The distance separating the particles is large in relation to their size; the particlescan therefore be represented by dots;

• The particles of a gas are in continuous and rapid motion: they collide constantlywith each other and the walls of their container, and this causes them to reboundin random directions;

• The particles of a gas do not attract each other, nor do they repel each other: theyare independent;

• The energy associated with the movement of the particles (kinetic energy), thatis, the energy that depends on their mass and speed, is a function of thetemperature of a gas.

The above statements include a number of the characteristics of our gas model. Infact, a gas is made up of independent molecules (or atoms) that are in perpetual motionand that are separated by large distances. The molecules collide continuously witheach other and the walls of the container. This statement agrees with our model:imagine a large number of molecules in constant movement in a limited volume—collisions are inevitable.


1.16

The kinetic theory also holds that the energy of the molecules depends on the

temperature of the gas, which does not contradict our model. We will discuss this

point in more detail later in the course.

Exercise 1.10

a) We have developed a model to describe the three states of matter. Are the properties

given for gases consistent with the kinetic theory?

b) Summarize the kinetic theory of gases in five points.

Ideal Gases?

Like our model, the kinetic theory of gases describes what is conventionally called

an ideal gas. The theory holds that gas molecules are as small as points, meaning

that they have no volume and are completely independent. This description

corresponds to an ideal situation, and is sufficient to explain the behaviour of gases

in most cases. In other words, while the model does not explain the reality entirely,

it is a close enough description of it. From now on, we will refer to the model that

corresponds to the kinetic theory of gases as the “model of an ideal gas.”

For the reasons we have just stated, the model of an ideal gas has limitations. For

example, a highly compressed gas is not “ideal.” Its molecules are very close together

and they cannot be considered as points with no volume. Furthermore, the shorter

distance between the molecules means that, while still very weak, the forces of attraction

are no longer negligible. Therefore, a compressed gas no longer behaves exactly like

an ideal gas. Its molecules, which are closer together, start to move rather like those

of a liquid. In fact, if it were further compressed, the gas would become a liquid.

Figure 1.7 compares a gas that is not compressed (a) with the same gas when it is

compressed (b).


1.17

?

Figure 1.7 - Ideal gases?

a) Low-pressure gas b) Compressed gas

The more the molecules of a gas are compressed, the less the gas will behave like an ideal gas, thatis, a gas whose molecules are completely independent and free to move around.

To sum up, the further apart the molecules of a gas are, the more independent they

become and the more closely the gas behaves like the ideal gas described in the kinetic

theory of gases (or in our model). By contrast, the more a gas is compressed, the smaller

the distance between its molecules and the greater the attraction between them.

Although the ideal gas model is limited, we should not forget that it describes the

reality in most cases. This is the model that will be used throughout this course, unless

otherwise indicated.

Exercise 1.11

a) What is meant by an “ideal gas?”

GWB


1.18

?

b) Give an example of a case in which a gas cannot be considered ideal.

While the kinetic theory of gases as we know it today may seem simple, it was developed

by the efforts of many scientists over a period of centuries. They first had to conclude

that matter was made up of atoms, discover the existence of molecules and, through

experimentation, develop their knowledge of gases. Well before the atomic theory

gained widespread acceptance, alchemists also wondered about the structure of matter.

1.2 HISTORY AND GASES: THE BIRTH OF MODERN CHEMISTRY

While today it is relatively simple to develop a model for the three states of matter,

it was not always so. Imagine for a moment that the concept of “molecule” does not

exist, that the atom is only a vague concept and that you are trying to understand

the nature of matter from what you can observe with your eyes. Not so simple! Research

on gases has contributed greatly to the evolution of chemistry. The great scientists

like Dalton, Lavoisier and Avogadro, all took an interest in gases.

THE FOUR MAIN ELEMENTS

Despite the fact that the concept of the atom is more than 20 centuries old, for a long

time it was believed that matter was continuous and unbroken, in other words, that

it was not composed of particles and that it was possible to continually cut a piece

of metal, iron for example, indefinitely. All the creation was made up of four main

elements2: air, earth, water and fire. Alchemy, the forerunner of modern chemistry,

held that these elements were the only basic constituents of all matter. Before they

accepted the atomic theory and turned to the study of gases, the alchemists of the

17th century grouped magic, astrology and science together.


1.19

2. At that time, the word “element” did not have the same meaning as it does today, and had no relation to modernatomic theory.

THE LATE 18TH CENTURY: CHEMISTRY BECOMES AN EXPERIMENTAL SCIENCE

It was only in the late 18th century that chemistry, like physics, became an

experimental science as we know it today. This was a veritable revolution started by

scientists who were studying gases.

In the late 18th century, chemistry had reached somewhat of an impasse because the

old theories were no longer consistent with a growing number of experimental

conclusions. The practice of chemistry became directed towards quantitative results:

scientists started to measure and weigh substances as they do today. Qualitative

information such as “longer,” “lighter” or “heavier” were no longer acceptable. The

chemist had become more than an observer with a trained eye; he now used

thermometers, calorimeters, aerometers and, above all, balance with high degree of

precision.

The discovery of carbon dioxide around 1750 was a key event. Researchers

concentrated their efforts and developed methods for collecting gases, separating them

and measuring their volume. They identified nitrogen, chlorine, carbon monoxide,

sour gas (sulphur dioxide) and nitrogen dioxide.

In 1784, Antoine Lavoisier, considered by many to be the founder of modern

chemistry, succeeded in determining once and for all the chemical composition of

water. By applying a rigorous quantitative method, he obtained water by burning a

mixture of gases composed of two volumes of “inflammable air” (hydrogen) for each

volume of “vital air” (oxygen). From that point on, water could no longer be

considered a principal element since it was known to be composed of hydrogen and

oxygen. Lavoisier studied air and established, using his usual rigorous method, that

it was a gaseous mixture. He also took an interest in combustion reactions and proved

that fire, like air, was not an element. The reign of the four principal elements, which

had prevailed since ancient times, had come to an end. Thanks to Lavoisier, the use

of the balance and the application of a rigorous quantitative method became the means

of controlling all chemical operations.


1.20

Figure 1.8 - Lavoisier’s gasometer

The gasometer is a balance for weighing gases. Lavoisier used it to determine the composition of air, among other things.

New scientific discoveries forced scientists to redefine the concept of “element”: from

that point on, the term “element” was used for any substance which could not be

decomposed into simpler substances. The composition of a large number of

substances was soon established. The method for naming them also had to change

because chemistry needed a rigorous language in order to progress. “Airs” soon

disappeared as did many colourful names (See “Change Is in the Air,” page 1.23).

Lavoisier and his colleagues established the foundations for the current nomenclature

by replacing the names of the different “airs” with other names, many of which are

still being used today.

Figure 1.9 traces two centuries’ worth of landmark discoveries that contributed to

the collapse of outdated alchemical principles.

GWB


1.21

Figure 1.9 - Historical facts and discoveries involving gases in the 17th and 18th centuries

A number of events and discoveries made it possible for science to advance and for chemistry to become a truly experimental science.

Historicaland

technicalevents

Historicaland

technicalevents

Historicaland

technicalevents

Historicaland

technicalevents

Historicaland

technicalevents

1603The Academy of Science

is founded in Rome

1600First discoveriesinvolving gases(Van Belmont)

1660The Royal Society

is founded

1661Manometer(Huygens)

1700The Royal Academy ofScience is founded in

Berlin

1714Fahrenheit

scale

1762Steam car(Cugnot)

1789The FrenchRevolution

starts

1784Synthesis of

water (Cavendishand Lavoisier)

1785Gas

lighting(Minckelers)

1786Steamboat

(Fitch)

1787Charles’s

law

1799Metric system

adopted inFrance

1789Lavoisier publishes seminal work on

chemistry entitled Traité élémentaire dechimie (start of chemistry as a science)

1794The École polytechnique

is founded in France

1766Cavendish identifies

hydrogen as a component of water

1774Oxygen and ammonia

gas identified(Priestley)

1777Composition ofair (Lavoisier)

1783Montgolfier brothers

operate the first hot-air balloon

1718Mercury

thermometer(Fahrenheit)

1742Celsiusscale

1754Carbon dioxide

identified(Black)

1730Alcohol

thermometer(Réaumur)

1737Kinetic theory of

gases (Bernoulli’sTheorem)

1776American Declaration of

Independence

1662Boyle’s law

1667Boyle’s study onthe expansion of

gases

1680First phosphorus

matches

1690Centrifugal

pump (Denis Papin)

1697Phlogiston

theory(Stahl)

1666The Académie royale

des sciences isfounded in Paris

1632First water

thermometer (J. Ray)

1644Torricelli: experiments onthe weight of air; the first

barometer is made

1647Pascal: experimentson vacuums (vacuum

distillation)

1650Vacuum pump

(O. von Guenicke)

Major discoveries

andchemists

Major discoveries

andchemists

Major discoveries

andchemists

Major discoveries

andchemists

Major discoveries

andchemists

1635The Académie

française is founded

1657The Accademia del Cimento

is founded in Florence


1.22

thermometerbarometerManometerexpansionPhlogiston theoryFahrenheit scale Kinetic theory ofgasesCarbon dioxide hot-air balloon

Scientists in the 18th century could not keep up with the growing family of “airs,” as gases were called in

those days. Each scientist had a different name for a newly identified gas so that what Joseph Priestley called

“dephlogisticated air” and Scheele called “fire air,” Lavoisier eventually called “oxygen.” It is still known as

such today. Below is a list of a few “airs” known at that time, and their current names.

When a gas like oxygen goes by six different names, you know it’s time to put an end to the confusion. Thus,

in 1787, a team of French scientists proposed a complete reform of chemical nomenclature. The nomenclature

method quickly became the bible of the new chemistry. Simple names were suggested for simple substances.

Thus, the numerous “airs” became “hydrogen” if the gas generated water (hydro), “oxygen” if it generated

acids (oxy), “nitrogen” if it was the portion of air (in its current meaning) that did not support respiration in

living beings (note that the French name for nitrogen is azote, from the Greek prefix “a” meaning “without”

and the Greek “zöein” meaning “life”), etc.

THE EARLY 19TH CENTURY: THE ATOM AND THE MOLECULE

In 1808, following his work on gases, John Dalton formulated his atomic theory. He

was particularly interested in meteorology, the composition of air and the properties

of gases in general. Using his new theory, he explained the behaviour of ammonia

gas and methane (the main component of natural gas).

Analytical studies on water and different gases caused chemistry to advance by leaps

and bounds. It was at this time that scientists formulated the fundamental laws which

are still being used in our day. For example, Avogadro introduced the concept of the

molecule by maintaining that simple gases are composed of diatomic molecules. Today,

we know of several diatomic gases, including oxygen (O2), nitrogen (N2) and

hydrogen (H2). Avogadro formulated the following hypothesis: equal volumes of

different gases at the same temperature and pressure contain the same number of

Current name

Hydrogen

Oxygen

Carbon dioxide (sour gas)

Nitrogen

Names used in the 18th century

Inflammable air, phlogiston

Vital air, dephlogisticated air, empireal air, vitriolated air, pure air, fire air

Inert air

Phlogisticated air

Change Is in the Air


1.23

molecules. This means that three containers of the same volume filled with three

different gases at the same temperature and pressure contain the same number of

molecules. This hypothesis is illustrated in Figure 1.10 and will be covered in detail

in Chapter 5.

Figure 1.10 - Illustration of Avogadro’s principle

One mole of gas at standard temperature and pressure occupies a volume of 22.4 litres, regardless of the type of gas.

Exercise 1.12

The idea that matter is composed of particles (atoms) goes back to antiquity, to about

the fourth century B.C.

a) Has this theory been in general use since that time to explain the composition of

matter and its behaviour? If not, which theory prevailed?

b) When did the atomic theory become prominent again?

H26.02 × 1023

molecules 22.4 L

N26.02 × 1023

molecules 22.4 L

O26.02 × 1023

molecules 22.4 L


1.24

?

c) Which great chemist debunked the theory of the four principal elements? How

did he do this?

Exercise 1.13

Chemistry became a truly experimental science in the late 18th century. What major

change can be noted in the attitude of the chemists of that time?

1.3 DIFFUSION AND BROWNIAN MOTION

As we saw earlier, modern chemistry uses the kinetic theory to describe gases. According

to the kinetic theory, molecules are in continuous motion and this has specific

consequences. For example, the movement of the molecules of a perfume allows us

to smell the perfume. We say that the perfume is “diffused” in the air.

DIFFUSION OF GASES

A keen sense of smell allows you to smell even the most subtle fragrances. You are

able to perceive smells because gaseous and odorous molecules disperse in the air

and reach the receptor cells in your nose.

Gas molecules are independent of each other. They move in straight lines and change

direction each time they collide with one another or with the walls of the container

in which they are confined. As a result of all these collisions, molecules move in a

zigzag pattern, not unlike as the metal ball in a pinball machine (Figure 1.11).


1.25

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Figure 1.11 - Diffusion

Like a pinball, gas molecules diffuse by moving in a zigzag pattern, which is the result of a series of collisions with other molecules.

Like pinballs, molecules that escape from a bottle of perfume move rapidly in all

directions, spreading the odour throughout the room. The natural tendency of gases

to spread in space is called diffusion.

Smells

The sense of smell, or olfaction, allows us to perceive the diffusion of odours. An odour

can only be smelled when there is direct contact between the nervous system and

the gaseous molecules emitting the odour. When inhaled, the molecules come into

contact with the olfactory receptors in the nose (Figure 1.12).

GWB


1.26

Figure 1.12 - The sense of smell in humans

Humans can appreciate the fragrance of a flower thanks to the olfactory receptors situated behind the eyes. Gaseous molecules emanate from the flower and enter the nose.

Those that reach the olfactory receptors create the sensation of smell.

In the case of a fire or a gas leak, the sense of smell can save our life. It is for this reason that a product

with a characteristic odour is added to odourless propane gas so leaks can be rapidly detected. It is easy to

see that the partial or total loss of the sense of smell can have fatal consequences. Such a loss can be

devastating to someone who loves good food!

It is common knowledge that dogs have a keen sense of smell. This is a dog’s best developed sense and it

allows it to detect very specific smells. Dogs are trained to sniff out drugs or to find victims in an avalanche.

Dogs have between 100 and 200 million olfactory cells whereas humans have only five million. You might

call dogs professional “noses!”

A keen Sense of Smell

Olfactory receptors

GWB


1.27

Have you ever wondered why people cry when they peel onions? This phenomenon

can be explained by the fact that onion juice contains propenyl sulfenic acid

(C3H6SO), which is an eye irritant. When you slice an onion, the molecules of this

acid escape from the liquid and are diffused in the surrounding air. Those that reach

your eyes irritate the mucous membranes and cause the eyes to produce tears whose

function is to protect the mucous membranes. To prevent your eyes from tearing,

you must prevent the gas from reaching them. An extreme solution would be to wear

an airtight mask, such as those worn by scuba divers. A simpler solution is to immerse

an onion in a bowl of water when you are peeling it so that the irritating substance

remains trapped in the water.

Odours and Toxicity

Often, if a substance has an unpleasant odour, it is wrongly thought to be toxic.

Although they are unpleasant, bad odours are not necessarily harmful. Minimal

quantities of a substance are sometimes sufficient to produce a very strong odour.

For example, the liquid that skunks spray to defend themselves smells terrible; however,

breathing the odour is not necessarily harmful. On the other hand, some gases have

a pleasant smell but are extremely dangerous. For example, hydrocyanic acid (HCN),

an extremely toxic substance, has a pleasant almond smell.

Nor are all odourless gases safe. Carbon monoxide (CO), a gas discharged from the

exhaust systems of cars, is highly toxic even though it is completely odourless. Methyl

alcohol (CH3OH) is another example. Also known as wood spirit, it is used as a solvent,

as a paint remover and as an additive in gasoline. It is a very volatile liquid with almost

no smell compared to other common solvents. When a bottle of methyl alcohol is

opened, the molecules of the liquid evapourate and diffuse in the ambient air. The

diffusion of this gas in a poorly ventilated area may cause eye irritation, headaches

and even fainting.

Speed of Diffusion

Not all gases diffuse as rapidly as a pleasant perfume or as the propenyl sulfenic acid

found in onions. Essential oils, for example, diffuse much more slowly because their

molecules are heavier. Two factors come into play here: 1) heavy molecules escape

from the liquid with more difficulty and 2) they move more slowly through the air.


1.28

The speed at which a gas diffuses is directly related to the speed of its molecules.

The faster the molecules move, the more quickly they diffuse. At the same temperature,

smaller molecules move faster than larger ones. Their small size also makes it easier

for them to negotiate a path around the air molecules. The heavy molecules move

less quickly than the lighter ones and are often more bulky.

The size of the molecules is related to the molar mass of a gas. The molecules of gases

with a large molar mass move more slowly than those of gases with a smaller molar

mass. Thus, gases with a small molar mass escape and diffuse more quickly than the

heavier gases. Figure 1.13 compares the movement of a heavy molecule with that of

a light molecule.

Figure 1.13 - Speed of diffusion of molecules

Heavy molecule Light molecule

At any given temperature, and in the same period of time, a light molecule moves faster and goesfurther than a heavy molecule. The diffusion of a gas with a small molar mass is faster than that

of a gas with a large molar mass.


1.29

Exercise 1.14

Knowing that the molecules of lighter gases diffuse more quickly, answer the

questions below about the following gases: H2, CO2, O2, He.

a) Using the periodic table, determine the molar mass of these four gases3.

b) Arrange the four gases by increasing order of speed of diffusion.

All gases diffuse into the air regardless of whether they smell good or bad, or have

no smell at all. Diffusion consists in the movement of gas molecules in every direction

in space. The smaller the molecules, the faster they move and the larger the

molecules, the slower they move. Gases diffuse into the air, which is itself a gaseous

medium whose molecules are far apart. What would happen in a liquid medium where

the distances between the molecules are much smaller?

DIFFUSION IN LIQUIDS AND SOLIDS

Diffusion is not limited to gases. It also occurs in liquids. If you pour a few drops of

grape juice or food colouring into a glass of water, you will see that the difference

between the colour of the juice (or of the food colouring) and the water, very distinct

at the beginning, will become less and less apparent until in the end the liquid is all

the same colour. Diffusion is slower in liquids than in gases because the molecules

are much closer together and their movements are much more limited. A molecule

that diffuses into a liquid can be likened to a person trying to move to the front of a

tightly-packed crowd in order to see the stage.


1.30

3. For more information on calculating molar mass, refer to Ouellet, Danielle, Ionic Phenomena: A Study of anEnvironmental Problem, Chapter 5, Learning Guide produced by SOFAD, or to “Appendix C - ChemicalPrerequisites.”

?

Exercise 1.15

The speed of molecular diffusion is greater in gases than in liquids. Explain this

phenomenon using the model that was developed for liquids and gases at the beginning

of the chapter.

Diffusion is almost nonexistent in solids because of the rigidity of the solid’s

structure. Remember that the molecules in a solid are tightly packed and occupy fixed

positions.

In practice, in order to have a substance diffuse in a solid, the temperature must be

raised until the solid becomes a liquid. Steel illustrates this well. Metallurgists make

steel by diffusing solid carbon into iron. The iron is heated until it becomes a liquid

and carbon is then introduced. The resulting alloy is harder than iron and less brittle

than carbon, making it a very resistant material.

Exercise 1.16

Why is there very little diffusion in solids?


1.31

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?

Diffusion has an important application in the nuclear industry. Natural uranium is a mixture of uranium

isotope 238 and uranium isotope 235. Since the isotopes of an element have the same chemical properties,

it is then not possible to separate them with the use of chemical reactions. Before using natural uranium in

the reactors, the mixture of uranium isotopes must be enriched with uranium 235. This procedure is achived

through diffusion.

As the speed of diffusion is faster in gases, natural uranium is first transformed into a gas called “Uranium

hexafluoride,” more precisely 238UF6(g) and 235UF6(g). These two gaseous isotopes are then diffuse through walls

that have different degrees of porosity. The gaseous molecules of the lighter isotope (235UF6 or uranium 235),

diffuses more quickly than the gaseous molecules of the heavy isotope (238UF6 or uranium 238). When the

gas is collected on the other side of the walls, the proportion of uranium 235 is greater than at the start.

We can then say that the combustible has been “enriched” with uranium 235.

BROWNIAN MOTION

We have seen that diffusion is faster in gases than in liquids because molecules move

more easily in gases. Diffusion is slower in liquids but it occurs just the same; it is

facilitated by the continual motion of the molecules. The direct consequences of the

movement of molecules in a liquid can be observed with the naked eye.

In 1827, less than 20 years after John Dalton formulated his atomic theory, the Scottish

botanist Robert Brown noted, under a microscope, that pollen grains suspended in

water displayed erratic motion. At first he thought that this movement was due to

the living nature of the grains of pollen; however, when he observed non-living dye

particles, he noted that they displayed the same motion. Figure 1.14 shows the

movement of fine particles of a solid (dust, pollen grains) in suspension in a liquid,

seen through a microscope.

Diffusion at the Service of Nuclear Science


1.32

Figure 1.14 - Brownian motion

Under the microscope, it can be seen that the fine particles of pollen move in a random fashion.This motion is the result of collisions between the liquid molecules (invisible) and the pollen.Brownian motion is therefore visible proof that the molecules of a liquid are in continuous

motion that is caused by different collisions with molecules in the surrounding environment.

This random motion is explained by the following hypothesis: the molecules of the

liquid collide continually with the pollen grains, in an erratic fashion. The pollen

particles move in a zigzag pattern due to bombardment from the molecules of the

liquid. The smaller the particles, the more pronounced the movement and the more

easily it can be observed. Such observations have given weight to the theory which

holds that the molecules of a liquid are in continuous movement. The motion of solid

particles, called Brownian motion, after the botanist, is concrete and visible proof

of molecular motion.

1.4 PHASE CHANGES

Regardless of whether snow, rain or extreme humidity is forecast, we can be sure

that the day will be wet. Water can take many forms and its molecular movements

and properties change depending on its state. You know from experience that water

can change from one state to another depending on the temperature. A day that starts

off snowy, with an outside temperature of –2°C can easily end with rain if the

temperature reaches 0°C during the day. Figure 1.15 summarizes the different phasechanges.

Fine pollen particlesseen through amicroscope

Invisiblemolecules


1.33

Figure 1.15 - Phase change triangles

In everyday language, certain phase changes are designated by terms other than those used in thefigure above. For example, condensation is often used instead of liquefaction, and evapouration or

boiling is used instead of vapourization.

ICE

Melting

Solidification

Sublimation

Crystallization

Liquefaction

Vapourization

WATER VAPOUR

SOLID

GAS

Melting

Solidification

Sublimation

Crystallization

Liquefaction

Vapourization

LIQUID

b) Phase changes of water

a) Spatial organization of the molecules


1.34

Melting

Solidification

Crystallization

Exercise 1.17

Using Figure 1.15, complete the following quiz to test your knowledge of this subject.

Give one point for each correct answer and count the total of the points.

Each of the following statements describes a phase change. Name the phase change.

a) A cube of ice in a glass of alcohol:

b) A puddle of water in the sun:

c) An iceberg floating on the ocean:

d) The surface of a lake in the fall, at –10°C:

e) The appearance of dew on grass:

MELTING AND SOLIDIFICATION

When a solid, such as ice, is heated, the molecules absorb the heat and thus acquire

energy. The greater the quantity of heat applied, the greater the amount of energy

stored by the molecules. This energy is manifested by the increased vibration of the

molecules. Thus, the molecules’ range of motion increases and the bonds that link

the molecules together weaken. At a sufficiently high temperature, the molecules vibrate

to the point of breaking, and the solid then collapses, somewhat like a house of cards.

The structure loses its shape and the molecules move past one another. The solid is

melting and becomes a liquid (Figure 1.6).


1.35

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Figure 1.16 - Melting of a solid

The molecules closest to the source of heat vibrate more rapidly (a). The vibrating motion is transmitted to the surrounding molecules (b). The vibrating movements spread to the interior of the solid (c). The structure of the solid collapses

and becomes liquid. This is the phenomenon of melting (d).

When this happens in reverse, that is, from a liquid state to a solid state, as in the

formation of ice on lakes, it is called solidification. The melting and solidification

of water are represented by the following equations.

Meltingsolid water + energy . liquid water

solid water + energy . liquid water

Solidification

VAPOURIZATION AND LIQUEFACTION

When you fill a kettle with water and plug it into an electrical outlet, the temperature

of the water rises to the point of boiling or vapourization, that is, to the point where

the water changes from a liquid state to a gaseous or vapour state.

Exercise 1.18

Referring to molecular motion, explain what happens when:

a) water is heated in a kettle.

a) b) c) d)


1.36

?

b) the temperature reaches the boiling point (100°C).

As the temperature of a liquid rises, its molecules become increasingly more active.

The molecules become more independent as the attractive forces between them weaken

and, when the temperature reaches the boiling point, the molecules escape from the

liquid in the form of vapour (Figure 1.17).

Figure 1.17 - Boiling water: the phenomenon of vapourization

At the boiling point, molecules have enough energy to become independent of one another and to change from a liquid state to a gaseous state.

The reverse transformation, from vapour to liquid, can be observed by placing a spoon

over the spout of the kettle. Upon contact with the cold spoon, the vapour condenses.

This phase change is the reverse of vapourization and is called liquefaction or

condensation.

Exercise 1.19

Name the phase change that is represented by the following equations.

a) Liquid water + energy water vapour

b) Liquid water + energy water vapour


1.37

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SUBLIMATION AND CRYSTALLIZATION

It is possible for substances to go directly from the solid state to the gaseous state.

This is demonstrated by dry ice (solid carbon dioxide), which is transformed directly

into a gas when heated. Dry ice is often used in the theatre or at rock concerts to produce

the effect of smoke or clouds at floor level, the density of this gas is higher than the

density of the air. White clouds are produced by the condensation of water vapour in

the air under the cooling effect of invisible CO2 gas. At –78°C, the pieces of solid CO2

are transformed directly into a gas, skipping the liquid state.

The same phenomenon occurs with naphthalene (mothballs) and solid deodorants.

This transformation is called sublimation; the reverse phase change, that is, the direct

passage from the gaseous state to the solid state is called crystallization.

Sublimation is also the phenomenon that allows us to see comets on a clear night.

In fact, comets are composed of dust and ice. When a comet approaches the Sun,

the water on its surface changes directly from a solid to a gas; the gas molecules diffuse

into space and crystallize behind the comet, creating a large tail which reflects the

light of the Sun. It is this tail that we admire when a comet passes close to the Earth.

In the winter, a number of amphibians and animals endure temperatures as cold as –40°C or more. Their

survival depends on the fact that their metabolism slows down considerably and their respiratory and circulatory

activities is suspended. However, certain insects, such as the larva of the gall wasp* and some animals,

such as the wood frog, actually freeze. The wood frog, for instance, controls the development of ice crystals

that form inside its limbs so that they do not cause irreversible damage to its cells.

Other insects, such as the spruce budworm, a defoliating insect** which ravages our forests, secrete an

antifreeze to maintain their body fluids in a liquid state when the temperature drops below the freezing point.

They can then survive without freezing, even at –45°C.

* A parasitic insect a few millimetres long which, when it lays its eggs on a plant, causes swelling or galls on theplant tissues.

** It destroys the leaves and needles of trees and vegetation. To find out more about these insects, consult Julien,Caroline, “Des animaux qui se congèlent pour résister au froid.” Québec Science, December 1995 - January 1996,pp. 30-33.

Animals That Freeze


1.38

THE HEATING CURVE FOR WATER

In the kitchen, a double boiler is generally used to heat certain foods such as chocolate.

The chocolate melts in a container placed over another which contains hot water.

Do you ever wonder why chocolate is not heated in a saucepan placed directly on

the burner?

Double Boilers

In order to understand the principle of the double boiler, we must first understand

what happens when water boils. Figure 1.18 shows a diagram of the variation in the

temperature of water in a double boiler.

Figure 1.18 - Heating water in a double boiler

The above graph shows the temperature of water as a function of time. Part a) of the graphcorresponds to the period that precedes boiling. Part b) shows a horizontal line. This part of the

graph is called a “plateau.” It indicates that the temperature of the substance remains stabledespite a constant supply of energy.

The graph in Figure 1.18 is composed of two parts. In the first, the temperature of

the water goes from about 15°C to 100°C, and the heat provided by the burner serves

to gradually increase the temperature of the water. In the second part, the temperature

of the water remains constant even though the burner continues to heat the double

boiler. This is because the water has reached the boiling point or the boilingtemperature. Let us take a look at what happens at this temperature.

BAKER

BAKER

BA

KE

R

BA

KE

R

BAKER

BAKER

100

50

0

0 1 2 3 4 5 6 7 8

Tem

pera

ture

(°C

)

Time (min.)

(b)

(a)


1.39

At the boiling point, all the heat is used to break apart the bonds between the molecules

of the liquid, thus changing the liquid into water vapour. This is why the temperature

of the water does not increase during boiling and we observe a plateau (the

horizontal part of the graph). If more heat or energy were to be provided, the water

would be transformed into vapour more quickly, but the temperature would remain

constant at 100°C until all the water had evapourated.

In a double boiler, the vapour escapes and is not recovered. Therefore, foods cook at

a constant temperature of 100°C. By contrast, if the chocolate were placed in a saucepan

directly on the stove burner, it would probably burn or stick to the pan since the

temperature of the burner is much higher than 100°C.

The Melting Point and the Boiling Point

Figure 1.19 shows the complete heating curve for water. It covers the three states and

contains two plateaus corresponding to the phase changes of melting and vapourization.

Figure 1.19 - Heating curve for water

The curve features two plateaus corresponding to the two phase changes, namely melting (0°C)and vapourization (100°C). The vapourization plateau is longer because

this phase change requires more energy than melting.

150

100

50

0

-50

Direction of the reaction

(a)(b)

(c)

(d)

(e)

Melting

Vapourization

Tem

pera

ture

(°C

)


1.40

Note that at the start of the curve, the temperature of the water is –50°C and it is in

the form of ice. When ice is heated (a), its temperature rises and the molecular

vibrations increase until the ice begins to melt. The temperature remains stable during

the entire melting phase until the ice is completely melted (plateau b). This

temperature, called the fusion point or the melting temperature or the melting point,is 0°C for water. At the melting point, the heat applied to the ice serves to break down

the ordered structure of the solid which then changes to a liquid.

Once the ice has melted, the heat raises the temperature of the liquid and thus increases

the intensity of the molecular movements (c). When the water has reached the boiling

point (100°C), vapourization begins (plateau d). The heat that is applied serves to

separate the molecules of the liquid; they become more independent and they form

a gas. When all the water has been transformed into vapour, the vapour in turn is

heated and its temperature increases as does the intensity of the molecular

movements (e). The latter operation must be carried out using a closed system (vessel

or container) from which no vapour can escape.

You may have noticed on the graph that the vapourization plateau is longer than the

melting plateau; this is because vapourization requires more energy than melting.

Exercise 1.20

Under highly controlled experimental conditions, a student obtains the heating curve

for a substance. He makes sure that he continually applies the same amount of heat

to the substance so that the supply of energy is constant. The resulting curve has two

plateaus, like the heating curve for water. The student notices that the first plateau

corresponds to the point at which the substance melts.

a) Did the temperature change in the period corresponding to the first plateau?

b) The student wonders what happened to the heat he applied during the melting

period, since the temperature of the substance remained constant. What would

you tell him?


1.41

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The heating curve of a substance makes it possible to determine the melting and boiling

temperatures of the substance, also called the melting and boiling points. These

temperatures correspond to the plateaus on the curve and are characteristic of each

substance. Chemists frequently use these temperatures to identify unknown substances.

The table in Figure 1.20 gives the melting and boiling temperatures of selected

substances.

Figure 1.20 - Table of melting and boiling temperatures of selected substances at

standard pressure (101.3 kPa)

* The melting and boiling temperatures of mixtures can vary slightly depending on the product brand or the sampletaken.

Note that the melting temperature of butter is lower than that of margarine. This is why butter melts when left on the counter on a hot summer day

whereas margarine remains solid under the same conditions.

The table shows that substances which are solid at room temperature, such as iron

and aluminum, have melting temperatures of more than 25°C. By contrast, gaseous

substances at room temperature have boiling temperatures of less than 25°C.

Substance Melting temperature (°C) Boiling temperature (°C)

Pure substances

Oxygen –219 –183

Nitrogen –210 –196

Water 0 100

Aluminum 660 2 467

Iron 1 535 2 750

Mixtures*

Corn oil –20 ––––––––

Olive oil –6 ––––––––

Peanut oil 3 ––––––––

Butter 32 ––––––––

Margarine 72 360


1.42

Exercise 1.21

Refer to Figure 1.20 to answer the following questions.

a) Based on the melting points of margarine and butter, which of the two substances

requires more energy in order to melt? Explain.

b) Which of the substances shown in the table are liquid at room temperature, that

is, at 25°C? Give their melting temperatures.

c) The liquid substances identified in b) have a melting temperature lower than 25°C.

What can you say about their boiling temperatures?

d) Which of the substances in the table are gases at 25°C? If possible, give the melting

and boiling temperatures of each substance.

e) Compare the melting and boiling temperatures of each of the substances named

in d), at room temperature.

f) One of the substances in the table undergoes a phase change when it is taken from

the cupboard (25°C) and placed in the refrigerator (2°C). What is this substance

and what state is being described?


1.43

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Exercise 1.22

The table below shows the melting and boiling temperatures of three substances named

A, B and C. For each one, indicate whether it is a solid, a liquid or a gas at room

temperature.

Expansion and Contraction

On the heating curve for water (Figure 1.19), ice (solid) heats up, that is, its temperature

rises before the first plateau is reached. After the plateau (melting), the temperature

rises again and the liquid heats up.

When a solid or a liquid is heated, molecular movement intensifies and the molecules

take up a little more space than at colder temperatures. Thus, most liquids and solids

expand when heated. Expansion is the property that causes substances to increase

in volume when their temperature is raised.

By contrast, when the temperature drops, most liquids and solids contract, that is,

their volume decreases. Contraction is the property of substances that causes them

to decrease in volume when the temperature drops. These properties of liquids and

solids are often used in technical applications. Examples of such applications are

thermometers, thermostats and electrical switches. Let us take a closer look at how

a thermometer works.

Substance Melting temperature (°C) Boiling temperature (°C) State

A –78 –33

B 1 064 2 807

C –117 79


1.44

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Thermometers

A thermometer is generally composed of a long hollow glass tube which has a small

bulb filled with liquid, usually mercury or alcohol, at one end. These two liquids are

used because they expand quickly and much more extensively than the glass that

contains them.

When a thermometer is placed in an environment warmer than itself, heat is

transmitted through the glass to the mercury inside the thermometer. The vibrations

of the mercury particles increase and the mercury expands. As the volume of the

mercury increases, the mercury rises in the tube. The level stabilizes when the

temperature of the mercury is the same as the temperature outside the thermometer.

By contrast, when the temperature drops, molecular movement decreases, the

mercury contracts and descends in the tube. The level stabilizes when the temperature

of the mercury is the same as the temperature outside the thermometer.

Anders Celsius, the Swedish astronomer, established the temperature scale that bears

his name by using a mercury thermometer and the transition temperatures for water.

In 1741, he suggested dividing the interval between the melting and boiling

temperatures of water, at standard atmospheric pressure, into 100 units. This

measurement scale was officially adopted in 1948 and is now the most commonly

used scale in the world.

Figure 1.21 - Thermometer graduated in degrees Celsius

The thermometer is an instrument that measures temperature. In 1742, the Swedish astronomer Anders Celsius arbitrarily set the melting

and boiling temperatures of water at 0° and 100°. He then calibrated the thermometer using these two reference points.

°C

40

0

20

30

70

100

80

6050

90

10


1.45

PRESSURE AND BOILING TEMPERATURE

Imagine that you are taking part in an expedition to Mount Everest. The closer you

get to the summit, the thinner the air becomes because the air pressure is decreasing.

Once you have scaled the summit, you decide to celebrate your victory by making a

nice cup of hot coffee to stave off the bitter cold. Surprise, surprise! The coffee is

barely warm even though the water is boiling vigorously. The explanation for this

lies in two concepts: pressure and boiling temperature.

Each substance has its own specific boiling temperature. Water boils at 100°C at

standard pressure. However, the boiling temperature varies according to the

ambient pressure. The boiling temperatures in the table in Figure 1.20 are accurate

provided the atmospheric pressure is standard (101.3 kPa).

In order to better understand the effect of pressure on the boiling temperature, let

us examine the following two situations: 1) how liquids behave at low atmospheric

pressure, as on Mount Everest and 2) how they behave at high atmospheric pressure.

These two situations are represented in Figure 1.22.

Figure 1.22 - Boiling temperature and atmospheric pressure

In Montréal, a temperature of 71°C is not sufficient to make water boil. However, on Mount Everest, water boils at 71°C where its low atmospheric pressure

offers little resistance to the vapour molecules that escape from the liquid.

Montréal: 71oC Mount Everest: 71oC


1.46


1.47

When a liquid boils, the molecules escape from the liquid and become gaseous. These

molecules then collide frequently with the air molecules that are close to the surface

of the liquid. At low atmospheric pressure, there are fewer air molecules at the surface

of the liquid and the liquid’s molecules need less energy to pass through them. They

therefore escape more easily. In this case, the boiling temperature is lower than it is

at standard atmospheric pressure.

Now imagine that the number of air molecules is doubled, that is, that the

atmospheric pressure is twice as high. The molecules attempting to escape from the

liquid come up against such a mass of air molecules that they can be forced to remain

in the liquid. The temperature must then be increased so that the liquid’s molecules

have sufficient energy to escape the liquid.

In summary, when the outside pressure decreases, the boiling temperature drops or,

in other words, the boiling liquid is cooler. Thus, the coffee prepared on the summit

of Mount Everest will be lukewarm whereas in Montréal or in New York, where the

atmospheric pressure is standard, the coffee will be very hot. Figure 1.23 shows the

boiling temperature of water at different locations on the globe. Atmospheric

pressure varies according to altitude.

Figure 1.23 - Variations in the boiling point of water according to altitude

The higher the altitude, the fewer air molecules there are. Atmospheric pressure decreases and so does the boiling temperature.

Dead Sea–393 m

Sea level

Mexico City, Mexico2 250 m

Quito, Equator2 849 m

Lhassa, Tibet3 684 m

Mount Everest8 847 m

101oC

100oC

92oC

90oC

87oC

71oC

Montréal0 m

1.5 TECHNICAL APPLICATIONS

It is believed that the popularity of large supermarkets is due in part to the high-

performance commercial refrigerators and freezers that make it possible to preserve

products over longer periods. These appliances reduce the losses of perishable foods,

and allow supermarkets to offer quality products at a lower price. Appliances such

as refrigerators and freezers are examples of the types of applications that involve

the phase changes of matter. For a phase change to occur, matter must either absorb

or lose energy. In other words, there must be an exchange of heat between matter

and the environment. Let us look at some applications of this type of change.

REFRIGERATORS, FREEZERS AND AIR CONDITIONERS

Freezers, refrigerators, air conditioners and heat pumps all function according to the

same principle: they use gases called “refrigerants” whose boiling temperature is around

–30°C. The most commonly used refrigerants are ammonia, freon and sulphur dioxide.

A short description of how these appliances work is given below.

Why do foods last longer when they are refrigerated? Because lowering the

temperature of foods reduces the activity of the micro-organisms and enzymes that

alter their quality. The temperature inside a refrigerator is lowered using a substance

called a “refrigerant” whose boiling temperature is lower than the temperature of the

space to be cooled. The most commonly used refrigerant is freon whose boiling

temperature is –30°C.

You may be wondering how refrigerants such as freon lower the temperature of

refrigerators to nearly 4°C (Figure 1.24). Freon is first compressed in a pump activated

by an electric motor. This pump is called a “compressor” (a). The increase in the gas

pressure causes the temperature to rise. The heated gas then passes through a coil,

called a condenser (b), situated outside the appliance, generally behind the refrigerator.

In the condenser, the gas is cooled by the ambient air and becomes a liquid

(liquefaction). As it leaves the condenser, the freon passes through an expansion valve (c)

with a twofold function. On the one hand, it regulates the flow of the refrigerant by

allowing into the evapourator only the quantity of freon needed to absorb the heat

from the air to be cooled. On the other hand, and this is a major point, the refrigerant

passes through a calibrated opening into a wide-diameter tube. This causes the

refrigerant to expand; in other words, it occupies more space, and its pressure drops

abruptly. As we have seen, any drop in pressure is accompanied by a drop in

temperature. A portion of the refrigerant vapourizes instantly and the cold vapours


1.48

lower the temperature of the remaining liquid refrigerant. The same phenomenon

occurs when deodorant is expelled from an aerosol can or air is released from an

over-inflated tire. Lastly, the cooled liquid freon reaches the evapourator where it

absorbs part of the heat from the air contained within the refrigerator. This heat

transforms the liquid freon into a gas (vapourization) (d). Thanks to this absorption

of heat, the temperature inside the refrigerator gradually decreases to approximately

4°C. Leaving the evapourator, the gas returns to the compressor and the cycle starts

again.

Figure 1.24 - The refrigerator

In the process of becoming a gas, a liquid absorbs heat. The refrigerator is based on this principle.In the compressor, the refrigerant gas is highly compressed. It then becomes hot but not yet liquid.

It is cooled by the ambient air as it passes through the condenser and becomes liquid at thatmoment, not before. As it leaves the expansion valve, a portion of the refrigerant is immediatelyvapourized and its temperature drops. It then flows to the evapourator where it becomes a gas

once again by absorbing the heat from the air in the refrigerator and the freezer. The temperaturewithin the refrigerator drops and the food is cooled.

Freezers and air conditioners function in exactly the same way as refrigerators except

that their cooling system is designed to attain different temperatures.

d) Evaporator

c) Expansion valve

b) Condenser

High-pressure section

Low-pressure section

Gas

Hot gas

a) Compressor

Cooled liquid

Simplified diagram of a refrigeratorRefrigerator

b) Condenser

a) Compressor

Freon gas

Liquid freon

d) Evaporator

c) Expansion valve


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Heat pumps are used to cool houses in summer and to heat them in winter. One heat pump can fulfill two

functions because the direction in which the refrigerant flows can be reversed, depending on the season.

The principle of the heat pump is similar to that of the refrigerator. The refrigerant, usually freon, flows through

a pipe linking a coil inside the house to a second coil outside the house. A compressor is placed between

the inside and outside coils.

In winter, the heat pump is used as a heating appliance because it compresses the freon, thus producing

heat. The outside air can also heat the freon at temperatures as low as –10°C since the boiling temperature

of freon is very low (–30°C). In summer, the direction of the refrigerant flow need only be reversed to remove

heat from the house. The cold refrigerant absorbs heat in the house as it passes through the inside coil. The

warm refrigerant, now a high-pressure gas, releases its heat as it passes through the outside coil.

Heat Pump

The high-pressure vapour leaves the compressor and passes through the inside coil where it givesup some of its heat as it changes to a liquid. A fan pushes the heat thus produced towards theinside of the house. When the liquid flows through the outside coil, its pressure drops and it

becomes very cold (–30°C) again. As the liquid is colder than the outside air (–10°C), it absorbsheat as it circulates in the coil, thus changing into a gas. The vapour then passes through the

compressor which increases its temperature and pressure.

Most heat pumps sold in Canada are of the “air-air” type. However, there are heat pumps that draw energy

from sources other than air. In some, for example, the outside coil is immersed in water or is buried underground.

These function according to the same principle, except that they use the soil or a layer of water as a source

of energy rather than air. These types of heat pumps are often more efficient but are also more costly.

Low-pressure and low-temperatureliquid (–30°C)

Coil Coil

Warmerinside air(21°C)

High-pressure and high-temperature liquid

(20°C)

High-pressure and high-temperature vapour

(30°C)

Reversiblevalve

Compressor

30°C–10°C

Low-pressure and low-temperature vapour

(–5°C)

Cold outside air

(–10°C)

Operation of an “air-air” heat pump in winter

Heat Pumps


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CFC’S: THE TWO SIDES OF THE COIN

Odourless, non-toxic and chemically inert—these are the properties of the gases

belonging to the category of chlorofluorocarbons (CFC’s), freon being the best known.

These gases are widely used as refrigerants in air conditioners and refrigerators because

they are very stable. This property is of particular interest to the refrigeration industry

because CFC’s do not break down inside the appliances.

By contrast, when an appliance is emptied of its refrigerant or there is a gas leak,

the CFC’s are released into the environment, and being very stable, can remain

suspended in the atmosphere for more than 100 years. Produced by humans in

enormous quantities (1 140 000 tons in 1988 alone), CFC’s remain intact for 10 to

15 years, long enough to enter the stratosphere. Under the action of the Sun’s rays,

they react with the ozone and damage this thin shield that protects terrestrial life

from harmful ultraviolet radiation.

To counteract this threat, 62 countries as well as the European Economic Community

(known today as the European Union) signed the Montréal Protocol on Substances

That Deplete the Ozone Layer (1987) in which they agreed to take appropriate measures

to eliminate the consumption of CFC’s by 1997. The agreement proposes replacing

CFC’s by HFC’s (hydrofluorocarbons), a group of refrigerants that do not contain

chlorine, to develop new technologies and to adopt good salvaging and recycling

practices.

In addition to being used in refrigerators and air conditioners, CFC’s serve as expansion

agents in the manufacturing of polystyrene glasses and plates, upholstery and

insulating materials. They are also used as propellants in aerosol cans, in solvents

used in electronics and in sterilizing agents used in hospitals.4


1.51

4. Based on an article by Lyne Lauzon entitled “Les ‘merveilleux’ CFC!”, in Franc-Nord, (Franc-Vert), November-December 1990, p. 24. Reproduced with the permission from Franc-Vert.

1.6 OTHER STATES OF MATTER

Throughout this chapter, we have discussed the three main states of matter. They are

by far the most common states encountered in daily life. There are, however, other

states we see less often but which are nonetheless important.

PLASMA

In addition to the solid, liquid and gaseous states, scientists often consider plasmaa fourth state of matter. Plasma is obtained by heating a gas to very high temperatures

of more than 5 000°C. At these temperatures, the energy supplied to the molecules

is so high and the collisions between the particles of gas become so violent that the

atoms are stripped of one or more electrons, which then move around freely. This

characteristic of plasma, to separate electrons from atoms, produces very strong electric

currents. Plasmas are therefore gases that have been heated to very high temperatures

and that are very rich in ions and free electrons.

Most of the universe is composed of plasma. Stars such as the Sun and even outer

space consist of plasma. The nuclear fusion reaction which occurs inside the Sun is

basically the same as the one we try to reproduce in fusion nuclear reactors.5

Hydro-Québec and its partners have been studying plasmas for a long time, with the

aim of eventually using fusion to produce nuclear energy. In the Tokamak de Varennes

nuclear fusion reactor near Sorel, the fusion reaction involves the use of plasma

composed of the two heavy hydrogen isotopes, deuterium (2H) and tritium (3H),

confined in a magnetic field.

AMORPHOUS SOLIDS AND LIQUID CRYSTAL

We have seen that solids have a definite shape, that liquids flow and take the shape

of their containers and that gases are compressible and expandable and have no definite

shape. Although the distinction between these three states may seem clear-cut, there

are cases where it is not.


1.52

5. The design of nuclear fusion reactors is not yet fully refined, unlike fission reactors which are used by many countriesto produce electricity.

For example, glass is considered a solid because at first glance it appears to have a

definite shape. However, over time, glass can become deformed and flows like a liquid.

Glass “slides” can be seen in old windows that have been in place for many years.

These changes in the glass distort the view when you look through the window. Glass

is one of those solid substances that share some of the properties of a liquid. They

are called amorphous solids. The molecules of these solids do not have the ordered

structure that normally characterizes a solid. Amorphous solids are more like

“frozen” liquids. Some familiar substances in this category are rubber, certain plastics,

asphalt, tar, volcanic rock, meteorites and glass.

While amorphous solids are solids that resemble liquids, liquid crystal is a liquid

substance that has some of the properties of a solid. The molecules in liquid crystal

are arranged in an ordered structure just as they are in a solid crystal.

Liquid crystal was discovered in 1888, but it remained a laboratory curiosity for more

than 30 years. The most spectacular application of liquid crystal is its utilization in

passive displays since it responds rapidly and reversibly to an electric field (see

Figure 1.25). A display cell contains a layer of liquid crystal between two glass plates

that have been slightly metallicized in order to make them conductive. When power

is supplied to the cell containing the liquid crystal, the molecules line up in the direction

of the current and the cell changes appearance—for example, the display changes

from clear to dark. If the power is switched off, the molecules return to their initial

position. The mobility of the liquid crystal is due to its liquid properties. To create

the letters and numbers that appear in the displays, small electric voltages are set up

across liquid crystal. This is how watch and calculator displays work and, in a more

sophisticated way, how the flat-screen displays of portable computers and televisions

work.


1.53

Figure 1.25 - Liquid crystal

The passive display on some wristwatches is achieved thanks to liquid crystal.

This chapter is drawing to a close. In it we covered the three major states of

matter that are by far the most commonly encountered in everyday life. The

rest of this course will be devoted to an in-depth study of the gaseous state of

matter and the behaviour of certain gases. In the next chapter we will look at

gases found in nature, their origin and their cycles as well as the broad range

of uses that humans make of gases.

Liquid crystal

Unenergized cell

Energized cell


1.54

Amorphous solid

Boiling temperature Brownian motion(boiling point)

Compressibility Crystallization

Diffusion

Expansibility

Gas

Ideal gas

Kinetic energy Kinetic theory of gases

Liquefaction LiquidLiquid crystal

Melting Melting temperature Model (melting point)

Phase change Plasma

Rotation

Solid SolidificationState Sublimation

Translation

Vapourization Vibration

A substance can exist in three forms: solid, liquid or gas. These are called the three

states of matter. The properties of each state can be explained using a model, designed

to describe what happens at the molecular level. The table below sums up the observable

(visible) properties of each state of matter and the characteristics of the model that

explains these properties. In most cases, the particles are molecules or atoms.

SUMMARY

KEY WORDS IN THIS CHAPTER


1.55

SUMMARY

KEY WORDS IN

THIS CHAPTER

The kinetic theory of gases, which is currently being used by scientists, describes

the model of ideal gases, and is consistent with the model used in the above table.

This theory can be summarized in five points:

• a gas is composed of particles: atoms or molecules;

• the particles are very far apart;

• the particles of a gas are in continuous motion;

• the particles are independent of one another and there is no attraction between

them;

• the kinetic energy of the particles is a function of their temperature.

GAS

Indefinite shape

Indefinite volume

Compressible

Particles relatively far apart

Vibration, rotation and translation movements

Independent particles (little or no attraction

for each other)

Particles randomly arranged

LIQUID

VISIBLE PROPERTIES

Indefinite shape

Defined volume

Very slight compressibility

MODEL

Particles tightly packed

Vibration and rotation movements

Strong attraction between the particles

Particles randomly arranged

SOLID

Defined shape

Defined volume

Resists compression

Particles tightly packed

Vibration movement

Very strong attraction between the particles

Particles arranged in an ordered structure


1.56

rotationtranslation

particles

The late 18th century marked a turning point in the history of chemistry. Chemistry

became a science in the modern sense of the word and has been experiment-based

ever since. Chemists abandoned obsolete theories in order to make way for the concept

of the atom, the elements (in the current sense of the word) and molecules. The study

of gases played a major role in the work that brought about this transition, which

was considered a veritable scientific revolution.

Diffusion is the natural tendency of a substance to spread as a result of molecular

movement. Rapid in gases, diffusion occurs slowly in liquids and is almost non-existent

in solids. Brownian motion refers to the motion of fine particles in a liquid; it is

visible proof of molecular movement in a liquid.

A phase change is a transition from one state to another (see Figure 1.15). The heating

curve of a substance (temperature as a function of time) features two plateaus that

mark the melting and boiling temperatures. During these periods, the temperature

remains constant since the heat provided serves to overcome the attraction between

the molecules. The increase in temperature causes liquids and solids to expand, that

is, to increase slightly in volume. The mercury thermometer is an application of the

expansion of liquids. The boiling temperature of a substance varies according to the

atmospheric pressure. For example, water boils at 71°C on Mount Everest and at 100°C

in Montréal.

The absorption of heat during the vapourization of a liquid is the basic principle of

refrigerators and other cooling appliances. The liquid refrigerant turns into a gas

thereby absorbing heat from the ambient air. It then releases this heat outside where

the appliance turns it back into a liquid.

Certain substances exist in states that are neither solid, liquid nor gaseous. Plasmais a highly ionized gas; amorphous solids appear to be solids, but their structure is

random like that of a liquid; liquid crystal flows like a liquid but its molecules are

ordered like those of a solid crystal.


1.57

Exercise 1.23

Using the model of matter developed in this chapter, explain why:

a) a solid is not compressible.

b) a liquid flows.

c) a gas occupies all the available volume.

Exercise 1. 24

Using the kinetic theory of gases, explain why the pleasant fragrance of a lily or a

rose spreads throughout the room in which it is placed.

REVIEW EXERCICES


1.58

REVIEWEXERCISES

?

?

Exercise 1.25

Which molecular movements are possible:

a) in solids?

b) in liquids?

c) in gases?

Exercise 1.26

Antoine Lavoisier is considered by many to be the founder of modern chemistry.

Explain why.

Exercise 1.27

Gases are everywhere in our lives. Methane (CH4) is the main constituent of natural

gas, one of the most popular fuels. Sulphur dioxide (SO2) and carbon dioxide (CO2)

are two major atmospheric pollutants. Nitrogen (N2) alone represents 80% of the air

we breathe. Helium (He) is a gas used to inflate balloons. Classify the five gases

mentioned in this paragraph by decreasing order of speed of diffusion.


1.59

?

?

?

Exercise 1.28

Compare the speed of diffusion in solids, liquids and gases. Briefly explain why there

is a difference in speed.

Exercise 1.29

The Scottish botanist Robert Brown observed that pollen grains which were

suspended in water moved erratically (Brownian motion). Why did the pollen grains

move?

Exercise 1.30

Identify the phase changes described by the following statements.

a) The melting of gold in the manufacture of gold ingots;

b) The transformation of sap into maple syrup;

c) The freezing of concentrated fruit juices before selling them;

d) Freon that becomes liquid when highly compressed;

e) The transformation of dry ice (CO2) to produce smoke.


1.60

?

?

?

Exercise 1.31

A container of refrigerated water (4ºC) is heated in a microwave oven for a few minutes.

When the container is removed from the oven, the water is boiling and rapidly giving

off steam.

a) Draw an approximate curve of the temperature as a function of time. Identify the

section of the curve that corresponds to boiling.

Heating curve for water

b) What is the boiling temperature of pure water at normal pressure? Indicate the

value at the appropriate place on the graph.

Exercise 1. 32

Refer to the table in Figure 1.20 and identify the state in which the following substances

are found at 100ºC.

a) Margarine:

b) Helium:

c) Aluminum:

Tem

pera

ture

(°C

)

Time (min.)0


1.61

?

?

Exercise 1. 33

Using Figure 1.23, compare the boiling temperature of water in Mexico City and in

Montréal. Explain this difference.

Exercise 1. 34

Freon is a refrigerant used in refrigerators. It circulates in a closed circuit, and goes,

in turn, from the liquid to the gaseous state and vice versa.

a) What phase change does freon undergo when the coil in which it circulates is in

contact with the air trapped in the refrigerator? Why does this phase change occur?

b) In the condenser (Figure 1.24), the gas is liquefied. Is the coil of the condenser in

contact with the air inside the refrigerator? Why?


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?

?

c) What are the two roles of the expansion valve?

Exercise 1.35

How can the use of refrigerators, air conditioners and other cooling appliances harm

the environment?

Exercise 1.36

a) Why is liquid crystal neither solid nor liquid?

b) Why is glass considered an amorphous solid?


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CHE-5041-2 Learning Guide - SOFAD Reactions 1: Energy and Chemical Dynamics Chemical Reactions 2: Equilibrium and Oxidation—Reduction The three learning guides are complemented by

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