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 The
Structure of the IPS 7th Edition Course
Link
to 8th edition structure
IPS has always had a central theme that is not confined within the boundaries
of individual disciplines, primarily chemistry and physics. As stated in
the preface to the text, the central theme is the study of matter leading
to the development of the atomic model. In broad terms, the course divides
naturally into three parts.
Chapters 1-6 provide the empirical framework without which the atomic model
becomes an answer in search of a question. The progression is from what
is around us in the greatest abundance, namely mixtures, to compounds and
elements. In the process, students learn about the characteristic properties
by which substances are recognized and separated. No distinction is made
between physical and chemical properties.
Chapters 7-9 introduce the atomic model. Radioactivity was chosen as the
vehicle because the discreteness in radioactive processes is clearly observable,
and because the subject, despite its importance, is widely neglected. Radioactivity
is one of the topics to which more time was allotted in the Seventh Edition.
Chapters 10-12 add the electric dimension to the atomic model, reinforcing
the material learned earlier. In the process, a valuable foundation is laid
for electrochemistry.
The division of the course along these lines provides natural breaking points
for teachers who wish to spread the IPS course over more than one year.
The Story Line of the Course
IPS has always had a central theme and a clear story line not confined to
a single discipline. As stated in the preface of the text, the central theme
is the study of matter leading to the development of the atomic model. The
story line is best described by the annotated table of contents that follows:
CHAPTER 1 - VOLUME AND MASS
| 1.1 |
Experiment: Heating Baking
Soda |
| 1.2 |
Volume |
| 1.3 |
Reading Scales |
| 1.4 |
Experiment: Measuring Volume
by Displacement of Water |
| 1.5 |
Shortcomings of Volume as
a Measure of Matter |
| 1.6 |
Mass |
| 1.7 |
Experiment: The Equal-Arm
Balance |
| 1.8 |
Experiment: Calibrating the
Balance |
| 1.9 |
Unequal-Arm Balances |
| 1.10 |
Electronic Balances |
| 1.11 |
Experiment: The Sensitivity
of a Balance |
As you introduce a class to IPS, try to set the tone for the entire year
on the first day. The short article "About the Course" (on page xii in the
text) and Experiment 1.1 will help you do that.
The purpose of the article is to alert the student that in IPS "throwing
around big words" is not synonymous with understanding. You may want to
read the article in class either before or after experiment 1.1.
The purpose of the experiment is to raise questions, some of which will
be answered later in this chapter. It also helps in developing laboratory
skills.
The last question in Experiment 1.1 provides motivation for the study of
volume and mass. We begin with volume, showing different methods of measuring
it and the need to say precisely what we mean by the volume of an object.
After pointing out the shortcomings of volume as a measure of the quantity
of matter, we then proceed to mass, which is operationally defined as that
property of matter that can be measured with an equal-arm balance.
The equal-arm balance is introduced to help students understand the basic
operation of a mechanical balance. The unequal-arm balance is introduced
next for schools equipped with this type of balance. For schools that have
electronic balances, the next section describes the use of this instrument.
CHAPTER 2 - MASS CHANGES IN CLOSED SYSTEMS
| 2.1 |
Experiment: The Mass of Dissolved
Salt |
| 2.2 |
Histograms |
| 2.3 |
Using a Computer to Draw
Histograms |
| 2.4 |
Experiment: The Mass of Ice
and Water |
| 2.5 |
Experiment: The Mass of Copper
and Sulfur |
| 2.6 |
Experiment: The Mass of a
Gas |
| 2.7 |
The Conservation of Mass |
| 2.8 |
Laws of Nature |
Although this chapter is among the shorter ones in the text, it is of prime
importance. Interwoven are two objectives: the development of the skills
related to the balance and the analysis of data, and the accumulation of
evidence leading to a fundamental law of nature, the law of conservation
of mass. It will take the entire chapter to reach the objectives.
Histograms, which are introduced in this chapter, will be used throughout
the course. The time you invest in teaching how to construct them will pay
handsome dividends later on. Once the students know how to construct histograms
by hand, we recommend that they use the software to save time and explore
various choices available to them.
Emphasize to students that a single experiment, involving only one kind
of change (such as dissolving salt), is not in itself very convincing evidence
for concluding that mass does not change when other changes take place.
This is why four separate mass-conservation experiments, all involving different
kinds of change, are included in this chapter. Do not omit any of them;
let your students do all of them to convince themselves of the plausibility
of conservation of mass.
CHAPTER 3 - CHARACTERISTIC PROPERTIES
| 3.1 |
Properties of Substances
and Properties of Objects |
| 3.2 |
Experiment: Freezing and
Melting |
| 3.3 |
Graphing |
| 3.4 |
Experiment: Boiling Point |
| 3.5 |
Boiling Point and Air Pressure |
| 3.6 |
Experiment: Mass and Volume |
| 3.7 |
Density |
| 3.8 |
Dividing and Multiplying
Measured Numbers |
| 3.9 |
Experiment: The Density of
Solids |
| 3.10 |
Experiment: The Density of
Liquids |
| 3.11 |
Experiment: The Density of
a Gas |
| 3.12 |
The Range of Densities |
| 3.13 |
Identifying Substances |
In the daily language one hears statements like "lead is heavier than iron."
Of course, lead is neither heavier nor lighter than iron, just as lead is
neither bigger nor smaller than iron. Mass, volume, and shape are properties
of objects. Properties that do not depend on the amount of a substance are
called characteristic properties.
The characteristic properties discussed in this chapter and in Chapter 4
have been selected for their usefulness in identifying substances and separating
mixtures. Hence, we concentrate on freezing point, boiling point, and density
in this chapter, and on solubility in Chapter 4.
CHAPTER 4 - SOLUBILITY
| 4.1 |
Experiment: Dissolving a
Solid in Water |
| 4.2 |
Concentration |
| 4.3 |
Experiment: Comparing the
Concentrations of Saturated Solutions |
| 4.4 |
Experiment: The Effect of
Temperature on Solubility |
| 4.5 |
Wood Alcohol and Grain Alcohol |
| 4.6 |
Experiment: Rubbing Alcohol
as a Solvent |
| 4.7 |
Sulfuric Acid |
| 4.8 |
Experiment: Two Gases |
| 4.9 |
Hydrogen |
| 4.10 |
Carbon Dioxide |
| 4.11 |
Experiment: The Solubility
of Carbon Dioxide |
| 4.12 |
The Solubility of Gases |
| 4.13 |
Acid Rain |
| 4.14 |
Drinking Water |
Solubility is a characteristic property of both the solute and the solvent.
It is expressed in a complex unit-grams of solute per 100 cm3 of solvent.
If we know the solubility of a substance in a given solvent and the quantity
we want to dissolve, we can calculate the minimum amount of solvent necessary.
Or, if we know how much solvent we have, we can use the solubility to find
the maximum amount of the solute we can dissolve in it.
Like density, solubility changes with temperature. However, the solubility
of some substances changes rather dramatically with temperature, whereas
the density of solids or liquids changes only slightly. The dependence of
solubility on temperature is very useful in separating substances in solution.
CHAPTER 5 - THE SEPARATION OF MIXTURES
| 5.1 |
Experiment: Fractional Distillation |
| 5.2 |
Petroleum |
| 5.3 |
The Separation of Insoluble
Solids |
| 5.4 |
Experiment: The Separation
of a Mixture of Solids |
| 5.5 |
The Separation of a Mixture
of Soluble Solids |
| 5.6 |
Experiment: Paper Chromatography |
| 5.7 |
A Mixture of Gases: Nitrogen
and Oxygen |
| 5.8 |
Low Temperatures |
| 5.9 |
Mixtures and Pure Substances |
As mentioned earlier, one of the criteria for selecting characteristic properties
for discussion was their usefulness in separating substances. Now we will
employ these properties for actual separations in the laboratory, describe
some applications of these methods in industry, and arrive at an operational
definition of a pure substance. Reading through this chapter, you may get
the impression that we are leaving students with a rather vague definition
of a pure substance. This is true. The boundary between a mixture and a
pure substance is not as well-defined as some textbooks may lead us to believe.
If your students realize at the end of this chapter that a pure substance
is something that cannot be broken up by any of the methods discussed, they
will have learned a valuable lesson.
CHAPTER 6 - COMPOUNDS AND ELEMENTS
| 6.1 |
Breaking Down Pure Substances |
| 6.2 |
Experiment: The Decomposition
of Water |
| 6.3 |
The Synthesis of Water |
| 6.4 |
Experiment: The Synthesis
of Zinc Chloride |
| 6.5 |
The Law of Constant Proportions |
| 6.6 |
Experiment: A Reaction with
Copper |
| 6.7 |
Experiment: The Separation
of a Mixture of Copper Oxide and Copper |
| 6.8 |
Complete and Incomplete Reactions |
| 6.9 |
Experiment: Precipitating
Copper |
| 6.10 |
Elements |
| 6.11 |
Elements near the Surface
of the Earth |
| 6.12 |
The Production of Iron and
Aluminum |
By definition, pure substances are not broken up into different components
by those separation methods used to separate mixtures. The aim of this chapter
is to show that, in general, pure substances can, nevertheless, be broken
up by other means, such as applying intense heat or an electric current.
Conversely, such pure substances (compounds) can also be synthesized from
other pure substances, but only by reacting in definite proportions.
A brief review of breaking down mercuric oxide and baking soda (part of
Section 5.9 and Experiment 1.1) using heat, leads to the breaking down of
water using electricity (Experiment 6.2). In each case, new pure substances
are produced that are quite different from the original substances. We then
reverse our method of attack and synthesize compounds. The examples used
are chosen to illustrate one of the basic differences between compounds
and mixtures: unlike mixtures, compounds can be synthesized only by reacting
them in definite proportions.
Early difficulties in the formation of the law of constant proportions sprang
in part from the difficulty of determining when a reaction was complete.
The reaction between copper and oxygen (Experiments 6.6 and 6.7) illustrates
this circumstance: The investigation into what has happened leads to an
understanding of complete and incomplete reactions.
Experiment 6.9 ends the sequence of experiments that started with Experiment
6.6 and continued in Experiment 6.7; copper was made to form a series of
pure substances and was then recovered, suggesting that the copper was there
all along. The section leads into the operational definition of an element
(Section 6.10). The reasoning used in the definition of an element is reinforced
with two historical examples. Be sure to spend enough time on this section.
Sections 6.11 and 6.12 balance the preceding discussion of scientific methodology
with a discussion of the abundance of elements near the surface of the earth,
and a description of the industrial processes for producing iron and aluminum.
You can assign these sections for reading, and follow up with a brief class
discussion.
CHAPTER 7 - RADIOACTIVITY
| 7.1 |
Radioactive Elements |
| 7.2 |
Radioactive Decomposition |
| 7.3 |
Experiment: Radioactive Background |
| 7.4 |
Experiment: Collecting Radioactive
Material on a Filter |
| 7.5 |
Experiment: Absorption and
Decay |
| 7.6 |
Radioactivity and Health |
| 7.7 |
A Closer Look at Radioactivity |
The objectives of this chapter are quite modest: learning the basics about
counting radioactive decays and noting the discreteness of the process.
Tying this discreteness to change of one element into another suggests a
particle model of matter. More than that, the counting of radioactive decays
provides us with a direct way of counting the number of atoms in a measurable
sample of an element, thereby providing a means of finding the mass of atoms.
Unlike in other chapters, the three experiments in Chapter 7 are to be done
by the class as a whole rather than by pairs of students. The reason is
simple: it is unlikely that you will have enough Geiger counters. However,
if you have more than one counter, divide the class into smaller groups
and have them work in parallel. The class will have the advantage of seeing
that while the details vary, the general trend is the same.
CHAPTER 8 - THE ATOMIC MODEL OF MATTER
| 8.1 |
A Model |
| 8.2 |
Experiment: A Black Box |
| 8.3 |
The Atomic Model of Matter |
| 8.4 |
"Experiment": Constant Composition
Using Fasteners and Rings |
| 8.5 |
Molecules |
| 8.6 |
Experiment: Flame Tests of
Some Elements |
| 8.7 |
Experiment: Spectra of Some
Elements |
| 8.8 |
Spectral Analysis |
| 8.9 |
"Experiment": An Analog for
Radioactive Decay |
| 8.10 |
Half-Life |
We now introduce the atomic model of matter, which will continue to be at
the center of our attention through Chapters 9, 11, and 12.
After a brief introduction to the meaning of a "model," the class applies
the idea to a Black Box, which provides an opportunity to make testable
predictions (Experiment 8.2).
Sections 8.3-8.6 sum up key observations made earlier in the course in the
context of the atomic model. The law of Conservation of Mass and the law
of Constant Proportions are given special attention.
The class performs experiments with spectra of atoms and is shown evidence
that the spectra present properties of the individual atoms rather than
properties of the elements in bulk.
Finally, the atomic model is used to predict the existence of a half-life
for radioactive elements.
CHAPTER 9 - THE SIZES AND MASSES OF MOLECULES AND ATOMS
| 9.1 |
The Thickness of a Thin Layer |
| 9.2 |
Experiment: The Thickness
of a Thin Sheet of Metal |
| 9.3 |
Scientific Notation |
| 9.4 |
Multiplying and Dividing
in Scientific Notation: Significant Digits |
| 9.5 |
Experiment: The Size and
Mass of an Oleic Acid Molecule |
| 9.6 |
The Mass of Helium Atoms |
| 9.7 |
The Mass of Polonium Atoms |
| 9.8 |
The Size of Atoms |
This chapter greatly strengthens the atomic model of matter introduced in
Chapter 8. The ability to ascribe a mass and a size to atoms in effect clinches
the argument for the acceptance of the model.
The chapter does make heavier demands on your students' mathematical skills
than other chapters in the book. Even though the steps of the experiments
are conceptually very simple, it takes a relatively long chain of operations
with powers of 10 to achieve the desired results. The details should be
treated only with a class that can follow the rather complex arithmetic
without losing sight of the physical content. With students who are weak
in mathematics, it may be advisable to treat the chapter lightly. Carry
them through a limited number of calculations on the chalkboard, with the
main goal being a basic understanding of the method used to find the masses
and sizes of atoms.
CHAPTER 10 - ELECTRIC CHARGE
| 10.1 |
Introduction |
| 10.2 |
A Measure for the Quantity
of Charge |
| 10.3 |
Experiment: Hydrogen Cells
and Light Bulbs |
| 10.4 |
Experiment: Flow of a Charge
at Different Points in a Circuit |
| 10.5 |
The Conservation of Electric
Charge |
| 10.6 |
The Effect of the Charge
Meter on the Circuit |
| 10.7 |
Charge, Current, and Time |
| 10.8 |
Experiment: Measuring Charge
with an Ammeter and a Clock |
In this chapter, we introduce a model of electric charge flow, the law of
Conservation of Charge, and two methods for measuring charge.
The emphasis throughout this chapter and the next is on electric charge
rather than electric current. It is important to keep this in mind, even
though from the end of this chapter on, we shall measure moving charge indirectly
by measuring current and time.
CHAPTER 11 - ATOMS AND ELECTRIC CHARGE
| 11.1 |
The Charge per Atom of Hydrogen
and Oxygen |
| 11.2 |
Experiment: The Electroplating
of Zinc |
| 11.3 |
The Elementary Charge |
| 11.4 |
The Elementary Charge and
the Law of Constant Proportions |
| 11.5 |
Experiment: Two Compounds
of Copper |
| 11.6 |
The Law of Multiple Proportions |
At the end of the preceding chapter, we established a method of measuring
electric charge with an ammeter and a clock. In this chapter, we use this
method to find the quantity of charge needed to plate out a single atom
of an element from a solution. The comparison of these charges will lead
us to the existence of a natural unit of charge, the elementary charge.
From this, the idea of "atoms" of electricity can be related very directly
to the law of constant proportions studied in Chapters 6 and 8.
CHAPTER 12 - CELLS AND CHARGE CARRIERS
| 12.1 |
Experiment: The Daniell Cell |
| 12.2 |
Experiment: Zinc and Copper
in Different Solutions |
| 12.3 |
Flashlight Cells |
| 12.4 |
Unintentional Cells and Corrosion |
| 12.5 |
The Motion of Electric Charge
Through a Vacuum |
| 12.6 |
Electrons |
| 12.7 |
Atoms and Ions |
| 12.8 |
The Motion of Charge Through
a Circuit |
| 12.9 |
The Direction of Electric
Current |
The first topic in this chapter is the Daniell cell and an investigation
of The basic reactions that make it work: dissolving zinc and plating out
copper. We then demonstrate that similar reactions occur in other cells,
both desirable ones and undesirable ones such as those causing corrosion.
By now we have added a great deal to the phenomenological knowledge of the
students. We have shown how to use that knowledge to expand the atomic model
of matter to relate charge per atom to constant composition, multiple proportions,
and simplest formulas. However, we have not shown how to connect the phenomena
involving electric charge with the mechanism of compound-forming (chemical)
properties of atoms. To do that we need electrons.
We use the passage of charge through a vacuum tube to introduce electrons.
Using the vacuum tube to make the existence of electrons plausible corresponds
to using radioactivity to make the existence of atoms plausible. Electrons
and atoms are then related through the introduction of positive and negative
ions to account for the movement of charge through a solution. |
