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 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.

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.

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.

 

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.

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.

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.

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.

 

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.

 

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.

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.

 

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.

 

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.

 

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.