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.

Chapters 7-10 use the discreteness of radioactive processes to motivate the development of the atomic model. The atomic model is shown to bring order to known facts and to allow us to make testable predictions. Simple methods are used to determine the sizes and masses of molecules and atoms. The periodic table is examined through a historical perspective and serves as a culmination of the course.

1.1 Experiment: Heating Baking Soda
1.2 Volume
1.3 Reading Scales
1.4 Experiment: Measuring Volume by Displacement of Water
1.5 Mass: The Equal-Arm Balance
1.6 Unequal-Arm Balances
1.7 Electronic Balances
1.8 Experiment: The Sensitivity of a Balance

As you introduce a new class to IPS, try to set the tone for the entire year on the first day. The short note “To the Student” (on page x in the text) and Experiment 1.1, Heating Baking Soda, will help you do that.
The purpose of the note is to alert students to the interplay of the three forms of active learning in the course: experimenting, reading, and solving problems. You may wish to read the note 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 provides the first steps in developing laboratory skills.

The last question in Experiment 1.1 serves as a 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 is measured with an equal-arm balance. However, in practice, the course no longer requires the use of the equal-arm balance.

Determining the sensitivity of a balance (Experiment 1.8) is written in such a way that it can be done with equal-arm balances, unequal-arm balances, or electronic balances.

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 of 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 KaleidaGraph 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 skip 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: Mass and Volume
3.3 Density
3.4 Dividing and Multiplying Measured Numbers
3.5 Experiment: The Density of Solids
3.6 Experiment: The Density of Liquids
3.7 Experiment: The Density of a Gas
3.8 The Range of Densities
3.9 Experiment: Freezing and Melting
3.10 Graphing
3.11 Experiment: Boiling Point
3.12 Boiling Point and Air Pressure
3.13 Identifying Substances

In 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 density, freezing point, and boiling point 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: Isopropanol as a Solvent
4.7 Experiment: The Solubility of Carbon Dioxide
4.8 The Solubility of Gases
4.9 Acid Rain
4.10 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 Mixtures and Pure Substances

As we 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 so sharp as may be believed from reading some textbooks. 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 their 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

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.

After recalling the decomposition of two pure substances by heating, we use electrolysis to break up water (Experiment 6.2) New pure substances are produced that are quite different from the original substance. 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 the components in definite proportions (Sections 6.3 – 6.5).

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 elements (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 balances the preceding discussion of scientific methodology with a discussion of the abundance of elements near the surface of the earth.

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 A Closer Look at Radioactivity
7.7 Radioactivity and Health

You may wonder why we proceed with the introduction of radioactivity at this point in the course. Here are the reasons:

• (i) It gives an excellent example of the surprises that nature has for us: Just as students learned how elements survive in the formation of compounds they find that some elements change into other elements all on their own.
  
  
• (ii) This change takes place in discrete steps, which can be counted.
  
  
• (iii) The combination of (i) and (ii) provides a motivation for the atomic model of matter and leads to a testable prediction (Chapter 8).
  
  
• (iv) Being able to count radioactive decays enables us to find the number of atoms in a given sample of an element. This, in turn, provides a conceptually simple way to find the mass of single atoms solely with the knowledge students gained in this course (Chapter 9).
  
  
• (v) Knowing the mass of atoms enables us to introduce the periodic table in a meaningful way (Chapter10).
  

In addition, it should be noted that radioactivity is largely ignored in the science curriculum. For many students, learning about radioactivity in IPS may be the only chance to do so.

Randomness, discreteness, and absorption can be demonstrated quite easily in the classroom. However, decay and the existence of a half-life require a source with a short half-life. The only practical way to get such a source is to collect it yourself. You can do this if there is a sufficient concentration of radon in the ground around your school, and if your school has a closed, unventilated room in the basement in which radioactive material can be collected from the air.

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 Chapter 10.

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 8.3–8.5 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.

In Sections 8.6 and 8.7, the class 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 Experiment: The Size and Mass of an Oleic Acid Molecule
9.4 The Mass of Helium Atoms
9.5 The Mass of Polonium Atoms
9.6 Atomic Masses and Molecular Formulas

One of the key ingredients of the atomic model introduced in Chapter 8 was that atoms are very light and small, or, equivalently, that there are many atoms in any sample of an element of measurable mass. Building the atomic model on this premise demands an answer to the question “What is the mass of a single atom of an element?” In this chapter we answer the question.

We arrive at our goal in stages, by preparing the students for the main experiment (Sections 9.4 and 9.5). We first find the thickness of an aluminum foil. We then apply the same approach to find the thickness of a layer of oleic acid (Sections 9.1–9.3).

Sections 9.4 and 9.5 run parallel to the film, “The Mass of Atoms,” which is now available in VHS and DVD formats and adds a lively dimension to the presentation in the text. Finally, the knowledge of atomic masses is applied to finding molecular formulas (Section 9.6).

10.1 Introduction
10.2 Classifying Elements
10.3 The Extraction of Similar Elements from Similar Compounds
10.4 Alkali Metals, Alkaline Earth Metals and Halogens
10.5 Activity: Atomic Mass and Other Properties of Atoms
10.6 The Elements in the Third Through Sixth Columns
10.7 Activity: Elements in the Fourth Row
10.8 The Fourth and Fifth Rows: A Historical Perspective

Chapter 10, which is new to the Eighth Edition, provides a climax to the year’s work by tying together many observations on the macroscopic and atomic levels. We begin by asking if there is any relation between atomic mass and various properties of the elements (Section 10.1). This leads us to the question of how to classify elements (Sections 10.2–10.4). Some of the historical comments in these sections relate directly to similar comment made in Chapter 6.

To drive home the point that any classification is a matter of judgment, we have the students do an activity with a set of 24 specially prepared cards (Section 10.5). The cards resemble entries in a periodic table of the elements. The activity has historical connotations, and raises the question of the order of potassium and argon. The activity is briefly extended in Section 10.7 to show the need for additional columns in the periodic table.

The chapter ends with an analysis of the periodic table as a model, and showing its success by highlighting the correct prediction of the properties of germanium by Mendeleev before the element was discovered.

Part 1 Scientific Notation
Part 2 Multiplying and Dividing in Scientific Notation: Significant Digits

This appendix provides instruction and practice for students who need to improve their skill calculating with numbers in scientific notation.