Part 1 - Properties of Matter

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.

Part 2 - Atoms and Molecules

Chapters 7–11 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. The motion of atoms and molecules is added to the model. Simple methods are used to determine the size and masses of molecules and atoms. The periodic table is examined through a historical perspective.

Part 3 - Energy and Forces

Chapters 12–16 introduce thermal energy as the basis for recognizing and measuring gravitational potential, elastic potential energy, and kinetic energy. The first two of these chapters investigate changes of energy from one form to another. Students quantitatively measure the change from gravitational potential energy to thermal energy, leading to the development of the law of conservation of energy. Three chapters on force and motion examine types of forces, force as a vector quantity, and the effects force on the motion of objects. In these chapters, students develop evidence for Newton’s laws of motion.

The division of the book into three sections provides natural breaking points for teachers who wish to spread the IPS course over more than one year or change the order of the sections.

1.1 Experiment: Heating Baking Soda
1.2 Volume
1.3 Reading Scales
1.4 Experiment: Measuring Volume by Displacement of Water
1.5 Limitations of Volume as a Measure of Matter
1.6 Mass: The Equal-Arm Balance
1.7 Single-Pan and Electronic Balances
1.8 Experiment: The Sensitivity of a Balance

The short note �To the Student� (on page xiii in the text) and Experiment 1.1, Heating Baking Soda and the short note �To the Student� (on page xiii in the text) set the tone for the entire year on the first day. 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 question of how to compare amounts of solids, liquids, and gases at the end of 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.

Experiment 1.8, in which students determine the sensitivity of a balance, has been written so 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 importance. Interwoven within it are two objectives: the development of 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 these objectives.

Histograms, which are introduced in this chapter, are used throughout the course. Time invested now in teaching students how to construct histograms will pay handsome dividends later on. Once students know how to construct histograms by hand, we recommend that they use KaleidaGraph software to save time and explore various choices available to them.

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.

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 Distinguishing 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 and Global Warming
4.10 Drinking Water

Solubility is a characteristic property of both the solute and the solvent. It is expressed as 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 needed. Or, if we know how much solvent we have, we can use the solubility to find the maximum amount of the solute that will 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 Mixtures Involving Gases
5.8 Mixtures and Pure Substances

One of the criteria for selecting characteristic properties for discussion is 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 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 Decomposing 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 cannot be 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, some pure substances (compounds) can also be synthesized from other pure substances, but only by reacting in definite proportions.

After describing the decomposition of two pure substances by heating, we use electrolysis to decompose 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 formulation of the law of constant proportions resulted, in part, from difficulty in determining when a reaction was complete. The reaction between copper and oxygen (Experiments 6.6 and 6.7) illustrates this. 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 used to form a series of pure substances and was then recovered, suggesting that the copper was there all along. This experiment leads to the operational definition of elements (Section 6.10). The reasoning used in the definition of an element is reinforced with two historical examples.

Section 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 provides an excellent example of the surprises that nature has in store for us: Now that students have learned how elements survive in the formation of compounds, they find that some elements actually change into other elements all on their own.

(ii) This change takes place in discrete steps, which can be counted.

(iii) Taken together, (i) and (ii) provide evidence for the atomic model of matter and lead to a testable prediction in 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 have gained in this course (Chapter 10).

(v) Knowing the mass of atoms leads to the introduction of the periodic table in a meaningful way (Chapter 10).

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

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

Unlike those 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 Constant Proportions and the Atomic Model
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 11. 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 display 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 Molecular Motion and Diffusion
9.2 Number of Molecules and Pressure of a Gas
9.3 A Prediction About the Relation Between Volume and Pressure of Gases
9.4 The Compressibility of Gases
9.5 Temperature and Molecular Speed
9.6 Avogadro’s Law
9.7 Masses of Atoms and Molecules
9.8 Behavior of Gases at High Pressures

The study of molecular motion involves the effect of the average motion of a large number of molecules. The constant pressure exerted by a gas on the walls of a container is an example of average behavior. The relationship among the number of molecules, pressure temperature and volume is explored through the use of the sphere gas machine. The fact that the compressibility, density and thermal expansion of a gas do not depend of what the gas is provides the basis for introducing Avogadro’s hypothesis—equal volumes of gas at the same pressure and temperature contain an equal number of molecules.

10.1 A Historical Sketch
10.2 Some Families of Elements
10.3 Activity: Atomic Mass and Other Properties of Atoms
10.4 The Elements in the Third Through Sixth Columns
10.5 Activity: The Elements in the Fourth Row
10.6 The Fourth and Fifth Rows: A Historical Perspective

Chapter 10 ties together many observations on the macroscopic and atomic levels. The chapter begins where Chapter 6 left off by discussing the discovery of a large number of elements. The similarity of properties of groups of elements (Section 10.2) introduces the idea of “families” of elements. Asking if there is any relationship between atomic mass and the properties of the elements leads us to the question of how to classify elements (Section 10.3)

To make the point that any classification is a matter of judgment, students do an activity with a set of 24 specially prepared cards. 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 10.4 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 shows its success by highlighting Mendeleev’s correct prediction of the properties of germanium before the element was discovered.

11.1 The Thickness of a Thin Layer
11.2 Experiment: The Thickness of a Thin Sheet of Metal
11.3 Experiment: The Size and Mass of an Oleic Acid Molecule
11.4 The Mass of Helium Atoms
11.5 The Mass of Polonium Atoms

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 that question.

We arrive at our goal in stages, by preparing the students for the main sections (Sections 11.4 and 11.5). We first find the thickness of a piece of aluminum foil. We then apply the same approach to find the thickness of a layer of oleic acid (Sections 11.1–11.3). Sections 11.4 and 11.5 parallel the film, “The Mass of Atoms,” which adds a lively dimension to the presentation in the text.

12.1 Introduction
12.2 Experiment: Mixing Warm and Cool Water
12.3 A Unit of Energy: The Joule
12.4 Experiment: Cooling a Warm Solid In Cool Water
12.5 Specific Heats of Different Substances
12.6 Experiment: Melting Ice
12.7 Heat of Fusion and Heat of Vaporization
12.8 Experiment: Heat of Reaction
12.9 Comparing the Energies Involved in Different Reactions

The usefulness of the idea of energy stems from the convertibility of energy into different forms. All forms of energy can be associated with a change in temperature. In this chapter, we define a change in some phenomenon as a change in energy when the original change is associated with a change in temperature.

Experiment 12.2 helps emphasize the difference between temperature and thermal energy. The results of the experiments are generalized in Sections 12.3 to 12.5, where specific heat is introduced. Melting ice in a calorimeter (Experiment 12.6) leads into the heat of fusion and heat of vaporization (Section 12.7). This is followed by the heat of reaction in Section 12.8.

13.1 Experiment: Heating Produced by a Slowly Falling Object
13.2 Gravitational Potential Energy
13.3 Kinetic Energy
13.4 Kinetic Energy as a Function of Speed
13.5 Experiment: Changing Gravitational Potential Energy to Kinetic Energy
13.6 The Law of Conservation of Energy<

Gravitational potential energy is introduced through the use of a slowly falling weight to raise the temperature of an aluminum cylinder. The dependence of the temperature rise on the weight of the falling body for a fixed distance is part of the experiment (Experiment 13.1); the dependence of the height is discussed in Section 13.2. The results are generalized in an end-of-chapter problem to show the change in gravitational potential energy depends only on the change in height and not the distance traveled along a slope.

The increase in temperature generated by stopping a wheel with a heavy rim is described in detail leading the definition of kinetic energy (Section 13.3 and 13.4). Once both gravitational potential energy and kinetic energy have been defined, the energy conversion between the two is measured in Experiment 13.5. The chapter is summed up with a discussion of the law of conservation of energy.

14.1 Introduction
14.2 Weight: The Gravitational Force
14.3 Activity: The Elastic Force: Hooke’s Law
14.4 Experiment: The Magnetic Force
14.5 Experiment: Sliding Friction
14.6 Friction and Weight
14.7 Newton’s Third Law

Although “force” is a common word, its use in science is specific and quantitative. This idea is made clear in the brief introduction and in Experiment 14.2, in which a spring is calibrated in arbitrary weight units to introduce the newton spring scale. The extension of the spring as a function of weight (Hooke’s Law) is used to develop the idea of proportionality. Further development of this topic appears in Appendix 2.

Experiment 14.4, in which the magnetic force between two magnets is measured as a function of the separation between the closest poles, introduces the idea of dependence of force on distance. Because friction plays a prominent role in daily life, we show that friction does not act on an object at rest unless there is also another force acting on the object. Students investigate the minimum force needed to keep a body moving under a variety of conditions (Experiment 14.5). This experiment, taken with Section 14.6, makes it evident that friction depends on weight, not weight per unit area. Finally, Newton’s third law is discussed for static conditions.

15.1 Balanced Forces on a Line
15.2 Representing Forces in a Plane
15.3 Experiment: Balanced Forces in a Plane
15.4 The Net Force
15.5 Forces and Their Components
15.6 Experiment: Forces Acting on Moving Objects
15.7 Newton's First Law

Through experience, students know that the effect of a force depends on its strength and direction. Since restricting force to a single line is quite artificial, we introduce vectors to represent forces. The usefulness of this representation is made clear in Experiment 15.3 in which students balance forces exerted by three spring scales in a plane. The relationship between the combination of two forces and a third force that balances them is discussed in Section 15.4.

A common misconception is that moving bodies move in the direction of the force acting on them. Students find out that this is not the case by blowing in various directions on a moving, low friction puck (Experiment 15.6). The chapter concludes with a review that includes the formulation of Newton’s first law.

16.1 Experiment: The Motion Detector
16.2 Velocity Graphs
16.3 Experiment: Motion Under a Constant Net Force: The Effect of Time
16.4 Experiment: Motion Under a Constant Net Force: The Effect of the Magnitude of the Force
16.5 Free Fall and the Effect of Mass: A Prediction
16.6 Experiment: Testing a Prediction: The Relation Between Mass and Change in Velocity
16.7 Newton's Second Law

The motion detector is introduced a device that, when coupled with a computer, can record the distance of an object as a function of time and produce velocity graphs that are analyzed in Section 16.2. The next two sections examine the motion of an object under a constant force; studying the effect of duration of the force in Experiment 16.3 and the magnitude of the force in Experiment 16. 4. The effect of mass on the velocity of an object during free fall is discussed in Section 16.5 and a prediction made which is investigated in Experiment 16.6. The chapter concludes with the formulation of Newton’s second law.

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.

This appendix amplifies the understanding and use of proportionality required in several chapters.